Volume 91, 1982
CONTENTS . (Chemical Sciences)
INORGANIC AND ANALYTICAL
Complexes of Cu(II), Ni(II) and Co(II) with isophthalic diliydrazide 1
C G Kumbhar * 1-6
Amorphous MoS 3 and A^MoSg (A = Li or Na : < x < 4)
T Murugesan and J Gopalakrishnan 7-13
Synthesis of l,5-disubstituted-2,4-dithiobiuret and l,5-disubstituted-2-
thiobiuret and their vanadyl (V) chloride complexes
K P Srivastava and I K Jain 15-19
Precipitation of uranium quinolin-8-olate from homogeneous solution
by urea hydrolysis
G Siva Reddy, A Varada Reddy and Y Krishna Reddy 21-30
Electrochemical studies on copper(II) glucuronate
R Payne and R J Magee 31-37
Molecular constants of PSF 8 and NSF 3
A Natarajan and S Somasundaram 39-45
Carbon-] 3 nuclear magnetic resonance studies on high spin iron
(III) porphyrins D V Behere and S Mitra 145-150
Study of mixed complexes by polarography : cadmium-glycine-methio-
nine and cadmium-glycine-ethylene diamine complexes
M Ramaiah, B G Bhat and R Sundaresan 151-156
In vitro antimicrobial activity studies on the mixed ligand complexes of
Hg(II) with 8-hydroxyquinoline and salicyclic acids Y Anjaneyulu,
R Prabhakar Rao, R Y Swamy, A Eknath and K Narasimha Rao 157-163
Synthesis and structural studies on Ni(II) chloride complexes of
N, N'-(substituted) formamidino-N'-(substituted) carbamides and
thiocarbamides K L Madhok 165-171
Magnetic susceptibility studies of Mu^,. M,O (M = Zn, Mg, Fe)
C E Deshpande, PP Bakare, M N S Murthy, NY Vasanthacharya
and P Ganguly 261-266
Liquid structure of vanadium tetrachloride from neutron diffraction
study R V Gopala Rao and B M Satpathy 267-277
tungsten(VI) complexes G S Sodhi, H S Sangari and N K Kaushik 289-294
Synthesis of some zero-valent complexes of iron via aryldiazenato
cationic complexes S Vancheesan 343-346
Spectrophotometric study of the formation of adducts between U(TTA) 4
and .some neutral organo sulphoxide donors
A Ramanujam, N M Gudi, M N Nadkarni, S K Patil andVV Ramakrishna 347-350
Chelation ion-exchange properties of salicylic acid-urea-formaldehyde
copolymers R M Joshi and M M Patel 351-358
Synthesis, identification and analytical properties of the 5,5'-methylene
disalicylhydroxamic acid (MEDSHA)
L F Capitan-Vallvey, F Salinas and DGazquez 399-407
Polar ographic study of Eu(III)-hydroxamic acid complexes
P K Bhansali and B Nemade 409-413
Redox reactions in non-aqueous media : determination of hydrazine
and its organic derivatives with lead (IV) acetate Balbir Chand Verma,
Saroj Chauhan, Jagmohan Entail and Rajnish Kumar Sood 415-419
Electrochemical studies on cobalt tris (acetylacetone)
R J Magee and Bookhari Annuar 491-498
Spectrophotometric studies on the formation of adducts in the extraction
of U (VI) by mixtures of HTTA with some neutral extractants
V V Ramakrishna, R Swamp and S K Patil 499-506
Syntheses and characterisation of ruthenium carbonyl clusters containing
phosphorous and arsenic bridging ligands K Natarajan and G Huttner 507-511
Spectral and magnetic studies of metal thiocyanate "complexes with
N-substituted thioureas S B Kokatnur and A S R Murty 513-519
ORGANIC
Synthesis and characterization of copolymers from 4-halo (chloro,
bromo) salicylic acid tfasmukh S Patel and Shanti R Patel 173-183
Micellar catalysed chlorination of acetophenone by chloramine-T
V Raghunathan, P S Raghavan, K Vaidyanathan and V S Srinivasan 241-246
Reactions of indoles with mercury(II) salts
Avijit Banerji and Manjusha Sarkar (nee Chaudhuri) 247-254
Quantitative structure activity relationships Part V. Release and
The study of cis- and trans-2 butene using mass spectrometry
Ezzat T M Selim 359-369
Kinetics of chlorination of ketones by 1-chlorobenzotriazole
P S Jayaraman, S Sundaram and N Venkatasubramanian 421-436
Isocoumarins : Part 7. New synthesis of 3-phenyl-3, 4-dihydroiso- .
coumarins and their antifungal activity B H Bhide and R C Shah 437-440
\A general synthesis of thiazoles. Part 4. Synthesis of 5-acyl-2, 4-
diaminothiazoles
S Rajappa, V Sudarsanam, V G Yadav and B V Gaikwad 441-444
r
A general synthesis of thiazoles. Part 5. Synthesis of 5-acyl-2-dialkyl-
aminothiazoles S Rajappa, V Sudarsanam, B G Advani and A V Rane 445-450
A general synthesis of thiazoles. Part 6. Synthesis of 2-amino-5-
heterylthiazoles S Rajappa, V Sudarsanam and V G Yadav 451-455
A general synthesis of thiazoles. Part 7. Direct generation of chloro-
methyl 5-thiazolyl ketones
S Rajappa, V Sudarsanam and R Sreenivasan 457-461
Nitroenamines. Part 9. The enaminic reactivity of 2-nitroinethylene-
thiazolidine S Rajappa and B G Advani 463-466
Tl(III) acetate oxidation of cyclanols and bicyclo (2,2,1) heptan-2-ols
Vangalur S Srinivasan and N Venkatasubramanian 467-473
On the mechanism of chiral aldol cyclization reaction
A Sarkar, H R Y Jois, T R Kasturi and D Dasgupta 475-481
Synthesis of 2-aryl-3-(3-disubstituted aminomethyl-2-thio-4-oxo thiazoli-
din-5yl)-methylenyl-indoles as CNS active agents
Chapla Chaudhary, Rajesh Agarwal and V S Misra 483-489
PHYSICAL AND THEORETICAL
High resolution electron microscopy of chloritoid minerals from diffe-
rent geological melieu G N Subbanna and G V Anantha Iyer 47-56
Strengths of some N-H...7E type of hydrogen bonds
G V G Krishna Mwthy and B Subrahmanyam 57-63
Preparation and molecular configurations of some salts of dipicrylamine
with organic and inorganic cations
M L Kundu. J N Kavoor and S K Ghosh 65-72
iv Intents
Periodic precipitation of-cobalt(II) oxiuate in'agar gel : Effect of parasitic
electrolytes on flocculation
TV Kanniah, S Ambrose, F D Gnanam and P Ramasamy 87-
Photoelectron spectroscopic studies of the adsorption of organic mole-
cules with lone pair orbitals on transition metal surfaces
& Yashonath, P K Bam, A Srinivasan, M S Hegde and CNR Rao 101-1
MNDO study of reaction paths : Hydroboration of carbonyl systems
Nab a K Ray and Rita Chadha 129-
Infrared and Raman spectra and thermodynamic functions of 4-methoxy
pyridine N-oxide K C Medhi 137-}
Dynamic distortion of QN '--skeleton in (CH 3 ) 4 NC1
H D Bist, Mahendra Pal, G S Raghuvanshi and V N Saiin 185-1
An ab initio molecular orbital study of thymine radicals
U Chandra Singh and A Muralikrishna Rao 193-2
EXAFS studies of cobalt oxides and oxide glasses
R Parthasarathy, Ravi V Prasad, P R Sarode and K J Rao 201-2
Electrical conductivity studies in sulphate glasses and the mixed alkali
effect H G K Sundar and K J Rao 207-^
Reactions of co-ordinated ligands : Kinetics and mechanisms in the
charge transfer interaction between dichloro or diaquo triethylene tetra-
mine Co(III) and ferrocyanide
G Visalakshi and K S Venkateswarlu 213-^
CNDO calculations of N-rnethyl substituted acrylamides
G Ramana Rao 223-:
Red edge excitation and proton association in the excited state of acridine
P Gangola, N B Joshi and D D Pant 229-:
A. theoretical study on the specific interaction of hexafLuorobenzene
with benzene and ^-xylene D V S Jain and F S Nandel 295-'.
Conformational behaviour and vibrational spectra of 3 -methyl 2-butane-
thipl. - s K Nandy and G S Kasfha 303-:
Determination of ionisation constants of nitrobenzidines
S Aravamuthan, C Kalidas and C S Venkatachalam "311-^
Sorption properties of oxides IX: Effect of anions on the sorption of
uranium(VI) on hydrous oxides
H S Mahal, B Venkataramani and K S Venkateswarlu 321-327
Electrochemical reduction of copper(II) galacturonate
Robert Payne and Robert J Magee 329-337
Metal complexes of isonicotinic acid hydrazide
H Sanke Gowda and R Janardhan 339-341
Kinetics of Tl(III) oxidation of hydroxylamine hydrochloride in aqueous
sulphuric acid Vangalur S Srinivasan and N Venkatasubramanian 371-375
Mossbauer studies on ferrous-zinc ferrites prepared by a novel technique
C E Deshpande, S K Date, M P Gupta and M N S Murthy 377-383
Chemical shifts of the x-ray K or L m absorption edges
K S Srivastava, M Husain, Kirti Sinha, Pratibha Gupta,
A K Srivastava, V Kumar and Shiv Singh 385-392
Isentropic compressibilities of ternary systems with 1-alkanol as non-
common component G Rajendra Naidu and P R Naidu 393-397
Optical absorption spectrum of thulium nitrate in solution
S V J Lakshman and C K Jayasankar 521-526
Studies on the electrochemical and thermodynamic behaviour of tin-tin
sulphide electrode in the presence of sulphide ions
Pushpa Sharma and Mukhtar Singh 527-533
Compliant fields for molecular interactions I: Lithium cation with
carbonyl donors A S N Murthy and Shoba Ranganathan 535-546
Complexes of Cu(O), Ni(II) and 0(11) with isophthalic
dihydrazide
C G KUMBHAR
Institute of Armament Technology, Girinagar, Pune 411 025, India
MS received 21 September 1981 ; revised 7 December 1981
Abstract. Complexes of isophlhalic dihydrazide (IPZ) of llic type MCl a (IPZ) H a O
[M - Cu(ll), Ni(ll), Co(ll)] and MSO 4 (1PZ) . H a O [M =- Cu(ll), Ni(ll)] have
been prepared and characterised from elemental analysis, magnetic moment, visible,
1R, and ESR spectra. Based on these data a polymeric octahedral structure has been
assigned to MCI 2 (IPZ) H a O complexes and sulphate bridged four coordinate
polymeric structure for MSO 4 (IPZ) HaO complexes. Thermogravimetric studies
of these complexes show that the thermal stability decreases in the order Ni(ll) >
Co(II) > Cu(II).
Keywords. Isophthalic dihydrazide ; metal chloride complexes; metal sulphate
complexes.
1. Introduction
Benzoyl hydrazide forms complexes with transition (Aggarwal and Narang 1976)
and non-transition metals (Aggarwal and Bahadur 1969 ; Aggarwal and Singh
1969} where it is observed that this ligand coordinates through both }C = O
O
and -^NH a of the -C.NH-NH 2 groups keeping the secondary amino group undis-
turbed. In nrjtal complexes of Co(ll), Ni(ll) and Cn(ll) with oxalic malonic
and terephthalic dihydrazides we have noticed (Kumbhar and Sadasivan 1976)
O
II
that these ligands function as bidentate and different -.0 NH-NH 3 groups of the
same ligand coordinate to different metal centres resulting in the formation of
polymeric complexes. Due to sterically unfavourable positions of hydrazide
groups on benzene ring to form chelates when ligand is bifunctional, isophthalic
dihydrazide (1PZ) is expected to form polynuclear complexes with metal salts
containing both Metal-Nitrogen and Metal-Oxygen bonds simultaneously. In
this paper synthesis and characterization of some metal complexes of isophthaiic
dihydrazide with MC1 2 [M = Cu(Il), Ni(H) and Co(ll)] and MSO 4 [M = Cii(II)
and Ni(ll)] are reported. Thermogravimetric studies of these complexes have also
2. Experimental
Isophthalic acid dihydrazide was synthesised as described in the literature fr
diethyl isophthalate [M.P. 221 C (Obs), 220 C (Lit)], (Voloviskii and Knoroz<
1964). Diethyl isophthalate was prepared after Majumdar and Sharma (19'
Syntheses of rmtal complexes were carried out by adding hot aqueous solutior
ligand to alcoholic solutions of respective metal salts in the molar ratio 1:1.'
coloured complexes which separated immediately were suction filtered, was
with water, alcohol and ether successively and air-dried.
IR spectra in nujol mull were recorded on Perkin Elmer Spectrophotom<
Model 457 using fCBr plates. Electronic spectra were obtained in the solid si
using nujol mull on Systronics Spectrocoiorimeter-103. Magnetic susceptil
ties were determined at room temperature by Gouy method using HgCo(CM
as caUbrant. ESR spectra were recorded on .Varian Microwave Spectrome
in the x-band region (9-1 GHz) using DPPH as internal standard. Thermogrc
metric study was carried out using MOM-BUDAPEST Model by heating
samples at a rate of 10 C per minute.
3. Results and discussions
3-1. Physical properties of complexes
The elemental analysis and physical properties of the complexes are listed in tabl<
The complexes are quite stable in air. They are insoluble in common orgai
solvents and water. However, they decompose in mineral acids and are sparin
soluble in DMF and pyridine. Molecular weights of these complexes could i
be measured as they are insoluble in suitable solvents. Elemental analysis lej
to the stoichiometry MCl 2 (IPZ)-H a O [M = Cu(II),Ni (II) and Co(II)] a
MSO 4 (IPZ).H 2 O[M = Cu(II) and Ni(II)] for the complexes. The associat
of one water molecule is confirmed both from IR spectra and thermogravime
data.
Table 1. Analytical data and physical properties of the complexes.
M
SI. Compound
No.
Colour
Found
Cal.
Found
Cal.
/*eff.
BM
values
1.
CuCl a (IPZ).
H a O
Green
18
00
18-34
19-90
20-46
2-00
2-19
2.
NiClgi
HgO
Light blue
16
90
17-20
21-00
20-76
2-82
2-18
3.
CoCl a
(IPZ).
H 2
Pink
16
90
17-24
19-90
20-73
4-42
4.
CuSO
4 (n>z)
H 2
Blue
17-
00
17-10
8-20
8-61
1-74
2-16
5. NiSO 4 (IPZ).H2O Light blue 15'90 16-01 8'22 8-53 2'98 2-18
valent to one water molecule is also evident in the TG and DTA curves at ~ 1 10 C.
Thus the presence of free water molecule is indicated. Coordinated water would
have lost at relatively higher temperatures. The U-.N-H frequencies are lower
by 50-100 crrr 1 as compared to their respective positions in the parent ligand.
The u-C = O is also shifted to lower region in the complexes t y 35-40 cm" 1 . The
shift to lower wave numbers of the arnino andcarbonyl vibration bands in all the
complexes is a clear indication of the involvement of both -NH a and >C = O
in coordination. The possibility of enol form is not indicated in the 1R spectra
of the complexes. The new strong band appearing in the region 1200- 1230 cm- 1
has been assigned to v-(C-O) (Aggarwal et al 1976). The medium to weak band
in the region 550-^00 cm- 1 is assigned to u-M-O following Adams (1967) who
has shown that u->M-O stretching frequency in the metal carbonyl complexes
occurs around 600 cm- 1 . It is observed that u-M-O is higher for MSO 4 complexes
than for MC1 2 complexes.
Electronic spectra of MC1 2 complexes give evidence for distorted octahedral
symmetry around metal centre. Thus CuCl 2 complex shows broad bands around
640 nm and 560 nm which are characteristic bands for octahedral environment
around Cu(II) (Billing and Underfill! 1968;. Mahapatra and Rama Rao 1971).
Ni(II) complex shows broad bands at 660nm and 580 nm and weak bands, at
520 nm and 490 nm. Manch and Fernelius (1961) have also observed such bands
for octahedral Ni(H) in the region 700-O50 nm. CoCl 2 complex shows broad
bands at 650 nm, 625 nm and 505 nm which are characterstic bands for octa-
hedral environment around Co([I) ion (Aggarwal et al 1976). Magnetic moment
values (table 1) of the complexes of CuCl 2 (2-OOB.M.), NiCl a (2-82 B.M.) and
CoCl 2 (4-42 B.M.) show one, two and three unpaired electrons respectively and
are well within the range required for octahedral environment around metal centre
(Figgis and Lewis 196.0).
The average 'g' value for CuCl 2 complex (g-2- 19) calculated from ESR spectra
of polycrystalline powder samples is comparable to those reported for Cu(Il)
complexes, with octahedral symmetry (Sadasivan and Arora 1976). Similarly 'g'
value for NiCl a complex (g = 2-18) is also consistent with those observed for
Ni(H) complexes having octahedral symmetry (Arora and KLumbhar 1977).
Thus in the MCl a -L-H a O complexes the stoichiometry and bidentate nature of
ligand suggest a polymeric structure as shown in figure la. The stoichiometry
and insolubility of the compounds in common organic solvents support this
contention.
Although electronic spectra for CuSO 4 (lPZ).H a O complex ex?ubit broad band
at 550 nm consistent with tetragonally distorted octahedral structure, the effective
magnetic moment (1-74 B.M.) value is comparable to the values reported for
square planar and tetrahedral copper sulphate complexes (Saconi and Ciapolini
1964 ; Beadle et al 1969). The tetrahedral Cu(II) complexes have the magnetic
' moment values within the range 1-89-1-92 B.M. and square planar Cu(II) in the
range 1 -83-.1-86 (Saconi 1966). The observed magnetic moment value for CuSO 4
(IPZ)-HaO is more close to square planar symmetry. The splitting of the strong
HH,
(a)
c
HjH
.NN.
"NN'
HH-j
(b)
Figure 1. (a) Proposed structure for MC1 3 complexes [M = Cu(II), Ni(II), Co{
(b) Proposed structure for [MSO 4 complexes M = Cu(II), Ni(II)].
sulphate bands in the infrared spectrum (bands at 1175cm- 1 , 1130 cm- 1 , 1050 cfl
is consistent with bidentate bridged SO 4 -~ion (Nakamoto et al 1957). Henc
square planar polymeric structure is suggested for CuSO 4 (lPZ)-H a O consist
of bidentate ligand and bridging sulphate ion as shown in figure It.
NiSQi. L-H 2 O shows broad band at 580 nm which is characteristic for tcl
hedral nickel complexes (Saconi 1966). Literature survey reveals (Yamada 19
that square planar nickel complexes are dtamagentic and red in solid state. Ma|
tic moment value for NiSO^ complex (2-98 B.M.) is within the range for te
hedral environment. The possibility of pseudo-tetrahedral environment is ral
less because such complexes possess magnetic moment in the range 3 -2-3 -3 B
(Suconi 1966). Hence we are tempted to suggest a tetrahedral polymeric sti
turc for NiSO 4 (LPZ).H a O (figure Ib).
3 . 3. Th&rmogravimetric analysis
Thcrmogravimotric study of all the complexes show almost similar decomposit
pattern. The DTA peaks are given in table 2. The cndothermic peaks
105-4 10 C in all the complexes on DTA curve is accompanied by weight 1
equivalent to one water molecule on TG curve.
In the case of MC1 2 complexes loss of chlorine occurs in the temperature rai
105-<320 C which is followed by elimination of amine groups in the temperat
range 300-450 C, with subsequent decomposition of the ligand and carbonizat
of the aromatic ring residue which spreads over the temperature range 450-600 C
The formation of respective metal oxides occurs at 600-^700 C. The then
Table 2. Thermal decomposition data. MC1 4 and MSO 2 complexes.
Compound
DTA peak temp.
(C)
CuCl a (IPZ).H a O
IIQ(-)
215 (-I-)
360 (+)
NiCl 2 (IPZ).H a O
105 (-)
330 (+)
425 (+)
CoC! a (lPZ).H 2 O
100 (-)
250 (-I-)
450 (+)
600 (H-) 640 (+)
CuS0 4 (IPZ).H 2 O
110 (-)
240 (+)
420 (-I-)
720 (-)
NiS0 4 (IPZ).H a O
105 (-)
280 (+)
340(-r-)
440 (+) 530 (+)
750 (-)
(-) and (+) signs after
tenvioratiiros represent ondotliorm
and oxotliorm respectively.
For MSO 4 complexes, after the elimination of one water molecule and decompo-
sition of amino groups, formation of anhydrous MSO 4 takes place at temperature
around 500 O. Anhydrous MSO 4 finally decomposes to respective metal oxide
via intermediate oxide formation as represented below. Formation of MO to
MO 2 is evidenced by peaks at 720 and 750 C on DTA curves for Cu(II) and
Ni(fl) complexes, which are also supported by the corresponding peaks on DTA
curves at the same temperature. Percentage residue observed on TG curve also
agrees well with calculated values for each intermediate stage of decomposition.
The decomposition pattern for MC1 2 and MSO 4 complexes can be represented
as
decomposition,
m MCI L HO-? 20 - MCI L carbonization
l MUa-L.HUoMU. L.
decomposition,
/\ nx-o/-k T TT r\ H 2 O T.,<y~ T carbonization
(n) MS0 4 - L . H 2 MSO d - L - MSO 4
oxidation . ... , -^
MO
500-600"C a 3 650-700 C a 700-900 C
Acknowledgement
The author thanks Dr Prem ,. warup, SPL, Delhi, for ESR spectral measurements.
Thanks are due to Dr S P Panda and Prof. N N Banerji, Applied Science
Faculty of this institute, for their keen interest and encouragement. The author
.rvuuiua i- AVI iy\Jt JKICIWI uguttu unu leuiieu viumuuri
Aggarwal R C and Bahadur A 1969 Indian J. Chem. 7 1039
Aggarwal R C and Narang K K 1976 Indian J. Chem. A14 66
Agganval R C and Singh S K 1969 Bull. Chem. Soc. Jpn. 42 3363
Aggarwal R C, Singh S K and Prasad L 1976 Indian J. Chem. A14 181
AWa R P and Kumbhar C G 1977 /. Armt. Studies 13 (1) 47
Beadle P J, Goldstein M, Goodgame M L and Grezeshkowiak R 1969 Inorg. Chem.
Billing D E and Underbill A E 1968 J. Inorg. Nucl. Chem. 30 2147
Figgis B M and Lewis J 1960 Modem coordination chemistry (eds.) J Lewis and R G
(New York : Interscience) p. 403
Kumbhar C G 1980 /. Armt. Studies 16 (1) 24
Kumbhar C G and Sadasivan N 1976 /. Armt. Studies 12 (2) 135 ; Chem. Abstr. 86
Mahapatra B K and Rama Rao D V 1971 Indian J. Chem. 9 715
Majumdar V R and Sharma B C 1978 Res. and Ind. 23 (2) 79
Manch W and Fernelius W C 1961 J. Chem. Ediic. 38 192
Marvel C S and Tarkoy N 1957 /. Am. Chem. Soc. 79 6000
Nakamoto K, Fujita J, Tanaka S and Kobayashi N 1957 /. Am. Chem. Soc. 79 4904
Sacconi L 1966 Coord. Chem. Rev. I 126
Sacconi L and Ciampolini M 1964 /. Chem. Soc.. 276
Sadasivan N and Arora R P 1976 Indian J. Pure Appl. Phys. 14 745
Voloviskii L W and Knorozova G V 1964 Zh. Obsrch. Khim. 34 343
Yamada S 1966 Coord. Chem. Rev. 1 415
Amorphous MoS 3 and A*MoS, (A = Li or Na ; < x < 4)t
T MURUGESAN and J GOPALAKRISHNAN*
Solid State and Structural Chemistry Unit, Indian Institute of Science,
Bangalore 560 01 2, India
MS received 12 June 1981
Abstract. Amorphous A.MoS 3 (A = U or Na; < x < 4) prepared by the re-
action of MoS 3 with rt-butyllithium or sodium naphthalide in. organic solvents have
been characterized by x-ray photoelectron Spectroscopy, infrared spectroscopy as well
as electrical and magnetic measurements. The results indicate that sulphur exists
as poly&ulphide species in MoS 3 and mainly as monosulphide in A,MoS a when
x ^-4; there is no discernible change in the Mo(3rf) binding energies of MoS 3 and
A f MoS a . Both MoS 3 and A.MoSa are diamagnetic and non-metallic at room
temperature. The data suggest that MoS 3 probably exists as Mo 2 " 1 ' (S 3 2 ~) with
Mo-Mo bonds, incorporation of alkali metal atoms resulting in the reduction of
proportion of polysv Iphide ions.
Keywords. Amorphous MoS a ; Li,MoS 3 ; Na,MoS 3 ; x-ray ph toelectron spectra.
I. Introduction
Among the transition metal trisulphides MS 3 (M = Ti, Zr, Hf, Kb, Ta, Mo and
W), MoS 3 and WS 3 can be prepared only in the amorphous state by low-tempera-
ture chemical or thermal decomposition of ammonium tetrathiometallates (Wilder-
vanck and Jellinek 1964 ; Diemann 1977). Until recently, the identity of these su 1-
phides as. true chemical compounds was somewhat in doubt. It has now been
established that these are definite compounds (not a mixture of disulphide and
amorphous, sulphur) possessing a chain-like structure similar to that of crystalline
trichaicogenides of other transition metals (Liang et al 1980a,b). It has recently
been found that MoS 3 can incorporate reversibly upto four atoms of alkali metal
per formula unit, A a MoS 3 (A = alkali metal ; < x < 4) making it a good
candidate for cathode material in solid state batteries (Jacobson et al 1979).
A^MoSs may be regarded as thioanalogues of the alkali metal oxygen bronzes of
molybdenum. We have investigated the structure and electronic properties of
MoS 3 and A^MoSs (A = Li or Na) by various physical methods in an attempt
to understand the nature of these solids.
| Contribution No. 125 from the Solid State and Structural Chemistry Unit.
* To whom all correspondence should be made.
2. Experimental
MoS 3 was prepared by thermal decomposition of (NH 4 ) 2 MoS 4 at 500 K in
stream of dry nitrogen (Jacobson et al 1979). The composition was found to t
MoS 2 . 98 from the chemical analysis of sulphur [S (found) = 49-73% ; S (caic.) :
50-06%]. Samples of Li,MoS 3 (x = 0-9, 2-2 and 3-7) and Ha.MoS 8 (* - 0-
and. 3-5) were prepared, as reported in the literature (Jacobson et al 1979), fc
reaction with n-butyllithium in rc-Jiexane and sodium naphthalide: in tetrahydr<
fiiran respectively. To prepare Li,MoS 3) a known amount of MoSs was treate
with a 1 M solution of #--butyllithhim in /nhexane in a flowing nitrogen atmospher
After the reaction, the solid was filtered and the concentration of n^butyUithmi
in the filtrate was determined by the addition of standard potassium hydroge
phthalate and back titration with standard potassium hydroxide. From the diffe
ence in concentration, the amount of lithium inserted into MoS 3 was calculate*
Similarly samples of Na,MoS a were prepared by reaction of MoS a with sodiui
naphthalide in dry tetrahydrofuran followed by determination of the concentratic
of sodium naphthalide in the filtrate as in the case of reaction with n- butyllithiun
Details are given in table 1 .
X-ray powder diffraction patterns, recorded with CuK tt radiation, showed
broad diffuse scattering with a maximum around 14 2J. The diffuse band becarr
sharper with increasing alkali metal content in A,MoSa. The absence of any oth<
discrete diffraction lines in the patterns indicates that the samples are x-rs
amorphous similar to MoSa-
X-ray photoelectron spectra (XPS) of the samples were recorded with a ESCA-
Mark II sepectrometer (VG Scientific Co. Ltd., UK) using A1K radtatio:
Infrared spectra were recorded with a Perkin-Elmer Model 580 spectromete
Electrical resistivities of the pelletized samples were measured by a two-pro!
technique. Magnetic susceptibilities were measured by Faraday method betwe<
150-300 K.
Table 1. Preparation of A a MoS 3 (A = Li or Na ; < x < 4).
<
Amount of
MoS a , g
(m moles)
2oTicerttration of n-butyllithium or
sodium naphtlialide (m moles)
Alkali metal in-
- sorted per mole Composition
of MoS 3 (moles)
Befove the
reaction
After the
reaction
2-88(15)
13-50
o-o
0-9 Li . B MoS s
2-88(15)
33-00
0-0
2-2 Li 2 . 2 MoS s
2-40(12-5)
50-00
3-75
3-7 Iii,.,MoS 8
3-84(20)
16-00
0-0
0-8 Na. 8 MoS 8
3. Results and discussion
We have studied the valence band and core level XPS of MoS 3 , Li . 9 MoS a ,
Li 2 . 3 MoS 3 and Li 3 . 7 MoS 3 to find out the nature of molybdenum and sulphur in
these compounds. The spectra are given in figures 1 and 2 and the binding
energies in table 2. For purpose of comparison, the spectra of MoS 3 are also
included in the figures. ,
The S(3.y) peak of MoS 3 occurs as a doublet at 12-7 and 16-6 eV binding
energies in contrast to a single (3s) peak at 14 eV in the case of MoS 2 (figure 1).
To account for the doublet structure, it was proposed in our earlier study from
this laboratory (Manthiram et al 1980) that two different kinds of sulphur are
present in MoS 3 : Mo*+ (S 2 2 ~) (S 2 -). Similar S(3s) doublet structure in MoS a
with a relative intensity of 2 : 1 has been found by Liang et al (1980a). They
proposed that MoS 3 consists of \ SI~ and 2S 2 ~ which requires that molybdenum
is present in 5+ formal oxidation state : Mo^SS") (S 2 ~) a . According to this
formulation, formation of A a MoS a with x up to four would imply a reduction
of Mo B + to Mo 2 + ;
(S 2 2 -) (S 2 -)o + 4 Li -> Li 4 +
Mo(3d 5/2 )
225 230 235
BINDING ENERGY (eV)
Figure 2. S (3s) and Mo (3rf s/a , 3<f s/a ) core level spectra of (a) MoS , (b) lii. % MoS 3 ,
(c) I4 r7 MoS, and (d) MoS>.
Table 2. XPS binding energy, electrical resistivity and magnetic susceptibility data
of MoS 8 and A.MoSj (A = Li or Na ; Q<x< 4).
Compound
S(3s) binding
energy (eV)
S(2s) binding
energy (cV)
Mo(3rf./.)
binding energy
(eV)
Electrical
resistivity p
at 300 K
(Ohm-crn)
Magnetic
susceptibility
% m x 10* at
300 K
(cgs emu)
MoS,
12-7, 16-6
226-7
229-1
3-8 xlO 4
-48
LI'D-B MoS,
11-2,14-8. 17
3 226-4
229-1
4-6 x 10*
-67
I i a . a MoS,
broad
226-4
229-1
3-6 X 10*
-86
Li s . 7 MoS (
14-8
226-3
229-1
4-0 Xl0*
-92
Na. . MoS.
3-4 x 10*
-70
in Li . 9 Mo:!s 3 ana LvsMoSs, the S(3.y) shows complex features and at the limiting
composition Li 3 . 7 MoS 3 , the S(3^) becomes a single band similar to that in MoS a .
In addition, the Mo(3f) binding energies remain almost constant [229-1 eV for
Mo(34/a)] in MoS 3 and LiMoS 3 . We also see a slight decrease in the (2s)
binding energy as we go from MoS 3 to Li 3 . 7 MoS 3 (table 2). The results seem
to indicate that incorporation of alkali metal into MoS 3 affects only sulphur and
not molybdenum.
If we assume that MoS a consists of a trisulphide ion, Sl~, and Mo- Mo chain,
Mo 2 +(S-), the experimental results can be explained as follows :
(i) incorporation upto a maximum of four alkali metal atoms without change
in the oxidation state of molybdenum,
Mo 2 +(S*-) + 4 Li -* Li 4 + Mo 2 +(S 2 -) 3 ,
(ii) presence of two different kinds of sulphur in MoS 3 in the ratio 2 : 1, and
(iii) the complex nature of S(3^) at intermediate values of x in A,,MoS a . In
these cases, the polysulphide ion bonds would have been partially broken result-
ing in Sl~, and S|j- species.
Infrared absorf tion spectra and electrical and magnetic proj erties of A.MoS a
are consistent with the above model. MoS 3 shows characteristic S-S stretching
vibration of the polysulphide ion at 515 and 540cm- 1 as shown in figure 3 (Rittner
et al 1979). The disappearance of these bands in Li 3 . 7 MoS 3 indicates that poly-
sulphide species is absent. In addition, a new band at 420cm" 1 appears in
Li,MoS 3 ; the band may be assigned to Li->S stretching vibration. Similar changes
in the infrared spectra of Li/TiS 3 have been reported by Chianelli and Dines (1975).
(a)
700 600 500 400 300
i) (cm" 1 )
Figure 3. Infrared spectra of (a) MoS,,, (b) Id,. 7 MoS, and (c) MoS (crystalline)
-50
-too
(4)
* 4 , *
(b)
9 i ^
(0
(d)
150 200 250 300
T(K)
Figure 4. x H ~T plots for (a) MoS n , (b) Li. MoS 3 , (c) Li a . a MoS a and (d) i,i 3 .,MoS 3 .
Room temperature electrical resistivity, /?, and magnetic susceptibility, # M , of
MoS s and A,Mo? 3 are given in table 2. It is seen that there is no significant
difference between the resistivities of Mo 3 S and A,MoS a . The magnetic suscepti-
bility data (bgure 4) show that the diamagnetic character of MoS 3 is retained in
A a MoS 3 albeit with increase in the magnitude of diamagentic / u . The results
support our formulation of MoSS 3 as Mo 2 +(St~), the diamagnetism being due to
Mo- Mo bonds as proposed by Liang et al (1980a). Insertion of alkali metal
does not seem to disrupt the Mo- Mo bonds in MoS 3 .
The resence of molybdenum in a formal oxidation state of 2 + in MoS 3 can be
understood in terms of Jellinek's (1968) model for transition metal sulphides.
Transition metal ions having large positive oxidation state such as Mo*+
and W 8+ would be unstable in the solid state in the presence of S 2 " ions because
the valence S(3p) states overlap with the empty Mo(4d) or W(5rf) states, result-
ing in electron transfer from S(3/>) to the M(d) until the metal Estates are lifted
just above those of 3(3^). In chemical terms, this would correspond to the reduc-
tion of the metal ion to lower oxidation states and oxidation of sulphide to poly-
sulphide :
MO<H-
2 - (3p)
A formal oxidation state around 2+ for molybdenum as well as Mo-Mo bonds
occur in molytdenum sulphides, e.g. Ohevrel phases, A,Mo 6 S 8 (Vvon 1978).
Acknowledgements
The authors thank Professor CNR Rao for suggesting the problem and taking
keen interest in the progress of the work. The authors also thank Dr MS Hegde
for XPS measurements and the UGC for financial support.
References
Liang K S, Cramer S P, Johnston D C, Chang C M, Jacobson A J, doNoufvilU) J P and
Chianelli R R 1980b /. Non-Cry st. Solids 42 345
Liang K S, doNcufville J P, Jacobson A J, Chianelli R R and Betts V 1980a /. Non-Cryst.
Solids 35 and 36 1249
Manthiram A, Sarode P R, Madlmsudan W H, Gopalakrishivwi J and Rao CNR 1980
/. Phys. Cheni. 84 2200 *
Rittnor W, Muller A, Neumann A, Bather W and Sharnia R C 1979 Angcw. Client., Int. Ed.
Engl 18 530
Wildervanck J C and Jollinck P 1964 2. Anorg. Allg. Ghent. 328 309
Yvon K 1978 Curt: Top. Mater. Sci:353
Synthesis of l,5-disubstituted-2,4-dithiobiuret and 1,5-disubstitutedL
2-thiobiuret and their vanadyl ( V) chloride complexes
K P SRIVASTAVA and I K JAIN*
Chemistry Group, Birla Institute of Technology and Science, Pilani 333031, India
* Present address : Geological Survey of India, Bhopal 462 002, India
MS received 10 July 1981 ; revised 24 November 1981
Abstract. Vanadyl (V) chloride forms VOC1 3 L (L is a bidentate ligand molecule)
type complexes with 1,5-disubstituted 2,4-dithiobiurets and l,5-disubstitutcd-2-
thiobiurets in carbon tetrachloride solution. Co-ordinations through sulphur atoms
in case of 1,5-diaryl substituted 2,4-dithiobiurets and sulphur and oxygen in case of
1,5-diaryl substituted 2-thiobiurets have been proposed for these complexes.
Keywords. Vanadyl (V) chloride complexes; 1,5-disubstituted 2,4-dithiobiurets ;
l,5-disubstituted-2-thiobiurets.
1. Introduction
Vanadyl (V) chloride forms VO (NHPli) 3 type complex (Nelson and McFadden
1933) with aniline. Addition and substitution complexes of VOCJ 3 are also known
(bunk et al 1958 ; Kraiiss and Gnatz 1962 ; Cazzi and Gecconi, 1953). Reduc"
tion of VOGla by carboxylic acids, pyridine and aliphatic amines and thioethers
with the formation of vanadyl (IV) carboxylates, VOCl 2 -3 C 5 H 5 N and VOC1 2 -2L
where L is CH a NH 2 ; (CH 3 ) 2 S and (C 2 H 5 ) 2 S respectively is also reported
(Selbin 1953 ; Paul and Kumar 1965 ; Baker et al 1967). Some unreduced
VOC1 3 - 2L type VOC1 3 complexes with aromatic amines are also said to have been
synthesized (Prasad and Upadhyaya I960).
From the literature it has been found that no work has teen carried out on com-
plexing ability of VOGL 3 with polydentate ligands. We have synthesized the
1,5 dtsubstituted-2,4-dithiobiurets and l,5-disubstituted-2-thiobiurets ligands and
their VOC1 3 complexes.
2. Experimental
All the chemicals used were of either BDH " AnalaR " or MERCK GR grade.
Solvents were purified by repeated distillation after appropriate drying. Phenyt
isocyanate (mustard oil) used was of E. MERCK. Other mustard oils viz
phenyl, ortho-tolyl and para-tolyl isothiocyanate were prepared in the laboratory
* To whoin all correspondence should be made.
15
K P Snvastava and I K Jain .
5el 1962). Vanadyl (V) chloride was prepared (H<xht et al 1947 ; Prandtl
Bleyer 1909) ty refluxing and distilling the mixture of V 2 O 5 and SOd* It
analysed to check its purity. Data are reported in table 1.
Synthesis of substituted dithtobiurets and thiobiurels ligands
ligands 1,5-dtphenyl 2,4-ditliiobiuret (DPDTB) ; 1,5-diphcnyl 2-thiobiuret
PS); l-phenyl 5-paratolyi 2,4 dithiobiuret (P.^-TDTB); 1-phenyl 5-para-
2-thiobiuret (P.^TTB) ; l-orthotolyl-5-phenyl 2,4-dithiobiuret (o-TPDTB ;
hotolyl-5-phenyl-2-thiobiuret (o-TPTB) ; 1,5-diparatolyl 2,4-dithiotiuret
TDTB) ; 1,5-diortho.tolyl 2,4-dithiobiuret (D-o-TDTB) and l-methyl-5-
tolyl 2,4-dithiobiuret (Me-_p-TDTB) were prepared by the known method
t 1962). Parity of ligands was checked by their sharp inciting points and
icntal analysis (table 1).
Preparation of complexes
Derations were carried out in a dry box. The ligand was dissolved in mini-
L of chloroform and then 0-02M solution was prepared by CC1 4 - VOCla
ion in CC1 4 (0-03 M) was added to the ice-cold ligand solution. The tempc-
e of reaction mixture was kept below 10 C. The coloured complex was
ed, washed, dried and then analysed for vanadium, chloride and sulphur
;el 1959).
Instrumentation
Magnetic measurements of the complexes were carried out on Gouy balance.
,r conductance was determined in N,N-dimothyl formamidc (10~^M) on
nctivity meter type LBR of Wissenschaftlich Technisch, Wersstatten, Germany^
;a dtp type cell The IR spectras were taken on Perkins Elmer Grating infra-
pectrophotometer model 237-B and 621. The important peaks of spectra
is ted in table 2.
lesults and discussion
/tical results (table 1) correspond to the empirical formula VOCl r L, where
i ligand molecule. The molar conductance values are in the range of 32-96-
> Mhos which are well below those for 1 : 1 electrolyte (Suttons 1971). Th e
a red conductance values seem to be due to the partial replacement of chlo-
fc>y solvent molecule. All the complexes are diamagnetic with a magnetic
ptibility (-0-19 to -0-43) x 1Q- 6 .
the ligands a medium broad band appearing around 3100-3200 may be due
U
r)
,2
o
U
U
o
'q.
s
NO c-i r-l
(N <M tN M iM 01
OO ON C O '/">
oor-ioomcoricoooOOvc '
01 oo oo r-l r-l
* r-l ON c^-i Tt r-l ON f"' T en '-' l~- *n
C*4 fo *~( t~ *"
cn^_vbr-'ro^<sbof^O fr)V i :i <'O
CM i-, _ r-l
ON ^5 vO
r <1 O . *
f*i *_* ^i t^ r^*i f^i *O
r-- . r~ t -* * -co
Q-, ,_ r-l
o ^- o o o ci
0-1 --
ri
rl o r* O o f>
-- . -O t~- -ON -O
r-l c-i *3-
^ <* m rl rl O r
VD rl 01
01 rl rl ol rl 01 ol
cr, _ -t -) o
oO <N ON ON o
<U fy-j r-i oo -t vo r-i -rf r1 "' rl rl
^ in * O m cS ''i' "*3" O ^~* * ' ^
_. c /i o "-O en
^ ^ * !> o! r^ O t^ *" * "^t" ^^* ^^ t^- VO
r-l r-< -> r-l
Th ^ fr t vi rp CM -n-
ON C3 1| ~'
^ r-1 O O O\ < ' co
U
O .-SJ
o o 15
u 5
IS O r- 5:
g, en S
?~1 G " O jj;
> S ^ * d ^ "^ ^ ^ S
^ js 1 -^ s--
,cf ^J to -55
32 -22 &
T3 4S ^j *^ ^ r O
, -2 5 ? !p
ftf ^ fQ ^ iJ >H
W 6D W ,S2 i- (D *U f t/J {^) ^o
. *?~* >-i ">T bfl >* 5^i f *^ >-
C ^r* > Oj_i*"C! -d
^'M^ib^^i'^S^^ S'Si'S
'-'j' l 3> :i !JJi>^-> < Ojc
: 3 S K a a
^ OO O r1 Ch r^* O ^i" *-h s ! G\ i/-^ c?
5-5
+ 4-
s
ll
i
_|- 0>
1 is
I
l!
-
i|
ONVO
l
"
in 5
O O
n C5
53
a
i
ompound
SI
et al 1967 ; Yarnaguchi et al 1958). Both these lands on compJexation arc shifted
down by 10-20 crrr 1 with reduced intensity which indicates the possibility of
thioketo sulphur atoms as co-ordination sites in complexes. The bands ~ 1450 crrr 1
in ligands are assigned to the NH-C = S group vibration which is the combination
of v (N-C-tt), v (C = S) and ~NH rocking vibrations (Yamaguchi et al 1958 ;
Randall 1949). Tne change in the nature of these vibrations on complexatjon
further confirms the ligand co-ordination to the metal atom through thioketo
sulphur atoms.
In DPTB, P-J7-TTB and 0-T-PTB ligands, the sharp peaks appearing around
1710cm.- 1 have been assigned to the vibrations of v (C = O) group (Srivastava
and Madhok- 1978). On complexation the v (C = O) absorption peak shifts
to a lower frequency by 10-20 cm- 1 and becomes medium in intensity. The V = O
stretching vibrations in the complexes appear ~ 1030 cm" 1 . These bands are weak
and medium in intensity and are in the region expected for vanadium oxygen
stretching frequencies (Miller and Cousins 1957).
All these observations show that co-ordination of DPDTB, P-^-TDTB, o-TPDTB,
D.p-TDTB, D-O-TDTB and Me.^-TDTB ligands to the metalatom is through
two thioketo sulphur atoms while in the case of DPTB, P./>-TTB and o-TPTB
ligands it is through thio-keto sulphur and C = O group oxygen atoms. The
six coordinated complexes fVOC! 3 (L)] so formed may have octahedral configu-
ration.
Acknowledgement
One of the authors (HO) thanks the UGC, New Delhi, for a fellowship.
References
Baker K L, Edwards D A, Fowles G W A and Williams R G 1967 /. Jnorg. Nucl. Own. 2V
1881
Cozzi D and Cecconi S 1953 Recerca Sci. 23 609
Dixit S N 1962 /. Indian Chem. Soc. 39 407
Funk H, Weiss W and Zeising M 1958 Z. Anorg. Allg. Chem, 296 36
Gosavi R K, Agarwala U and Rao CNR 1967 /. Am. Chem. Soc. 89 235
Hecht H, lander G and Schlapmann H 1947 Z. Inorg. Chem. 254 255
Krauss H L and Gnatz G 1962 Chem. Bcr. 95 1023
Miller F A and Cousins L R 1957 J. Chem. Phys. 26 329
Nelson R E and McFadden A H 1933 Pftc. Indian Mad. Scl. 42 119
Paul R C and Kumar A 1965 J. Inorg. Nucl. Chem. 27 2537
Prandtl W and Bleyer B 1909 Z. Anorg. Chem. 65 153
Prasad S and Upadhyaya K N 1960 J. Proc. Inst. Chem. 32 265, 270
Randall F and Fuson Dangl 1947 Infrared (let of organic structures (van Nostrand)
Scheinmann F 1970 Nuclear magnetic resonance and PR spectroscopy Vol. I (Oxford, New York :
Pcrgamon Press) p. 176
Sclbin J 1953 Chem. Rev. 65 153
Srivastava K P and Madhok K L 1978 /. Inorg. Nucl. Chem. 40 1821 .
Suttons G J 1971 Aust. J. Chem. 24 919
Precipitation of uranium quinolin-8>olate from homogeneous solution
by urea hydrolysis
G SIVA REDDYt, A VARADA REDDY and
Y KRISHNA REDDY*
Department of Chemistry, SV University, Tirupati 517 502, India
t Present address : SVUPG Extension Centre, Cuddapah, India
MS received 7 February 1981 ; revised 16 October 1981
Abstract. A very crystalline precipitate of uranium quinolin-8-olate (uranyl oxinate)
has been obtained from homogeneous solution by utilising urea hydrolysis to raise
the pH of the initially acid solution of uranium (VI) in the presence of sodium
acetate and the precipitant. The precipitate so obtained is stoichiometric unlike
the one obtained in conventional precipitation and possesses the same composition
(uranium : oxine ratio) irrespective of the pH of the final solution unlike that
obtained in earlier PFHS methods. The results obtained show that the determi-
nation is not affected by the presence of other metal ions when the precipitation
is carried out in the presence of EDTA and the precipitate obtained by PFHS is also
more thermally stable.
Keywords. Uranium(VI) complex ; qutn.olin-8-ol ; precipitation from homogeneous
solutions ; urea hydrolysis ; thermal analysis.
1. Introduction
Precipitation from homogeneous solution (PFHS) technique is often employed
to produce pure and crystalline precipitates, and the precipitates so obtained also
possess a better stioshiometry when compared to the precipitates ottained by the
conventional method (Gordon et al 1959 ; Cartwright et al 1967). For example,
copper cupferrate obtained by PFHS irothod can be weighed directly for the
gravimetric determination of the metal (Heyn and Dave 1960) whereas the same
obtained by the conventional method is to be ignited to oxide before weighing.
Corsini and Abraham (1968) reported that the uranium quinolin-8-^olate (uranyl
oxinate) obtained by direct addition of the reagent was deficient in oxine. This
red compound UO a (C 9 HeON) 2 C 8 H 7 ON, which is usually employed for the ana-
lysis of uranium (VI) solutions was mentioned several times in literature (Hecht
and Reich-Rohrwig 1929 ; Frere 1933 ; Fleck 1937 ; Classen and Visser 1946 ;
Moeller and Wilkins 1953 ; Wendlandt 1956 ; Van Tassel and Wendlandt 1959,
1961 ; Bullwinkel and Nobel 1959 ;' Horton and Wendlandt 1963 ; Tackett 1964 ;
Majeeand U-oroon lyoij Majee ana wooawara iyoo; .Fleming ana Lynton
Milner et al I960), but the deficiency of oxine was not observed. Similarly, the
workers who studied the PPHS determination of uranium qiunolin-8-.olate by the
hydrolysis of 8-acetoxyquinoline (Bordner et al 1961) and also by evaporation
of mixed solvents (Howick and Rihs 1964) were not aware of this problem as
Corsini and Abraham's (loc. tit.) work was published much later. On the other
hand it wa^ reported that the complex obtained by PPHS method had different
compositions depending on the pH of the solution, a dark red compound UO a
(C H B ON) 2 C g H 7 ON was obtained at pH 5-0 and an orange compound (UO 2
(G a H ON) 2 ) 2 C 9 H 7 ON resulted when the pH was increased to 6-8 (Bordner et al
1961). Hence the earlier observation of deficiency of oxine in the red compound
may be d.ie to the coprecipitation of the latter compound with the former.
Since PFHS methods produced stoichiometrjc precipitates which could be
weighed directly unlike in conventional methods (Cartwright et al 1967 ; Siva
Reddy and Krishna Reddy 1980) in the present investigation, uranium uinolin-
8-olate has been precipitated from homogeneous solution by urea hydrolysis so
as to remove the deficiency of oxine. Crystal sizes and thermal behaviour of the
precipitates obtained by PFHS and conventional methods are compared.
2. Experimental
2- 1. Reagents
AH the chemicals used were of AnalaR grade, supplied by BDH. Uranium (VI)
solution: About 8- 5 gin of uranyl nitrate, UO 2 (NO a ) 2 6H 2 O was transferred to a
1-litre flask, dissolved in water and diluted to the mark. The solution was stan-
dardised gravi metrically using diammonium uranate procedure (Vogel 1975).
Qninol'm~$>-^l (oxine) solution : 2% solution (wfv) in acetone.
Dilute nitric acid : Aqueous solution, 1 : 1 (v/v) of HNO 3 .
Dilute ammonia solution : Ammonium hydroxide solution, 1 : 1 (v/v).
2.2. Apparatus
Elico pH-meter model Ll^lO was used in the investigation. The sensitivity of
the instrument is 0-05 pH units. The particle sizes were measured using Leitz-
Wetzler microscope and photomicrographs were also taken using the same micro-
scope fitted with camera. The thermograms were recorded using a thermobalance
supplied by Stanton Redcroft. The crucible used was made of alumina and the
thermocouple employed was constructed from platinum and rhodium. The
temperature was raised at the rate of 6 C/min. The samples were dried under
vacuum for several days before recording the thermograms.
2-3. Procedure
An aliquot containing uranium (VI) was transferred into a clean 250ml beaker
and to it were added 10 gm of sodium acetate and 15 gm. of urea. The initial pH
of the solution was adjusted to 2-0 with 1 : I nitric acid. Quinolin->8-.ol solution
5 hr. Tne solution was cooled to room temperature ana. me precipitate was miere a
through a weighed sintered glass- crucible of medium porosity, washed with warm
water (ca. temperature 40 C) and dried for 1 hr at 1 15-120 C and weighed.
3. Results and discussions
3. 1. Determination of uranium
Determination of uranium in the range 20-300 mg was successful and the results
are presented in table 2. For a set often measurements of 81-25 mg of uranium
(VI), the average amount and standard deviation were found to be 81-23 and
0'07mg. For the same amount of uranium (VI), the amount found by bromo-
metric procedure was 81-22 mg with a standard deviation of 0- 12 mg. The micro-
analysis of the precipitate indicated the composition which correspond to
U0 2 (CoHaONJa C H 7 ON (table 1).
Table 1. Microanalysis of the complex.
Element Theoretical Percentage
percentage obtained
C
46-03
45-95
H
2-72
2-68
N
5-97
5-97
Table 2. Determination of uranium.
Amount of
Amount of
Difference
uranium taken*
uranium
mg (n = 3)
mg
found,
mg (n = 3)
8-1 3
8-l
0-0 3
20 -3j
20- 3 2
O-Oi
40-6 2
40 -5 8
0-0 4
60 -9 a
60 -9
o-o 3
81 -2s
81 -2 3
0-0,**
243-8
243-3
0-5
325-0
324-2
0-8
* Based on gravimetric (ammonium diuranate) procedure.
siiuw mat
the pH is increased beyond 9-0, it was observed that only dark red compound
formed unlike in the earlier PPHS and conventional methods (Corsini and Abra-
ham 1968 ; Bordner et al 1961 ; Howick and Rihs 1964). Moreover the com-
plex possesses a good stoichiometry (no deficiency of oxine is observed). Addition
of 2 to 3 fold excess of reagent did not affect the determination, though a slight
excess ensured complete precipitation.
3.2. Panicle size
The precipitate ol tained is very dense and crystalline and higher amounts of uranium
could be handled very easily. The distributions of particle sizes are given in
table 3. The median value was obtained from the plots of particle size vs. loga-
rithm of cumulative number. The median value calculated was 192 x 112/j. The
precipitate obtained by conventional method is amorphous. The photomicro-
graphs of the precipitate obtained are shown in figure 1.
3.3. Thermal behaviour
The thermograms are shown in figure 2. The precipitate obtained from homo-
geneous solution did not lose weight up to 200 C whereas the same obtained by
conventional method was. stable up to 180 C. The precipitate yielded unsolvated
chela te on further heating due to the sublimation of supplementary quinolin->8-.ol
molecule and the fo>complex, UO 3 (CaHfjON^ was stable in the range of 240-
310 C. At 440 G, the weight of the precipitate corresponded to U 3 O 8 and from
this temperature no weight loss was recorded. For the precipitate obtained by
conventional method, the formation of 6/>complex and oxide were observed at
220 C and 420 C respectively. The thermograms recorded in the present
procedure agreed fairly well with the earlier ones but not with the pyrolysis curve
recorded by 0uval (1953) who stated that formation of U 3 O 8 did not take place
before 940 C.
Table 3. Distribution of particle sizes of the precipitate obtained by PFHS.
Total number of particles counted = 500
1 microscopic division = 16/<.
Lengthwise distribution Breadthwise distribution
Microscopic
Percentage of
Microscopic
Percentage of
divisions
distribution
divisions
distribution
1-10
64-7
1-4
52-6
11-20
16-8 v
5-8
18
21-30
11
9-12
15
31-40
5
13-16
6-6
41-50
2
17-20
5
Precipitation of uranium qiiinolin-8-olate
25
Figure 1. Photomicrographs of uranium quinolin-8-olate (a) from homogeneous
solution, (b) by conventional method.
precipitate obtained by PFHS
-- Precipitate obtained by convent ionai.
method
33-9 mg.
39-9 mg.
. 85-2.mg,.
J
200 *00
TEMPERATURE
Figure 2. Pyrolysis of uranium quinolin-8-olate.
600
3.4. pH increase in presence of various anions
Tne function of the anion in urea hydrolysis is three fold ;. (i) buffer action,
(ii) complex formation and (iii) incorporation of anion in the precipitate and
formation of basic salt. In the precipitation of hydroxides and basic salts, the
anions help in the formation of basic salts though other reasons are not ruled out
(Gartwright 1967). Similarly in the precipitation of neutral salts or complexes,
the anion serves mainly as buffer, though complex formation is not ruled out (Siva
Reddy and Krishna Reddy 1979). This can bs understood only from the plots,
of pH increase with time and percentage recovery of the metal with j H in presence
of various anions. pH increase in the presence of various anions is shown in
figure 3 and the percentage recovery of uranium (VI) with pH is given in figure 4.
pH increase is very rapid in the presence of ammonium chloride, ammonium sul-
phate, tartaric acid and succinic acid and slow in the presence of acetate and for-
mate buffers. A very crystalline precipitate is obtained in the presence of sodium
acetate and ammonium acetate. The pH increase with time shown in figure 3
is for nearly 80 mg of uranium (VI), 15 ml of quinolin-8-ol solution and 10 gm
of buffer. Each value represented in the figure is the average for six experiments.
Prom the plot of percentage recovery of uranium with pH (figure 4), it is seen that
the initial precipitation process is delayed in the presence of succinic acid and EDTA
due to complex formation other buffers appear to control pH raise only.
3.5. Effect of diverse ions
Wnen precipitation was carried out in the presence of EDTA, 100 mg each of
copper (II), iron (III), zinc (II), cadmium (II), lead (II), aluminium (III) and
28
G Siva Reddy, A Varada Reddy and Y Krishna Reddy
) 2 SO tt
Tarlaric or succinic cid
Sodium acetate
Ammonium ace tat* + 2g Na->
EDTAor2g tartaric
Ammonium acid
tat
468
TIME IN HOURS
10
Figure 3. pH increase with time in presence of various buffers.
3 10
=> 80
z
60
I
In presence of :
Na^EOTAand sodium
tartarale
Succinic acid
other buffers
Initial molarity of EDTA in the solution = 0-04 M
Amount of cation added = 100 mg.
Amount of uranium
Cation added found for 81-2 mg
of uranium taken mg
Difference
mg
Copper(II)
81 -2
0-0 5
Nickel(II)
81-1 8
0-0 7
Iron(nt)
81-3 9
0-0 7
Manganesc(tl)
81 -2
0-0 S
Zinc(II)
81 -2o
0-0 5
Cadmium(II)
8M 8
0-0 7
Lead(ri)*
8l-3 4
0-0 9
AluminiumCni)
81 -3,
0-0 7
Thorium(IV)*
81 -2 B
o-o a
* Added in the form of nitrates.
Other salts are added in the form of sulphates.
4. Conclusion
Prom the above discussion, it is clear that the PFHS determination of uranium
by urea hydrolysis is accurate. NO deficiency of quinolin-8-ol (oxine) or differ-
ences in compositions has been observed unlike in earlier conventional and PFHS
methods. The highly crystalline precipitate obtained in the present method
facilitated easy filtration and less interference from impurities.
Acknowledgements
The authors thank Prof. N Appala Raju for his interest in the work and Prof.
G Aravamudan for providing facilities for thermal analysis. The financial
assistance of University Grants Commission, New Delhi, is gratefully acknowledged.
References
Bordner J, Salesin E D and Gordon L 1961 Talanta 8 579
Bullwinkel E P and Nobel Jr P 1959 /. Am. Chem. Soc. 81 813
Cartwright P F S 1967 Talanta 14 690
Cartwright P F S, Newman E J and Wilson B W 1967 Analyst 92 664
Corsini A and Abraham J 1968 Talanta 15 562
Duual C 1953 Inorganic thermogravi metric analysis (Amsterdam : Elsveier) 510
Fleck H R 1937 Analyst 62 378
Fleming J E and Lynton W 1967 Can. J. Chem. 45 1637
Frere F J 1933 /. Am. Cfiem. Soc. 55 4362
Gordon L, Salutsky M L and Willard H H 1959 Precipitation from homogeneous solutic
(New York : John Wiley) 2.
Hecht F and Reich-Rohrwig W 1929 Monatsh. Cfiem. 53-54 596
Heyn A H A and Dave N G 1960 Talanta 5 1 19
Horton G R and Wendlandt W W 1963 /. Inorg. Nucl. Chetn. 25 247
Howick L C and Rihs T 1964 Talanta 11 667
Majee R J and Gordon L 1965 Talanta 12, 441, 445
Majee R J and Woodward I 1966 Talanta 13 209
Milner G W C in Wilson C L and Wilson D W 1960 Comprehensive analytical chemistry
Vol. 1C (Amsterdam : Elsveier) 610
Moeller T and Wilkins D H 1953 Inorganic synthesis 4 (New York : McGraw-Hill) 101
Siva Reddy G and Krishna Reddy Y 1979 Talanta 26 245
Idem 1980 Analyst 105 391
Tackett J E and Sawyer D J 1964 Inorg. Chem. 3 692
Van Tassel J H and Wendlandt W W 1959 /. Am. Chem. Soc. 81 813
Van Tassel J H and Wendlandt W W 1961 ./. Am. Chem Soc. 83 810
Vogel A I 1975 A text-book of quantitative inorganic analysis (London: ELBS and Longmans
p. 539
Wendlandt W W 1956 Anal. Chem. 28 499
Electrochemical studies on copper(II) glucuronatc
R PAYKE and R J MAGEE*
Department of Inorganic and Analytical Chemistry, La Trobc University, Buiidoora
Melbourne, Victoria, Australia 3083
MS received 16 October 1981
Abstract. The complexes formed in the interaction of the copper(H) ion with
glucuronic acid over the pH range 4-0-11-0 wore investigated using d.c. polaro-
graphy, cyclic voltammetry and chronoamperometry. It was found that below
about pH 6-1 no complex forms, while in the pH range of approximately 6-2-7-4
and again 7-5-9-8 stable complexes were formed in solution. At high. pH values,
the complexes appear to break up. The complex formed hi the pH range 6-2-7-4
was studied and stability constants determined by two different methods.
Keywords. Electrochemical investigation ; polarogtaphy, cyclic voltammctry ;
stability constants ; copper(II) glucuionatc.
1. Introduction
In investigations on the uptake of copper by certain bacteria (Payne at al 1981)
a copper complex of glucuronic acid was isolated and the structure compared
with that of model compounds.. As a result of these investigations, it became
necessary to investigate the interaction of copper(II) ions with glucuronic acid
over a wide pH range. This study was carried out using electrochemical techniques.
While a number of reports have appeared (Biswas et al 1978 ; Rajan and Martell
1967) on the polyhydroxy acids and their complexes with Cu(II), particularly
citric and tartaric acids;, little work has been reported on the copper complexes
of the uronic acids in general and glucuronic acid in particular. Makridou
et al 1977 have studied the formation of complexes of the types MA and MA'
between different metal ions, includingCu(II), and glucuronic acid and galacturonic
acid by a potentiometric method. They determined stability constants and
concluded that the metal complexes of galacturonic acid are more stable than
those of glucuronic acid.
In the present paper, the results of a study on the formation of complexes
between the Gu,(II) ion and gtucuronic acid, using polarography methods, arc
presented.
2. Experimental
The polarographic, cyclic voltammetric and chronoamperometric studies were
carried out on the AMBL47l^Multipolarograph System and the Princeton Applied
*To whom all correspondence should be made.
31
Research (PAR) 170 Electrochemistry System. Results were plotted on a Hewlett-
Packard 7040A X^Y Recorder. pH was recorded on an ETI572 Digital pH
meter. The polarographic cell had a three electrode configuration consisting
of a saturated calomel reference electrode and a platinum counter electrode. For
the d.c. polarographic measurements a glass capillary dropping electrode (DME)
was used : for the cyclic voltammetric and chrouoampero metric measurements
the hanging mercury drop electrode (HMDE) was used. Experiments were also
carried out using a glassy carbon electrode. All polarographic and cyclic voltani-
rnetric data were obtained at 25 + 0-02C, solutions being deoxygenated with
p re-dried, oxygen-free nitrogen.
Sodium glucuronate was either prepared by direct titration of the acid with
sodium hydroxide solution or was purchased directly (Sigma Chemicals). In
both cases, the product was recrystallised. Copper nitrate was used as the source
of Cu(H) ions. Stock solutions of copper nitrate were standardized by titration
with EDTA using a potentiometric end-point determination. All results were
obtained at an ionic strength of 0-74 M OSIaClO 4 ). pH values were checked
before and after recording voltammograins.
3. Results and discussion
3. 1. Polarographic investigations
3 -la Effect on E l/z of variation in pH : With the polarographic cell containing
6-25 x 10~ 3 mol dm~ 3 Cu 2 +, a ligand concentration of 237 mol dm- 3 and
0-5 rnol dm- 3 of NaClO 4 , the pH was varied over the range 4-0-9-8. In each
case, polarograms were obtained in the potential range +0-2 to -0-700 volt.
In the range pH 4-0 to about 6-3 or 6-4, one wave (wave I) was obtained with
an jE'i/a value around -0-02 to 0-03 volt and was clearly indicative of Cu(II)
in a 2-electron reduction step. Around pH 6-4, a second wave (wave II) began
to appear in addition to and following wave I. As the pH was increased, wave II
increased steadily in height, while wave I decreased in height. Maximum deve-
lopment of wave II appears to be around pH 7-4. At this pH, a very small
residual first wave (wave I) at JS 1/3 = 0-036 volt still persisted with wave II
showing an 1/2 around 0-14 volt. However, beyond pH 7-4, the residual
first wave disappeared and a third wave (wave 111) now began to develop more
negative than wave II ( 1/2 ^ 0-36 volt). Both wave II and wave III existed
together up to about 9-8 : however, wave II gradually decreased in height, while
wave III increased in height with increasing pH. At high pH values, waves II
and III disappeared and a new wave (wave IV) with E i/z around 4-0-102 deve-
loped. Examples of the development of the waves, mentioned above are shown
in figure 1 (a), (b), (c), (d)
In the pH range 4 -0-6 -3, the single wave (wave 1) present was shown to be
due to the reduction, of free Cu 2+ ion (figure la) i.e., no complex is formed in acidic
solution between Cu 2+ and glucuronic acid. In the range pH 6-4-7-4, the first
wave is due to reduction of Cu 2+ , while the second wave (wave II) is indicative
+2 +-1 -1 -2 -3-4 -5
E(V)vs.SCE
Figure 1. D.C. polarograms at DME at varying pH values of Cu(II) glucuronate
complcxation in aqueous solution, 0-5 M NaCIO 4 . (a) pH - 5-74 (b) nH = 6-4 I
' 0-004% Triton
appearance of wave IV appears to be indicative of the ligand itself. To check
tins assumption, polarograms of the ligand were obtained. At pH 11 -0 a wave
was obtained identical to that obtained for the solution which had the copper
complex present (figure id). It is concluded that this wave may be due to the
reduction of the ligand or a mercury complex formed by the ligand It would
appear then that at high pH values any copper complexes formed at lower pH
values break up releasing the ligand.
3.2. Wave II
3-2a. Variation of limiting current with height of Hg column : Using the same
concentration of On*, ligand and NaClO 4 as for the investigation of the effect
ol pH on . E i/2 , the effect of the height of the mercury column on the limiting
f C< * r C mP , leX ^^ ****** pH 7 ' 4 was mined. Limiting
currents for wave II were found to be proportional to rf indicating diffusion
control under the polarographic conditions. - "uiu^on
E vs log (ijid 1} gave a linear plot from which the slope (average of a series
of results) was found to be 87 mv, indicating a quasi-reversible reduction.
Assuming a value of a not too different from 5, gives a value of n 2 for the
uuinter of electrons involved in the reduction.
32c. E 1/z vs log C (x): With thepHat 7-4, the effect of change in the concen-
tration of ligand on E i/z was examined. A linear plot was obtained indicating
the presence of a single comi lex at this pH. From the relationship
. 0-0591, ^ 0-0591 , ,, % N
A^i/a = log -- j- P ' log (C (x)).
the value p (number of ligands coordinated to metal ion) and log /? were deter-
mined using the slope and intercept of the straight line plot.
The value obtained for p was approximately 2 (1-85). For the stability
constant determination (logy?), a standard pH titration method was also carried
out for comparison with the polarographic method and the copper complexes
of glutamic acid were also determined by both methods. The basis of the pH
titration method, as given by Albert and Sargent (1971) is that the average
irumter of ligands bound by one atom of the metal is defined as
of loimd ligand [Cu (glue)*] + 2 [Cu (gluc) a ]
~ total moles of Cu 2 + " [Cu 2 +] + [Cu (gtac)+] + [Cu (gluey '
This may be re-written in terms of stability constants as
K, (glue-) + 2JC a JC, (glue-)* A(g^-) + 2y? 2 (glac-) 3
1 + K, (glue-) + K& feluc-) 3 1 + /?! (glue-) + 0Y(gluc~p '
On re-arranging we get
__ _ - a . (2 - ) (glue-)
(1 - ) '(glue-) ~ ^ ^ (T=^J / ; 2-
A plot of n/(L -/i)(gluc-) v^ (2 -) (gluc-)/(l n) should give a straight line of
intercept & and slope /? 2 .
In the present work, a computer was used to determine the stability constants
instead of using the graphical plot. Results are shown in table 1.
Table 1. Stability constants of Cu(II) glucuronato and glutamato.
Complex
Method
Literature
velues
Polarography
pH titration
Cu glucuronato
log/?!
log$j 4-10G
1-01 1-48
4-103
Makridou
et al 1977
Cu glutamate log/Jj .. 8-314 8-20
uncut 01 pn on J^ia v i uu 01 WU.YU JLL
pH
-i/a
6-3
7-05
-0-147
6-5
-Q-Q94
7-15
-0-150
6-7
-0-136
7-25
-0-154
6-85
-0-141
7-5
-0-161
6-95
-0-143
8.1
3- 2d. E i/2 vs pH : For wave II, the pH was varied from the value at which
the wave first appeared until its disappearance at the higher j.H. For each pH
the JBi/z va lue was determined. Trie variation oE i/Si with pHis shown in table 2.
From the plot of E 1/2 vs pH a straight line resulted. Now,
Thus, from the slope of the plot which is equal to -0-59 i/n m, nt = number
of hydrogen ions involved may be determined. The average of a number of
experiments gave a value of m % 2.
3.3. Wave III
As indicated earlier, this wave which appears above pH 7-5 may possibly be
the result of the formation of a second complex. Attempts were made to apply
the same sort of tests to it as applied to wave II. However, difficulties were
always experienced in deciding the point at which the wave began as it followed
so closely on wave II. Log plot analyses indicated a value between 1 and 2
electrons for the number of electrons involved in the reduction and there was
evidence for considerable irreversibility.
4. Cyclic voltairimetry
Cyclic voltammograms were obtained at varying pH values, under the same
conditions as used for the d.c. polarographic studies discussed above. Figure
2 (a), (b), (c) shows the results obtained. Below pH^6-l, a voltamrnogram
with one cathodic peak and one anodic peak (figure 2a) was obtained. This
voltammogram showed clearly that only free Cu 2+ ions were present in the
solution. Above pH 6-3 a second cathodic peak appeared and by pH 7-4 this
was the only cathodic peak ;. the first wave due to Cu 2 * had disappeared. Figure
2b shows the voltammogram at pH 7-4. The cathodic peak at E 9 approximately
0-14 volt corresponds to reduction of the same copper complex observed in
the polarographic study at this pH. The anodic peak (!) corresponds to the
ECV1 vs SCE
Figure 2. Cyclic voltaramograms at HMDE (scan rate 20mV/sec, r=25C)
in aq. solution, 0-5M NaClO 4 in presence of nitrogen at varying pH values,
(a) pH = 5-74 ; (b) pH = 7-40_; (c) pH = 8-70. ,
for which there is no corresponding cathodic wave was found and is thought
to be due to oxidation of the ligand or a mercury complex of free ligand. Accor-
dingly, cyclic voltammograrm were obtained for the free ligand. As expected, a
peak was obtained exactly at the same peak potential as' the anodic peak in
figure 2b, confirming that this peak is caused by oxidation of the ligand.
The data for peak l po at different scan rates is shown in table 3. Figure 2c
shows the cyclic voltammogram. at pH 8 -5 of the copper-glucuronic acid system.
T*wo cathodic peaks are present, the first, very small one at E f 0-14 (J^)
representing a residual part of the complex which forms between pH 6-1 and 7-4.
The second cathodic peak with E 9 around 0-4 (ii po ) is the major peak in the
voltamrnogram and represents the complex which forms at pH values greater
than 7-4. The absence of a corresponding anodic peak for the cathodic peak
11^ indicates that the electrode reaction is irreversible. For the small cathodic
peak, / Po , a corresponding anodic peak is discernible (/ Po ) as a shoulder on a larger
anodic peak (III fe ). This again indicated the quasi-reversible nature of the
electrode reaction of this complex. The large anodic peak (in pa ) found in the
voltammograms is of interest.- There appears to be no corresponding cathodic
peak at any scan rate. Initially it was supposed that it was due to oxidation of
free ligand or a mercury complex of the ligand. However, experiments carried
out with the free ligand indicated that it was not a ligaiid peak. On further
investigation,- it was found that this peak of sharp symmetry increased in magni-
tude with increasing scan rate and decreasing concentration, thus showing charac-
Scat rate
V sec- 1
E fe v s SCE
0-005V
*~
(mV)
i,
'..
i
0-002
-0-144
-0-064
80
0-425
0-689
1-62 *
0-005
-0-148
-0-064
84
-453
Q-7Q9
1-57
0-01
-0-156
0-068
88
a -709
0-709
1-00
0-02
-0-154
-0-062
92
0-866
0-778
0-898
0-05
-0-158
-0-065
93
0-90Q
0-720
0-800
Q-1Q
-0-160
-Q-066
94
1-024
0-787
0-762
0-70
-0-180
-0-060
118
1-732
1-339
0-76Q
5. Chronoamperometery
Chronoamperograms were recorded for the complex formed around pH 7-4 by
applying a voltage on the plateau of wave II to the HMDE. The derived
"current versus time- 172 plot " was a straight line. This linear /versus t~ 1/z plot
shows that the electrode reaction responsible for wave II is diffusion controlled
and that there is no preceding chemical reaction coupled with the electron transfer
process.
References
Albert A and Sargeant E P 1971 The determination of iontsation constants (London : Chapman
and Hall)
Biswas S P, Krishnamoorthy T S and Venkateswarlu Ch 1978 Indian J. Chetn. A16 972
Makridou C, Cromer-Morin M and Scharff J P 1977 Bull. Soc. Chfm. Fr. 1-2 Part 1 59
Payne R, Magee R J, Sarode R and Rao CNR 19.81 Inorg. Nucl. Chem. Lett. 17 125
Rajan K S and Martell A E 1967 /. Inorg. Nucl. Chem. 29 463
Shah S K, Suyan K M and Gupta CM 1980 Talanta 27 455
Molecular constants of PSF, and NSF,
A NATARAJAN* and S SOMASUNDARAM
Department of Physics, Autonomous Post-Graduate Centre, TiruchirapalU 620020,
India
MS received 28 July 1981
Abstract. A complete vibrational analysis of PSF 3 and NSF S molecules is described
in this paper. Urey-Bradley and General valence Force Fields have been computed
for these molecules, belonging to C 3e symmetry using the fundamental frequencies
obtained from infrared spectra. The mean Square amplitudes, Coriolis coupling
coefficients and centrifugal distortion constants have also been calculated and
presented here.
Keywords. Symmetry; vibrational frequencies; normal coordinate analysis ; mole-
cular constants.
Ik Introduction
The molecules PSF 3 and NSF 3 possess C& symmetry. Their fundamental
frequencies are distributed accordingly as 3j + 3^.
Considerable amount of work has been done on the spectra of PSP 8 . Normal
coordinate analysis has been carried out by Shurvell (19.69) and Koniger and
MJuller (1977). Recently fresh vibrational assignments have been made for PSF S
and NSF 3 by Koniger et al (1979) and they are used for the present work. The
internal coordinates, numbering of atoms and the orientation of Cartesian
coordinate axes of these molecules are shown in figure 1.
Normal coordinate analysis has been performed by Koniger et al (1979) using
the general valence force field (GVFF). However, the GVFF has been repeated
and in addition Urey-Bradley force flelc* (UBFF) has also been worj^ed out. The
mean square amplitudes of vibration, the generalized mean square amplitudes of
vibrations, shrinkage constants, Coriolis coupling cpejjicients #n4 centrifugal
Distortion constants have been computed and reported here.
2. Theoretical considerations
2.1. Molecular force field
UBFF and GVFF are used to obtain the force constants. A reliable set of force
constants has been ob.tained by Wilsqn's (1939) F-G matri? formalism.
Figure 1. The orientation of Cartesian coordinate axes, the numbering of atom
and the i.itcmal coordinates of XY^Z (pyramidal) ty.ic of molecules.
2.2. Mean square amplitudes of vibration
The symmetrised mean square amplitude matrix S has been evaluated using the
relation (Cyvin 1960)
where the 's are the characteristic vectors and A is a diagonal matrix consisting
of the mean square values of the normal coordinates given by Block (1932). The
generalised mean square amplitudes (Morino and Hirota 1955) {(AZ) 2 )>
and {(A^) 2 ) have been evaluated using the 2 matrix from the relation
where
= AS,
,4 ==M- 1 .B'G- 1 .
2 . 3. Cor folis coupling coefficients
''
active couplings for these molecules are a x e and e x e belonging to *
and " respectively. The Coriolis coupling coefl5:ients have been calculated from
the relation given by Meal and Polo (1956)
where C a matrix is obtained from the geometry and atomic masses.
2.4. Centrifugal distortion constants. , : .
The centrifugal distortion constants have been computed using the relations given
by Kivelson and Wilson (1952, 1953), and by Cyvin. e/ al (1968).
To check whether the chosen set of symmetry coordinates -contributes maximum
to the potential energy associated with the normal coordinates of the molecule,
the P.E.D. was calculated using the relation
100
It is found (table 1) that the contributions to the main diagonal elements are
predominant which confirms that the chosen coordinates form a normal mode
representation.
3. Results and discussion
The molecular parameters, the observed and calculated frequencies of the funda-
mentals are given in table. 1. The valence force constants in GVFF and the
Urey-iBradley force constants are presented in table 2. The values agree very
well with those calculated by Koniger et al (1979). The H-S force constant
in NSF 3 molecule is found to be 12-683 m dyne/A in general valence and 12-369
in UBFF. This indicates the doiible bond nature of the N-S bond.
It is slightly larger than ^_ s in other molecules such as NSC1 [/ N _ s : 10 57 m dyne/
A, Namasivayam and Nair 1978; 10 '15m dyne/A, Peacock ef al 1969; 10 -41m
dyne/A, Nagarajan at al 1967]. This may be due to the delocalisation of electrons
in the N-'S bond region and consequently enhancement of nitrogen-'SUlphur back
donation of electron density. This results in the reduction of such back donation
from fluorine to sulphur. Hence the f s - p is comparatively small.
The mean square amplitude quantities at 300 K calculated using L matrix
elements are presented in table 3. The mean amplitudes of vibration at 300 K
for bonded and non-bonded distances are given in table 4, and the generalised
mean square amplitudes (both parallel and perpendicular) of vibration along with
the shrinkage constants are given in table 5. The mean amplitudes of vibration
for non-bonded distances are larger than those for bonded distances. The Coriolis
coupling coefficients of PSF 3 and NSF 3 molecules are listed in table 6. The
Coriolis coupling between the two degenerate modos of the fundamentals v 4 and
v s are fairly strong for rotation about the Z axis. The Coriolis coupling coeffi-
cients satisfy the sum rule
2/ B .
where / x and 7 B are the principal moments of inertia whose calculated values are
also given in table 1. The calculated values of * 4 , 6 and S OG agree very well
with the experimental values obtained by Koniger et al (1979) in the band
contour analysis of the spectra and also with the experimental values investigated
in the microwave analysis by Small and Smith (1976). These values are also given
(2 <s ON 7 s T t r~
Ai ^ VO T--
-* "* oo ^
+ t + t t t
to j* * C^j to C/5
g
05 CTv .. 01 VO ^J"
vo a\ L~ o >n n
* T^ ,1,
AH
ri* Q^ O4 ^^ gf\
^ -i -1- -1- 4-
V
o? to ^ to* 3 ?
i
_M t: t^. 00 ON O
ON ^ oo O *?
M
ro "* ^
pi*t
M
r- ^,00
CO
p-l
CO
P-
3
rn r-1 fi T 1 T
C/5
"ca
U
^ S 't ^" 2 M
vo as ^. o, * M
t*-<
O
,
|
.'
' o<^ ^ ?n oi ON < i
t f>9ooO t T".vD^ Q (!, ^j. vb
O
vO ^^ ^r OS J
o
1
i
cG
tQ
2' co- ^* co ^*
O.
r S SS 2 ?
Cj
.'
+ * + "H ?
jl
. a
o<
^ ^ S " ^o
g
vJ3 oo Q ,1, J_, o
s3
-1- 4- -|- + 4-
5"
CO" - rf ^ *
o
r;
1
>n O <^> r'l o\ O
s
M
r<^ . ^ o >n *p
fa
ON O o cij ^| m
u
r-5 O
00 sO VO o .
o
1
;' 3
'-', i ^ ^ ?a ^
g 8 53 * cn
d!
u
fc
o ^5 9O
a
8
F-1
1
r-o_._ J .ootM VD rt ?^ f^Os^
. T* lr> ^ AJ
pj <N r ,j ^ rl ^. . f ^. o) os 5
^.^^SS'*^^^ *
3
'
at
H
* *-- ^ "* ** o>f*
^ pk pk. >. ^ S^ ^^i
=J
"N ^ ^ "N "^
1
r> li<' K. uLl G tl
x -' cS s, S. H
Valence force constants Urcy-Bradlcy force constants
PSF 3 - NSF 3 PSF 3 NSF 3
/R
5-8521
12-6830 2$
: K 5-8269
12-3692^
/r
6-2036
4-8850
r 6-2181
4-5541
/
0-3410
0-2620
f a 0-6759
0-6043
/Rt
0-4363
0-1218 13
5j 0-4287
0-2532
/cT/cm
0-5390
0-7210 /,
, y 0-2349
0-2252
fp-fpp
0-7338
0-3190 f s
z 0-1762
0-1410
/R/S
0-7243
0-3372
/r/J
0-6362
0-4131
f ta
0-0411
0-1862
f
-0-0237
0-0460
Table 3; Mean square amplitudes- of vibration of PSF 3 and NSF 3 tit 300 K
(10- 1 A 2 ).
PSF 3 NSF a
<r R 16-0720
11-5866
a t 12-9593
17-9935
cr rr - 0-7306
- 0-4385
<r Rr - 1-2124
- 1-2126
o- -cr tta 94-3682
119-6820
OOURR 81*3021
131-1066
<r rtt - 8-1070
- 3-3928
<r r jj - 5-2334
- 7-9621
a K p ~ 2-3510
- 5-9198
a aB 17-6240
8-7983
Table 4. Mean amplitudes of vibration (A) of bonded and jioubojidod distances
of PSF 3 and NSF 3 at 300 K.
PSF a NSF
0-0401
-0340
L
0-0360
0-0424
'*
0-0723
0-0764
/...'
0. ; Q427
0-0716
rerpenaicuiar
j.vMjivw>uiti ^viuJ" 1 "I* laiaiici OJllluKagO
F S
16-0712
27-3312
27-3312
P-F
12-9521
29- 6321
41-2213
t
PSF F...F
52-2830
16-3217
37-9424
0-0026
S...F
18-2137
23-1522
36-5628
0-0082
N S
11-5832
21-7372
15-4000
S F
17-9922
53-5862
68-2162
, ,
NSF 3 F...F
58-3631
134-6021
154-2133
0-0052
N...F
51-2216
72-1325
83-2165
0-0093
Table 6. Coriolis coupling coefficients of PSF ( and NSF 4 molecules.
Coupling 01 X e ;
PSF a
NSF 3
&4
0-6724
0-3453
l5
-0-6631
-0-3112
l8
0-0492
0-5138
r
-0*2961
-0-2191
^26
0-1276
0-8465
26
0-3149
0-1049
34
0-3911
0-1739
36
0-3462
0-0332
Csa
-0-5228
-0-5082
Coupling e x e
*
0-5813
0-5612
(0-57)*
(0-5)*
(0-59)**
'55
-0-5891
-0-2330
(-0-56)*
(-0-2)*
(_0-49)**
(-0-2256)*
yn
<i 6
0-2592 .
0-1482
(0-26)*
(0-15)*
(0-16)**
(0-1567)f
C*45
-0-6525
0-4070
C*48
0-5441
0-6066
b 56
0-4631
-0-0177
* band contour analysis : Koniger et al (1979), ** spectral analysis : Clark and Ellestad (1976)
f microwave analysis : Small and Smith (1976)
PSF. NSF 3
1-0678 0-2971
(1-082)* (0-2742)*
-1-4735 -1-1863
1-9763 1-9363
(1-9047)* (1-961)*
Band contour analysis : Koniger et al (1979)
The centrifugal distortion constants are presented in table 7. Since these mole-
ilesare symmetric tops, the centrifugal stretching coefficients R 5 , R Q and &j vanish
ie other coefficients D J9 D K and D JK have been evaluated. These values agree
:ry well with those obtained by Koniger et al (1979) in the band contour
la lysis.
Conclusion
is expected th<tt the analysis presented here would help us in knowing the
2ctro5coplc properties of these molecules. The mean amplitude calculations
e iiseful in the interpretation of electron diffraction studies in the molecular
-ucture determinations and the shrinkage constants are helpful in the refinement
1 bond lengths obtained experimentally. The Coriolis constants arc used in the
lerpretation of vibration-rotation spectra of these molecules. The centrifugal
Portion constants calculated here for NSF 8 will be useful in the study of
icrowave spectra of the molecule.
jferences
ock F 1932 Z. Phys. 74 295
ark J H and Ellestad OH 1975 /. Mol. Speetnvc. 56 3P6
vin S J 1960 Acta Polytech. Scand. Ph 6 279
vin S J, Cyvin B N and Hagen O 1968 Z. Naturforsck. A23 1649
velson D and Wilson Jr E B 1952 /. Chem. Phys. 20 1575
velson D ai d Wilson Jr E B 1953 /. Chem. Phys. 21 1229
miger F and Muller A 1977 Spectrosc. Chim. Acta A33 971
miger F, Muller A and Blom C E 1979 /. Mol. Spectrosc. 77 76
sal J H and Polo S R 1956 /. Chem. Phys. 24 1119
orino Y and Hirota E 1955 /. Chem. Phys. 23 73?
igarajan G, Muller A, Glemser O and Cyvin S J 1967 Spectrosc. Chim. Acta A23 2863
iraasivayam R and Viswanatha Nair 1978 Z. Phys. Chem. Leipzig 259 1117
acock C J, Heidborn U and Muller A 1969 J. Mol. Spectrosc. 30 338
urvell H F 1969 Spectrosc. Chim. Acta A25 973
High resolution electron microscopy of chloritoid minerals from -
different geological melieu
G TS SUBBANNA and G V ANANTHA IYER*
Materials Research Laboratory and Department of Inorganic and
Physical Chemistry, Indian Institute of Science, Bangalore 560012, India
MS received 14 December 1981
Abstract. Structural defects of three chloritoid minerals from distinct geologic
melieu have been investigated by high resolution electron microscopy. X-ra> powder
and electron diffraction patterns indicate that the chloritoid from one geological
source (A) is 2Mi +-2M 2 monoclmic variant while those from another geological
source (B) are 2M Z monoclinic variants. In a typical one-dimensional lattice image
of a crystal from source A, the 2M 2 matrix is broken by insertion of triclinic inter-
growths. Another crystal with the 2M 2 matrix showed single, triple, quadruple and
quintuple layers displaying an unusually high degree of disorder. Lattice images
of 2M 2 monoclinic variants from source B yielded more homogeneous micrographs.
The important rinding from the present Studies is that the chloritoid from source A
is a Severely disordered low-temperature intermediate phase in the conversion of the
triclinic cHloritoid to the high-temperature ordered monoclinic variants of source B.
Severely disordered chloritoids, marking the beginning of low grade metarnorpliism,
are generated as intermediates between the state of complete disordered arrangement
towards the end of low grade metarnorphism within the narrow stability range of
40a-500 C.
Keywords. Chlaritoid mineral; electron microscopy; lattice images; different
geological melieu.
1. Introduction
Pine details of mineral structures have long been successfully obtained \ y x-ray
diffraction techninu.es. In spite of the unquestioned power of x-ray crystallography,
there is no way we can obtain direct information on the ultramic restructure of
solids by this technique. High resolution transmission electron microscopy is
well suited for the study of deviations from perfect periodicity in minerals because
it provides direct interpretable lattice images down to 2 A resolution. Thus,
column defects, stacking faults and intercalation of different structural units have
been examined by the lattice imaging technique (Anderson 1978) During the
last few years there have been many studies utilising high resolution electron micro-
scopy to study minerals (e.g. see Hutchison et al 1977 ; Thomas et al 1979 -
Buseck 1979). We considered it worthwhile examining lattice images of a mineral
* To whom all correspondence should be made.
47
iiUiU UUJ.GJ.C1LL bUULCCSi LU 1IUU UULUUW LUG g
structure and whether the observed ultramic restructure can reveal the mechanism
of reactions occurring during metamorphism. For this purpose we have investi-
gated samples of chloritoid occurring in two different parts of Karnataka South
India, having distinct geological environments. An attempt has been made to
correlate the observed structure with the thermal history.
2. Experimental
Of the three chloritoids selected for the present study one (KB 7543) is from
Kibbanahalli belt east of Banasandra (13- 15', 76- 41'), Tumkur district of Karna-
taka and the other two samples (HN 7801 and HN 7808) are from Shingarana-
halli (12 52' : 76 15') and HaradanahalU (12 51' : 76 14') both situated north
east of Holenarasipur belt in Hassan District of Karanataka. We designate
Kibbanahalli and Holenarasipur belts as sources A and B respectively. KB 7543
chloritoid occurs in a greenish buff coloured highly fissile rock as dark thin
prismatic crystals along restricted foliation planes. The coexisting minerals in
the rock are chlorite, mica and quartz. HN 7801 occurs as dark stubby porphyro-
blastupto a cm in size, in a light green foliated rock with phengite and chlorite.
HN 7808 chloritoids are dark stubby porphyroblast crystals occurring in a matrix
of olive green chlorite and with almandine garnet. The chloritoid crystals from
the rock specimens could easily be separated. They were crushed, powdered
and pure fractions were obtained using heavy liquids like bromoform and methylenc
iodide. The chemical analyses data of the samples are given in table 1. X-ray
powder diffraction analysis were carried out with Philips x-ray diffractometer
using CoK radiation. The unit cell constants determined from x-ray studies are
comparable with the cell constants reported by Jefferson and Thomas (1978).
Table 1. Chemical composition of the chloritoids.
Sample No. KB 7543 HN 7801 HN 7808
SiOs
25-41
26-49
24-08
TiO z
1-02
1-23
1-88
A1 2 O 3
33-89
39-98
35-03
Fc 2 O a
4-01
3-12
6-04
FeO
26-68
19-52
21-63
MnO
1-04
0-84
0-92
MgO
0-95
1-82
2-89
CaO
0-07
0-05
0-08
N^O
0-14
0-19
0-30
K 2
0-05
0-04
0-01
H 2 0-i
6-58
6-66
6-98
Ho~
0-05
0-05
0-04
A double tilt holder was used for orienting the crystals such that C* axis is per"
pendicular to the electron beam. 70 pm as well as. 50 /mi objective apertures
were used for forming the image and through -focus images were recorded at various
focussing conditions for all specimens. Generally an underfocus of about 80nm
relative to the Gaussian image is. necessary to obtain images of optimum contrast
which could be directly interpretable in terms of the structure of the crystal and
of the defects within it. Since the main aim in this study is to determine the devia-
tions from perfect periodicity in chloritoid minerals obtained from different geo-
logical melieu by direct imaging of non-periodic features to characterise the struc-
ture type and their intergrowths, one-dimensional lattice images are sufficient to
investigate the stacking disorders. Chloritoids cleave on (110) as well as (001)
planes, ?nd a crystal could be easily oriented with one of the principal axes normal
to the beam. With the goniometer stage ( 25 max tilt) fitted to the Philips EM
301 electron microscope used for the study it was possible to record images of only
one section even though in principle it is possible to obtain images, from the next
section by tilting the crystal through 30. This limitation is partly due to insuffi*
cient tilting and also due to increase in crystal thickness after tilting it through 30.
The imaging code was. established experimentally from the analysis of characterised
species following the method of Jefferson and Thomas (1978).
3. Results and discussion
3.1. Structure of chloritoid
Close association of rnonoclinic and triclinic chloritoid with the ideal structural
formula
(Fe+ 2 Mg) 2 A1(OH) 4 Al a O 2 (SiO 4 ) 2
is frequently described in the literature and specimens of chloritoid giving x-ray
powder patterns characteristic of both structure types have been reported
(Halferdahl 1957). The rnonoclinic chloritoid structure has been solved and
refined (Harrison and Brindley 1957; Hanscom 1975). The triclinic structure is
topologically very similar to monoclinic polymorph and the structure has been
recently resolved by Hanscom (1980),
The monoclinic and triclinic chloritoid structures consist of alternating brucite
and corundum-type octahedral sheets joined by isolated SiO 4 tetrahedra. Inter-
layer hydrogen bonding occurs in the monoclinic structure (Hanscom 1975).
Because the triclinic structure is so similar to monoclinic chloritoids, interlaye r
hydrogen bonding is assumed (Hanscom 1980). Halferdahl (1957) suggested
that the monoclinic chloritoid structure resulted from twinning of triclinic on a
unit cell scale. One possible mechanism of twinning of the triclinic structure on a
unit cell scale is rotation of one unit cell relative to another (Hanscom 1980). If
two triclinic chloritoid unit cells are stacked one on top of another so that their
c-axes are collinear, and the upper unit cell is rotated clockwise by 60, the result,
ing unit cell is very similar to that of a inonoclinic polymorph.
x-ray diffraction, studies have proved the existence of two monoclinic variants of
cMoritoid designating them as 2M a and 2M Z , in accordance with the notation
used for micas. 2Mj_ is the new monoclinic structure, the cell dimensions of this
variant in contrast with those of the 2Af 2 structures is that x and y axes are
interchanged, with c-repeat and angle being completely different. Wnen viewed
down on (110), the 2M Z and 2M l chioritoid structures are identical. Unit cell
dimensions for the three structural types of chioritoid by x-ray diffraction methods
are as follows :
A 2091
Triclinic (ITc) = 9-43 A a = 9603'
b = 5-48 A = 101 52'
c = 9 14 A y = 90 00'
A 2082
Monoclinic 2M 3
(2M a + 2MJ a= 9 -47 A
b = 5-48 A /? = 101*39
e = 18-14A
fl= 5-47 A
Z> = 9 -47 A p~ 97 24'
c= 17-90 A
A 2092 a= 9-47A
Mouoclinic2M 2 b= 5-48 A /? = 101 39'
tf=18-14A
Jefferson and Thomas (1978) conclude that the 2Afj variant is possibly an inter-
mediate phase in the conversion of the triclinic to monoclinic (2Af 2 ) structure. In
addition they have observed three layer longer period chioritoid structure in the
lattice images of 2M 2 regions of trie lime chioritoid. KB 7543 corresponds to
A 2082 while HN 7801 and HN 7808 arc similar to A 2092 in their cell
dimensions.
3-2. Lattice images ,
Electron diffraction patterns of chloritoids (KB 7543) from Kibbanahalli belt
(source A) had predominant streaking parallel to (001) direction indicating high
disorder. Figure 1 shows a typical one-dimensional lattice image of KB 7543
wherein the 2Afi matrix is broicen by insertion of triclinic (ITc) intergrowths
characterised ly the halving of the fringe spacing. The lattice image shown in
figure 2 is also from KB 7543 chioritoid, principally of the 2M Z matrix with inter-
growths of single, triple, quadruple and quintuple layers indicating unusually high
degree of disorder not observed by Jefferson and Thomas (1978) in (A 2092)
monoclinic (2Mi + 2M^ chioritoid samples. In their study, lattice images of
High resolution electron microscopy
51
fs
'I
nl
Tj-
m
c--
PQ
a -is
M IU
fi .3
High resolution electron microscopy
<*>
I
s
a
S
C4
c3
*-.
O .
'/l
r? S
1 1
I
O
s
VJ
O *
'So w
o ^J
.
ll
Wl *O
s .s
a iar more Homogeneous piuiiue man me cmumum xso
of disorder, has been observed by Jefferson and Thomas only in the triclinic chlori-
toid (A 2091). In figure 2, two 4-layer intergrowths appear side by side flanked
on either side ty 3-layer intergrowths, within the 2M Z matrix. In addition,
1-layer and unusual 5-layer intergrowths are also present in the micrograph. The
3 and 4-layer intergrowths cannot be considered as genuine polytypes because
they do not repeat three times in a sequence according to Buseck and lijima (1974).
According to Jefferson and Thomas (1978) and Jefferson (1980) the 3 and 4-layer
intergrowths are genuine polytypes if they repeat twice in a sequence.
Electron diffraction patterns of HN 7801 chloritoid crystals from Holer.arasipur
belt (source B) indicate the principal structure to be 2M Z monoclinic. The lattice
image shown in figure 3 is predominantly of the 2M Z monoclinic variant interrupted
by strips of 3-layer intergrowths with rare triclinic intergrowths.
Electron diffraction patterns of HN 7808 chloritoid crystals from Holenarasipur
belt (source 3) showed the principle structure to be 2M Z monoclinic variant. The
electron micrograph given in figure 4 clearly shows the true structural periodicity
of 2M Z with triclinic (ITc) intergrowth characterised by halving of the fringe
spacing. In contrast to KB 7543, HN 7801 and HN 7808 display a far more
homogeneous picture. The three layer intergrowths observed in HN 7801 were
also present in some of the examined crystals of HN 7808.
33. Geologic implications
Normally triclinic chloritoid appears at the beginning of low grade metamorphism.
The lower thermal stability limit for low grade metamorphism given by Wir.kler
(1974) is 400 C. Jefferson and Thomas (1978) have observed in triclinic chlori-
toid a kind of unusual and severe degree of disorder. In the present work such
an unusual degree of disorder is observed in the monoclinic chloritoid KB 7543
(2Mj + 2M 2 ) variant which appears at the legmning of low grade metamorphism.
The 2M 2 variant of chloritoid HN 7808, which displays a homogeneous picture,
coexists with chlorite and almandine garnet suggests, that the stability limit of
chloritoid corresponds to almandine 1 eld. Their thermal stability limit of alman-
dine, according to Winkler (1974) is 500 C at 4 kb pressure. The thermal stabi-
lity limit of the monoclinic chloritoid HN 7801 coexisting with phengitc and chlo-
rite can be inferred to be intermediate between KB 7543 and HN 7808, i.e. 450 C.
4. Conclusions
The present study of the chloritoids from different geological melieu shows that
the monoclinic chloritoid KB 7543 (from source A) is a severely disordered inter-
mediate phase in the conversion of the triclinic chloritoid to the 2M Z monoclinic
variant. It is clear that the severely disordered 2M + 2M Z variant marking the
beginning of low grade metamorphism from Kibbanahalli belt (source A) is gene-
rated as an intermediate between a state of almost complete disorder and a com-
pletely ordered arrangement as in 2M Z monoclinic chloritoid, HN 7508 (from source
3) coexisting with almandine garret from Holenarasipur belt. Furthermore,
fields of the entire range of low grade metamorphism. The lattice image studies
of the chloritoids indicate increase in ordered arrangement with progressive meta-
morphism.
Acknowledgement
The authors thank Professor C H R Rao for suggesting the problem and
Dr D A Jefferson for his advice in obtaining lattice images.
References
Anderson J S 1978 Proc. Indian Acad. Sci. A87 295
Buscck P R 1979 Proc. EMAG 93
Buseck P R and lijima S 1974 Am. Mineral. 59 1
Halfcrdalil L B 1957 Carnegie lust. Washington, Pap. Geophys. Lab. p. 200
Hanscom R H 1975 Acta Crystallogr. B31 780
Hanscom R H 1980 Am. Mineral. 65 534
Harrison F W and Brindley G W 1957 Acta Crystallogr. 10 77
Hutchison J !>, Jefferson D A and Thomas J M 1977 In Surface and defect properties of solids
Vol. 6 (London: The Chemical Society) p. 420
Jefferson D A 1980 Proc. EMAG 113
Jefferson D A and Thomas J M 1978 Proc. R. Soc. London A361 399
Thomas J M, Jefferson D A, Nfellinson L G, Smith D J and Crawford ES 1979 Cham. Scipt.
14 167
Wiriklcr H G F 1974 P&trogeneaia of metatnwphic rocks (New York; Springer-Verlag)
Strengths of some N-H ... * type of hydrogen bonds
G V G KRISHNA MURf Y and B SUBRAHMANYAM*
Department of Chemistry, Osmania University, Hyderabad 500 007, India
MS received 25 May 1981 ; revised 4 September 1981
Abstract. Hydrogen bond formation between N-H of acetanilide and -n electrons
of some arenes has been studied by infrared spectroscopy. In the mixtures of CCU
and arenes new N-H bands appeared at lower frequencies in addition to the N-H
band observed in pure CCU. The lower frequency bands are assigned to the N-H
bonded to it electrons of the arenes. Formation constants of these complexes have
been determined with different arene concentrations. The frequency shifts of the
N-H band and the formation constants were found to increase with increase in
alkyl substitution in the benzene ring. The relative frequency shifts and logarithmic
values of formation constants bear linear relationship with each other.
Keywords. N-H ... u hydrogen bonds ; formation constants; ionisatiorj p6ten-
tials ; frequency shifts.
1. Introduction
Extensive studies were made on hydrogen bonds formed between proton donor
groups and Lewis bases containing non-bonded or w-electrons (Searles and Tamres
1951 ;. Pultin and Werner 1965). In the early fifties evidence was shown for the
formation of hydrogen bonds with n electrons of aromatic compounds (Jones and
Badger 1951 ; Tamres 1952 ; Josien and Sourisseau 1959). Infrared studies
involving frequency shifts and intensity changes of the proton donor groups like
hydrogen halides, OH and OD in aromatic hydrocarbons was subsequently carried
out (Cook 1956 ; Basila 1961).
Association constants of the complexes between aromatics and OH were deter-
mined by Josien et al (1958), Basila et al (1965), Golinska et al (1968) and
Yoshida Zenichi and Ishibi Nobuyuki (1969). The complexes between N-H and
it electrons did not receive much attention. Josien et al (1958) determined the
association constants of the complexes formed between N-H of pyrrole and a few
aromatic compounds. Hence it was considered worthwhile to determine the
association constants of N--H. . . complexes using different types of N-H com-
pounds. We report in this paper the association constants of complexes formed
between acetanilide and some alkyl benzenes.
* To whom aN correspondence should be made.
58 G V G Krishna Murty and B Sitbrahmanyam
2. Experimental
2. 1. Materials ;
Acetanilide was recrystallised twice from 200 ml of distilled water containing 4 ml
of methylated spirit. The crystals were dried in an Abderhalden drier and kept
in a vacuum desiccator. Benzene, toluene, xylene, mssitylene were purified by
the nuthods suggested in the literature (Vogel 1971 ; Oilman 1932). They were
kspt overnight on the drying agent and distilled. The distillates were collected
at the appropriate boiling points. They were again dried over sodium wire to
remove last traces of moisture, and then distilled twice before use, collecting each
time the middle fractions.
2-2. Measurements of spectra
As acetanilide is known to exhibit self-association in solutions of 0-01 M and
above, the infrared spectra oT this compound were recorded at 24->25 C in dilute
solutions (0-002 -fl/ to 0-006 M") of CC1 4 on Perkin Elmer 337 grating spectro-
photom^ter using matched quartz cells 0-5 cm path length. The N-H band of
acetanilide was found to obey Beer's law at these concentrations (figure 1). For
obtaining the free and bonded IShH bands the spectra of acetanilide were recorded
in CC1 4 containing varied amounts of alkyl benzenes (0-5 M to 6M). Solvent
mixtures of the sams composition were placed in the reference beam. The fre-
quencies of free and bonded IShH stretching bands, the absorbances of the free
N-H bands and the concentrations of acetanilide and alkyl benzenes are given in
table 1.
055
0-40
0'20
ws 4
d
"I-
J3
| ^2
f\ oo r-l oo M
">
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crt
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5
O ^
o
IS
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r\ r \ rO j-r, ^j-
'5
o "5
5
Cu %
o
c
"S ""*
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"o
.D
* 1 r
^ fq *->
35 S C: M n
si 1 s s
.2
d ^
t^T
a
S ^
^
c?
a 1 S^
11
-S
o o i^
S ^ o o
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rj ' T;' .* "\5 fO
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sSi
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S
an
'
5
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d
r 1
cCJ
Si
s?
o
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'S S "*
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a a o -s
ia>
MO o in r^
bo <b 06
&
13
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^a S c
t i TA r4
: S S * ?? : ?f^^^S
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W
bb bo bb 6b bb
.1
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g
o
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83
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Sg 12 ggs g gg
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obb bo bob bb bb
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rt *S f-T*
e
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^ -r 3 -s 2
si o v
jn<n>n <nv> oooooo ooo roin
i
es
-j Mt *> -25 /s
C 55 (D rt
o o o is r ^
v ^r ^^ ^5j> nf ^j- ^
M
cS
r "^ +3 cS c ^*^
2
w -^x
<a
_o
, s
*"*
o
|dll
OO Oo OQ a o ^_
t>m Os r v5 ac_I "^oo
oo-* Si^ Sso^ S? 2 v^--
' ? T ^T 1 IVT o^ 5-58
I
8
S '^ 1 1
COD (Nl-* rMvH ,_ ^, ^^
3
o
cu
u ^
I
iH
J
d
g .
i|
*o ^
1
1 I 8 f !l
I 1 ! 5 I 1
1
1
7<
-i* ri n rt v>
M D (a G)/m
[M A - M D + (a - c)lm] (a - c)/m *
(1)
where M D and M A are the molar concentrations of the proton donor and proton
acceptor respectively and a is the absorbance of the free N-H stretching band
in the presence of proton acceptor, m and c are the slope and intercept of the
Beer's law plot. In. the present investigation the intercept c is zero. The values
of K thus calculated are presented in table 1, together with the frequency shifts
and the values of ionization potentials of the alkyl benzenes.
3. Results and discussion
Infrared spectrum of acetanilide in dilute CC1 4 solution exhibited only one N-H
band at 3430cm" 1 . In the presence of alkyl benzene a new band appeared at a
lower frequency the intensity of which increased with increase in concentration
of alkyl benzene with simultaneous decrease in the intensity of the band at
3430 cnr 1 (figure 2). The band at lower frequency is therefore attributed to N-H
bonded to the n electrons, of the aromatic hydrocarbons.
Further the frequency shift which is the difference between the frequencies of
the free and bonded N-H stretching vibrations generally depends upon the dielec-
tric constant of the solvent and local association or hydrogen bonding. Absence
of linear relationship between the relative frequency shifts and (D~l)/(2D + 1)
(figure 3) suggests that the dielectric constant of the arenes plays less significant
part than hydrogen bond formation between N-H and n electrons of the alkyl
benzene.
0-00
0-10
z
<t
CD
o 0-20
CO
0-30
3800 3400 3000
FREQUENCY (CM* 1 )
Figure 2. N-H band of acetanilide in (a) CC1 4 , (b) CC1 4 containing benzene
(c) CC1 4 containing toluene,
240
b
a
o. 236
o
X
: l
MS
' 'i
232
* o
?pft
1 1 1 1 I . j ,
0-60
00
Figure 3. Plot of AV/V vs. ( D - 1)/(2D -I- 1) : (a) benzene, (b) toluene, (c) w-xyieiie,
(d) 1,3,5-mesitylene.
CM
O
60
20
<j
0-80 -
0-50
8-00 10-00
potential (e.v.)
Figure 4. Plot of A/v of N-H strctcliing band of acetanilide v.y. ionisation poten-
tial of the alkyl beazcao, (a) benzene, (b) toluene, (c) m-xylene, (d) 1, 3,5-mcsitylene,
(e) hexa methyl benzene.
0-60 -
0-400 0-600
1+ log k
Figure 5. Plot of Av/v vs. (1 + log K).
0-800
Tlie magnitude of the frequency shift is usually taken as a measure of the strength
of the hydrogen bond formed (Gordy 1939 ; Gordy and Stanford 1940, 1941).
The small frequency shifts observed in the present study indicate weak hydrogen
bonds. Although the frequency shifts are small, they are sufficiently pronounced
to distinguish among the basicities of aromatic compounds. A linear plot is
observed between Av/v and IP of the aromatic base (figure 4). The shifts are
found to increase regularly with increase in alkyl substitution in the benzene ring.
Introduction of alkyl group in benzene ring decreases the IP and enhances the
electron donor ability of benzene and consequently its proton accepting ability.
Hence the strengths of hydrogen bonds formed with aromatic compounds studied
vary in the order
hexa methyl tenzene > mesitylene > m-xylene > toluene > benzene.
Formation constants also exhibit the same trends. There is a regular increase
in the K value as the number of methyl groups in the benzene ring increases. As
both frequency shifts and formation constants determine the strength of hydrogen
bonds, a correlation between the two is hopefully expected. Correlation of this
type fails when there are steric effects hindering the formation of hydrogen bonds
because frequency shifts are presumed to depend upon the energy of the hydrogen
bond whereas the formation constants depend upon energy as well as entropy.
The linear plot of Av/v vs (1 + log K) (figure 5) shows that there are no steric
effects involved to decrease the chances of hydrogen bond formation.
Acknowledgements
One of the authors (GVGK Murty) is thankful to the University Grants Commis-
sion for the award of a fellowship. The authors wish to express their grateful
Basila M R, Saier E L, and Cousins L R 1965 /. Am. Chem. Soc. 87 1665
Britz R J N, Devris M J and Robenheimer H G 1968 J.S. Afr. Chem. Inst. 21 183 (Eng.)
Cook D 1956 /. Chem. Phys. 25 778
Gordy W 1939 /. Chem. Phys. 7 93
Gordy W and Stanford S C 1940 /. Chem. Phys. 8 170
Gordy W and Stanford S C 1941 /. Chem. Phys. 9 204
Oilman H et al 1932 Organic syntheses collective volume (New York : John Wiley) p. '.
Goliaska M, Mikola J J and Szezcpaniak K 1968 Acta Phys. Pol. 34 421 (Eng.)
Jones L H and Badger R M 1951 J. Am. Chem. Soc. 13 3132
Josien M L, Nelson F and Paul P 1958 /. Chim. Phys. 55 454
Josien M and Sourisseau G 1959 Hydrogen bonding (New York : Paragon) p. 129
Kartha V B, Norman Jones R and Robertson R E 1963 Proc. Indian Acad. Sci. ASS 21<
Pullin J A and Werner R L 1965 Spectra Chim. 21 1257
Searles S and Tamres M 1951 /. Am. Chem. Soc. 73 3704
Tamres M 1952 /. Am. Chem. Soc. 74 3375
Vogel A I 1971 Text-book of practical organic chemistry 3rd ed. (London : Longman) p. 172
Yoshida Zenichi and Ishibi Nobuyuki 1969 Bull. Chem. Soc. Jpn. 42 3254
Preparation and molecular configurations of some salts of
dipicrylamine with organic and inorganic cations
M L KUKDU, J N KAPOOR* and S K GHOSH
Physical Research Wing, Fertilizer (Planning and Development) India Ltd.,
Sindii 828 122, India
* Chemical Research Wing
MS received 7 July 1981
Abstract. An investigation has been made on the structural characteristics of a
class of salts of dipicrylamine with organic and inorganic cations. A general con-
clusion regarding the mechanism of salt formation, nature of bondings and molecular
configurations of this class of salts has been suggested from the studies on infrared
spectra of these salts. The main features of the spectra of all the salts are almost
similar showing similar nature of bondings and molecular configurations for all the
salts. In this class of salts either organic or inorganic cations are linked with
dipicrylamine anion by van der Waals forces through oxygens of one of the nitro
groups of the dipicrylamine. The nitro groups are twisted out of plane of
the benzene rings.
Keywords. Molecular configuration ; infrared spectra ; cholinium salt with
dipicrylamine ; acctylcholinium salt with dipicrylamine ; trimethylammonium salt
with dipicrylamine.
1. Introduction
Dipicrylamine i.e., 2, 4, 6-2',4',6'-hexanitro diphenylarnine is an organic analyti-
cal reagent which due to its acidic character, on treatment with sodium, ammonium^
calcium and magnesium hydroxides forms corresponding soluble salts. Kertes
(19561) was the first to investigate that dipicrylamine can react with organic bases
like alkylamines, pyridine, piperidine etc. and studied the possibility of methods
for colorimetric estimation of these organic bases. An investigation on the salt
formation of dipicrylamine with organic and inorganic bases has been carried
out in this laboratory and the method has been successfully utilized in the gravi-
metric estimation of the organic and inorganic bases in presence of common
contaminants. The mechanism of salt formation and the structural details of
these salts were not known. The structural characteristics of guanidinium
(Ghosh et al 1969) and pyridinium (Kapoor et al 1972) salts with dipicrylamine
have already been reported. In the present discourse the structural characteristics
mode of salt formation of dipicrylamine with organic and inorganic bases.
2. Experimental
2. 1. Preparation of the salts
To prepare the salts, f rst mexan (magnesium salt with dipicrylamine) was prepared
by treating a mixture of magnesium oxide ?nd dipicrylamine in the ratio 0-416 : 1
(by weight) in a given quantity of water with constant stirring and the precipitate
(mexan) thus obtained was used as reagent for other salts of dipicrylamine due
to its higher solubility in water compared to that of dipicrylamine.
The cholinium, acetylcholinium and trimethylammoniiim salts with dipicrylamine
were prepared by adding dropwise the aqueous solution of the respective chloride
salts in 3% aqueous solution of mexan with constant stirring at room temperatvre
(28-30C). The reactions were instantaneous giving red crystalline precipitates
of the cholinium, acetylcholinium and trimethylammoaium salts with dipicrylamine.
The precipitates were washed with ice cold distilled water and dried to constant
weight at 100 C. In the case of trimethylammonium salt, the trimethylammonium
chloride was first prepared by neutralizing an aqueous solution of trimethylamine
(40% wt/vol) with requisite amount of JV/10 hydrochloric acid using methyl red
as indicator. The neutral aqueous solution of trimethylammonium chloride was
allowed to concentrate on a water bath to obtain crystalline mass of trimethyl-
ammonium chloride.
Potassium, rubidium and cesium salts with dipicrylamine were prepared by
reacting the respective chloride salts with mexan. The precipitates of the
respective salts, were washed with saturated solutions of the salts to make the
precipitates free from mexan, if any, and dried at room temperature. The
chemical composition of the products are shown in table 1.
2-2. Apparatus
2-2a. Infrared spectra : The infrared spectra of the salts along with the parent
materials were recorded on a Perkin Elmer 421^model dual grating infrared
spectrophotometer in the frequency range from 4000 to 550 cm.- 1 using a scanning
speed of 17 min for the entire range. The pure and crystallised samples were
subjected to infrared recording in KBr matrix. About 2 mg of the samples
were mixed thoroughly with 250-^300 mg KBr powder. The mixture was placed
in a stainless steel vacuum die and pressed under hydraulic press for 10 min at
15 tons/sq. inch pressure. The pellets of dimension 13 mm thus formed were
subjected to infrared studies.
2-2b. Dipole moment: The dipole moments of dipicrylamine and potassium
salt with dipicrylamine were measured by the principle of heterodyne beat
method in a dipole meter (WTW, Germany) with the accuracy of order 10~ 4 .
Measurements were done at a frequency of 1 Me/sec, and at 30 C in benzene
solution of weight fraction not exceeding 0-05. Halverstadt-^Kumlers method
n + ~1 1Q/1TV
J _ | -
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Table 1. Chemical
Resulting salt
ium Salt with dipicrylamin
[ l4 NO)+(C ]2 H4N 7 O ia )-
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3. Results ana discussions
The infrared spectra of all the resulting salts are almost similar. The represen-
tative spectra of one salt with organic cation (viz., cholinium. salt) along with
its parent materials are shown in figure 1 and the spectra of the salts with
inorganic cations are shown in figure 2. The important absorption frequencies
which have been affected on salt formation are listed in table 2.
The main features of infrared spectra of all the resulting salts being similar, it
is expected that the nature of bonding, mode of salt formation and the mole-
cular configurations will be same for all the salts. So, the structural characteristics
with respect to cholinium salt with dipicrylamine have been discussed in detail.
140 tOO <JQO BOO (300 laOOSTOQ 2700 2TO03ICO 3300 i300JJOO
r*60JENC<r (CM' 1 )
Figure 1. Infrared spectra of (A) Dipicrylamine, (B) Cholinium chloride,
(C) Cholinium salt with dipicryJamine,
io too oo .100 two ooo itoo "ZTOO woo itoo
K 53 1
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Figure 3. Resonating structures of dipicrylamine.
It may be mentioned here that dipicrylamine has a number of resonating forms
of which two forms contribute predominantly to the resonance as shown in
figure 3. Comparing the occurrence of absorption bands in diphenylamine and
related compounds (Bellamy and Beecher 1952 ;. Bellamy 1958 ; Hadzi and
Skrblzak 1957) in the region 1655-1600 cm- 1 , it appears that the band at 1720-
1740cm- 1 region in the case of organic salts and at 1670- 1665 cm.- 1 in the case
of inorganic salts is too high in frequency to be assigned to aromatic ring vibration
only. However, in the type of compounds having structural element Ph -N=Ph
which of course, contains C=N link, a strong band at higher frequency around
1670cm- 1 has been observed by Marion et al (1951) and this is the characteristic
of anilino structures, in general, and indolenines in particular. Accordingly, it is
reasonable to assign the band at 1720-1740 cm- 1 region to C=N vibration
coupled with aromatic ring vibration (Ph-N=Ph). This band has consistently
appeared in other salts of dipicrylamine and this bonding is also in conformity
with the structure of guanidinium (Gupta and Datta 1975) and potassium (Kundu
and Ghosh 1980) salts with dipicrylamine determined by x-ray method. Further,
the N-H deformation mode (Rud'Co et al 1969) which has appeared at 1505 cm- 1
in the spectrum of dipicrylamine is absent in the spectra of the salts. These
suggest that choiinium chloride reacts with resonating form (II) of the dipicrylamine
in preference to that with other resonating form (I) (figure 3). Tnis is also
supported by the faster rate of reaction of choiinium chloride with dipicrylamine
to form the choiinium salt with dipicrylamine. Since the reaction rate of both
types of salt (organic and inorganic) is similar, it may be assumed that the band
due to structural element Ph N=Ph for both types of salts appear in the region
mentioned above.
The band at 1530cm- 1 as observed in the spectrum of dipicrylamine conforms
to the normal asymmetrical C-NO 3 stretching vibration (Bellamy 1958). In the
spectra of the salts the strong band in the region 1565-1555 cm- 1 may be presumed
to be the displaced frequency of the nitro groups which are twisted out of plane
of the aromatic rings due to the steric effect and also due to the presence of
choline ion in the vicinity of the nitro groups which take part in the salt formation
by van der Waals forces. The displacement of nitro groups to higher frequency
due to the presence of a strong electronegative group in the para position or of
a large group in the ortho position have been stated in the literature (Conduit
1959). The same phenomenon has been observed in the structure of guanidinium
salt with dipicrylamine. Similar is the case with the symmetrical vibration
has been shifted from 1345cm- 1 to 1338-1333 cm-i
Figure 4. Possible molecular configuration of cholinium salt with dipicrylamine.
way as in the case of organic salts. In the structure of rubidium hydrogen di-o-
nitrobenzoate and potassium hydrogen di-p-nitrobenzoate, Shrivastava and
Speakman (1961) have shown that nitro groups are twisted out of plane of the
benzene rings due to the linkage of metal ions with oxygen, of NO 2 group. Banerjoe
et al (1969) have also shown that in the case of alkali metal complex with
o-nitrobenzoic acid, N = O frequency is metal sensitive. The twisting of the
NO^ groups out of plane of the aromatic rings due to the linkage like K ONO
is further confirmed from the crystal structure of the potassium salt with dipicryl-
amine. Again, the symmetrical deformation mode (CH^N) at 14 10- 1400 cmr 1
of cholinium chloride is shifted in the salt, there is a decrease in the intensity of
the asymmetrical deformation mode around 1472cm.- 1 and enhancement of the
intensity of the absorption band in the region 1380-1375 cmr 1 due to the salt
formation. The enhancement of intensity is probably due to overlapping of
CH 3 -N symmetrical deformation mode in the region 14 10-- 14-00 cm- 1 which
has been lowered at 1380-1375 cmr 1 due to the salt formation, since the absorp-
tion band at 14 10- 1400 cmr 1 is absent in the spectra of the resulting salt. The
same phenomena have been observed in the case of acetyl cholinium and trimcthyl-
ammonium salts with dipicrylamine. The possible molecular configuration of
the cholinium salt with dipicrylamine is as shown in figure 4.
4. Conclusion
Dipicrylamine exists in resonating forms involving structural elements, ph-N-Ph
and Ph-N=Ph which are not equivalent. When dipicrylamine reacts with
organic and inorganic bases, the first resonating form is restricted and second
resonating form takes part in the reaction to form the respective salts. In the
structure of the salts of dipicrylamine, resonance occurs between the structures
involving structural element Ph-N=Ph and Ph=N-P>., which are equivalent.
The presence of dipole moment in dipicrylamine (3-68 x 10~ 18 esu) and in the
potassium salt (4-96 x 10~ 18 esu) also indicates that both structures are reso-
nating in character which is further confirmed from the structure of the potassium
salt (Kundu and G.iosh 1980). In the potassium salt the dipole moment is
lowered because of loss resonance energy compared to that of dipicrylamine.
This suggests that the potassium salt has equivalent resonating structures while
dipicrylamine has non-equivalent resonating structures (Pauling 1960).
It may therefore be concluded that dipicrylamine forms salts with both organic
and inorganic cations in similar fashion. The nature of bondings and linka|e
oi dipicrylamme ED ions with organic and inorganic cations are same in both types
of salts although the individual crystal structures may be different. Ghosh
et al (1968) have shown that when potassium nitrate reacts with the guanidinium
salt with dipicrylamine in water, potassium can replace guanidine ion forming the
potassium salt with dipicrylamine. The exchageability of organic ion by inorganic
ion also corroborates the similar molecular configurations of both types of salts
of dipicrylamine. This class of salts may be regarded as a loose molecular
complex type.
Acknowledgement
The authors are thankful to Dr B J Ansari and Sri A K Chakraborty for
recording the IR spectra. Our sincere thanks are also due to Dr S C Chakra-
borty of Burdwan University for helpful discussions.
References
Banerjee A B, Layton A J, Nyholm R S and Truter M R 1969 J. Chem. Soc. A 2536
Bellamy L J 1958 Infrared spectra of complex molecules (London : Methuen)
Bellamy L J and Bcecher L 1952 /. Chem, Soc. 1701
Conduit C P 1959 /. Chem. Soc. 3273
Ghosh Pasupati, Guha A, Chakraborty A K, Ghosh S K and Banerji K C 1969 Technol.
6 129
Ghosh Pasupati and Barker J M 1968 Technol. 5 229
Gupta M P and Datta B P 1975 Ada Cryst. B3l 1272
Hadzi D and Skrblzak J 1957 /. Chem. Soc. 843
Halverstadt I F and Kumbler W T 1942 /. Am. Chem. Soc. 63 2988
Kapoor J N, Kundu M L, Chakraborty A K and Ghosh S K 1972 Technol. 9 234
Kertes S 1956 Anal. Chem. Acta 15 73
Kundu M L and Ghosh S K 1980 Acta Cryst. B36 941
Marion L, Ramsey D A and Jones R N 1951 /. Am. Chem. Soc. 73 305
Pauling L 1960 The nature of chemical bonds, Ithaca, Cornell University Press
Rud'Co A P, Chernyuk I N, Ruzum Y S and Pilyugin GT 1969 Chem. Abstr. 70 11 0146 e
Shrivastava H N and Speakman J C 1961 /. Chem. Soc. 1151
Spectral and fluorimetric studies on the effect of surfactants on
thionine
S N GUHA, p N MOORTHY and K N RAO*
Chemistry Division, Bhabha Atomic Research Centre, Trombay, Bombay 400 085,
India
MS received 29 June 1981 ; revised 10 December 1981
Abstract. The effect of surfactants on the absorption and emission properties of
thionine (TH + ) have been studied in detail. Among the various surfactants investi-
gated sodium lauryl sulphate (SLS) has marked effect on these properties. Changes
in the absorption spectrum and the decrease in fluorescence intensity at [SLS] below
the critical micelle concentration (CMC) are attributed to the formation of a dye-
surfactant complex. At [SLS] above CMC, the restoration of dye spectrum with
increased extinction coefficient at the A max and a small but definite red shift of the
/Ijnax are interpreted as due to the incorporation of the dye into the SLS micelle.
The absorbance and spectral shift data suggest the thionine cation to be localized
near the micelle Stern layer in the case of SLS micelles but completely outside the
micelle in the aqueous environment in the case of CTABr. From the absorbance
and fluorescence data, the association constant for the formation of the TH+-SLS
complex in the premicellar region, and the binding constant for the incorporation
of the dye into the micelle in the micellar region have been computed. The values
of both these constants were found to increase markedly in the presence of
electrolytes.
Keywords. Thionine ; dyes ; surfactants ; micelles ; critical micelle concentra-
tion ; fluorescence.
1. Introduction
The behaviour of dyes in the presence of surfactant molecules is important for
understanding the thermal and light-induced reactions in biomembranes (Singhal
et al 1970; Hevesj et al 1970). Such reactions occur through the mediation of
excited and free radical species whose behaviour in a micellar medium can be
significantly different from that in a homogeneous aqueous medium. Our main
interest was the study of the photoredox reactions of the cationic dye thionine,
TH+. Our work on this dye in homogeneous aqueous solutions has been reported
earlier (Guha et al 1979). Before studying the photoredox chemistry in micellar
systems we considered it worthwhile to study its location and interactions in
such systems. We have therefore investigated the absorption and emission
characteristics of this dye in the presence of various surfactants. The thionine-
SLS system was studied in detail and the results are reported here.
* To whom all correspondence should be made
73
cae rea colour due to impuriues disappeared in ine organic pnase. inis was
followed by twice crystallization from an aqueous HC1 solution . SLS (Fluka, pract.
grade) and sodium dodecyl benzene sulfonate, SDDBS (Fluka, Tech. grade) were
purified by repeated washing with diethyl ether followed by drying over fused
calcium chloride in a vacuum desiccator. Triton X- 100 (Koch- Light, scintil-
lation grade), Brij-35 (Pierce Ohem. Co., specially purified grade) and Cetyl
pyridinium chloride, CPC (E. Merck) were used as such. Cetyl trimethyl ammo-
nium bromide, GTABr (Hopkin and Williams) was purified by dissolving the
substance in the minimum quantity of methan'ol, precipitating with diethyl ether
and drying in a vacuum desiccator over fused CaCl a . All other chemicals were
the purest commercially available. Solutions were prepared in triply distilled
water. Requisite volumes of stock solutions of thionine (10~ 3 mol dm- 3 ) and
the surfactant (10" 1 mol dor 3 ) were diluted together to give solutions containing
the two at the required concentrations. Absorption spectra were recorded on a
Hitachi Perkin Elmer IlV-visible spectrophotometer employing appropriate
blanks, and fluorescence measurements were carried out using an Aminco-.
Bowman spectrophotofluorometer.
3. Results
3.1. Sodium lanryl sulphate
The results of both spectral and fluorescence measurements for 10~ 5 mol dmr 3
thionine at different SLS concentrations are summarized in figure 1. Similar
behaviour was observed at other dye concentrations. It is seen that with increas-
ing SLS concentration the absorbance at 597 nm (^ max of thionine) as well as
fluorescence at 622 nm (for ^citation = 597 nm) both first decrease, reach a mini-
mum at 10~ 3 mol dnr 3 SLS, steeply rise between 1-3 x 10~ 3 mol dmr 3 SLS and
level off thereafter. In the low surfactant concentration region where the absor-
bance of the 597 nm dye band and fluorescence intensity decrease, two new bands
at 515 nm and 635 nm appear (figure 2) whose absorbances increase to a maximum
at 1 x 10~ 3 mol dmr 3 SLS and sharply decrease to zero in the 1-3 x 10~ 3 mol dnr 3
SLS region. Similar changes in absorption spectra and fluores.cer.ce of dyes in
presence of surfactant molecules of opposite charge have been reported in the past
(Corrin and Harkins 1947; Mukherjee and Mysels 1955; Malik and Cnand 1972;
Hwesi and Roz'.a 1971). In the case of the cationic dye pinacynol and anionic
surfactant SLS, a new absorption band has been observed (Mukherjee and Mysels
1955), which is r.ot present in the pure aqueous solution. For 3,3-diethylthia-
carbocyanine iodide, a cationic dye, significant change in the absorption spectrum
in presence of SLS has been reported (Sato et al 1980). The disappearance of
the dye absorption band and formation of new bands known as metachromism
generally ot served when the dye and surfactant bear mutually opposite
charges. In the thiohine-SLS (Hevesi and Rozsa 1971) and thionine- Rh6G- SLS
systems (Lohoczki and Hevesi L72) additional band at 465 nm has been assigned
Jo a dye-surfactant complex. However, this band was observed by us only when
0-8
fSLS] (mol Am' 3 )
igiirc J. Effect of SLS on the absorbancc and fluorescence of ihionine solutions.
0-6 -
0-8
Q-.3-
Q-t
0-0
550
A (nm)
700
Figure 2. Absorption spectra of thionine (10"~ 5 mol dm~ 3 ) : (a) in absence of SLS ;
and in the presence of (b) unpurified SLS, 10~ 3 mol dm" 3 , (c) purified SLS, 10 ~ 3
mol dm- 3 , (d) purified SLS 0-09 mol dnr 3 .
unpurjfied SLS was used (Curve I , figure 2). By mass spectral analysis the impu-
rity recovered by evapora tion of the ether extract obtained during the purification
of SLS was found to be dodecanol. Since the solubility of this impurity in water
bands at 515 and 635 nm. Decanol similarly added was found to show the same
effect.
The colour changes observed at low surfactant concentration have teen
variously attributed in the past to formation of ion pairs (Colichman 1950), com-
plexes (Malik and Ciiand 1972), insoluble complex salts (Klevens 1947) and dye
aggregates (Corrin and Harkins 1947). The fact that such a behaviour is
characteristic of oppositely charged dye and surfactant molecules agrees with the
first three possibilities. Mukherjee and Mysels (1955) have in fact ch?racterised
and isolated a 1 : 1 dye-detergent complex salt in the case of methylene blue. In
the pinacyanol-SLS system, a highly insoluble salt was found to form a stable
suspension in the presence of somewhat more than stoichioinetric amounts of the
detergent (Mukherjee and Mysels 1955). In the present thionine-SLS system, a
precipitate was observed only at 3 x 10~ 5 mol dm- 3 TH+ and 10~ 4 mol dm.- 3 SLS.
At other compositions the solutions were optically clear and precipitation could
not be induced by any means. In the transition region 10- 5 -^10- 4 mol dnr 3 SLS
where the dye band intensity diminishes abruptly and new bands at 515 and
635 nm appear, thionine spectra were found to exhibit isobestic points at 530 and
620 nm (figure 3). Such a spectral behaviour and also the change in fluorescence
intensity can be interpreted as due to an equilibrium involving association of
the dye cation (D+) and the lauryl sulfate auions (S~) :
D+ + S- ^ DS. (1)
At [TH+] = 5 x 10- 6 mol dm- 3 , [SLS] = 3-5 x 1Q- 3 mol dor 3 , Balint et al
(1977) observed bands at 465, 515 and 635 nm in addition to the monomer
band at 597 and the dimer band at 656 nm. They assigned the 465 nm band to
the dye-surfactant complex and the 635 nm band to higher dye aggregates. The
515 nm band was not discussed. As mentioned before, the 465 nm band is due
to interaction with the dodecanol impurity in SLS. Formation of dye aggregates
in other dye-detergent systems have also been reported (Matagaand Koizumi
1954; Sato et al 1980). To explain the behaviour of anthraquinoid acid dyes
in the presence of surfactant molecules Datyner (1961) assumed that the dye
surfactant complex may aggregate to form larger particles. Except at one compo-
sition as noted above there was no precipitate formation in the thionine-iSLS
system and hence such aggregation of the complex to larger particles does not
seem to be favoured in this system.
From figure 1 it may be seen that the disappearance of the thionine band at
597 nm and formation of the new bands at 515 and 635 nm are accompanied by
decrease in thionine fluorescence. In fact at 10~ 3 mol dnr 3 SLS when the thio-
nine band has virtually disappeared and the atsorbance at 515 and 635 nm are
maximum no fluorescence is observed with /U, 515, 597 and 635 nm. Hence, it is
to be concluded that the TH+^SLS complex is non-fluorescent. This must be due
to rapid degradation of excitation energy via internal conversion facilitated by
the long hydrocarbon chain in the SLS moity in the complex.
Both absorbance and fluorescence data can be used to compute the association
constant (^between the dye cation and surfactant anion involved in equilibrium 1.
400 420 440 460 480 SOO 620 540 560
A tnm)
S80 600 620 640 660 680 TOO
Figure 3. Absorption spectra of thioninc (3 x 10- 5 mol dm~ 3 ) at different concen-
trations of purified SLS (0)0, (l)2x 10~ 5 , (2) 3 X 10~ 5 ,(3)4x 10~ G ,(4)5 X 10~ B ,
(5) 6 X 10- 5 and (6) 7 X 10~ B mol dnT 3 .
.is if the extinction coefficients of the free (or aqueous) and associated (or
nplexed) thioninc species at a given wavelength are respectively e a and e c , the
asured absorbance is given by:
A = * w (l -/) [DW + ./.PU (2)
vhere f e , the fraction of the dye present as complex is given by
fe
==l l
the cell pathlength and [D] t and [S] t are the total thionine and surfactant
.centrations. Substituting the value of f e in (2), the measured absorbaoce
i be shown to be related to the surfactant concentration according to:
A ~\-l
I / \ 1 i // \ -rn tr fit f f JY1 \\ 1 (A\
lilarly the variation in the observed fluorescence intensity 7 obs should follow
: eouation :
fla "" vs {[ "" or vs ""
should be linear and K a can be calculated from the intercept and slope of such
plots. As/o is not known, first an approximate plot can be constructed using {S] t
instead of [S] t f a [D] t and the approximate K a so evaluated then used to compute
f e at each surfactant concentration from (3). A more accurate plot is now con-
structed using these f e values and this successive approximation procedure is
repeated until the K a and/ e values become invariant. From the intercept of the
plot corresponding to (4), e can be calculated as e fl(Z is known. The K a values
so obtained from the measured absorbances at 597 nm and the fluorescence
intensities at 622 nna agree with each other within 10%, the average value being
2-11 x 10 l dm 1 mol- 1 . K a values were similarly computed from the absorbances
at 515 nm (2-33 x 10 1 dm 3 mor 1 ) and 635 nm (1-8 x 10 4 dm 3 mol" 1 ).
The extinction coefficient of the complex computed from the intercept is 0-41 x
10 l dm 3 mol-^cnr 1 which is about a factor of three smaller than that correspond-
ing to the measured absorbance at 10~ 3 mol dmr 3 SLS wherein the absorbance
in the 597 nm band is at a minimum. The discrepancy is attributable to an
appreciable contribution from the micelle bound thionine which, as will be seen
later has an extinction coefficient at 597 nm even higher than the aqueous thionine
(monomeric) species.
The sharp changes in the absorbance and fluorescence occurring in the region
of SLS concentration 1-3 x 10- 3 rnol dmr 3 are related to the formation of sur-
factant micelles in which the dye is incorporated. Such sharp changes have in
the past been made use of for the determination of the CMC of surfactants
(Gorrin and Harkins 1947; Mukherjee and Mysels 1955). The CMC of SLS
evaluated from the inflexion points of curves in figure 4 are summarized in table 1.
As observed by Mukherjee and Mysels (1955) the CMC values so obtained are
lower than the ones obtained by light scattering, conductivity and viscosity
measurements. Also they increase and approach the latter with increasing
thionine concentration.
As mentioned before, thionine is present almost exclusively as the complex
at 10- 3 mol dmr 3 SLS and hence at this surfactant concentration the absorbance
at 597 nm and fluorescence are at a minimum and the absorbances at 515 nm and
63 5 nm are at a maximum. Addition of electrolytes such as Na 3 S0 4 , H a SO 4 and
NaCl was found to restore fluorescence and absorbance at 597 nm and bleach the
515 and 635 run bands. Measurements made at a xed concentration of Na a SO 4
and varying concentrations of SLS revealed that the CMC of the latter is lowered
in presence of the electrolyte (figure 4). As a result KHmol dmr* of SLS is
well above the CMC in presence of 0-02 mol dm- 3 Na a S0 4 and hence the
changes observed on addition of electrolytes to a thionine solution in 10-* mol
dmr 3 SLS can be attributed to micellization and incorporation of the dye in the
micelle (figure 5). This lowering of CMC on addition of electrolyte is in
agreement with previous reports in the literature (Corrin and Harkins 1947-
-06
- 0-Qii
10 10
[SLS] (mol dm")
Figure 4. Dependence of absorbance (at ^ mas = 597 nra) on SLS concentration at
different thionine concentrations and the effect of added Na 2 SO 4 .
Table 1. Effect of thionine concentration on the CMC of SLS micelles.
[TH+] x 10* CMC of SLS x 10*
(mol dm~ a ) (mol dm- 3 )
1-0
5-0
10-0
30-0
1-7
2'3
2-6
3-3
Muto et al 1973). The log-log relationship between CMC and electrolyte
concentration generally observed in the case of ionic surfactants (Corrin and
Harkins 1947;. Schick 1964; Birdi et al 1980) holds good.
The restoration of the 597 nm band and the characteristic thionine fluorescence
at SLS concentrations well above the CMC would indicate that the dye surfactant
comt>lex is unstable in the micellar environment. At high surfactant concentra-
Figure 5. Effect of Na 2 SO 4 on the absorption spectrum of thionine (10~ 5 mol dm~ 3 )
in presence of SLS (a) neat aqueous solution, (b) 10~ 3 mol dm~ 3 SLS, (c) 10~ 3
niol dm~ 3 SLS and 0-02 mol dnr 3 Na 2 SO 4 and (d) 0-09 mol dor 3 SLS.
(figure 2). Similar behaviour in other dye-detergent systems had been attributed
(Kapoor and Mishra 1976) to the disaggregation of dye aggregates in the surfac-
tant micelles to give the fluorescent monomeric species. In the thionine-SLS
system, however, comparison of the absorption spectrum in the abser.ce of surfac-
tant and at high [SLS] revealed the presence of the dimer band in both cases
although somewhat reduced in intensity in the latter (figure 2). If dye disaggre-
gation was solely responsible for these changes then fluorescence yield after correc-
tion for reabsorption in the system should be constant, but was found to increase
with increasing [SLS] above the CMC. The ^ mas of thionine monomer band also
exhibited a small but definite red shift. A more plausible explanation of the
change in absorption and fluorescence in the micellar system would be that the
dye in the micellar environment has a different extinction coefficient and radiative
life time as compared to the pure aqueous environment. Since the dye is present
as the complex in the premicellar region, the equilibrium
K
\"/
would be established in the micellar region, and the following equations can be
derived :.
j m is me ucicuun ui me u,yo iu tnc nu^ciicir iuim u^ M, ana L^Jt JS
total micelle concentration given by
[M\ t ~([S], - CMC)/aggregation no. . (9)
From the above equation it is possible, as before, to compute the dye-micelle
"binding constant K* by successive approximation. The values computed from
absorbance and fluorescence data agree with each other. The average value was
found to be 1-37 x 10 dm 3 mol" 1 . From the intercept of the plot corresponding
to (7) the extinction coefficient for the micelle-bound thionine at 597 nm was
found to be 6-9 x 10 4 dm 3 mol- 1 cnr 1 which closely agrees with the value corres-
ponding to the measured absorbance at the highest SLS concentration.
3.2. Effect of electrolyte addition
Both the dye-surfactant association constant K tt and the dye-micelle binding con-
stant K b were found to appreciably increase on addition of electrolytes sudh as
Na 2 SO 4 . The values of these constants for a few typical concentrations of Na a SO 4
are summarized in table 2. In the region of electrolyte concentration employed,
the micelle structure is not altered. In SLS, for example, the transition to rod
like micelle occurs at electrolyte concentration above 0-45 mol drrr 3 (Ikeda et al
1981). However, at lower electrolyte concentrations there is a small and gradual
increase in. micelle molecular weight and hence aggregation, number. In the
computation of K t we have ignored this. The effect of this would be to give a
value of K* lower than the true value. Therefore the observed increase in binding
constant with increasing concentration of electrolyte is inferred to be genuine and
not an. artefact of neglecting the increase in aggregation number in presence of
electrolytes.
3.3. Medium polarity effects
It has been mentioned earlier that the absortance of thionine solutions at 597 nm
which reaches a minimum at ~ 10~ 8 mol dm~ 3 steeply rises near the CMC and
levels off to a plateau of small positive slope beyond 3 x 10~ s mol dm~ 8 SLS.
There is also an appreciable red shift of the ;i m , e.g., in 0-09 mol drrr 3 SLS the
maximum is shifted to 602-5 nm. At the respective maxima the extinction coeffi-
cient in 0-09 mol dnr 3 SLS micellar medium is about 20% higher as compared
to the homogeneous aqueous medium. The red shift and increase in extinction
coefficient both reflect a decrease in the polarity or the dielectric constant of
Table 2. Effect of electrolyte addition on K and K t values
NaCl K a K>
(mol dm- 3 ) (dm 8 moH) (dm 3 mol' 1 )
NCI 2'11 X10* 1-37 xlO 8
(0-02) 4-8 Xl0* 7-5 X 10 6
the medium around the probe molecule. Thus, for example, in water-alcohol
mixtures, the extinction coefficient increases linearly with decreasing dielectric
constant (figure 6). From this plot the dielectric constant experienced by thionine
in the SLS micellar system can be read off as ~ 56 against the ot served Ae w of 1- 1
X 10* dm 3 moHcnr 1 . Similarly, as shown in figure 6 the A^ values also follow
linear variations with the solvent polarity parameter, E T (see Reichard 1965 for
E T values). Deviations from the linear correlation may be noted in the case
of DMSO and dioxane. Ignoring this deviation, the polarity of the environment
of thionine in the 0-09 mol dnr 3 SLS micellar system can be inferred to corres-
pond to an E T value of 57-5, i.e., the polarity is between that of water and
methanol. The red shift of the thionine absorption maximum also shows a
correlation with the solvent polarity parameter, but as the shifts are rather small
their accurate measurement is rather difficult and hence no attempt has been made
to evaluate the solvent polarity parameter from this correlation.
o-o
10 20 30 40 50 6 70 80
DIELECTRIC CONSTANT/* SOLVENT POLARITY -
3.4. Other surfactants
Among the other surfactants investigated, the behaviour of thionine in the anionic
surfactant SDDBS closely paralleled the behaviour -in SLS : a decrease in the
absorbance at 597 nm and the fluorescence intensity with increasing [SDDBS]
up to ~ 5 x 10-* mol dm~ 3 followed by a sharp increase in the region 1 2 8 x
IQ- 3 mol dmr 3 culminating in a plateau beyond ~ 3 x 10~ 3 mol dnr 3 . It may
be noted that the CMC of SDDBS is 1 -2 x 1CT 3 mol dnr 3 . The initial decrease
was accompanied by the appearance of new bands at 550 and 635nm. As in
the case of SLS the new bands disappeared at [SDDBS] > CMC. These obser-
vations are subject to the same interpretation as in the case of SLS. Association
and binding constants were not calculated.
In sharp contrast to the case of SLS and SDDBS the behaviour in presence
of the neutral surfactants Triton X-.100 and Brij-35 and the cationic surfactants
CTABr and CPC was very different. There was no decrease in the absorbance
at 597 nm nor did new bands appear at surfactant concentrations below the CMC.
The ion^pair type of complex is obviously not possible in the case of the catio-
nic surfactants as thionine and surfactant head groups bear like charges. Although
the neutral surfactant Triton X-100 is known to form charge transfer complex
with strong electron acceptors such as TCNQ (Muto el al 1970) thionine seems
to be too poor an electron acceptor to form such a complex with the rather
poor electron donors, v/z., Triton X-^100 and Brij-35. Beyond the CMC there
was a small but definite red shift of the thionine absorj. tion maximum and also
an increase in e*. These data together with the inferred values for the dielectric
constants and solvent polarity parameters as deduced from the linear plots of
figure 6 are given in table 3,
4. Discussion
Several publications have appeared on the study of micelles ever since Me. Bain
suggested aggregation of surfactant molecules atove a critical concentration and
Hartley's (1935) model of micelle as a tiny oil droplet in an ionic coat of
hydrated ions. However many aspects of micelles pertaining to micelle shape,
Table 3. Solvent polarity and dielectric constant of different surfactant micelle*
as probed by TH.
Surfactant
(well above CMC)
A*
SP
DC
SLS (+0-8 M NaCl)
1-5 xlO*
56-0
47-0
SLS
M XlO*
57-5
56*0
SDDBS
1-15 XlO 4
58-0
56-0
Tritonx -100
1-0 XlO 4 ,,,
58-5
57-0
Brij-35
0-6 XlO 4
60/5
67-0
CPC
0-4 xlO*
6J-5
71-5
CTABr
0-2'* XlO 4 t
62-5
76-0
water penetration, surface roughness, adsorption, sites, interior viscosity and chain
conformation are not yet fully unctetood (Msngsr 1979). The main point of
controversy in the recent past, with which of course the other questions above
are to some extent tied uj. , is water penetration into micelles (See Wennerstrom
and Lindman 1979 and Menger and Bonicamp 1981 for two divergent view-
points). Whereas one extreme, the " reef " model views the micelle interior as
completely dry with all the surfactant methylene groups lying entirely within the
ionic coat, the opposite extreme, the "Fjord" model allows water percolation
nearly to the micelle centre. Experimental results on micelles are often interpreted
as evidence to suj port either of these extreme view points or some median view-
point. Such evidences are invariably deduced on the basis of certain precon-
ceived notions, nota le among them being (i) the micelle is a closed entity
(Franses et al 1981) with an impervious (but imaginary) boundary separating the
lipopoic moieties in the interior from the hydrophilic head groups on the exterior
and (ii) a solubilizate used as a probe molecule is localized in the lipopoic interior
and does not jerturb the micelle structure. As pointed out by both Menger
(1979) and Lindman and Wennerstrom (1981) such questionable conclusions are
due to the methodology employed in experiments designed to prole the micelle
interior. In this one compares the spectroscopic properties of \ robe molecules
in the micelle and different solvents or solvent mixtures and presupposes a corre-
lation between such properties and the polarity of the environment indicated in
terms of the dielectric constant or some empirical solvent polarity parameter-
Although such correlations do exist (figure 6) what is in doubt is whether any
solvent or solvent mixture c?n te assumed to simulate a micellar environment.
Also, as is evident from figure 6, beyond a certain pote rity, the measured spectro-
scopic parameter may exhibit an opposite trend.
In connection with this controversy regarding water penetration into micelles
may be mentioned a recent neutron scattering study (Hayter and Penfold'1981)
the results of which reconcile the two extreme viewpoints by supporting the idea
of ' a little water penetration ' into the paraffin core due to entrainment of water
by the bound counterions.
Although the question regarding the nature of micelle interior is still unresolved,
there is some agreement on the location of solubilizates (Menger 1979; Lindman
and Wennerstrom 1981,). The majority of solubilizates, including the water
insoluble compounds such as benzophenonc, bromobenzene, pyrine, etc., prefer
the highly aqueous micelle surface to the lipopoic interior. Such being the case,
the location of an ionic highly water soluble compound such as thionine used
in the present study should unquestionablybethe highly aqueous micellar surface.
In experiments employing such probe molecules, the question one has to ask is
not how much water-like the micelle interior is, but how much water-unlike the
surface region is. From the data summarized in table 3 it is evident that in all
cases the environment around thionine is highly polar, the polarity being some-
where between that of water and methanol. The highest polarity is observed in
GTABr micelles and the least in SLS micelles. Due to electrostatic repulsion of
likecharged ions of TH+ and the headgroivpin CTABr, the location of TH+ in
the surface region m these micelles as compared to water is accountable on the
basis of the surface roughness of micelles arising from their dynamic nature, i.e.,
monoraeric units constantly enter into and exit from micelles. On an average,
there is considerable protrusion of the methylene groups into the headgrcaip
region (Aniasson 1978). This at once rules out the picture of a micelle as. a
closed impervious compartment and would obviously allow for penetration of
water into the micelle. The water so penetrated cannot be expected to behave
entirely like bulk water, but somewhat like a less extensively hydrogen bonded
structure filling the crevices between the paraffin chains.
It is known (Ikeda et al 1981) that in the presence of high concentration of
electrolytes larger rodlike aggregates of surfactants are formed. In such larger
aggregates one can expect more methylene protrusion per micelle. This is reflec-
ted in the lower value for the polarity around thionine in SLS micellar solutions
'containing 0-8 rnol dnr 3 NaCl (table 3). The gradual increase in the absorbance
of thionine solutions beyond the CMC (figure 1) is also attributable to the same,
as it is known that larger aggregates are favoured with increasing surfactant
concentration.
References
Aniasson BAG 1978 /. Phys. Chem. 82 2805
Balint E, Hevisi J and Vass I 1977 Ada Phys. Chem. 23 123
Birdi K S, Dalsagcr S U and Backlund S 1980 /. Chem. Soc, Faraday Trans. I 176 2035
Colickman E I 1950 /. Am. Chem. Soc. 72 1384
Corrin M L and Harkins W D 1947 /. Am. Chem. Soc. 69 683
Datyner A 1961 /. Soc. Dyers Colour. 77 304
Franses E I, Davis H T, Miller W G and Scrivcn L E 1981 J. Phys. Chem. 84 2413
Guha S N, Moorthy P N and Rao K N 1979 Mol. Photochem. 9 183
Hartley G S 1935 Traits. Faraday Soc. 31 31
Hayter J B and Pinfold J 1981 /. Chem. Soc. Faraday Trans. I 77 1851
Hcvesi J, Lohoczki E and Balint E 1970 Zh. Prikl Spektrosk. 13 440
Hevesi J and Rozsa Zs 1971 Act a Phys. Chem. 17 127
Ikeda S, Hayashi S and Imae T 1981 /. Phys. Chem. 85 106
Kapoor R C and Mishra V N 1976 J. Indian Chem. Soc. 53 965
Klevens H B 1947 /. Phys. Chem. 51 1143
Liudman B and Wennerstrom H 1981 in Topics in Current Chemistry Series, Vol. 87. Micelles
(New York : Springer-Verlag)
Lohoczki E and Hcvesi J 1972 Dokl. Akad. Nauk SSSR 206 1158
Malik W U and Chand P 1971 /. Electroanal. Chem. 40 385
Mataga N and Koizumi M 1954 Bull. Chem. Soc. Jpn. 27 197
Menger F M 1979 Ace. Chem. Res. 12 111
Monger F M and Bonicamp J M 1981 /. Am. Chem. Soc. 103 2140
Mukherjee P and Mysels K J 1955 /. Am. Chem. Soc. 77 2937
Muto S, Dcguchi K, Kobayasht E, Kancko E and Meguro K 1970 /. Colloid Interface Set.
33 475
Muto S, Aono Y and Meguro K 1973 Bull. Chem. Soc. Jpn. 46 2872
Reichardt C 1965 Angew. Chem., Int. Ed. Engl. 4 29
Sato H, Kawasaki M, Kessatavik K, Kusumoto Y, Nakashima N and Yoshihara K 1980 Chem
Lett. p. 1529
Schick M J 1964 /. Phys. Chem. 68 3585
Periodic precipitation of cobalt (II) oxinate in agar gel :
Effect of parasitic electrolytes on flocculation
N KAKNIAH, S AMBROSE, F D GHANAM and
P RAMASAMY*
A.C. College of Technology,
Perarignar Anna University of Technology, Madras 600 025, India
MS received 13 March 1981; revised 5 November 1981
Abstract. The experimental conditions for the periodic precipitation of cobalt(II)
oxinate in agar agar gel have been extensively studied- When different cobalt salts
are taken over the Set gel impregnated with oxine, different parasitic electrolytes
are formed along with cobalt(II) oxinate. These parasitic electrolytes affect the
solubility and hence the periodic precipitation of cobalt(ll) oxkate. The effect of
the parasitic electrolytes on the flocculation value has been studied on the basis of
Shinohara's revised flocculation theory. The spacing law of Jablczyiski and the
time law have been experimentally verified.
Keywords. Periodic precipitation; oobait(II) oxinate; parasitic electrolyte;
flocoulation value; agar gel.
1. Introduction
Periodic precipitation of many sparingly soluble substances has been reported in
the literature (Stern 1967). Stern (1954) has reviewed the different theories and
factors influencing the periodic precipitation. Shouji Shinohara (1970) has
revised the coagulation theory of Dhar and Chatterjee (1922) to explain the perio-
dic precipitation in a quantitative manner. We have recently reported the periodic
precipitation of cobalt(II) oxinate in agar agar gel (Kanniah et al 1981). In this
paper the influence of various parasitic electrolytes on the periodic precipitation
of cobalt(II) oxinate has been discussed in detail.
2. Theory
As the outer electrolyte diffuses into the gel impregnated with the inner electrolyte,
the inner one itself diffuses in the opposite direction, both obeying Pick's law of
diffusion given as
^-zx21 l m
dt - dx z ' (i)
* TO wliom all correspondence should be made,
87
at
fa*
(2)
where C\ and C a are the concentrations of the outer and inner electrolytes and D
and D z are the diffusion coefficients of the outer and inner electrolytes respectively.
The reaction occurs only on the boundary which is a plane perpendicular to the
axis of the tube. The outer electrolyte is on one side of the boundary and the
inner on the other side. The sparingly soluble reaction product remains as a sol
along with the parasitic substance at the boundary, known as 'sol front'. The
advancing speed of the sol front is given by
where D to is the diffusion coefficient of the outer electrolyte at infinite dilution
calculated using TSJernst equation.
The value of k which controls the behaviour of the sol front is determined
by Adair using the boundary condition as
/) 2 ^-? /boundary = 0.
This can be conveniently written as
SQ Alfc i /2ft == '
where
(4)
(5)
boundary,
^ = ii 2 / boundary.
ox I
If we assume that the diffusion coefficients are constant, then
Di = Ao and D z = X> 20 .
f _ G(k)-G(k,)
G (k,) '
=
where
(6)
= erf
snce
V 7 *
Substituting in (5)
exp (~fc 2 /2) _ exp ( - fc 2 ^ 2 /2)
sq G(k) ~ l/2-G(ks)
exp(-FW2]
erf(/c/V2) ~ 1 -
using error function
_ exp(-Jc 2 * 2 /2) .
~ erfc (Jb/V2) ' ^ }
Equation (7) is known as Adair's equation. The front constant k is calculated
for different valu.es of q using (7).
The sol front advances by forming sol of the sparingly soluble reaction pro-
duct. The substance is in a state of supersaturated solution, before the formation
of sol. The concentration of this supersaturated solution just before the forma-
tion of the sol, known as the reduced concentration (C^ is given by
From the above equation it can be noted that the reduced concentration
is constant everywhere, though the speed of the sol front slows down gradually
as it advances.
As the diffusion proceeds the concentration of the outer electrolyte varies with
the distance from the largest value (Cj ) at the gel boundary to zero .at the sol
front. The flocculation of the sol is caused by the ions of the outer electrolyte
and the parasitic electrolyte formed during the sol formation. As the ionic con-
centration of the outer electrolyte reaches a value /which is characteristic of the
sol, flocculation occurs. This characteristic value F can be defined from (6) as
r= c
where
k = k/P and p = x n ^/x n . (11 '
2.1. Flocculation value
The sol of the sparingly soluble reaction product is formed with the concomitant
formation of a soluble parasitic electrolyte which produces ions on ionization.
The influence of the parasitic ions depends on the valency of the outer electrolyte
and the parasitic electrolyte. That is whether the outer electrolyte and the para-
sitic electrolyte are monovalent, divalent or trivalent in action. The flocculation
solution. To this solution, 25% ammonia solution was added drop by drof until
a faint but permanent turbidity was obtained. A few drops of 2N acetic acid
were addsd to produce a clear solution. The pH of this solution was adjusted
to 4-25. This oxine solution was mixed with agar agar gel solution and the gel
solution with oxine was made upto 300 ml with hot double distilled water. This
gave 0-03 M oxine in 1% agar agar. Similarly 1% agar agar solutions containing
oxine of concentrations 0-06 M and 0-05 M were prepared. 50 ml of these clear
solutions were poured into a corning tube of 20 mm diameter and allowed to
set. After 3 hr 10 ml of cotalt(II) sulphate solutions of 1-031, 0-859, 0-687,
0-515 and 0-344M concentrations were carefully taken over the set gel.
The experiments were carried out with cobalt(II) chloride and cobalt(II) bromide
as the outer electrolytes. In all these cases the concentrations of the oi'ter electro-
lyte were 1-031, 0-859, 0-687, 0-515 and 0-344 and the concentrations of the
inner electrolyte were 0-06 M, 0-05 M and 0-03 M. The experiments were carried
out at room temperature (30 0). Sharp brown coloured disc like precipitate
rings demarcated by clear void spaces were oltained (figure 1). The distance
measurements were made with ? cathetometer. The IBM 1130 computer was
used for the calculation of floeculation values using Shinohara's revised coagulation
theory.
4. Results
4.1. Verification of time law
According to Shinohara the movement of the sol front can be expressed as
~ K.
IS is known as the velocity constant. For a given pair of concentrations of the
inner and outer electrolytes, K is found to be a constant, thus verifying the time
law. The velocity constant increases with the increase in the concentration of
the outer electrolyte. Figure 2 shows the dependence of K on the concen-
trations of OoSO 4 , OoOl a and OoBr a , for the same inner electrolyte concentration.
4.2. Verification of the spacing law
According to Table zynski's (1923) spacing law
x* - SoP"*
where p is known as the spacing coefficient, x m ^ and x are the positions of the
(n + l)th and rtth rings from the gel boundary and XQ is a constant. Figure 3
shows a plot of the numerical order of the ring (n) against log * for the same
Periodic precipitation of cobalt (If) oxlnate
Figure 1. Periodic precipitation of cobalt(ll) oxinato in a gar a gar gel. The con-
ccntr?tion of oxine is 0-05 mole/litre. The concentration of cobalt Sulphate is
1-031 mole/litre and 0-515 mole/litre for tubes 1 and 2. respectively.
0-6
0-3
o COClj
* COBr.2
cosa
I .. 2 .3 4
# ( mole" 1 lit )
/C 10
Figure 2. The dependence of ^ on the concentration, of the outer electro iytes.
1-0 -
0-9
0-8
5 10 15
n
Figure 3. Verification of Jablczjrnsfci's spacing law.
concentration of the different outer electrolytes. The linear variation verifies the
spacing law of Jablczynski.
4.3. The dependence of spacing coefficient on the concentrations of the outer
electrolytes
0.12-
Figure 4. The dependence of spacing coefficient on the outer electrolyte
concentration.
observation is true for all the outer electrolytes taken over the set gel. The
dependence of the spacing coefficient on the outer electrolyte concentration has
been represented in figure 4.
4.4. Calculation of flocculation value
The reduced concentration of the sol is calculated using the equation,
required to initiate
The characteristic concentration of the outer electrolyte,
flocculation is calculated using the equation
jo
Then the flocculation value of the sol
F Cao * F.
The flocculation values are computed for various concentrations of the different
outer electrolytes. Table 1 shows the computed values of C^, -T, Fand G(k),
when OoSO 4 is used as the outer electrolyte. Tables 2 and 3 show the corres-
ponding values for cobalt chloride and cobalt bromide respectively. The com-
puted flocculation value is the highest when cobalt sulphate is used as the outer
electrolyte. The concentration of the inner electrolyte being constant, the floccu-
lation values of cobalt(II) chloride and cobalt(II) bromide are almost equal for a
particular concentration of the outer electrolyte. As the concentration of the
outer electrolyte decreases, r value decreases for all the three outer electrolytes,
o
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x 2"
O
X
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a particular concentration 01 me inner eiectrojyie.
5. Discussion
5.1. Dependence of the velocity constant
The diffusion coefficients of CoBr a (1-2904 x iO~ 5 cm 2 sec- 1 ) and CoCl a (1-2763 x
10- 5 cm 2 sec- 1 ) are almost equal. The velocity of the sol (v) front is the same for
both the outer electrolytes. Jfcr.ce the values of the velocity constant do r.ot show
any appreciable difference. As the diffusion coefficient of CoSO 4 is very low
(0-8674 x 10" 5 cm 2 sec" 1 ) the velocity of the sol front and hence the velocity
constant are less than those for CoBr a ar.d CoCl 2 .
5.2. Effect of concentration of the outer electrolyte on spacing coefficient
The relation, between the spacing coefficient (p) and the concentration of outer
electrolyte (C ao ) is given by Matalon and Packter's (1955) equation
(p _ i) = A 4 B/Cio,
where A and B are constants. The plot of (p 1) versus 1/C 10 gives a straight
line (figure 4). The slope of the line gives B which is directly dependent on
supersaturation . Since the same straight line represents both CoCl a and CoBr 2
the supersaturation and the solubility of cobalt oxinate are the same when HCl
and HBr are formed as the parasitic "electrolytes. The slope of this line is greater
than that of the straight line for CoSO 4 . This imans that the degree of super-
saturation is less when CoSO 4 is the outer electrolyte. In other words the presence
of H 2 SO 4 as the parasitic electrolyte increases the solubility of cobalt(II) oxinate.
5.3. Flocculation value
When the outer electrolyte is allowed to diffuse into the gel impregnated with
oxinc, cobalt oxinate is formed as a sol with the concomitant formation of a
soluble parasitic electrolyte.
CoSO d -+ 2C H 7 ON -* Co (C H ON) a 4 H a SO 4 ,
Co(N0 8 ) a + 2C 9 H 7 ON-* Co (C H 6 ON) 2 + 2HN0 3 ,
CoCl 2 4 2C 9 H 7 ON -> Co (C 9 H 6 ON) 2 4- 2HC1,
CoBr 2 + 2C 9 H 7 ON - Co (C 9 H 6 ON) 2 + 2HBr.
These parasitic electrolytes have, different characteristics and produce different
ions on ionization. They will have different effect on the flocculation of the posi-
tively charged cobalt oxinate sol formed due to the adsorption of excess of cobalt
ion. The flocculation of the sol is caused by the counter ions supplied by the outer
electrolyte and the parasitic electrolyte. So it may appear that the divalent anion
sulph.at' 5 should be more effective in flocculation than the monovalent anions,
chloride pp.'d bromide. But the analysis of the tables 1, 2 and 3 shows that
the flocculation value of sulphate is greater than that of others. The concentration
of the supersaturated solution (C^) of cobalt oxinate formed before the sol forma-
lion is greater lor (josu 4 . in other words only alter tne auainmem 01
C 3o , the sol formation is possible. More outer electrolyte has to diffuse to satisfy
this, condition. It is very clear from the foregoing arguments, that the presence
of parasitic electrolyte H 2 SO 4 increases the solubility of cobalt oxinate. It takes
longer to exceed the solubility product. Hence the flocculation value (F) and
the spacing coefficient Q;) are greater than those for other outer electrolytes. This
result is in accordance with Wagner's (1950) theoretical prediction. Similar results
have been reported for calcite (Gnanam et al 1980) and silver chromate (Verma
and Ghosh 1953).
We have recently reported the periodic precipitation of cobalt oxinate formed
by diffusing cobalt(Il) nitrate (Kanniah et al 1981) into agar gel containing oxine.
The computed values of the concentration of the supersaturated solution (Cao)
are identical in the case of cobalt bromide, chloride and nitrate. Hence the
parasitic electrolytes HBr, HCl and HKO 3 aftxct the solubility of cobalt oxinate
to the same extent. Until the formation of the sol the conditions are identical
in the case of cobalt nitrate, chloride and bromide. However a close analysis of
the flocculation values will reveal the fact that the flocculation values are in the
following order :
NO a - < Cl- < Br-.
Hencc the flocculating capacity is in the reverse order.
6. Conclusion
The periodic precipitation of cobalt(II) oxinate obtained by diffusing different
electrolytes into the agar agar gel impregnated with oxine reveals the influence
of the parasitic electrolyte on the solubility ofcobolt(II) oxinate. The flocculat-
ing capacities of different aiuons arc compared.
Acknowledgement
This work is supported by CSIR and UGC, New Delhi. One of the authors (NK)
thanks UGC for the award of a fellowship.
References
Dhat H R and Chatter jeo A C 1922 Kolloid. Z. 31 15
Gnanam F D, Krishnan S, RamaSamy P and Laddha G S 1980 J. Colloid Interface Set. 73 193
JablczynSki K 1923 Bull. Soc. Chim. France 33 1952
Kanniah N, Gnanam F D and RamaSamy P 1981 J. Colloid Interface Sci. 80 377
Mataloii R and Packtcr A 1955 J. Colloid Sci. 10 46
Shouji Shinohara 1970 /. Phys. Soc. Jpn. Z9 1 073
Stern K H 1954 Ckem. Rev. 53 79
Stern K H 1967 Bibliography of Liesegang rings (USA : National Standards)
Photoelectron spectroscopic studies of the adsorption of organic
molecules with lone pair orbitals on transition metal surfaces!
S YASHONATH, P K BASU, A SRMVASAN,
M S HEGDE and C N R RAO*
Solid State and Structural Chemistry Unit, Indian Institute of Science,
Bangalore 560 012, India
MS received 18 February 1982
Abstract. Ultraviolet and x-ray photoelectron spectroscopy have been employed
to investigate the adsorption of methanol, ethanol, diethylether, acetaldehyde,
acetone, methyl acetate and methylamine on surfaces of Fe, Ni and Cu. All these
molecules adsorb molecularly at low temperatures (^lOOK). Lone pair orbitals
of these molecules are stabilized on these metal surfaces (by 0-4-1-OeV) due to
molecular chemisorption. The molecules generally undergo, transformations as
the temperature is raised to 120 K or above. The new species produced seems to
depend on the metal surface. Some of the product species identified are methoxy
species, formaldehyde and carbon monoxide in the case of methanol and methyl
acetate, ethoxy species in the case of ethanol and 2-propanol in the case of acetone.
Keywords. Photoelectron spectroscopy ; adsorption of organic molecules ; UVPS
and XPS studies of adsorption.
1. Introduction
Adsorption of molecules on metal surfaces is fruitfully studied by techniques of
electron spectroscopy (Rao and Hegde 1981 j Rao 1981 ; Thomas 1974). Ultra-
violet photoelectron spectroscopy (UVPS) has been found to be specially useful
in investigating electron states of adsorbed molecules and in characterizing
adsorbed species (Spicer et al 1975 ; Lloyd et al 1977). In this laboratory, we
have investigated adsorption of CO, N 2 and O^ on transition metal surfaces by
employing uvps and related techniques (Kamath et al 1982a ; Rao et al 1982 ;
Jagannathan et al 1980). We considered it most worthwhile to systematically
investigate the adsorption of several organic molecules possessing lone-pair
orbitals on the surfaces of a few transition metals by employing UVPS. This is
because such molecules would be expected to chemisorb on metals through their
lone-pair orbitals and UVPS should directly give information on the nature of
bonding (Luth et al 1977). The molecules we have examined are methanol,
| Contribution No. 166 from the Solid State and Structural Chemistry Unit.
* To wham correspondence should be made.
101
ethanol, diethyl-ether, acetaldehyde, acetone, methyl acetate and methylamine
and the transition metals employed are Fe, Ni and Cu. It was our purpose to
compare the electron states of such a related series of adsorbate molecules on
the three metals and to study the thermal transformations of the adsorbate
molecules. We have obtained quantitative information on the stabilization of the
lone-pair orbitals of the different molecules due to chemisorption on metals by
matching the experimental difference uv photoelectron spectra with the gas
phase spectra (Rao et al 1979 ; Turner et al 1970) of the free molecules. We
have employed x-ray photoelectron spectra in the C (ls) t O (Is) and N (Is)
regions to study the nature of the adsorbed species. By means of the changes
observed in both UVPS and XPS, we have attempted to characterize the species
r esulting from the transformations of the adsorbed molecules. It has thus
been possible to show that all the molecules studied adsorb molecularly at low
temperatures, but undergo transformations at higher temperatures.
2. Experimental
All the spectra were recorded on the BSCA spectrometer of VG Scientific|Limited,
UK, fitted with a sample preparation chamber and a gas handling manifold.
Specpure strips of Fe, Ni and Cu were used. The metals were etched with argon
ions under UHV (~ 5 x 1(H torr) conditions to obtain atomically clean surfaces
(Rao et al 1980 ; Jagannathan et al 1980). All the organic compounds were
purified by fractionation. The metals were exposed to the adsorbate vapours in
the sample preparation chamber to the desired extent. Exposures are referred
to in Langmuirs, L (1L = 10~ 6 torr sec). The temperature of the sample could
be varied by using a special probe designed for the purpose, uv photoelectron
spectra were recorded with Hell radiation (40-81eV).
Difference spectra of adsorbed molecules were plotted with the aid of a DEC-'
1090 computer system wherein the spectra of the metals were subtracted from the
observed spectra after multiplication with an appropriate attenuation factor.
3. Results and discussion
3.1. Methanol
At low temperatures (~ 80 K), methanol is found to adsorb molecularly on Fe,
Ni and Cu surfaces. Thus, the Hell uv photoelectron spectra of methanol
adsorbed on these metal surfaces (figures 1^3) show features very similar to those
of methanol in gas phase. Difference spectra due to the adsorbed species are
also compared with the gas phase spectrum in figures 1-3. The binding energy
of the 6a" + la' band in the gas phase could be matched with the third band in
the difference spectra to obtain satisfactory electron states of methanol molecularly
adsorbed on the three metals. The energies and assignments are summarized in
table 1 . We see that there is a shift of the lone-pair orbital towards higher binding
bV about fl-6ieV fine to o.hfimisorntinn Tn -jr-rav rT-ir>trp1pntr/vi", ortra
(d)-(a)
(b)ht to 373 K
Figure 1. Hell spectra of methanol adsorbed on Fc at different temperatures and
exposures. Difference spectra are also shown in the figure along with the positions
of bands in the gas phase.
Significant changes occur both in the Hell spectra and the x-ray photoelectron
spectra on progressively heating the sample from 80 K to 300 K. The major
changes observed on the surfaces of the three metals are as follows :
3. la. The Hell spectrum of methanol adsorbed on Fe at 223 K (figure 1) shows
only two bands corresponding to 2a" and 6a" + la' orbitals. The bands corres-
ponding to la' and 5a r orbitals of methanol are absent in the spectrum. We attri-
bute this to the formation of the methoxy species, CH 3 O. Methoxy species is
known to be formed on heating methanol on Ni surface (Kojima et al 1981 ;'
Rubloff and Demuth 1977 ; Demuth and Ibach 1979) and Cu (Bowker and
Madix 1980; Ryberg 1981 ; Sexton et al 1981 \ Steinbach and Spenglet 1981;
Table 1. Electron state of molecularly adsorbed organic molecules.
Gas phase
Fe
Ni
Cu Assignment
Methcaiol
10-8
6-1(0-6)
5-8(0-6)
6-5(0-5) 2"( >
12-7
7-6
7-0
8-1 7o'( )
15-2
10-1
9-8
10-6 6a"+la'(<r co ',it ca }
15- 6
11-1
17-7
12-4
12-3
12-8 5a'
Dietliylether
9-61
6-2(0-0
4-8(0-0)
5-6(0-0)
11-08
. .
. .
8. Q
7-5*
7-7*
O
9-9
10-2
10-4
16-23
12-0
11-7
12-3
14-5
15-3
15-9
Acetaldehyde
10-3
5-5
a' ( )
13-24
8-0*
;'W
14-15
ct' (a)
15-34
a' (a)
15-6
10-5*
"(*<*,>
16-47
a'
Acetone
9-7
4- 6(0.- 5)
5-5(0-4) a' 00
12-6
. .
8-5*
9-1* *(>)
14-0
, .
a"(CH s )
15-7
10'2
11 '2 a"(n wa )
18-15
12-65
13-7 a'^cH,)
Methyl acetate
10-5
6-4(0-4)
5-8(0-3)
6-5(0-3)
11-3
^0=0
12-9
8-4
7-7
8-5
14-05
10-1
9-4
10-3
14-9
16-3
12:0
11-35
12-2
14-3
13-6
14-5
Methylamlne
9-6
6-0(1-0)
, .
5- 8 (Q- 6} H
13-2
8-55
. .
8> ^CH
14-3
'*
1
l"H
III
||.s
<D r CJ
1 a s
I
s IS "3
B 5?
ll-
o o< u
o s: o
~* ro **:
^3
-- 1
--388
u
8 B
oj co ^"^
I &3
p. tl -
-
rt {A
ll
I
o a
d d
15
d
rO
a
si
ll
x u
o -^
IS
3.1b : In the case of Cu, warming to 123 K results in the appearance of two
bands at 5-3 and 9-3 eV in the difference spectrum (figure 3). Formaldehyde
adsorbed on Cu gives a uv photoelectron spectrum (Kojima et al 1981) very
similar to that shown in figure 3 with bands at 5-4 eV and 8-5 eV (weak) as well
as a broad band extending between 8-9 and U-OeV with the maximum around
10 eV. The C (Is) and O (1 s) bands in XPS at 123 K are shifted towards lower
binding energies suggesting a decomposition of methanol (figure 5). It, therefore,
appears that methanol dissociates to give formaldehyde on the Cu surface studied
by us at 123 K. This observation differs from that of Bowker and Madix (1980)
who found the methoxy species on Cu (110) surface when methanol adsorbed
at 140 Kl was subsequently heated to 270 K ; these workers employed much
higher exposures than in the present study.
3. Ic : In the case of Ni, warming to 123 K results in drastic changes in the Hell
spectrum (figure 2). The C (Is) and O (Is) bands in XPS are shifted towards
lower binding energies indicating new species drue to decomposition of methanol.
We attribute the complex Hell spectra to the presence of more than one type of
species, probably CH 3 O, H 2 CO and other products ; such products are known
to be formed on Ni surface (Kojima et al 1981 ; Demuth and Ibach 1979).
3. Id : On warming to 300 K or higher, we observe considerable changes in the
UVPS of the adsorbed species on all the three metal surfaces. The C (Is) and
O (\s) bands in XPS are also further shifted to lower binding energies (figures 4
and 5) and appear around 284 -OeV (282- 7 eV for Cu) and 530-6 eV respectively.
These changes suggest the formation of CO and carbide species on the metal
surfaces due to the decomposition of the intermediate species observed at 123/
223 K.
Our observation of the occurrences of different types of transformations of
methanol on the three metal surfaces studied by us is interesting indeed. In table 2
we show the various adsorbed species formed on Fe, Ni and Cu surfaces at
different temperatures. We can understand the formation of CH 3 O and H 2 CO
from CH 3 OH on metal surfaces in terms of the sequences shown in chart 1. Thus,
the formation of methoxy species on Fe surface could occur as shown in sequence
(a) of chart 1. Formation of H 2 CO can occur either by sequence (b) or (c) in
chart 1. Sequence (c) has been noticed on Ag metal by Wachs and Madix (1978).
However, in the case of Cu, sequence (b) is more probable since the heat of chemi-
sorption of hydrogen (A# ) is somewhat low. In table 3 we have listed the
heats of chemisorption of hydrogen taken from Stevenson (1955). We see that
on Cu, A# a is lowest suggesting thereby that adsorption of hydrogen is not as
favoured. Formation of H 2 CO may therefore not occur by sequence (c) which
would require chemisorption of hydrogen on Cu.
3-2. Ethanol
Figure 6 shows the ultraviolet photoelectron spectrum of ethanol adsorbed on
Fe at 273 K. The spectrum shows four distinct bands at 5 5, 7- 8, 9 -2 and 10 5 eV.
On comparison with the gas phase spectrum of ethanol, we see that bands due
and imthylamirte adsorbed on metal surfaces.
Approximate
temperature
(K)
Fe
Ni
Cu
Metltanol
80
123/223
300
Ethanal
273
Diethylether
80
223-323
Methyl acetate
80
123/173
Metltylamine
80
123
173
300
GH 3 OH
CH a O
CO
C 3 H 5
(chemisorbed)
Transformation
products
CH 3 COOCH 3
CH,O
CH a OH
Transformation
products
CO
(physisorbed)
Transformation
products
CHgCOOCHa
Transformation
products
CH 3 NH 3 -f decomposed
products
Decomposition
products
CH 3 OH
H 2 CO
CO
(physisorbed)
Transformation
products
CH 3 COOCH 3
H 2 CO
CH 3 NH 2
Partial decomposition
Decomposition
products
(1981). Ethanol is apparently adsorbed molecularly only at low temperatures just
as methanol.
3.3. Diethylether
Figures 7-^9 show the UVP spectra of diethylether adsorbed at 80 K on Fe, Ni
and Cu surfaces. Difference spectra and the gas. phase spectrum are also shown
in these figures. In table 1 we have shown the positions and assignments of
chemisorbed diethylether. On matching the most intense band of the difference
spectra with the 17-2 eV gas phase band, we find that the lone-pair orbital shows
a shift of about 0-8 eV on Fe due to chemisorption ,' no such shift is observed
I H
M M
I.
I "7
M M
H
H f
M M
(a)
H
H \!/ H
M M M M
I *H H'
MM MM
(b)
M M M
2H 2 (g)
I - x - >
<f H
M M A .ii i A
Chart 1. Mechanism of formation of methoxy species and formaldehyde from
methanol adsorbed on metals.
Table 3. Heat of chemisorption of hydrogen (\H a ) on metals and the products
formed by methanol on their surfaces.
Metal
kcal/mole
Product on warmin g
to 123/2 23 K
Fe
31-6
CH 3
Ni
28-9
CH 8 0/H 3 CO
Cu
25-6
H ? CO
* From Stevenson (1955).
finding this difference in behaviour amongst the three metals ; errors in matching
the gas phase and difference spectra could partly contribute to this difficulty
since there are only three bands in the gas phase spectrum, xp spectra in the
C (Is) and O (Is) regions of Fe and Hi surfaces are shown in figures 10 and U.
The C (hr) region shows two distinctbands in the case of Fe at 286 5 and 289 6 eV
while in the case of Ni (and Cu, not shown in the figure) a single band at 285-9
is observed. The O (Is) band appears around 533-7 eV. XPS and UVPS data
suggest that molecular diethylether interacts in a distinctly different way with Fe
as compared to Ni and Cu.
On warming, drastic changes are observed in the UVP snectra ffi<mrfic 7_oA
I L 273 K
OL 80 K
5-0 "'" 15.0
Figure 6. Hell Spec tra of ethanol adsorbed on Fe at 273 K.
80ft hi to 223 K
E F =O
5.0
10.0
BE.eV
15.0
8
!-< **>
X o
o -*
o *
TJ -<
<o <a
-P -
P.
a
W
I
^ J
g a
B!
o
rt ti
J3 *
<o
a
s
11
&B
o
*2 P)
w fl)
w &
o
-p.
o Sj
*& p^ p^
lsl>
t3 52
59 .3
'
5
o S*.2
t ( pj ^1^
^ d 52
WrH
t/3 <->
fl S -S
.
its
-J 60
^
^!5
rt 'Q o
sil
~-< C3
a 2
'
a
appears from the difference spectra at these temperatures (223-323 K), that
diethylether undergoes different transformations on these metal surfaces. We
are n ot able to assign the spectra to any definite species at this stage.
3.4. Acetone
Hell spectra of acetone adsorbed on Ni and Cu at 80 K are shown in figures 12
and 13. Difference spectra are also compared with the gas phase spectrum in
these figures. It is clear that at 80 K acetone is molecularly adsorbed on both
Ni and Cu surfaces. The highest lying lone-pair orbital shows a chemisorption
shift towards higher binding energy by about 0-5 eV. Electron states of mole-
cularly adsorbed acetone and their assignments are listed in table 1. The O(lj)
bands in XPS (figures 14 and 15) appear around 532- 5 eV in both the metals ;
the C (Is) spectrum shows two distinct bands on Ni at 285 and 288 eV corres-
ponding to the two types of carbons (methyl and carbonyl respectively). These
bands appear at 291-23 and 293-88 eV in the spectrum of acetone in the gas
Cds)
(a) ht to 273 K
283
288 293
BE.eV
0(1s)
530
535
_ (a)ht to 273 K
) ht to 123K
(a) 0.35L 80K
540
PES study of adsorption of organic molecules on metals
115
COs)
A.OL 80K ht to173K
2.51 80K ht to 123K
0(1s)
Figure 15. Carbon Is and oxygen Is bands in xps of acetone adsorbed onCu at
different temperatures and exposures.
phase (Bafcke et al 1980). On Cu, the two bands in the C (Is) spectrum are not
resolved (figure 15).
We notice several changes in UVPS of adsorbed acetone on warming. The
resultant changes are noted below :
3-4a : In the case of Cu, on warming to 123 K, uvps shows definitive changes
as shown in figure 13. The resulting spectrum is similar to the gas phase spectrum
of 2-propanol (figure 13) reported by Katsumata et al (1973). The formation of
2-propanol implies surface reduction of acetone. The C (Is) and O (1*) bands
in XPS show a shoulder on the lower binding energy side at 123 K.
1L 213K
8.0
4.0
12.0
8.0
BE.eV
B e,v
< b) (a) hi to 298K
(a)
/
I
x 1L 213K
X
GAS PHASE
1
1
1
12.0
Figure 16. (A) Difference spectra of acetone adsorbed on Fe at different conditions,
(B) Carbon Is bands in xps and (C) Oxygen Is bands of acetone adsorbed on Fe.
On further warming to 300 K, we notice considerable changes in UVPS ; the
C(ls) and O(ls) bands are shifted to lower binding energies and suggest the
presence of adsorbed CO.
3.4c : Adsorption of acetone on Fe at 213 K shows broad features quite unlike
the gas phase acetone spectrum (figure 16). This and the appearance of C (Is)
and O (Is) bands in xps at lower binding energies (figure 16) indicate dissociative
adsorption of acetone at this temperature. Molecular adsorption occurs only
at low temperatures ( ~100 K). Further heating to 300 K results in the appearance
of C (Is) and O (Is) at still lower binding energies; these along with the changes
observed in UVPS are indicative of the formation of CO species.
3.5. Acet aldehyde
uv photoelectron difference spectra of acetaldehyde adsorbed on Fe are shown
in figure 17 along with the gas phase spectrum. Although we could match the
observed bands in the difference spectrum against the gas phase spectrum of
acetaldehyde, we find that the bands in the difference spectra are much too broad
vj ttwd
12
(c)
(b)ht to 323 K
(a)ht to 223K
0.5L 123K
GAS PHASE
16
12
BE.eV
Figure 17. Difference Hell spectra of acetaldehyde adsorbed on Fe at different
temperatures and exposures; gas phase spectrum is also shown.
to be due to molecularly chemisorbed species alone. It is possible that there is
some decomposition over the entire range 123 K-323 K. However, if we decide
to match the 13 -2 eV gas phase band with the second band in the difference spectra,
we find a lone-pair shift of 0-4 eV due to the molecularly chemisorbed species.
3.6. Methyl acetate
Hell photoelectron spectra of methyl acetate adsorbed on Fe, Ki and Cu at 80 K
are shown in figures 18-20. Difference spectra are also compared with the gas
phase spectrum (Sweigart and Turner 1972) in these figures. The 12-9 eV band
in the gas phase spectrum could be matched with the second band in the difference
spectra. The difference spectra show a single band corresponding to the 10-5 eV
(rt ) and ll-3eV (TT CO ) bands in the gas phase. The highest energy band in the
region 13 -5-14- 5 eV that we see in the spectra of the adsorbed species is not
found in the gas phase spectrum, since the latter was obtained with Hel excitation.
The lone-pair orbital (n ) shows a chemisorption shift towards higher binding
energy of ~0-3 eV on all the three metal surfaces. In table 1 we have listed the
energies and some of the assignments of UVPS bands of chemisorbed methyl
acetate. We see that the lone-pair shift is much lower in methyl acetate than
in methanol or acetone.
XPS spectra (figures 21-23) in the C (Is) region show two bands around 285-7 eV
and 289-7 eV (290-0 eV for Hi) ; the band at 285-7 eV is due to methyl carbons
while the carbonyl carbon is responsible for the band near 289-0 eV. The band
near 285-7 eV is about twice as intense as the 289-7 eV band as expected. This
is similar to the two bands observed in the C (Is) and O (Is) regions by Bowker
and Madix (1981) and Edwards (1976) in their studies on acetic acid. The O (Is)
n-^A r.f-1-nar.f,-. O nnl-.^i A C1 3 . 1 a.\T n. oil til a. +1-if>a
UJ
'O OQ
CJ O
3g
8 v
Ctf ;3
S
Ctf Qj rt
a & o
a\ g g
rH . nj >
| &^>
WO S O
|i4 -M ctf
"cS O
O
C/3
T3
^J . w
> p w TJ
N .g (U 2
uT 3 .8
QQ 5 ,
SIM
rv ra r,
U.
JLJ
SI
"5 a
a.
<+-< <D
o g
</j *i
CL
a t-<
X 03
Its
5
II
^
is a
1
la
u s
*a .
03 t
O 3
^ O
and XPS studies at 80 K clearly establish the presence of molecularty adsorbed
methyl acetate on the three metals. We, however, notice significant changes in
the spectra on warming the samples.
The main results of thermal effects are as follows :
3.6a : On the Fe surface, heating to 173 K causes drastic changes in the uvps
(figure 18). The resulting spectrum is identical to that of the methoxy species
obtained on adsorbing methanol at 223 K. The C (Is) region in XPS shows a
band at 284 8 eV while in the O (Is) region we see a band at 531 3 eV. In the case
of methanol, we found the C (Is) and O (Is) bands at 284-7 and 531 1 eV respec-
tively. It therefore appears that methyl acetate decomposes to give methoxy
species on Fe when warmed to 173 K.
3. 6b : On the H surface, warming to 173 K results in diffuse bands with features
similar to the 80 K spectra in the uvps (figure 19). In the C (Is) region, we see
bands at 285-7 and 289-0 eV and a new band on the low energy side at 284-0 eV.
The O (Is) region shows bands at 530-5 and 532-6 eV. These suggest that part
of the adsorbed methyl acetate has undergone decomposition.
3-6c : On warming to 123 K, the UVPS of methyl acetate adsorbed on Cu
shows drastic changes (figure 20). We see only two bands in the difference
spectrum. The difference spectrum is similar to that obtained on warming
methanol adsorbed on Cu from 80 K to 123 K wherein we found formation of
formaldehyde on the surface. The C(ls) band in XPS appears at 284- 6 eV at
this temperature and the O(ls) band is broad with the maximum at 530- 6 eV
and a shoulder at 532-4 eV. At 123 K, methanol adsorbed on Cu shows the
C (Is) band at 284- 5 eV ; the O (Is) band is broad with a maximum at 532-3 eV
and a distinct shoulder at 530- 6 eV. The C (Is) and O (Is) bands are as expected
of methyl acetate adsorbed at 123 K except that in the latter case the 530-6 eV
band is more intense indicating decomposition of methyl acetate. Based on these
XPS and uvps results, we suggest that formaldehyde is formed on warming
methyl acetate adsorbed on Cu just as in the case of methanol. It is interesting
that adsorption of methyl formate (HCOOCH 3 ) on Cu at 295 K also yields
formaldehyde (Kojima et al 1981).
In the adsorption of methyl acetate, only the lone-pair orbital, rt , shows a
relatively small chemisorption shift, but the orbital, TT CO is unaffected. Methyl
acetate, therefore, seems to interact with the metal through the lone-pair orbital
of the ether-type oxygen rather than the carbonyl oxygen as shown in chart 2.
The sequence of reactions taking place on metal surfaces is also shown in this
chart. Formaldehyde could be formed as per sequence (a) on Cu, while methoxy
species could be produced as per sequence (b) on Fe. It is interesting that, the
transformations found in methyl acetate are exactly the same as those found with
methanol. In view of this observation (and also because of the small shift of
the carbonyl lone-pair orbital energy), we seem to be justif ed in suggesting bonding
through the ether-type oxygen of the ester to the metal in chart 2.
3-7. Ammonia
HoC
M
9 -H- (a>
H/ X f M
I +
M OTHER PRODUCTS
,CH 3
H 3 C'
4 OTHER PRODUCTS
M .
Chart 2. Mechanism of formation of methoxy species and formaldehyde from
methyl acetate adsorbed on metals.
10L 60Khlto298K
5L 80K
2L 80K
CLEAN
5.0
10.0 . 15.0
B.E.cV ;
Figure 24. Hell spectra of ammonia adsorbed on Fe at different temperatures
ajtd exposures,
PES study of adsorption of organic molecules on metals
123
pond to the two bands at 10-2 and 14-9 eVin the gas phase spectrum of ammonia.
On warming to 298 K, we notice significant changes in the spectrum similar to
those reported by Kishi and Roberts (1977) suggesting dissociation of ammonia.
The NCls) signal of molecularly adsorbed species appears around 400 eV (figure
25). On warming this to 298 K, we see a weak band around 397 eV corres-
ponding to a nitrogen species (N or NH) formed due to dissociation of ammonia.
3.8. Methylamine
Hell spectra of methylamine adsorbed on Cu at 80 K and Fe at 173 K are shown
in figures 26 and 27. The difference spectra are compared with the Hel gas phase
spectrum in these figures. The gas phase spectrum was shifted by ~ 4-4eV
(for Cu) and ~4-6eV(forFe)so as to obtain the best matching of all the bands
between the gas phase and the difference spectra. The highest occupied orbital
(lone-pair orbital of nitrogen) on methylamine is shifted towards higher binding
energy by 1-0 eV on adsorption on Fe surface and by 0-6eV on Cu surface.
Electron states of adsorbed CH 3 NH a are tabulated in table 1. XP spectra in the
C(\s) region at low temperatures show bands at 286-0 eV (for Cu) and 285-0 eV
(for Fe) ; in the N(lj) region, bands are found at 399- 6 eV (for Cu), 399-1 eV
(for Fe) as shown in figures 28. and 29. The relatively lower binding energy of
the C(lj) peak on Fe could be due to partial decomposition of methylamine at
173 K on this surface. Such decomposition is also indicated by the broad maxima
in the uvp difference spectra.
N Os)
10L &OKhtto298K
B.E,eV
(b)-(a)
hl to 298K
B.E,eV
Figure 26. Hell spectra of methylamine adsorbed on Fe at different temperatures
and exposures. Difference spectra are also shown along with the positions of bands
in the gas phase.
UVPS of Cu + CH 3 NH 2 show considerable changes on warming to 123 K
and C(ls) and N(lj) bands in XPS are shifted towards lower binding energies
appearing at 285-0 eV and 399 -leV. These observations suggest occurrence
of decomposition of CH 3 NH 2 to a nitrogenous species on the metal surface (just
as in the case of NH 3 ). On warming to 300 K, drastic changes are seen in the
UVPS on both Cu and Fe surfaces with almost complete disappearance of the
bands due to molecular methylamine. The C (Is) band in XPS is shifted to
284-0 eV and the 1S( (Is) band shows little or no shift, but the intensity is appre-
ciably lowered.
4. Concluding remarks
(a) All the organic molecules containing lone-pair orbitals that we have examined
(d) - (a)
(c)-(a)
(b)- (a)
(b)ht toSOOK
(b)ht to123K
OL 80K
5.0
10.0
BE.eV
15.0
20.0
Figure 27. Hell spectra of methylamine adsorbed on Cu at different temperatures
and exposures. Difference spectra are also shown along with the positions of bands
in the gas phase.
(b) We are able to estimate the stabilization of the lone-pair ortitals due to
chemisorption in terms of the increase in its binding energy. These data are
presented in table 4 where we have listed the molecules in the older of increasing
ionization energy of the lone-pair orbital. The data suggest that the shift generally
shows an increasing trend with the decreasing first ionization energy as expected.
Furthermore, the stabilization of the lone pair seems to vary as Fe > Ni > Cu,
decreasing with increasing number of ^-electrons, a trend that is not readily under^
standable.
(c) Chemisorption of molecules gives rise to variation in the core level binding
energy of the atom containing the lone pair (Is binding energy of oxygen or
COs)
282 286
BE.eV
(O
397 401
BE.eV
Figure 28. Carbon 1* and nitrogen Is bands in XPS of methylamine adsor
Cu at different temperatures and exposures.
Table 4. Stabilization of the lone-pair orbitals (it eV) ori molecular chemis<
Molecule
Lone-
pair
IE
Lone-pair shifts
on
Fe Ni
Cu ,
Methanol
10-8
0-6 0-6
0-5
Methyl acetate
10-5*
0-4 0-3
0-3
Acetone
9-6
0-5
0-4
Diethylether
9-6
0-8
Methylamine
9-6
1-Q
0-6
NOS)
(b)
(a) ht to 298K
397
402
COs)
BE.eV
407 *V
ht to 298 K
(a)
__-/
1
^ ^
t
-5L 173K
1
283
288
' BE.eV
293 V
Figure 29. Carbon Is and nitrogen Is bands in XPS of methylamine adsorbed on
Fe at different temperatures and exposures.
(by matching C (Is) binding energies of adsorbates with those of free molecules),
we find the O(ls) binding energy increases by about 0-6 0-2 eV in the case
of methanol and acetone adsorbed on the three metal surfaces studied.
(d) Warming the substrate to temperatures above 120 K generally results in the
transformation of the adsorbate molecules in all the cases. The nature of the
species produced appears to depend on the metal surface and temperature as can
be seen from table 2. Methanol gives methoxy species or formaldehyde depending
on metal. These transformations are not only reflected in changes in the valence
band region (in uvps)but also in the C (Is) and O (Is) binding energies in XPS.
It is interesting that methyl acetate gives the same transformation products on
metal surfaces as methan ol (table 2). In the case of methylamine we fin d formation
of a nitrogenous species as in ammonia.
(e) The results of our studies on the adsorption of organic molecules find support
from Auger studies carried out for the first time in this laboratory (Kami th et al
1982b). Thus, Auger spectra also show the transformation of methanol to
formaldehyde on Cu surface.
Acknowledgement
"U.R.B.^ .rx n., v^uou n. vv uuu. juuy vv i-. i?ou ./. JZieciron SpeclrOsc. K61. rnenom. zu jjj
Bowker M and Madix R J 1980 Surf. Set. 95 190
Bowk&r M and Madix R J 1981 Appl. Surface Sci. 8 299
Carlson T A, Agron P A, Thomas T M and Grimm F A 1981 /. Electron Spectrosc. Rd.
Phenom. 23 13
Demuth J E and Ibach H 1979 Chem. Phys. Lett. 60 395
Edwards D 1976 Inorg. Chiin. Acta 18 65
Jagannathan K, Srinivasan A, Hegde M S and Rao CNR 1980 Surf. Sci. 99 309
Karaath P V, Yashonath S, Srinivasan A, Basu P K and Rao CNR 1982a /. Indian Chem.
Soc. 59 153
Karaath P V, Yashonath S, Srinivasan A and Rao CNR 1982b Appl. Surf. Sci. (in print)
Katsumata S, Iwai T and Kitnura K 1973 Bull. Chem. Soc. (Jpn.} 46 3391
Kishi K and Roberts M W 1977 Surf. Sci. 62 252
Kojiraa I, Sugihara H, Miyazaki E and Yasumori I 1981 J. Chem. Soc. Faraday Trans. 1 77
1315
Lloyd D R, Quinn C M and Richardson N V 1977 in Surface and defect properties of solids
(eds) M W Roberts and J M Thomas Specialist Periodic Reports, Chem. Soc. London 6
179-227
Luth H, Rubloff G W and Grobraan W D 1977 Surf. Set. 63 325
Rao CNR 1981 Proc. Indian Nat. Sci. Acad. A47 111
Rao C N R, Basu P K and Hegde M S 1979 Appl. Spec. Rev. 15 1
Rao CNR and Hegde M S, 1981 in. Preparation and characterization of materials (eds) J M
Honig and CNR Rao (New York : Academic Press)
Rao CNR, Kamath P V and Yashonath S 1982 Chem. Phys. Lett. 88 13
Rao CNR, Sarma D D and Hegde M S 1980 Proc. R. Soc. (London) A370 269
Rubloff G W and Demuth J E 1977 /. Vacuum Sci. Technol. 14 419
Ryberg R 1981 Chem. Phys. Lett. 83 423 -
Sexton B A 1981 Surf. Sci. 102 271
SpicerWE,YuK Y, Lindaul, PianettuP and Collins D M 1975 in Surface and defect properties
of solids (s^ltf. W Roberts and J M Thomas Specialist Periodic Reports Chem. Soc. London
5 103-161
Steinbach F and Spengler H J 1981 Surf. Sci. 104 318
Stevenson D P 1955 /. Chem. Phys. 23 203
Sweigart D A and Turner D W 1972 J. Am. Chem. Soc. 94 5592
Thomas J M 1974 in Progress in surface and membrane science (eds) J F Danielli and
D A Cadenhead (New York : Academic Press) 8 49
Turner D W, Baker C, Baker A D and Brundle C R 1970 Molecular photoelectron spectroscopy
(New York : Wiley Interscience)
Wachs I E and Madix R J 1978 Surf. Sci. 76 531
MNDO study of reaction paths : Hydroboration of
carbonyl systems
NABA K RAY* and RITA CHADHA
Department of Chemistry, University of Delhi, Delhi 110007, India
MS received 7 January 1982
Abstract. The hydroboration reactions of acetaldehyde and acetone have been
investigated by the MNDO method. The reactions have been shown to be two-
step reactions involving an intermediate adduct. This adduct subsequently under-
goes hydrogen rearrangement. The hydroboration reactions of acetaldehyde and
acetone have been compared with the corresponding reaction of formaldehyde.
The charge transfer effects accompanying these reactions have also been discussed.
Keywords. Hydroboration of acetaldehyde and acetone; effect of methyl substitution;
charge transfer effects.
1. Introduction
In the year 1939, Brown et al demonstrated that diborane reacts rapidly with
simple aldehydes and ketones, such as acetaldehyde and acetone, to produce the
corresponding dialkoxyboranes (Brown et al 1939). Since these substances are
readily hydrolysed to form acid and the corresponding alcohol, it is evident that
the procedure offers a promising route for the reduction of carbonyl groups. For
this reason, we have undertaken a theoretical study on the hydroboration of
carbonyl systems.
The results of our study on the hydroboration of formaldehyde (Ray and
Chadha 1981) had shown that the reaction is a two-step reaction involving the
initial formation of an adduct, followed by its rearrangement to the product via
a four-centre-like transition state. This latter step was shown to be the rate-
determining step. In the present work, we have carried out calculations on the
reaction of borane with acetaldehyde and acetone.
2. Method of calculation
The Modified Neglect of Diatomic Overlap (MNDO) method with the usual
parameters (Dewar and Thiel 1977: Dewar and McKee 1977) was employed for
all the calculations reported here. For the methyl group, C 3l> symmetry was
maintained throughout the calculations. All other geometrical parameters were
completely optimized. The transition states were located by the energy
minimization method (Rothman and Lohr 1980). The incipient C-H bond length
was employed as the reaction coordinate and the heat of formation of the
systems was plotted as a function of this reaction coordinate. Rothman and
Lohr (1980) have shown that the maxima obtained on such reaction surfaces are
transition states, provided that the reaction pathway is continuous. The force
constant matrix at such points also has only one negative eigenvalue.
3. Results and discussion
The reaction profile for the hydroboration reaction of acetaldehyde with borane
(figure 1) indicates two transition states and one intermediate along the reaction
pathway. Hence, like the corresponding reaction of formaldehyde, this is a two-
step reaction, involving the initial formation of an adduct as a stable intermediate.
The activation energy required for its formation is 26 -5 kcal/mol, a value much
higher than that found (Ray and Chadha 1981) for formaldehyde (9 7 kcal/mol).
Thus, the substitution of a hydrogen atom in formaldehyde by a methyl group
results in a considerable increase in the activation energy required for the first step.
The adduct is more stable than the reactants by 9-5 kcal/mol (c.f. 9-2 kcal/mol
for formaldehyde). Methyl substitution, therefore, has only a small effect
on the relative stabilities of reactants and adduct. The activation barrier for the
second step increases to 28 -2 kcal/mol. The reaction enthalpy is < 43 1 kcal/mol.
Figure 2 gives the optimized geometries of the four stationary points on the
reaction surface. As in the case of formaldehyde, the carbon-oxygen bond
distance increases as the reaction proceeds. The B~O bond distance decreases
to 1 50 A in the adduct, after which it increases to a value of 1 53 A in T2, and
(R)
U820
Figure 2. The optimized geometries (in angstroms and degrees) for the transition
state (Tl) for formation of the adduct, the adduct (A), the transition state (T2) for
the formation of product (?) for the hydroboration of acetaldehyde.
then decreases again to a value of 1 34 A in the product. The COB bond angle
decreases from its value of 136-5 in T\ to 87-2 in T2, after which it increases to
127-2 in the product. In this case, too, T2 has a four-centre-like structure.
The hydroboration reaction of acetone, which is the simplest ketone, has also
been studied. The profile of this reaction is similar to that of the corresponding
reaction of acetaldehyde (figure 3). The first step, the formation of adduct,
requires an activation energy of 26 -4 kcal/mol. Therefore, the activation barrier
for the first step is almost the same for the hydroboration of acetaldehyde and
acetone. The second step proceeds with an activation energy of 31-3 kcal/mol.
The reaction enthalpy is < 36 -2 kcal/mol.
Figure 4 gives the optimized geometry of each stationary point on the reaction
surface. As expected, the carbon-oxygen bond distance increases as the reaction
proceeds. The B-O bond distance is 1 -49 A in the adduct, after which it increases
to a value of 1 52 A in T2, and then decreases to 1 34 A in the product. The
COB bond angle also decreases from its value of 144-7 in Tl to 89-2 in 72,
after which it increases to 126-9 in the product. T2 again has a four-centre-
like structure. Comparison of figure 2 with figure 4 indicates that the geometries
at the stationary points are essentially similar for both the reactions studied here.
(Tl)
(T2)
(R)
(P) /
Figure 3. The energy profile (kcal/mol) for the reaction of acetone with borane.
Figure 4. The optimized geometries (in angstroms and degrees) for the stationary
points on the reaction surface for hydroboration of acetone.
4. Charge transfer effects
The charges on various atoms in the stationary points on the potential surface for
Af nrn Z>
Charge
R
Ti
A
T2
P
C 0-24
0-27
0-29
0-41
0-17
O -0-28
-0-31
-0-03
-0-24
-0-25
B 0-24
0-23
-0-21
-0-15
o-u
H x -0-08
-0-13
-0-12
-0-11
Q-00
H 2 -0-08
-0-09
-0-06
-0-03
-0-05
H 3 -0-08
-0-04
-0-07
-0-03
-0-02
H 4 0-01
0-03
0-09
0-06
0-02
CH 3 0-03
0-04
O'll
0-09
0-02
a see figure 1
* see figure 2.
Table 2. Charges on various atoms
surface for hydroboration of acetone.
in the
stationary points on
the reaction
Atnm 8
Charge
R
TI
A
T2
P
C 0-19
0-23
0-25
0-39
0-13
O -0-29
-0-31
. -0-04
-0-26
-0-24
B 0-24
0-23
-0-20
-0-13
0-10
^ -0-08
-0-13
-0-12
-0-14
0-02
H 2 -0-08
-0-09
-Q-07
-0-03
-0-05
H 3 -0-08
-0-05
-0-07
-0-03
-0-02
(CH 3 )i 0-05
0-06
0-11
0-10
0-03
(CH 8 ) a 0-05
0-06
0-14
0-10
0-03
see figure 3 6 see figure 4
Table 3. Heats of formation at stationary points on the reaction surface for
hydroboration of carbanyl systems.
Hoat of formation (kcal/mol)
Point
Acctaldohyde Acetone
R
-30-6
-37-1
Tl
-4-i
-10-7
A
-40-1
-45-6
T2
-11-9
-14-3
P
-73-7
-73-3
/l D (lccal/mol)
28-2
31-3
" sec figures 1 and 3.
D Energy of activation for tho rate-determining stop.
The progress of the reaction from Tl to A is accompanied with an increase in
charge density on the boron atom by 0-44 units. The total amount of charge
transferred from acetaldehyde to borane, in the adduct, is 0-46 units. In T2,
the amount of charge transferred reduces to 0-32 units, due to back donation of
charge to the oxygen atom.
The charges on various atoms in the stationary points on the potential surface
for hydroboration of acetone are listed in table 2. In this case, the amounts of
charge transferred from acetone to borane in Tl, A and T2 are, respectively,
0-04, 0-46 and 0-33 units.
In all the cases, therefore, the adduct is a charge transfer complex and the
amount of charge transfer from the carbonyl system to borane is nearly constant.
5. Conclusions
The present study has indicated that the reaction path for the hydroboration of
aldehydes and ketones includes an intermediate charge transfer adduct and two
transition states. The pronounced effect of substitution of one hydrogen atom
of formaldehyde by a methyl group on the activation barrier to the first step is
probably due to steric factors. Unlike the case of nitriles (Chadha and Ray 1982)
the transition state for the first step is more susceptible to steric hindrance.
The results of the present work are in agreement with the mechanism proposed
by Brown and Subba Rao (1960). As expected, methyl substitution increases
the stability of the adduct (table 3). The electron donating property of the methyl
nrrmm 1C rocr\/-moiMr ff\n fTlio /aflFVrt T il/-/* 4-l^a, Tvtr/1-r/Vl-wvrSiti/Yii ra.vr'+in-n n.f f~.---.o1
he authors (RC) thanks csm, New Delhi, for a Senior Research Fellow-
lanks are also due to the Staff of the Computer Centre, Delhi University,
operation.
C, Schlesinger H I and Burg A B 1939 J. Am. Chem. Soc. 61 673
C and Subba Rao 1960 J. Am. Chem. Soc. 82 681
J S and McKee M L 1977 /. Am. Chem. Soc. 99 5231
J S and Thiel W 1977 /. Am. Chem. Soc. 99 4899, 4907
and Chadha R 1981 Indian J. Chem. A20 492
and Chadha R 1982 Theor. Chim. Acta 60 451
M J and Lohr L L Jr 1980 Chem. Phys. Lett. 70 405
Infrared and Raman spectra and thermodynamic functions
of 4-methoxypyridine N-oxide*
K C MEDHI
Department of Physics, Gauhati University, Gauhati 78 1 014, India
MS received 24 February 1982
Abstract. The infrared spectrum of 4-methoxypyridine N-oxide in the region
4000-30 cm" 1 in the solid and liquid states and the polarized laser Raman spectrum
of the molecule in the liquid state have been investigated. A vibrational assignment
of the observed frequencies based on the state of polarization of the Raman lines
and comparison with the related molecules is presented. Ideal gas state thermo-
dynamic functions of the molecule are calculated in the temperature range 273-15-
1500 K.
Keywords. Infrared spectrum ; Raman spectrum ; thermodynamic functions ;
4-methoxypyridine N-oxide.
1. Introduction
As a part of the earlier investigations on the vibrational assignments for some
substituted pyridine molecules (Medhi et al 1965 ; Medhi 1965, 1972, 1977), the
infrared and Raman spectra of 4-methoxy pyridine N-oxide have been studied in
the present work. The infrared spectrum of this molecule in a limited frequency
range was reported by Hideyo Shin do (1958), Costa and Blasina (1955) and
Ghersetti et al (1973), and only a few frequencies were assigned. There is, how-
ever, no earlier study on the Raman spectrum of 4-methoxypyridin.e N-oxide.
The present paper gives a detailed account of the infrared and Raman spectra
of the molecule of 4-methoxypyridine N-^oxide, and a vibrational assignment of the
observed frequencies, based on the polarization of the Raman lines and comparison
with the assignments proposed for other related molecules (Green 1962 ; Green
et al 1963 ; Long and George 1963 ;. Allan et al 1971 ; Berezin and
Elkin 1973 a) is presented. The ideal gas state thermodynamic properties of
the molecule are also calculated on the basis of this vibrational assignment.
2. Experimental
The sample of 4-methoxypyridine N-oxide was supplied by Aldrich Chemical
Company, U.S.A- *t was purified by repeated sublimation in vacuum.
equipped with an argon ion laser. The 514- 5 nm line (200 mw) was used to
excite the spectrum. The polarization of the Raman lines was measured by the
same method as described previously (Medhi 1977). A Carl-Zeiss Specord IR
75 spectrophotometer was employed to measure the mid infrared spectrum in the
range 4000-400 cmr 1 either in hexachlorobutadiene (HCB) or in nujol mull as well
as in thin liquid 1 Im prepared from the melt of the substance. The far infrared
spectrum from 500 to 30cm- 1 was obtained in nujol mull using a Perkin-Elmer
Model 180 spectrophotometer. Because of the high boiling point of the sub->
stance, the vapour-phase infrared spectrum could not be obtained.
3. Results mid discussion
Table 1 gives the details of the infrared and Raman spectra of 4-methoxypyridine
N-oxide together with the probable assignments for the observed frequencies.
The fundamental frequencies are summarized in table 2. The calculated values
of the thermodynarnic functions are listed in table 3.
Table 1. Observed infrared and Raman frequencies and assignments for 4-methoxy-
pyridine N-oxide.
Infrared (cm" 1 )
R?man (cm" 1 ]
f
gnmont
HCB or nujol mull
Melt
Melt
3112(s)
3065 (mw)
3041 (s)
3026 (ras)
2997 (mw)
3108 (s)
3060 (mw)
3025 (s)
3092 (3) p
'3022(l)p
Vi (a 1 )
v 2 (a 1 )
V4 (a 1 )
v 29 (V
-CH 3 asymmetric
2932 (w)
2897 (mw)
2847 (ms)
2812(w)
2785 (ins)
2581 (w)
2235 b (ms)
2045 (mw)
2011 (w)
1941 (w)
1903 (raw)
1889(sh)
2943 (mw)
2895 (mw)
2840 (s)
2769 (ms)
2571 (w)
2509 (w)
2038 (ms)
1927 (w)
I877(mw)
1755 (w)
Stretching
2948 (1 ) p v s (a'), - CH 3 asymmetric
stretching
' 2 x v u U')
2901 (0) v u + v lz (A')
2845 (2) p v (a'), -CH 3 symmetric
stretching
v 9 4"V ia U')
V 8 +'V 20 U')
v 9 4- V M U')
Vu 4- v ?4 (A')
V H 4' v a4 U')
v! 9 4'vjlu')
2 4'V 2 [U')
I + v 31 (A')
Table 1 (Contd.)
Infrared (cm" 1 )
Raman (cm." 1 )
A on5rt***v.f
HCBornujolmull
Melt
Melt
1623 (s)
1625 (s)
1624 (3) p
V 7 (')
1566(sh)
I560(sh)
Vj.8 "t-Va,,^')
1 563 (ms)
1 557 (ms)
1559(l)p
v 8 (#')
1520(sh)
1520 (sh)
v 22 4- v 25 (A 1 )
I509(sh)
I510(sh>
2 x v 21 (,4M
1496(vs)
1490(vs)
1.490 (0) p
V B (a')
1471*(s)
1466t(ms)
VIQ (<*') and v 30 (a"), CH 3
asymmetric deformations
1462* (vs)
1455t (ms)
1458 (0)p
Vu (a'), CH 3 symmetric
= 0-
73 deformation
1439 (s)
1441 (s)
1442(sh)
VH (a')
1415(0)p
Vi 7 +V M U')
1319 (ms)
13lO(sh)
1310 (sh)p
V 13 (fi')
= 0-
63
1298(vs))
I291(vs)|
1289(vs)
1290(2}p
Vu (a')
1233(msO
Vl 5 (ft'}
1205 (vs)
1227(vs)
Vj 6 (a')~N-O stretching
1188 (sh)
1184(sh>p
V35 + Vw(A'}
1180 (s)
1172 (s)
1175 (4) p
V J7 ('}
H24(w)
v 2B + v 2o (A')
1111 (w)
1099(s)
1102 (0)p
Vis (')
1058(0) dp
Vai (a"), CHaOut-of-plane rocking
1033 (s)
Vao(a')
1014(vs)
1022(vs)
1027 (l)p
Vao (a')and V2a(fl').-O--CH 3 strctching
962 (w)
955 (w)
v 32 (a")
900 (w)
897 (w)
v aa (a)
856 (sh)
855 (sh)
858(10)p
v 22 (ff'),-N-O in-plane deformation
850 (vs)
v 34 (a")
836 (sh)
840 (vs)
842 (sh) p
Va. 3 (a'), CH a in-plane rocking
811*(ms)
~810(sh)
v 35 (a")
756 (vs)
758 (vs)
758(0)p
v w (a')
690 (mw)
704 (w)
708 (0) dp
v 36 (a*)
657 (raw)
657 (mw)
661 (l)p
v 2B (a')
= o-
74
586b(s)
?
542 (w)
540 1 (w)
541 (0) dp
Va? (") -N-O out-of-plane
deformation
522(s)
526(ms)
530 (0) dp
V 38 (")
462 (ms)
463 (ms)
465 (0)p
Vaa(^') Z.^^^ deformation
420 (sh)
v 39 (a"}
405 (ms)
406 (l)p
v 27 (a')
376 (mw)
370 (0) dp
v 40 (a"), /.COG deformation
252 (mw)
244 (0) p
v 28 (a 1 )
198 b (ms>
v 42 (a"}
148 (mw)
f
85 (raw)
< lattice modes
47 (w)
I
* Frequency observed in solid film prepared from the melt.
Symmetry Mode No. Fund.
species (Wilson 1934)
Wavertumber
(cm- 1 )
a' 20b Vi
3092
2 v
3060
20a v 3
3041
7b v 4
3022
V G
2948
V
2845
8a v 7
1624
8b v 8
1559
19a v
1490
V JO
1466
v u
1458
19b v
1441
14 v 33
1310
13 v u
1290
3 V,
. u
1.233
v l6
1227
9a v, 7
1175
15 v, 8
1102
18a v l9
1033
1 v ao
1027
v ai
(1027)
v aa
858
v aa
840
12 v a4
758
6b v as
661
V-ifl
465
6a v n
406
18b v 2B
244
a" v
2997
V 30
(1466)
V 31
1058
17a v 33
955
5 v 33
897
10a v si
850
10b v 35
811
4 v
708
Vy,
541
11 V a8
530
1 6a v 89
420
V 40
370
16b v
198
( ) Frequency assigned more than once.
? Frequency not assigned.
(O T/A
IVJ
V
Cal/degree/
mole
T
Cal/degree/
mole
Cal/degree/
mole
r
Cal/degree/
mole
273-15
26-93
15-45
81-58
66-13
298-15
29-43
16-52
84-04
67-53
300
29-61
16-60
84-22
67-63
400
39-06
21-06
94-06
73-01
500
47-03
25-48
103-67
78-19
600
53-46
29-63
112-83
83-20
700
58-65
33-42
121-47
88-06
800
62-89
36-84
129-59
92-75
900
66-40
39-94
137-20
97-27
1000
69-35
42-74
144-36
101-62
1100
71-83
45-27
151-09
10.5-82
1200
73-93
47-58
157-43
109-86
1300
75-73
49-67
163-42
113-75
1400
77-27
51-59
169-09
117-50
1500
78-60
53-35
174-47
121-12
* Ideal gas state at standard pressure of 1 atom.
The molecule of 4-methoxypyridine N-oxide is assumed to have the C t symmetry
with 28a' and 14" normal modes of vibrations predicted from group theoretical
considerations. All the forty-two normal vibrations are active both in the infrared
and Raman spectra. The Raman lines of a' species should be polarized, and
those belonging to the a" species depolarized.
3.1. Pyridyl vibrations
3. la Class a' : There appear four frequencies in the C-H stretching region.
These are readily assigned to such valence oscillations. The polarized Raman
lines observed at 1624, 1,559 and 1490cm.- 1 and the strong infrared band at 1,439
cm" 1 may be attributed to the modes primarily derived from the stretching of the
ring. The infrared spectrum shows very strong absorption at 1291cm.- 1 ;. its
Raman counterpart located at 1290cm- 1 is polarized. This may be confidently
assigned to the mode characterised as C-iO stretching in agreement with the assign-
ments made in the case of other molecules (Katritzky and Coats 1959 ; Briggs
et al 1957;. Sax et fl/1960; Spinner and White 1962; Katritzky 1959). Two other
substituent-sensitive modes belonging to the a! class may be identified with the
polarized Raman lines at 758 and 244 cm- 1 . Infrared bands, of medium strength
are observed at 657 and 405cm.- 1 . The corresponding Raman line at 661 cm.-*
appears to be depolarized and that at 406cm" 1 is polarized. These may be rea-
sonably assigned to the planar ring angle deformation modes (Green et al 1963 ;
Berezin and Elfcin 1973a). The remaining modes belonging to the a 1 class may be
and the moderately intense infrared band at 1233 cm- 1 .
3-lb Class a" : The Raman lines observed at 708 and 530cm- 1 are depola-
rized and are assigned as a" fundamentals arising principally from the out-of-plane
bending motions of the ring carbon atoms and the ring hydrogen atoms respec-
tively. The remaining a" modes are identified with the observed infrared fre-
quencies at 962, 900, 850, 8U, 420 and 19 8 cm- 1 in agreement with the assignments
proposed for related molecules (Medhi 1977 ;. Green et al 1963 ; Long and
George 1963).
3-2. N-0 vibrations
The infrared spectrum shows very strong absorption at 1205 cm- 1 , this may be
attributed to the mode that involves the stretching of the NO bond in agreement
with the previous assignment (Costa and Blasina 1955). The in-plane N-O defor-
mation mode is unambiguously assigned to the intensely strong polarized Raman
line at 858 cur 1 , whilst the out-of-plane bending mode may be responsible for the
depolarized Raman shift observed at 541cm.- 1 (Berezin and Elkin 1973 b ;
Katritzfcy and Coats 1959 ; Wiley and Slaymaker 1957 ; Colthup et al 1975).
3-3. OCHz vibrations
The two components of the out-of-phase CH 3 stretching mode may be assigned
to the polarized Raman shift observed at 2948cm.- 1 and the moderately intense
infrared band at 2997cm- 1 , while the symmetric stretching mode is identified
with the polarized Raman line at 2845cm"" 1 (Green 1962 ; Badger and
Moritz 1959 ; Pozefsky and Coggeshall 1951 ; Wiberley et al I960 ; Henbest
et al 1957 ; Seth-Paul et al 1974). It is known that the symmetric CH 3 defor-
mation mode, being somewhat sensitive to the electro-negativity of the attached
oxygen atom in the OCH 3 group, shifts to higher frequency (Sheppard 1955 ;
Bellamy and Williams 1956; Wilrmhurst 1,957). Thus the infrared band appear-
ing with considerable intensity at 1462cm" 1 is reasonably assigned to this
mode, The CH 3 asymmetric deformations may be responsible for the observed
frequency at 1471cm- 1 . While the assignments of the CH 3 rocking modes are
less certain, the torsional mode is not located in this case.
In methoxy compounds usually a strong band appears in the interval 1050--
1010cm" 1 which has been interpreted as O-CH 3 stretching fundamental (Kat-
ritzfcy and Coats 1959; Briggs et al 1957; Sax et al 196.0; Spinner and White
1,962). In agreement with this the strong infrared band observed at 1014cm- 1
is assigned to such mode. The in-plane and out-of-plane O-CH 3 deformation
may be represented by the polarized Raman line at 465cm- 1 (Allan et al 1971)
and the depolarized Raman shift observed at 370cm- 1 , respectively.
In addition to these, the infrared spectrum shows strong but somewhat broad
absorption at 586cm- 1 in the solid state ; but the band completely disappears on
melting the substance by raising the temperature-, or when the compound is dissol-
ved in methylene chloride. Therefore, it cannot be assigned as a fundamental
Besides the above fundamental frequencies, the infrared and Raman spectra of
4-methoxypyridine N^oxide show some weaker bands which are interpreted as
overtones and combination bands.
4. Thermo dynamic properties
Ideal gas state thermodynamtc functions of 4-methoxypyridine N-oxide were deter-
mined at several temperatures between 273-15 and 1500 K using the fundamental
frequencies given in table 2 and the following structural parameters and relative
atomic masses :
C-C = I -397 A, C->H = 1-084 A, NO = 1-37 A, Z.CCC = /.CNC =
Z.CCH = L CNO = /.COG = 120.
For OCH 3 group -CXH = 1-0936 A, C-O = \ -426 A, and all angles are
tetrahedral.
H= 1-0079, C = 12-011, O = 16-0, N = 14-01 amu.
The rotational constants and the reduced moment of inertia of the methyl group,
as calculated from the structure defined above, were found to be A = 0- 14860cm- 1 ,
B = 0-03347 cm- 1 , C = 0-02746 cm- 1 and I m = 5-2668 x 1(H gm cm 2 . The
thermodynamic functions were calculated by assuming a rigid rotor, harmonic
oscillator approximation and free internal rotation of the methyl group.
Acknowledgements
The author is thankful to Professor I Haque for his kind permission to record the
mid infrared spectrum in his laboratory. The author is also thankful to
S K Barooah for his assistance in taking some infrared records.
References
Allan A, McKean D C, Perchard J P and Josien M L 1971 Spectrochim. Acfa A27 1409
Badger G M and Moritz A G 1959 Spectrochim. Acta 15 672
Bellamy L J and Williams R L 1956 /. Chem. Soc. London p. 2753
Berezin V I and Elkin M D 1973a Opt. Spectrosc. (USSK) 34 395
Berezin V I and Elkin M D I973b Opt. Spectrosc. (USSK) 35 49
Briggs L H, Colebrook L D, Fales H M and Wildman W C 1957 Anal. Chem. 29 904
Colthup N B, Daly L H and Wiberley S E 1975 Introduction to infrared and Raman spectro-
scopy (New York : Academic Press) p. 272
Costa G and Blasina P 1955 Z. Phys. Chem. 4 24
Ghersetti S, Giorgianni S, Capricci P L and Spunta G 1973 Spectrochim. Acta A29 1207
Green J H S 1962 Spectrochim. Acta 18 39
Green J H S, Kynaston W and Paisley H M 1963 Spectrochim. Acta 19 549
Uo.rJiPct M "R A/r^oVmQ d n NTif.Tinlls PI anH Wnorland A A 1Q<J7 T rlien, Knr T n-nJn-n n 1<dfi9
Medhi K C 1965 Opt. Spektrosk. 19 49
Medhi K C 1972 Indian J. Phys. 46 300
Medhi K C 1977 Indian J. Phys. A51 399
Medhi K C and Mukherjee D K 1965 Spectrochim. Acta 21 895
Pozefsky A and Coggeshall N D 1951 Anal. Chem. 23 1611
Sax K J, Saari W S, Mahoney C L and Gordon J M I960 /. Org. Chem. 25 1590
Seth-Paul W A, Tollenaere J P, Meeusen H and Ho'fler F 1974 Spectrochim. Acta
Sheppard N 1955 Trans. Faraday Soc. 51 1465
Spinner E and White J C B 1962 J. Chem. Soc. p. 3115
Wiberley S E, Bunce S C and Bauer W H i960 Anal. Chem. 32 217
Wiley R H and Slaymaker S C 1957 J. Am. Chem. Soc. 79 2233
Wilmshurst J K 1957 /. Chem. Phys. 26 426
Wilson E B 1934 Phys. Rev. 45 706
Carbon-13 nuclear magnetic resonance studies on high spin iron(III)
porphyrins
D V BEHERE and S MITRA*
Chemical Physics Group, Tata Institute of Fundamental Research, Colaba,
Bombay 400 005, India
MS received 10 February 19S2
Abstract. Carbon-13 NMR studies on a series of high spin iron(UI) porphyrins,
namely tetraphenylporphyrin iron(ni) halides [Fe(TPP) X, X = Cl, Br, I] in
CDC1 3 solution are reported. As expected the 13 C shifts are found to be an order
of magnitude larger than the corresponding proton shifts. The dipolar contri-
bution, which is quite important for the proton NMR, becomes much less signi-
ficant for the 13 C shifts. No systematic variation in the 13 C shift across the series
is observed, except for the meso-carbon which shows a small but gradual decrease
in going from the chloro to the iodo complex. The 13 C shift for the various carbon
atoms of the porphyrin ligand shows interesting pattern which is discussed in terms
of spin delocalisation mechanisms.
Keywords. Nuclear magnetic resonance ; high spin iron(IH) porphyrins ; C-13
shifts.
1. Introduction
Nuclear magnetic resonance (N;MR) studies on synthetic iron porphyrins have
been very useful in understanding the magnetic and electronic properties of iron
in haem proteins (Wuthrich and Baumann 1973a,b, 1974; Goff 1978, 1981; LaMa r
and Walker 1979). Most of the NMR studies on iron porphyrins have been
done on proton nucleus. 13 C NMR is, however, a more direct, sensitive and accU"
rate probe for paramagnetic complexes. Besides being a direct probe to the
distribution of unpaired spin across the porphyrin skeleton, 13 C NMR has an
additional advantage over- proton NMR. Since the 13 C shifts are generally an
order of magnitude larger than the corresponding proton shifts (Horrocks 1973 j
Doddrell and Gregson 1974 ; Mitra 1977) the dipolar contribution to the 18 C
shifts is expected to be usually negligible, which makes the interpretation of the
data easier.
We have recently reported (Behere et al 1982) a detailed proton NMR study on
a series of five coordinated high spin tetraphenylporphinato iron(III) halides,
[Fe(TPP)X, X = C1, Br, I (figure 1)]. These studies confirmed the dominant
* To whom correspondence should be made.
145
Figure 1. Molecular geometry of Fc(TPP) X.
influence of the porphyrin ligand on the shift pattern., since the changes in the
axial halide ligand were found to have minimal effect on the shifts of the various
protons. Based on the pattern of the unpaired spin density distribution across
various proton sites, a mechanism for the delocalisation of the unpaired spin
across the porphyrin ligand was suggested (Behere et al 1982). The present 13 C
NMR study is an extension of our proton NMR study (Behere et al 1982) and
forms a part of our research programme on metalloporphyrins (Behere and
Mi tra 1979, 1980; Behere et al 1977, 1979, 1981, 1982). 13 C NMR on Fe(TPP)Cl
had been reported earlier by Ooff (1978) but a subsequent study by Mispelter
et al (1979) proved that the previous assignments were grossly in error.
2. Experimental
The Fe(TPP)X samples were prepared by the previously reported methods (Adler
et al 1,970). 13 C NMR in natural abundance was recorded at 67- 89 MHz on
Br-uker FT NMR spectrometer, Deuterated chloroform solutions having solute
concentrations in 30->60mM were used. Since aggregation effect often compli-
cates the 13 G NMR results on metalloporphyrins, a study at different concentrations
was done to determine the optimum concentration range. Spectral width of about
50,000 Hz were employed. About 30,000-40,000 transients were collected. Pulses
of 15 /A sec width were applied at a rate of 0-4 sec repetition. The spectra were
r\A A
V
believed to be shifted to 1000-1400 ppm down-field and hence could not be
observed simultaneously along with other resonances. As in the proton NMR
study the two ortho and meta carbons of the phenyi ring show inequivalence with
respect to the iron atom which lies out of the mean porphyrm plane. Table 1
summarises the relevant 13 C shifts on the three complexes ; the data were corrected
for diamagnetic shifts using the corresponding values of ZnTPP (Wuthrich and
Baumann 1973a). The proton NMR shifts are included in table 1 for comparison.
The 13 C shifts consist of both dipolar and contact terms. The dipolar contri-
bution is given in axial symmetry by
H
3N
(1)
where (K^ K\\) is the paramagnetic anisotropy and f and r are the structural
parameters as defined by Horrocks (1973). Both these informations are available
on the Fe(TPP)X series (Behere et al 1982). The dipolar term can therefore be
easily calculated for pyrrole, meso and quaternary phenyi carbon atoms, but for
the other phenyi carbons slight complication is involved due to the rotation of
the phenyi ring in solution, which will affect the evaluation of r and G. Hence
we calculated r and at an interval of every 10 by rotating the phenyi ring
through 40 with respect to porphyrin plane and used an average value for the
calculation of the dipolar terms. The dipolar and contact terms so obtained are
included in table 1.
TPPPe
Figure 2. A typical l3 C spectrum of Fe(TPP)Br at room temperature. All the
carbon resonances are labelled.
00 VO VO T-H o^ ^ o
13
10 O\ O r^l r^, ts
O-,
f~: O t^ ^J CM _, T i
H-t
&
i + 1 1 1 1 1
S7
o<
b
t-. tp , <N ^ vo ^
(-1
iS?
fe
So
ro <~H
5
1 ! 1 1 1 I 1
8
a
K5
V) f} +^ "! O OO Oo
2
"eo
t-s t^ 00 <0 t^ O
^ s fc ^ 7 -
,g
1 + 1 1 1 1 1
'i
2
o o n fn vo
CS CN OO o
"u
1
01
O
? ?T 7 +
tl
i^ "^ f") C^\ CM ( VO
a
53
(SI ^, vb O^ ^| <^j ^H
53
Uj
O t>- JR Tj r-J r ( r 1
c/)
oT
oo
GO
1 ? 1 1 1 1 1
ha
9
o
<a
'O CN OO (N OO Tf O
*3
fc
P<
bD
So
7 7 i i i i i
0>
CJ
3
t- vo T-H Y-H o n vo
t^ tN CM *O i/j 5 ^j-
c^i QQ t"* V) (-4 cN *^
t?
2o
eo
7 + i i i i i
<u
4^
to
u
Wl
rf ^
1
o
' . 'T 1 ^ T 1
oW
* * " "^t" CO CM
s
CO
1 1 +
s
X
p7
Q> Tj- Q\ . _^ _.
c
fe
i?
U
ST
fc
So
7 ^ 7 7 i i 7
5
s
3
d
o
w
fi
<^ M r- <S VO <T) <ri
I
"r*
&
""^ V% C~4 C-l i I T i -lj
s
W
So
1 1 I 1 1 1 1
o
o
u
1 1 1 1 1 1 1
a
H
o> ^H ^ ON m co
cS
1H
4)
?S S ^ ^ S 2
H
1
T + 7 7 i i i
H
05
3
.1
_
3
^>
o
(H
s
<u
a
!3
'S* > ""^ '^
to
o S o "o *^ *s ^**
s
1 1 II 1 5 t
s o* 5 5 s S p2
s
2
4. Discussion
Table 1 shows several interesting results. As expected 13 C shifts are much larger
than the corresponding proton shifts. This is true not only for pyrrole but for
phenyl carbon shifts as well. In view of the large shifts, the dipolar contributions
which lie in the range of 1-40 ppm appear insignificant. This is an encouraging
result as it allows the neglect of dipolar contribution for the analysis of 13 C shifts.
This is contrary to the situation for proton NMR where the dipolar contribution
plays a significant role in the interpretation of the data (Mitra 1977 ; Behere
et al 1982).
The variation in the 13 C shifts across the Fe(TPP)X does not show any definite
trend, though the systematic decrease in the meso-carbon shift from chloride to
the iodide c omplex appears to be real. It is interesting that a very recent theoretical
calculation on five coordinated high spin iron(III) porphyrin predicts just a similar
variation in the meso-carbon shift (Mun et al 1981). Nevertheless the effect of
variation in the axial fclide on the 13 C shift of the basal porphyrin ring is small
and for most cases difficult to discern.
We shall now discuss the spin delocalisation mechanisms responsible for the
observed contact shifts of the various carbons and protons in this series. We
observe from table 1 that the meso-carbon shows a large down -field shift while
the phenyl quaternary carbon bonded to it is considerably up-field shifted. The
shift pattern of the phenyl ortho, meta and para carbons and protons is quite
interesting. The 13 C shifts for these nuclei show a sharp decrease in magnitude
in going from ortho to meta to para but the sign of the shift remains the same.
This is in contrast to the situation for the corresponding proton shifts which show
alternation in sign but no attenuation.
The ferric ion in the Fe(TPP)X series has unpaired electrons in all the five
d*orbitals of n and a symmetry. The unpaired electrons can therefore delocalise
over the porphyrin ring through its n and a molecular orbitals. It has recently
been shown that the meso-carbon shift arises mainly through the unpaired spin-
density in the n molecular orbitals (Mispelter et al 1981). This unpaired spin-
density in the Tr-MO can induce an unpaired spin density at the phenyl quaternary
carbon either through n-a correlation (Carrington and McLachlan 1967 ; LaMar
1973) or through direct p w -ff v interaction (LaMar 1973). The latter contri-
bution is expected to be small because of the orientation of the phenyl rings.
Nevertheless both these mechanisms will induce at the phenyl quaternary carbon
a spin density opposite in sign to that at the meso-carbon, as is experimentally
observed. The Unpaired spin density induced at the phenyl quaternary carbon
in both a and n MO propagates over the phenyl rings. The proton contact shifts
found earlier (Behere et al 1982) are consistent with the spin delocalisation in
predominantly rc-^MO of the phenyl rings. The 13 C contact shifts of the phenyl
carbons however show typical variation expected of <r-delocalisation resulting
from the unpaired spin density in the cr-atomic orbitals of these carbon atoms.
While these simple arguments explain qualitatively the unpaired spin density at
various carbon sites, a quantitative description may be much more complicated
Carrineton and McLachlan 1967).
experimental worK reponea in mis paper was aoae at
facility. The authors thank the staff members of the facility for their help.
References
Adler A D, LOngo F R, Kampas F and Kim J 1970 /. Inorg. Nucl. Chem. 32 2443
Behere D V, Date S K and Mitra S 1979 Chem. Phys. Lett. 68 544
Behere D V, Marathe V R and Mitra S 1977 /. Am. Chem. Soc. 99 4149
Behere D V and Mitra S 1979 Inorg. Chem. 18 1723
Behere D V and Mitra S 1980 Indian J. Chem. A19 505
Behere D V, Birdy R and Mitra S 1981 Inorg. Chem. 20 2786
Behere D V, Birdy R and Mitra S 1982 Inorg. Chem. 21
Carrington A and McLachlan A D 1967 Introduction to magnetic resonance (New York : Harper
and Row) p. 82, 94
Doddrell D M and Gregson A K 1974 Chem. Phys. Lett. 29 512
Goff H M 1978 Biochim. Biophys. Acta 57 336
Goff H M 1981 /. Am. Chem. Soc. 103 3714
Horrocks W Dew (Jr) 1973 In NMR of paramagnetic molecules (eds) G N LaMar, R H Holm
and W Dew Horrocks (New York : Academic Press) Chap. 4
LaMar G N 1973 In NMR of paramagnetic molecules (eds) G N LaMar, W Dew Horrocks and
R H Holm (New York : Academic Press) Chap. 3
LaMar G N and Walker F Ann 1979 in The porphyrins (ed) D H Dolphin (New York : Academic
Press) Vol. 4, p. 61
Mispelter J, Momenteau M and Ihoste J M 1979 /. Chem. Soc. Chem. Comm. p. 808
Mispelter J, Momenteau M and Ihoste J M 1981 /. Chem. Soc. Dalton p. 1729
Mitra S 1977 Progr. Inorg. Chem. 22 307
Mun S K, Mallick M K, Mishra S, Chang J C and Das T P 1981 /. Am. Chem. Soc. 103 5024
Wuthrich K and Baumann R I973a Ann. N. Y. Acad. Sci. 222 709
' Wuthrich K and Baumann R 1973b Helv. Chim. Acta 56 585
' Wuthrich K and Baumann R 1974 Helv. Chim. Acta 57 336
Study of mixed complexes by polarography : cadmium-glycine-
methionine and cadmium-glycine-ethylenediamine complexes
M RAMAIAHt, B G BHAT* and R SUNDARESAN^t
Chemistry Department, Indian Institute of Technology, Powai, Bombay 400076,
India
t Present address : Chemistry Department, Regional Engineering College,
Warangal 506 004, India
tt Analytical Chemistry Division, Bhabha Atomic Research Centre, Trombay,
Bombay 400 085, India
MS received 24 May 1980 ; revised 23 January 1982
Abstract. The mixed complexes of cadmium with glycine and methionine and
glycine and ethylenediamine have been studied by polarography and the stability
constants of the various species formed have been evaluated.
Keywords. Mixed complexes ; cadmium ; glycine ; methionine ; ethylenediamine ;
polarography.
1. Introduction
Though Schaap and Me Masters (1961) pioneered the extension of the polarographic
method of DeFord and Hume (1951) to the study of mixed complexes, not much
work has since been reported in the literature. This paper presents a polarographic
study of the mixed complexes of cadmium with glyciiie-methionine and glycine-
ethylenediamine.
2. Experimental
Glycine (E Merck, pro analysi) and DL-methionine (E Merck, LR) were used
without purification. Ethylenediamine (E Merck, LR) was standardised against
hydrochloric acid using methyl orange as the indicator. A stock solution of
cadmium was prepared from cadmium sulphate (E Merck, GR) and standardised
with EDTA. Potassium nitrate, used as the supporting electrolyte, was of BDH
AnalaR grade. The solutions were made in double distilled water and pH was
measured with a Philips pH-meter (pp->9040). Polarograms were taken on a
manual set-up using a H*cell with an agar plug and a saturated calomel electrode
(SCE) served as the reference electrode. Currents are reported after correcting
for the residual currents. Correction for the IR drop in the potentials was not
necessary.
* to whom correspondence should be made,
151
separately in 1-OM potassium nitrate at different pH. The electrode reaction
was reversible in both the cases. The glycinate [G] and the methioninate [Me]
concentrations were calculated from the pH of the solution and the pK a of the
ligands determined as 9 -76 for glycine and 9-10 for raethionine. The halfrwave
potentials were measured as a function of log [G] or log [Me] from which the
stability constants were calculated as log fa = 9-40 for cadmium-glycine and
log ^ = 3-80, iogj? 2 = 6 -35 and log /? 3 = 8-19 for cadmium-methionine systems
by the methods of Lingane (1941) and DeFord and Hume (1951) respectively
In the investigation of cadmlunvglycine-methionine system three series of
measurements were made under the same experimental conditions as for the
'simple' systems, keeping [Me] constant and varying [G]. The polarograms
were well defined and the reduction was reversible. The half-wave potentials
(table 1) were used to calculate a function, F 00 , using the relationship (Schaap
and McMasters 1961)
=antilog [0-4343 n
+ log {z*<) /*<)
(1)
where the symbols have the usual meaning. This may be written, at constant
methioninate concentration, as
(2)
where
F 00 (Me,G) = A 4- B[Q] + C[G] 2 + D [G] 3
A = [1+ /? M c lGo [Me] + /? Mca , o [Me] 2 4- J? MflsGo [Me]']
= [14- Ao IMe] 4- Ao [Me] 2 4- jffao [Me] 3 ],
B = [jS MeoGi 4- MeiGl [Me] 4- y? M02Gl [Me] 2 ]
=[#*4- jff a [Me] 4- #u [Me] J,
C = [j? M6oGa 4- J^c, [Me]]
-CjSoa + A. [Me]] and
re f er to t^ 6 stability constants of the mixed complex species CdMe.0,,^
The constants A, B, Cand D, evaluated by a graphical procedure, are reported in
table 1.
The values of A agreed with the calculated values based on the stability constants
obtained from the ' simple ' system, fa and /? 2 i were calculated from the values
of 3 using (3) as 1-0 x JO 6 and 1-26 x 10 9 respectively. Similarly fa z and $t a
were determined from C as 3-98 x 10 7 and 3-16 x 10 9 respectively. The average
value of D corresponds to fa and agreed with that obtained from cadmium
glycine system.
The relative stability of a mixed complex over the parent binary complex ancl
the compatibility between the ligands, indicated by the "mixing constant" k^
sec- 1 / 2 ; 1/a(|) = ~Q-58G VvsSCE; i, (|) 2'350A ! '*(.),, =2-00 M
[G] x 10 a
[M]
1/3 F QQ X 10"*
- FvsSCE
F JO x 10-
fa x:a-a
[M.--] =3-58 x 10- a M
1-42
0-701
1-25
1-74
0-66
2-84
0-705
1-69
2-43
2-75
5-67
0-710
2-49
2-62
1-71
9-93
0-717
4-25
3-27
1-63
14-19
0-725
7-84
4-82
2-24
21-29
0-733
13-94
6-08
2-08
28-37
0-738
21-24
7-13
1-93
42-57
0-743
31-15
7-08
1-28
56-74
0-759
106-1
18-53
2-97
yl=l-00 X
10*; 5 = 1-65 X 10 8
; C-l-X
>xlQ 8 ; -2-30. x 10 9 .
[Me] =7- 15 X 10-'
>M
1-42
0-725
7-74
5-25
..
2-84
0-727
9-03
7-14
3-32
5-67
0-729
11-36
7-68
2-61
9-93
0-735
17-31
10-38
4-21
14-19
0-738
20-97
9-84
2-57
21-29
0-745
37-25
14-21
3-76
28-37
0-749
50-60
15-37
3-23
42-57
0-759
108-9
23-93
4-16
56-74
0-764
159-7
26-91
3-65
A =7-00 x
10*; 5 = 6-20 X 10 6
; C = 2'7;
X1Q8; D =2-20 Xl0 9 .
[Me] = 0-107 M
1-42
0-739
22-93
10-07
..
2-84
0-739
22-93
5-04
. .
5-67
0-743
31-16
17-04
3-93
9-93
0-746
39-12
17-83
3-05
14-19
0-749
47-48
18-31
2-47
21-29
0-754
72-37
23-89
4-27
28-37
0-760
110-3
31-30
5'?2
42-57
0-765
168-1
34-44
4-61
56-74
0-772
287-4
46-86
5-65
/<=2-15xl0 5 ; 5 = 1-48x10'; (7 = 3-75 x 10 8 ; Z>=2-70xlO.
and the enhanced (or sometimes decreased) stability due to factors other than
statistical, given by the "stabilisation constant" k a are calculated from the
expressions (Marcus and Eliezer 1962)
and log k, = log k a - log (w!/ x ! yl). (5)
The values of k M and k t for this system, calculated in this manner, are given in
table 2.
3.2. Cadmium-'glycine-*ethylenediamine complexes
A preliminary investigation of cadmium-ethylenediamine complexes in 1-OM
potassium nitrate indicated reversible reduction. The dissociation constant, pK.
of ethylenediamine was determined as 10-10 and the stability constant /? 3 was
calculated from the half-wave potential data as 10 11 ' 98 . In the study of cadmium-
glycine-ethylenediamine system, three sets of data were obtained at three concen-
trations of glycine, varying that of ethylenediamine. The reduction was reversible
and the half-wave potential data (not presented here for the sake of brevity) were
solved for the stability constants of the various complex species as discussed
earlier. These values as well as log k m and log k t are given in table 2.
Table 2. Stability constants and log& M and log/c s values.
Complex
log (1
11-23
log k
a
log k s
a)Jl<U'ti3
Present
work
Literature
values
Present work
Cd (Me)
3-80
3-81
Cd(Me) 3
6-35
6-24
Cd(Me) a
8-19
8-32
. >
Cd(G)
4-54
Cd(G) a
7-60
8-08
Cd(G) 3
9-40
9-78
Cd(Me) G
6-00
-0-98 -1-28
Cd(Me) 2 G
9-10
-HO -51 -HO -03
Cd(Me) G 2
9-50
HO -SO -HO -02
Cd(En)
5-60
Cd(En} 3
10-63
Cd(En),
11-98
12-10
Cd(En) G
9-11
-0-01 -0-31
+0-97
+0-50
1.0
Figure 1. Distribution of cadmium as complexes.
Table 3. Complex equilibria and equilibrium constants.
Equilibria
logK
l-G^ Cd(Mo)(G)
Cd -!- 2 Mo -h G ^ Cd(Mo) 3 (G)
Cd + Me +-2G ** Cd(Mo) (G) s
Cd(Mc) (G) -i~ G ^ Cd(Me) (G) a
Cd (Mo) (G) -I- Me ^ Cd (Mo) 2 (G)
Cd(Mc) a (G) + G ^ Cd(Mo) (G), -I- Me
Cd(G) 2 + Me ^ Cd(Mc) (G) 2
Cd(Mo) -|- G ^ Cd(Mo) (G)
Cd(Me) a + G ^ Cd(Mc) (G) + Me
Cd(Me) 3 + G ^ Cd(Mc) 2 (G) + Me
Cd(Mc) (G) 8 -I- G ^ Cd(G) 3 -I- Me
Cd + En + G ^ Cd(En)(G)
Cd H-En -H2G ^ Cd(En)(G),
Cd+.2En-l-G^Cd(En) 3 (G>
Cd(En)(G) -I- En ^ Cd(En), (G)
Cd(En)(G) + G ^ Cd(En)(G) a
Cd(En) (G) a -I- Ep. ^ Cd(En) a (G) 4~ G
Cd(G 2 4- En ^ Gd(En) (G) + G
6 -00
9-10
9-50
3-50
3-10
Q-40
1-90
2-20
-0-35
0-91
-0-10
9-11
11-23
12-02
2-91
2-12
0-79
1-51
.. \ t/~i\ i n.\
It is seen from table 2 that log k M for the 1, 2 and 2, 1 complexes of both the
systems are positive indicating the compatibility between the ligands. Glycine
is more compatible with ethylenediamine probably because the chelating power
of ethylenediamine is more due to the presence of two nitrogen donors. The
co-ordination unsaturated 1, lis 'not important 'in both the systems as inferred
from the negative values of log k s and log k M . It is, therefore, apparent that the
species present in solution in. the concentration ranges studied are mostly 20, 30,
21, 12 and 03 complexes. The distribution of cadmium as these complex species
at [Me] = 0-05 M is depicted in figure 1 as a function of glycinate concentration
as an example. The equilibria between the different complex species ?re given
in table 3 from which the facility with which a ligand adds on to or substitutes
another ligand may be deduced.
References
DeFord D D and Hume D N 1951 /. Am. Client. Soc. 73 5321
Lingane J J 1941 Cheat. Rev. 29 1
Marcus Y and Eliezer E 1962 J. Phys. Chem. 66 1661
Scliaap W B and McMasters D L 1961 J. Am. Chem. Soc. 83 4699
In vitro antimicrobial-activity studies on the mixed ligand
complexes of Hg(II) with 8-hydroxyquinoline and
salicylic acids
Y ANJANEYULU 1 ", R PRABHAKAR RAO, R Y SWAMY,
A EKNATH* and K NARASIMHA RAO*
Department of Chemistry, Nagarjuna University, Nagarjunanagar 522 510, India
* Government Medical College, Guntur 522004, India
MS received 4 December 1981
Abstract. A series of mixed ligand complexes of Hg(II) with the general formula
Hg (OX) (SA) (where OX : S-hydroxyquinoline, SA : salicylic, 5-chloro-, 3,5-
dibromo, 3,5-diiodo, 3,5-dinitro, acetyl thiosalicylic acids) are isolated in pure
state and characterised by elemental analysis and infrared data. The low molar
conductance of the complexes in dimethylformamide indicates non-electrolyte nature.
The antimicrobial activity of these complexes against various bacteria and fungi
is studied which indicates that in several cases, the mixed ligand complexes possess
fairly highly antimicrobial activity than the binary mercury-oxinate. The lipo-
philic tendency of these complexes and its influence on the antimicrobial activity is
critically examined. A probable mechanism for the toxic action of these complexes
against various organisms is discussed.
Keywords. Antimicrobial activity ; mixed ligand complexes of Hg(II) ; 8-hydroxy-
quinoline and salicylic acids.
1. Introduction
Though. 8-hydroxyquinoline (oxine) and its divalent metal chelates are known
to possess fungicidal and bactericidal properties, the high cost of 8-hydroxyquino-
line limits their applicability. Albert et al (1953) explained the antimicrobial
activity of copper-oxinate assuming that the Zw-chelate due to its high liposolubility
penetrates the cell, reaches the site of action and there it undergoes dissociation
into 1 : 1 complex and free 8-hydroxyquinolme. The 1 : 1 charged complex
thus formed will become the toxic entity by combining with and blocking the
metal-binding sites on enzymes. The same mechanism may equally apply well
for explaining the antimicrobial activity of all divalent metal oxinates.
It has been observed that the cost factor can be minimised by replacing one
oxine molecule in the divalent metal oxinates with low cost fungicides like salicylic
hetero ligands. The results of our study on the mixed ligand complexes of Hg(II)
with 8-hydroxyquinoline and salicylic acids are presented in this paper.
2, Experimental
All the chemicals used are analytical grade (BDH) reagents.
2.1. General method for the preparation of the complex
Equimolar solutions of salicylic acid or substituted salicylic acids (0-2M),
8-hydroxyquinoline (0-2M) and Hg(II) acetate (0-2M) in 80% aqueous methanol
are mixed. After stirring for half an hour, the product is removed by filtration,
washed with several volumes of water and boiled in acetone and filtered. The
complexes are dried at 70 C for 12hr. Metal and nitrogen are estimated by
standard methods.
2.2. Physical measurements
Infrared spectra are recorded by using Perkin Elmer model 577 spectrophotometer
(4000 cm" 1 to 200 cnT 1 ) by KBr disc technique. The conductivity of the complexes
in DMF (10~ 3 M) is measured at 27 C by systronics conductivity bridge 305.
2-3. Antimicrobial activity
The antimicrobial activity of the compounds in dimethyl formamide (DMF) are
examined in vitro by serial dilution method (Schaub et al 1958) against various
bacteria and by paper disc method (Jasper etal 1958) against fungi. All the stock
cultures were supplied by the Department of Microbiology, All India Institute
of Medical Sciences, New Delhi, India. Peptone water and saline water is used
for making the inoculum for bacteria (18hr culture) and fungi respectively.
Nutrient broth and Saboround's dextrose agar (M/s. Hindustan Dehydrated Media,
Bombay) are used as test media for bacteria and fungi respectively. The minimum
inhibition concentration (MIC ywg/ml) of the compounds against bacteria and
average zone of inhibition (mm) of the compounds at 1000^g/ml against fungi
is given in tables 2 and 3. All the tests are carried out in duplicate.
3. Results and discussion
With 5d l configuration Hg(II) forms tetrahedral complexes using 6s 6p 3 hybrid
orbitals for bonding leaving a completely non-bonding shell (dyd%) which can
cause least perturbation to preferred stereochemistry. The elemental analyses
of these complexes (table 1) show that Hg(II) forms mixed ligand complexes
which can be represented as shown below (figure 1).
o
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W 'o
it?
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r-H CM
s
oo
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Q ~
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< 60
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1
s
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el
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SB
s
s
00 CM
Table 2. Antibacterial activity MIC (^g/ml) of mercury complexes with 8-hydroxy-.
quinoline and substituted salicylic acids at 37 G after 18 hours in nutrient broth.
SI.
No
Compound
Gram-positive
Gram-negative
.1
2
3
4
5
6
7
8
9
10
1.
Hg (OX)
3-1
50
25
50
50
50
25
10.0
12-5
25
2.
Hg (OX) (SA)
6-2
50
25
50
50
50
25
50
25
25
3.
Hg (OX) (CH5A)
12-5
50
3-1
12-5
6-2
100
12-5
100
12-5
50
4.
Hg (OX) UBr-SA)
50
12-5
12-5
12'5
25
100
25
50
25
25
5.
Hs (OX) (21-SA)
ioa
50
100
100
100
>too
>100
>100
>100
>100
6.
Ha(OX)(2NO a -SA)
100
100
25
12-5
25
100
25
100
25
50
7.
H S (OX) (Ace-SA)
6-2
12-5
12-5
3-1
6-2
25
12-5
50
12-5
12-5
8.
Hg (OX) (Thio-SA)
50
12-5
6-2
6-2
12-5
100
12-5
100
6-2
6-2
(1) Staphyloco'cciis albtis, (2) Staphylococcus aiireus, (3) Schigella schmitzi, (4) Pseudomonas
pyogenes, (5) Shigella sonnei, (6) Kle'osilla asrogenes, (7) Shigella flexneri, (8) Vibrio chokrae,
(9) Salmonella typhi, (10) Salmonella paratyphi-'B.
Table 3. Atitifungil activity of the mercury complexes with 8-hydroquinoline
and substituted Salicylic acids at 1000 ^g/ml after 48 hours at 30C.
SI.
No.
Compound
Zone of inhibitions
at 1000 ^
ig/ml
. in mm
Fungi 1 2
3
4
5
1.
Hg (OX),
10 9
9
1Q
8
2.
Hg (OX) (SA)
9 8
10
10
12
3.
Hg (OX) (Cl-SA)
8 10
9
9
10
4.
Hg (OX) (2Br~SA)
7 12
9
9
5.
Hg .(OX) (21-SA)
8 7
10
8
8
6,
Hg (OX) ( 2N0 2 -SA)
8 9
8
8
12
7.
Hg (OX) (Aca-SA)
8
9
11
8.
Hg (OX) (Thio-SA)
7
I tie low molar conductance values of these complexes in DMF indicate that
they are of non-electrolyte type.
3-1. IR data
The important absorption peaks of the IR spectra of these complexes agree with
their structure. In all the mixed ligand complexes the symmetric and asymmetric
vibrations of (O-C-O) group (Bellamy 1956) of salicylic acids are observed at
~ 1420 cm" 1 and ~ 1570 cm" 1 respectively, while the carbonyl stretching frequency
which appeared in the free salicylic acids between 1650-1 670 cm" 1 disappeared.
This clearly indicates metal-carboxylate linkage. Charles et al (1956) reported
that in several 8-hydroxyquinoline complexes of divalent metals, the v (c _ o) appeared
at ~ 1120 cm" 1 region and the position of the band slightly varies with the metal.
The v (c _ o) which appeared in the free oxine molecule at 1090 cm" 1 is found to be
shifted in all the mixed complexes giving a strong absorption band at 1110cm" 1
which clearly indicates the coordination of 8-hydroxyquinoline in the complexes.
In all the mixed ligand complexes, the observed sharp peaks between ~ 540-
560 cm"" 1 and ~ 340-400 cm" 1 may be assigned to the M-N and M-O stretching
frequencies respectively (Nakamato 1970).
3 2. Antimicrobial activity
In many cases the toxic effect of the mercury-oxine-salicylic or substituted salicylic
acid mixed ligand complexes against various bacteria and fungi is found to be
either equal or slightly greater when compared to the bis (8-hydroxyquinolinato)
mercury(II) complex. Salicylic acid or substituted salicylic acids and; their
mercury chelates are found to have measurable activity against these bacteria
and fungi at relatively very high concentrations (for bacteria > 1 00 /*g/ml, fungi'
> 2000 ^g/ml). This may be due to their higher water solubility. '..'...
In explaining the antimicrobial activity of bis (8-hydroxyquinolinato) copper(II),
Albert et al (1953) believed that the 1 : 2 chelate due to its liposolubility is neces-
sary to transport the toxic moiety, i.e., 1 : 1 chelate to the site of action. The
assumption was supported by the fact that antimicrobial activity of these complexes
was reversed in the presence of excess of copper. This may be due to the inability
of the ionically charged 1 : 1 chelate (which is produced in the presence of excess
of metal) to penetrate the cell membrane. Block (1955) proposed that the, natural
chelators within the cell were poisoned by removing copper from Cu(II)-oxine,
eiuu o-J4 i yuJ-UA i y4u-jiij,vjijijuc JLCJULUJV^ mt. ^UJJJJCJL cum JLUIJLU njjiu. -suj.ui-'J.c uucj.ai.coi.
Esposito and Fletcher (1961) proposed that the activity of copper(II)-8-hydroxy-
quinoline was due to the 1 : 1 complex which could bind with an enzyme site
that was involved in the biosynthesis of pteridines. This was based on the reversal
of inhibition by several pteridines and precursors. It was also believed that a
similar mechanism may be working well in explaining the toxic action of all
other bivalent oxinates.
According to Overton's concept of cell permeability the lipoid membrane
surrounding the cell favours the passiage through that membrane of lipid-soluble
materials and liposolubility is considered as one of the important factors which
control the antimicrobial activity of any toxic agent. The partition of the toxic
agent between olyl alcohol or chloroform and 7-4 pH phosphate buffer (pH of
the biological medium) system is considered as a good model to understand the
Hpophobic or lipophilic tendency (Dweyer and Mellor 1964). So we have deter-
mined distribution of all these complexes in between chloroform and 7-4pH
buffer and the results are given in table 4. As expected the mercury-oxine-sali-
cylic or substituted salicylic acid mixed complexes have lower partition coefficient
in chloroform when compared to the binary mercury oxinate. However, in
Table 4. Percentage of extraction of metal into chloroform at 7-4pH.
SI. Complex
No.
% of mercury
extracted into
chloroform
1.. Hg(OX) a
80
2. Hg(SA)(OX)
23
3'. Hg (Cl-SA) (OX)
28
4. Hg (2Br-SA) (OX)
76
5. Hg(2I-SA),(OX)
38
6. Hg (2NO,-SA) (OX)
34
7. Hg.(Ace~SA) (OX)
44
8. Hg (Thio-SA) (OX)
62
The extraction of mercury (II) with various salicylic acids into chloroform are found to be less
than 20% at 7-4 pH.
many cases the mixed complexes have equal or slightly more toxic effect against
various bacteria and fungi in comparison with binary complex. This indicates
that in the mixed complexes not only the 1 : 1 mercury-oxine complex is acting
as toxic agent but also the released salicylic acid may be playing an important
role in the antimicrobial activity through a different mechanism. The salicylic
acids or mercury salicylate chelates, though possess toxic effect, due to their higher
water solubility cannot go to the site of action as much as the mixed complexes
can penetrate. If in the mixed complexes also the 1:1 mercury-oxine is the
only toxic moiety then the antimicrobial activity of the mixed complexes should
increase with increasing pk : values of the salicylic acids and mercury-oxine-diiodo-
salicylic acid must have maximum activity. But no such relation is found to
exist from their antimicrobial activity screening studies (tables 2 and 3). It is also
believed that if the geometry and charge distribution around the molecules are
incompatible with geometry and charge distribution around the pores of the
fungal or bacterial cell wall, penetration through the wall by the toxic agent
cannot take place and toxic reactions within the spore do not occur. This may be
one of the reasons for certain mixed ligand complexes showing less effective
antimicrobial activity than the corresponding Hg (OX 2 ) complex.
Acknowledgements
The authors wish to thank Prof. L N Mohapatra for the supply of stock cultures
of bacteria and fungi. The authors also wish to thank Prof. Govardhan Mehata,
Dean, School of Chemistry, University of Hyderabad, for his help in getting the
IR spectra. The financial support given by the Council of Scientific and Industrial
Research, India, for carrying out this work is gratefully acknowledged.
References
Albert A, Gibson M I and Rubbo S D 1953 Sr. J. Exptl. Pathol 34 119
Bellamy L J 1956 The infrared spectra of complex molecules (London : MethUen) p. 279
Block S S 1955 /. Agrtc. Food Chem. 3 229
Charles R G, Freiser H, Friedel R, Hilliard L E and Johnston W D 1956 Spectrochim. Acta
8 1
Dweyer F P and Mellor DP 1964 Chelating agents and metal chelates (New York ;
Academic Press) p 407
Esposito R G and Fletcher A M 1961 Arch. Biochem. Biophys. 93 369
Jasper C, Maruzzella and Henry P A 1958 /. Am. Pharm. Assoc. 47 471
Nakamato K 1970 Infrared spectra of inorganic and coordination compounds (New York 5
Wiley Interscience)
Schaub I G, Foley M K, Scott E G and Bailey W R 1958 Diagnostic bacteriology (St. Louis q
The C V Mosby Company)
Zentmyer G A, Rich S and Horsfall J G 1960 Phytopathology 50 421
rnthesis and structural studies on Ni(II) chloride complexes of
,N'-(substituted) formamidino-N'-(substituted) carbamides and
iocarbamides
K L MADHOK
Centre for Rural Development and Appropriate Technology, Indian Institute of
Technology, Delhi Hauz Khas, New Delhi 110016, India
MS received 16 October 1981 ; revised 10 December 1981
Abstract. Nickel(II) chloride reacts with N,N'-djarylformamidina-N*-arylcarba-
mides, thiocarbamides and N-benzoylformamidino-N'-arylcarbamides, thiocarba-
mides forming complexes of the gsneral formula [Ni (Ar.NH.C (NH) NH.C.X.
NH.R) a ] Cl a (R = phenyl, orthotolyl and paratolyl ; Ar == benzoyl ; X = S, O)
and [Ni (R.NH.C.X.NH.C (N.Ph.).NH.Ph.)a] C1 3 (R= phenyl, orthotolyl,
paratolyl ; Ph = phenyl ; X = S, O). The ligands when oxidized with iodine
undergo ring closure to related 3,5-diarylamino 1,2,4-thiadiazolidines and 3,5-diaryl-
amino-l,2,4-diazolidines, while the complexes are not susceptible to oxidation.
This confirms the binding in complexes is through sulphur and oxygen of the ligands.
Keywords. Nickel(II) chloride complexes ; potentiometric oxidations ; 3,5-diaryl-
amtno-l,2,4-thiadiazolidines ; 3>5-diarylamino-l,2>4-diazQlidines,
Introduction
etal complexes of sulphur donor ligands have received great attention during
:ent years (Horsfall and Rich 1951) because of their versatile use as antifungal
d antibacterial agents. A survey of literature reveals that compounds containing
>th C=O and C=S groups possess significant fungicidal activity (Horsfall and
ch 1951). Despite the fact that a variety of sulphur donor ligands have been
idied (Akbar Ali and Livingstone 1974 ; Mishra 1980 ; Srivastava and Madhofc
78 '; Madhok and Srivastava 1,980), for the synthesis of metal complexes it
ipears that N,NXsubstituted) formamidino-N"-(substituted) carbamides
>SFSC) and N,N'-(substituted) forma mi din o-N'Xsubstitu ted) thiocarbamides
JSFSTC) having both C=O and C=S with =NH groups, has not been used,
continuation of our earlier work (Srivastava and Madhoto 1,978 ; Madhok and
ivastava 1980) on metal chelates of substituted thioureas, thiobiurets, the
esent paper describes the studies of Ni (II) chloride complexes of the title ligands.
experimental
The ligands N 5 N-dipnenylforrnamidmo N -phenylcarbarmde (DPFPC), N,N,
diphenylformamidino N"-phenylthiocarbamdie (DPFPTC), N,N'-diphenylforma
midino N"-orthotolylthiocarbamide (DPF.o.TTC), N,N'-diphenylformamidino
N"-paratolylthiocarbamide (DPF.^.TTC), N-benzoylforma midino N'-phenyl-
carbamide (BFPC), N-benzoylformamidino N'-phenylthiocarbamide (BFPTC),
N-benzoylformamidino N'-orthotolylthiocarbamide (BF.o.TTC) and N-benzoyl-
formamidino N'-paratolylthiocarbamide (BF.^.TTC) were prepared by the
method as described earlier (Srivastava and Madhofc 1978). The purity of ligands
was checked by sharp melting point and elemental analysis.
2.1. General method of the preparation of complexes
Standard alcoholic solution (0- 1 M) of nickel chloride (200 ml) was mixed together
with 200 ml of (0-2 M) alcoholic solution of the ligands and refluxed at 70 C
for about 2 hr. It was then allowed to cool during which a silver grey precipitate
was obtained which was analysed after being dried in vacuo. Melting points of
these complexes were determined in open capillary tubes on a unimelt tempe-
ra tiire apparatus and are uncorrected. In all the complexes, nickel was estimated
as dimethylglyoximatonickel(II). Sulphur and chlorine were estimated by
standard methods (Clarke 1960; Erdy 1965). The analytical results are
recorded in table 1.
Table 1. Analytical results of Ni(II) chloride complexes.
Compound
found
(calc.)
found
(calc.)
%s
found
(calc.)
Molar conductance
Molarity
(M)
Conductance
[Ni(DPFPTC) a ]Cl a
7-153
(7-134)
8-662
(8-642)
7-692
(7-789)
o-io xio- 2
150
[Ni(DPFPC),]Cl,
7-399
(7-432)
8-980
(8-991)
-
O'lO xlO" a
148
[Ni (DPF.o.TTC)J d a
6-955
(6-907)
8-420
(8-356)
7-555
(7-532)
0-12x10-?
152
[Ni (DPF.jj.TTC)J Cl,
6-955
(6-907)
8-486
(8-356)
7-700
(7-532)
0-11 xio-
158
[Ni(BFPTC),]Cl,
8-100
(8-089)
9-820
(9-786)
11-155
(11-100)
0-12 X lQ-
160
[Ni(BFPC) a ]Cl,
8-488
(6-463)
10-300
(10-230)
O'lO x 10-'
157
[NUBFATTCHiGl,
7-890
(7-787)
9-550
(9-421)
8-932
(8-892)
0-11 xio->
149
rvr: /TST? _ TTT \ 1
ine magnetic susceptibility 01 the chelates was determined by Gouy s magnetic
alance applying a field strength of about 4- 5 x 10 3 gauss. Mercury(II) tetra-
iiocyanateocobaltate(II) [Hg Co(CNS) 4 ] was used as the standard.
The infrared spectra of the ligands and complexes were recorded in KBr pallets
Q Perkin Elmer grating infrared spectrophotometer model 237-B in the range
f 4000-650 cm" 1 using the pallet technique. The spectra are complicated and
ifficult to interpret, however only those peaks that could be assigned with reason-
ble certainty are listed in table 3.
The solubility of all these complexes is high in dimethylformamide. Thus the
Dnductance measurements were carried out in freshly distilled dimethylformamide
>lution, on conductivity meter type LBR of Wissenschafflich Technische,. Werk-
latten, Germany, with dip type cell. The solutions of the complexes were prepared
mmediately before use.
Studies on oxidation of ligands and their complexes were carried out with
)dine solution in tetrahydrofuran using calomel and platinum electrodes. 20 ml
f Af/500 solution in THF of the ligands and their metal complexes were titrated
ith M/50 iodine solution in THF.
Absorption spectra of Ni(II) chloride complexes were measured by the standard
lethod using Perkin-Elmer UV-VIS spectrophotometer model 139. Ethanol and
lethanol used were of BDH AnalaR quality and distilled before use. The
bsorption bands of Ni complexes in ethanol and rnethanol are represented in
ible 2. Transition energy E T was calculated from the relation
_ 2-859 x 10 s
* ;U(inA) '
nd the oscillatory strength / was calculated from the following equation :
/ = 4-32 x W- l9 $Edv,
tilizing
# m Av = J Edv
rhere At> is the wrve number of the half band width. All the data are recorded
a table 2.
:. Results and discussion
m the complexes are coloured. Complexes are insoluble in most of the common
rganic solvents but are soluble in excess of alcohol, tetrahydrofuran and dimethyl-.
srmamide. All the complexes decomposed on heating above 160 C, complexes
re also decomposed by mineral acids.
Magnetic susceptibility of the complexes is found in the range of -0-285 to
-0-500 x 10~ 6 g. The negative susceptibility values are indicative of the
limagnetic nature of the complexes.
The observed values of molar conductance in DMF are in the range of 148-
60 mhos. The molar conductance results indicate the electrolytic nature of the
1
.a
.M
O Q O OOOO'-'OOO
S 0\ ^- OCMOQ\7SOQ <N ^o
^H CM oc r imi-^^}-'-2aQ'-<
OOOOQ
P.OOOC30
r~ot~>c^ (
oo rn o\ -i o\
O U
Table 3. IR spectral data of HgandS and their Ni(II) chloride chelates (in cm" 1 ).
=NH C=O N-H C-H C-N Stretch (C=S) C=S
Compound stretch stretch bend stretch 4-N-H stretch Stretch
bend 4-
(C=S)bend
PFPTC
3385s
1641yw
1540s
144lvw
1225m
729m
:Clo.2DPFPTC
3280mb
1635s
1585S
1485s
1225m
718m
PFPC
3400s '
1735\v
1440m
1582m
1660m
1260m
9
iC! 2 .2DPFPC
1720w
1630m
1470w
1422m
1222m
* *
PF.o.TTC
3402s
1660m
1540m
1430s
1 270m
760b
;Cl a .2DPF.o.TTC
1648s
1552m
1405s
1258m
745m
PF.^.TTC
3400s
1600s
1566m
1441m
1220m
756m
iCla^DPF.p-TTC
1625W
1575m
146Qw
1205b
740m
'PTC
3400s
1700m
1680s
1550s
1413s
1227s
71 6w
C1 3 .2BFPTC
3310b
1700m
1630s
1595m
1480w
1215m
702m
'PC
3360s
1725m
1685s
1550m
1440s
I280w
. .
iCJ 2 .2BFPC
1710m
1635vw
1590m
1440ms
1255ms
. .
7.0.TTC
3400ms
1698m
1625sb
1525w
1400m
I775w .
750m
C1 2 .2BF.0.TTC
3300m
1698m
1605m
I520w
1475w
HSOws
730w
7./J.TTC
3400s
1700m
1625s
1563s
1440m
1282s
799w
C1 2 .2BF./>.TTC
3270m
1700m
1620m
1570m
1445\vb
1266m
778w
= Strong, m = medium, b = broad, w = weak.
On an examination of UV spectra of the ligands and complexes in alcohol, it is
)served that the absorption band of the ligands (DPFPTC), (DPF.o.TTC, DPF.p.
1C, BFPTC, BF.o.TTC and BF.p.TTC), 240-265 nm in ethanol and methanol
is been shifted to 305 to 347 nm in complexes. This shift is attributed to the
ctthat during the complex formation n* energy level is longer due to stabili-
.tion of the excited state, so the n-n* transition is shifted to lower wavelength,
5., lower frequency and consequently lower energy. This also accounts for the
ct that thiocarbonyl group is acting as donor in the complex formation. In
tse of DPFPC and BFPC complexes the rc-rc* transitions (247-265 nm) band
shifted to lower wavelength 307 to 347 nm.
From the infrared spectra of the ligands and their metal chelates, it can be seen
.at the ligands exhibit a C=O stretching band of medium intensity in the region
198-1 1735cm.- 1 which is observed to be stronger than the usual ketone C = O
ind (Schcinmann 1970). On dictation with metal the carbonyl absorption
;ak is shifted to lower frequency (IScrrr 1 ) of comparatively low intensity. As
.ere is almost no change in the benzoylic carbonyl frequencies on complexation,
e benzoylic C =O group cannot be considered as a site for coordination. A peak
; 1400- 1490cm- 1 is due to mixed bend of (C-N) stretch, N-H bend and C = S
md. The strong bands at 3355 to 3400cm.- 1 in the case of di substituted forma-
Potentiometric titrations of the ligands indicate that an equal amount 01 iodine
is consumed in the oxidation reaction. It is due to the oxidation of ligands to
thiazoles. The interaction of N,N'-diphenyl guanidines or asymmetric guani dines
with arylisocyanates and arylisothiocyanates have been shown to afford N,N'-
diarylforrnarnldino N"-arylcarbamides and thiocarbamides(I)andN-benzoylforma-
midino N'-arylcarbamides and thiocarbamides(II). The compounds I and II
when oxidized undergo ring closure to III and IV, the related 3,5-diarylamino-
1,2,4-diiazolidines and 3,5-diarylamino-l,2,4*thiadiazolidines (Dixit 1961).
/ NH V HN C=N' ph
RNHCT s O=N ( Ph oxidation I i
|| | ^RN=C NH + Phi
X NH Ph . X
I III
N,N'-diarylformamidmo-N''-arylcarbamidcs, 3,5-diarylamino-l ,2,4-diazoli dines and
thiocarbamides thiadiazolidines.
Ph = phenyl ; R = phenyl, orthotolyl, paratolyl ; X= O, S
HN C=N'Ar
*NHC X C NH Ar oxidotion RN = C NH -f- HI
X NH X
It - IV
N-benzoylformamidino N'-arylcarbamides, 3,5-diarylamino-l,2,4-diazolidiites and
thiocarbamides. thiadiazolidines.
R = C 6 H 6 CO ; Ar = phenyl, orthotolyl, paratolyl ; X = O, S
On titrating the metal complexes with iodine almost constant values of potential
are obtained indicating that the complexes are not being oxidized by iodine.
This may be due to the fact that the sulphur and oxygen atoms are already bonded
to the m?tal in the metal chloride complexes and are not free to form the thiols.
All these observations show that coordination in the case of DPFPTC, DPF.o.
TTC, DPF.p.TTC, BFPTC, BF.o.TTC and BF.p.TTC is through sulphur and
nitrogen (of the =NH group) while in the case of BFPC and t)PFPC complexes
the coordination is through oxygen and nitrogen atoms.
Acknowledgements
The author is thankful to Dr K P Srivastava, Professor of Chemistry,, Birla
Institute of Technology and Science, Pilani, for providing necessary facilities and
to the CSIR, Hew Delhi, for financial assistance.
cbar AH JVL ana Livingstone a Ji iyi<* uoora. ^nem. Kev. ij iui
arke H T 1960 A handbook of organic analysis (London : Advard Arnold) p. 311
xit S N 1961 /. Indian Chem. Soc. 38 221
dey L 1965 Gravimetric analysis part II (London : Pergamon) p. 532
jrsfall J G and Rich S 1951 Contrib. Boyce Thompson Inst. 16 361
adhok K L and Srivastava K P 1980 Indian J. Chem. A19 808
ishra H P 1980 Ph.D. thesis (Sambalpur University) p. 6
cholson A R and Button G J 1969 Aust. J. Chem. 22 1543
ivastava K P and Madhok K L I978a Indian J. Chem. A16 359, 990
ivastava K P and Madhok K L 1978b /. Chem. Eng. Data 23 256
ivastava K P, Srivastava G P, Arya S K and Madhok K L 1980 /. Chem. Eng. Data 25
173
pnthesis and characterization of copolymers from 4-haIo(chloro,
omo) salicylic acid
HASMUKH S PATEL and SHANTI R PATEL
Department of Chemistry, Sardar Patel University, Vallabh Vidyanagar 388 120,
India
MS received 29 May 1981; revised 2 January 1982
Abstract. Copolymers have been prepared by condensing a mixture of either
4-chloro Or 4-bromosalicylic acid and any one of the comonomer like salicylic acid,
^-hydroxybenzoic acid, ^-aminosalicylic acid, ^-aminobenzoic acid, jj-cresol and
^-halo(chloro, bromo)phenol with formaldehyde in the presence of 5M H a SC>4.
Copolymer composition of each of the copolymer has been estimated on the basis
of halogen content and/or on the basis of results of non-aqueous titrations of the
copolymer against standard sodium methoxide and/or tetra-n-butylammonium
hydroxide. The IR spectral characteristics of copolymers have been noted. The
viscometric and thermal studies of copolymers have also been carried out.
Keywords. Copolymers from 4-halo(chloro, bromo)salicylic acid.
Introduction
^polymers from a mixture of salicylic acid and a phenolic or amino derivative
.cTTormaidehyde are reported to find many interesting applications (Rozhanov-
aya and Zasova 1969 ; Isura Shigeru et al 1971 ; Kimura and Kashiha 1972 ;
asyuk and Natanson 1976). Some of the copolymers synthesized from a mixture
resorcinol, formaldehyde and either salicylic acid (Ward et al 1975), /?-resorcylic
id (DeGeiso et al 1963) or ^-aminosalicylic acid (Dyaltova et al 1964) have
oved to be specific ion exchangers. Copolymers from salicylic acid, urea and
mialdehyde are reported to yield soluble polyelectrolytes (Makhmudov et al
71). In the light of the reports about the promising applications of these
polymers, it was thought that copolymers from a mixture of 4-halo(chloro,
omo)salicylic acid and a phenol or an amine and formaldehyde may possess
teresting properties. Hence the work described in the present paper dealing
th the synthesis and characterization of such copolymers was undertaken. Out
the two comonomers condensed jointly with formaldehyde one is either 4-chloro
!$) or 4-bromosalicylic acid (BS) and the other comonomer is either salicylic
id (SA), p-hydroxybenzoic acid (PHBA), ^-aminosalicyh'c acid (PAS), p-amino-
ST\/-\\ _ -i-i__ __ i-_ ! /n/-<'n\ .,. _ t __ __ .1 ____ i
The spectral characteristics of the copolymers have been noted. The viscosity
of the copolymers have been measured in DMF and in 80 : 20 (v/v) DMF: water
containing 1-0% KBr. Solutions in DMF exhibited polyelectrolyte behaviour
in viscometry experiments. The thermal properties of the copolymers have been
examined by thermogravimetry.
2. Experimental
Copolymerization of a mixture of 4-chloro (CS) or 4-bromo (BS) salicylic acid
and a comonomer [salicylic acid (SA), ^-hydroxybenzoic acid (PHBA), /?-amino-
salicylic acid (PAS), j7-amhio benzole acid (PABA), /?-cresol (PC), j9-chlorophenol
(PCP) or p-bromophenol (PBP)] with formaldehyde (F) was carried out in presence
of an acid catalyst. In all the syntheses of the copolymers, the molar proportions
of CS (or BS), the comonomer and formaldehyde were 1:1:2 and 5 M H 2 SO 4 was
used as a catalyst. A typical copolymer synthesis is described here. Other co-
polymer samples listed in tables 1 and 2 were prepared similarly. However, each
copolymer sample was treated typically to free it from the polymer formed by
separate condensation of each of the two comonomers with formaldehyde. This
treatment is based on the diiference in the solubility behaviour of the copolymer
and the corresponding homopolymers. The solubility behaviour is summarised
in table 1.
Table 1. Solubilities of polymers and copolymers.
Polymer
PAS-F"
PABA-F"
PC-F d
PCP-F'
Solubility in"
Copolymer
aq. Solvent Acetone 1,4- DMF Pyridine
NaHCO 3 ether dioxane
CS (BS)-F D
+ -44-
+ +
SA-F'
. i . _j_
T* T^
+ +
PHBA-F"
+ +
+ +
CS (BS).SA.F
-J-.
+ +
CS (BS).PHBA.F
+
+ +
CS (BS).PAS.F
CS (BS).PABA.F
CS (BS).PC.F
CS.PCP.F
BS.PBP.F
+
+
+
+
+
+
+
+
+
+
+
+
+
* W 8
U
O W
*n *^
\ooooo<~'O CN ' l n^o
*""" r'
CO CO CO "^J" Tl" "*^* ""vf" "^J"
V 'O
[ST
-a O
o .0 o o
>n Vi ro c'l
S OH
-H CM CO *
*0 ' P3 ^
t~-t-~ oooooooooooo
(4-1
o 5 I,
OJ H +3 II
O
^jj j? c8
8 S -
O '43 aT
T: .2 w
T3 ft ^
&;5
*-* pu
o "^
"C PQ
H
S 5 *-
13 "eO O .O <tt
" "ce tfl ^ ffl | ^
S < +-> ^ g O
s ^ *^
o*nv*i *noC>>no
cooo'^-o\t-^tNo i n
vo */*j \o *n *o *n vo i n
2 C3 w
O pj
o ,? 'S
fl 13
SOW)
Q (D *M .*
13 ft eg
*^J "f"i ^> s?
rt o
O
"3 " iS i ' e S ^
CD *n C5 *^ c^ *n in G?
ON oo oo i~~ ^ ^ c^ r^
U
a to *.. s S o
a,
a % * ^
1
a
tt
.2
CO Co co CO; co *n in C5
cB
' rt
.'.' S3
in co t*^ "^ ^o oo ON DO
^<*rt''TcocMcvies
o
S
S 1 *0 4->
&5 a
OC3OOOOOO
TJ
9 r2 rt
m
o ^ _
Sa
u O
8
noocoo^om
^(nvo'nfNcMocM
"o
rt
5
a*
U
OOoCO'ooOO
$ s
Xfl *
o
fJ V0
U
rf
1 ^
M
co co i/% in cs
II 1 (
c;
(S^r-. o r~ oo I s * o
^ M
So
vo ^ i *n f^ r^ t*- o o
T 1 r| t 1 CS
81
1
s 1
&
s
I
oo ^h co ON
g ^
*3 *tz*
Tt* ^* ^3" CO
^^ TO
1
Q J~
55
S.O >3
oN f&
CM P
c8 ~>T
O -JS-
rt
fe
LLt tLi IY. rr.
1 3 ^2 r*^ r~i
'o
ft
. ., <! i <Cftft'<J'<l
Pi fe . ft ft w w pq ffl
<-<fr;!il-<<;-<:-<
lg
r o
S
O
OJ O5 OJ OJ t/3 to Ol CO
e ibi
umufflumow
sample* uo.o^i.^.
A mixture of CS (3-45g, 0-02 mole), SA (2-76 g, 0-02 mole), 37% formaldehyde
solution (3 -45 ml, 0-04 mole) and 40ml 5M H 2 S04 was refluxed with good
stirring at 130 C for 6 hr. During this time, the solid product separated out. It
was filtered and washed with hot water. The air-dried polymer sample was
Soxhlet-extracted with benzene to remove unreacted monomers. The copolymer
sample was dissolved in ethanol and reprecipitated as a pasty mass by adding
distilled water. The pasty mass was washed by decantation and collected. It
was dissolved in dilute alkali and reprecipitated by gradual addition of dilute HC1
with stirring. The solid was allowed to settle and filtered by decantation. It
was washed with water and allowed to dry in air.
The dried sample was powdered and stirred with acetone (30 ml) for about half
an hour. The solid was filtered, washed and retreated in the same manner with
acetone. This was repeated till the soluble portion was completely removed. The
residue was copolymer. It was soluble in 1,4-dioxane, THF and DMF. It did
not melt up to 360 C, yield 2-5g.
2. la. The copolymer samples labelled as BS.SA.F, CS.PHBA.F and BS.PHBA.F
were prepared and treated in the manner described above.
2- Ib. Treatment ofcopolymers designated as CS.PAS.F, BS.PAS.F, CS.PABA.F and
BS.PABA.F: Each of the copolymer samples CS.PAS.F, BS.PAS.F, CS.PABA.F
and BS.PABA.F prepared by the method described above was washed repeatedly
with acetone. The residue was extracted with DMF leaving a small amount of
insoluble material. The latter was rejected. The copolymer was precipitated
from the DMF extract by diluting it with water. The solid, obtained on filtration
and washing, was dried and treated again with acetone. The residue was the
required copojymer. The yields are: CS.PAS.F, 3-5 g; BS.PAS.F, 3-2g;
CS.PABA.F, 3-Og and BS.PABA.F, 3-Og.
2 1C; Treatment of copolymer samples designated as CS.PC.F and BS.PC.F : The
copolymer samples CS r PC.F and BS.PC.F were treated repeatedly with 2% aq.
sodium hydrogen carbonate solution (50ml) and then washed with 1,4-dioxane
(20 ml). The residue was dissolved in DMF, filtered and diluted with water to
precipitate the copolymer. The yields of copolymers are : CS.PC.F, 3 2 g,
BS.PC.F, 3-Og.
2 Id. Treatment of copolymer samples designated as CS.PCP.F and BS.PBP.F :
Each of the copolymer samples CS.PCP.F and BS.PBP.F was treated with 2% aq.
sodium hydrogen carbonate solution (50ml) and then dissolved in solvent ether
(30ml) and filtered. The residue was the homopolymer of j7-chloro or bromo-
phenol. The solution of the copolymer in ether was evaporated to dryness.
The solid was collected. It was soluble in ethanol, 1,4-dioxane, THF and DMF.
The yields of copolymers are; CSJPCP.F, 3-Og; BS.PBP.F, 2-5g.
.uaiogen content in ail copoiymer samples was estimated by me uanus metttoa.
In the case of nitrogen containing copolymers, nitrogen content was estimated by
Duma's method. The results are presented in table 2.
3-2. IR spectra of all copoiymer samples were taken in KBr on Beckmann IR-5
spectrophotometer.
3.3. Number average molecular weights (Mri) of soluble copolymers were estimated
by vapour pressure osmometry (VPO) in 1,4-dioxane at 51 C and also by non-
aqueous conductometric titration.
Conductometric titration of each of the copoiymer samples was carried out in
pyridine against standard tetra-K-butylammonium hydroxide (TBAH). In case
of copolymers containing phenolic OH groups conductometric titration was also
carried out in pyridine against standard sodium methoxide (NaOMe). The
details of the procedure are reported in the literature (Chatterjee and Gupta 1971,
1974, 1977). The results are presented in tables 2 and 3.
3-4. Viscometric measurements of the solutions of the copoiymer samples were
carried out using Ubbelohde viscometer. Viscosity of all copoiymer samples was
measured in DMF and 80 : 20 (v/v) DMF/water containing 1 % KBr. Intrinsic
viscosities of all the copolymers are presented in tables 2 and 3.
3.5. Thermogravimetry of all the copoiymer samples was carried out on "Du
Pont-950 thermogravimetric analyzer " in CO 2 atmosphere at a heating rate of
10 C per min.
4. Results and discussion
The polymer samples, salicylic acid -formaldehyde (DeGeiso et al 1962), ^-amino-
salicylic acid-formaldehyde (Patel et al 1981), p-hydroxybenzoic acid-formaldehyde,
p-aminobenzoic acid-formaldehyde and ^-cresol-formaldehyde (Chatterjee 1970,
1971), polymers are all reported to have linear structures. As both the
comonomers employed in the present copoiymer synthesis are bifunctional, these
copolymers would be linear. The distribution of the comonomer units along
the polymer chains would be random. On the basis of this the following structure
can be assigned to the copolymers.
COOH
CH*>
Ci/Br
<2 H 2
I : R COOH or CH 3 or Cl or Br~, R! = H and R 2 = OH or NH 3
II ; R = COOH, Ri = OH and R ? == H or|NH g .
el
H
L
1
X
?>**
IT
^Ir3
00 ^ tS (N
"cS
h_J
o
a
c
a ^
^i Jrt
f j V
9
M o oo vc vo g
O * '
. . ^*
SO
"cS ^
r~ >o t-i o vo oo
cS
B 8
11 CS CS fO CO
I
J> KS o o a
1
?rt O O o O
C
i, gn
Jj ft ^ Jj
o
<O Q
i,-^
O O O O O rt
w
*<*5 O
'o
It
rt m /^ m o o
'' O N OO t~~ r-i
ri
ri
fl
*N CN i 1 T | cf} ^
03
o b o o 6 o
,0
1K>
c3
o
'i*
M
13
>
S3 T-I
T3
tO
u
a
S
-^
p,
Q ,(* "-3
^
S
t||
A
i- t^ rt 10
II
*lll/
5S ff o o
O\ O\ <N o
O S.o
g oX
CJ *^
^ji S ^zJ tM fl3 ^j ^ ^*^
"
<O T ~ l
g S
*2j ^
M
SH
^ "fl
o" K.
II -3-
| S g
^|
g g |
*s s a 4
J "
^ fe
OH c H
S
1 Is
a _ "
3 1 1 ^
"^f" V% Vj r*i
OO * vr> oo
^> VO ff>
|l
a "r*
55 g S M_ S
VO V"> 00 O
^3 t^ ^O t (
A O
5 -SO
t-< to
c^ ^H fn co
CM
en ys a g JS g
O
'i!
wllllT 5
^'lj -i a
g I> A O
00 00 ON (^
f
+3
U
<*i, * H
'
a o
83 a,
o P<
a
S "E
S o
"O c3
ft - * S B
8 S s
B g B S
*From com
copolyraer
spectra of the copolymer samples.
Examination of the IR spectra shown in figures 1 and 2 reveal that they comprised
a broad absorption band from 3500 to 2600 cm" 1 due to OH and/or NH stretch-
ing vibrations. Such bands comprised distinct ' inflections at 3030, 2930, and
2850 cm' 1 and are expected due to aromatic C-H and bridge-methylene C-H
stretching. Separate carbonyl bands corresponding to two different C=O could
not be observed in the spectra of copolymers obtained from comonomer acids.
Only a slightly broad band covering expected vicinal positions is observed. In
some spectra at least one inflection is observed in this broad band. Hence the
average position of v c=0 in the spectrum of each copolymer is noted in tables 2 and 3.
There is no clear indication of the presence of bands due to - CH 2 O CH 2
in the IR spectrum. Because of the relatively small proportion of this type of
bridge and because of poorly resolved IR spectrum it is likely that the presence of
the band due to CH 2 O CH 2 grouping may not have been observed. Hence
no definite conclusion can be drawn about the presence or absence of such bridges
in the polymer chains.
The average degree of polymerization (DP) of copolymers estimated by non-
aqueous conductometric titration is shown in tables 2 and 3. The values varied
40OO 250O 2000 1500 1300 1100 1000 9OO
f Wa ven umber (cm" 1 )
800.
180
Hasmukh S Patel and Shanti R Patel
4000 2500 2000
1500 1200
Wavenumber (crrf 1
1000 900 800
Figure 2. IR spectra of copolymer samples : 1. CS.PABA.F, 2. BS.PABA.F,
3. CS.PC.F, 4. BS.PC.F, 5. CS.PCP.F, 6. BS.PBP.F.
from 5 to 8 and are nearly consistent with the number average molecular weight
(Mri) of dioxane soluble copolymers estimated by VPO method. These lower
values are expected as all the comonomers employed in the copolymerization
reaction except /?-cresol are much less reactive than phenol in phenol-formaldehyde
reaction. The lowering of reactivity in the electrophilic substitution reaction is
due to the presence of deactivating COOH and/or halogen.
On the basis of the percentage of Cl or Br in the copolymer sample it is possible
to calculate copolymer composition of a copolymer if the other comonomer does
not contain halogen. On the basis of percentage of Cl the total number of moles
of 4-chlorosalicylic acid repeating unit (x) in 100 g of copolymer is given by
x = percentage of Cl -r 35- 5.
fhr
nf 4-rhlr>rnfia1iryv1ir. ar.irl
unit t
/. Copolymer composition will be given by the molar ratio
4-Chlorosalicylic acid repeating unit : Comonomer repeating unit
=x :y
In the case of copolymer obtained from nitrogen containing comonomer the
composition based on halogen content of the copolymer can be used to calculate
the expected N %. These values along with those of estimated N % are shown in
table 2.
In the case of copolymers prepared from Cl or Br containing comonomers
the data obtained from titration were employed to calculate copolymer composition.
The titration against standard sodium methoxide in pyridine furnished information
about total number of moles of both COOH and phenolic OH groups. The
titration against standard TBAH would furnish information about only COOH
groups (Chatterjee 1970). It can be shown that if the amount of 1 M TBAH
required for 1 g of copolymer is equal to a ml and amount of 1 M NaOMe
required for 1 g of copolymer is equal to b ml, the ratio a : (b Id) = 4-chloro-
salicylic acid repeating unit : phenolic comonomer repeating unit.
The copolymer composition of copolymer containing />-cresol comouomer has
also been estimated by a similar method. It was found that the copolymer compo-
sition estimated by this method compared well with the copolymer composition
calculated on the basis of halogen content. The results are shown in table 3.
It is worth noting that the copolymer composition is highly dependent on the
accuracy of the measured value of appropriate characteristic of one of the two
comonomers. In the present cases there would be uncertainty in the measured
property on the basis of which the copolymer composition is calculated. Hence
information about the copolymer composition will not be that accurate as to
permit a further detailed analysis. The predicted relative reactivity of 4-halo-
(chloro, bromo) salicylic acid and the other comonomer in normal to homo-
polymerization is retained even in its copolymerjzation reaction. This order can
be predicted on the basis of the structure of the comonomer. A phenol con-
taining electron attracting groups shows lower reactivity in its reaction with formal-
dehyde. This is found to be true even in copolymerization reaction.
It was observed that all the copolymer samples showed polyelectrolyte behaviour
in their solutions in DMF. This type of behaviour of a typical copolymer sample,
CS.SA.F, was suppressed by increasing polarity of the solvent by using solvent/
water mixture or by adding electrolyte to the solvent. By carrying out some
experiments using various DMF/water mixtures as solvents and using different
.amounts of KBr, it was found that polyelectrolyte behaviour is suppressed in
80 : 20 (v/v) DMF/water containing 1-0% KBr by weight. Hence the viscometric
study of all copolymer samples was carried out in 80 : 20 (v/v) DMF/water
containing 1 % KBr by weight. From the plots of reduced viscosity vs. concen-
tration the values of intrinsic viscosities of all copolymers were estimated. These
values are shown in tables 2 and 3.
Table 4. Thcrmogravimetric analysis of copolymors.
Percentage weight-loss at temperature C T Temp. Weight-loss at first
Copolymer range stage degradation %
samples 100 200 300 400 500 600 700 for first
step Calcd. Observed
decom.
GS.SA.F
4
6
18
34
60
94
250-330
27-
50
24
BS.SA.F
4
6
19
42
68
96
250-330
25-
52
23
CS.PHBA.F
4
7
18
36
68
98
250-320
27-
72
21
BS.PHBA.F
2
8
20
49
67
100
25.10-320
25-
52
21
CS.PAS.F
6
9
26
36
74
100
250-300
25-
52
26
BS.PAS.F.
5
7
20
44
70
96
250-300
22-
66
20
CS.PABA.F
7
10
28
48
78
94
250-300
26-20
28
BS.PABA.F
4
6
20
43
72
98
250-300
2.3-
32
20
CS.PC.F
1
10
16
23
40
100
250-300
9-
02
10
BS.PC.F
5
10
18
32
68
98
250-300
8-
60
10
CS.PCP.F
1
4
17
21
31
42
52 250-300
16-
20
17
BS.PBP.F
4
16
26
46
57
64 250-300
13-
65
16
Examination of TG analysis results reported in table 4 reveals that each copolyrner
sample undergoes degradation in two steps. It is reported that salicylic acid-
formaldehyde (SA-F) (DeGeiso et al 1962), p-aminosalicylic acid-formaldehyde
(PAS-F) (Patel et al 1981) and p-chloro (or bromo) salicylic acid-formaldehyde
(CS-F or BS-F) (Patel and Patel 1982) polymers undergo degradation in two
steps. The first step in degradation of the above-mentioned SA-F type polymers
and which appears from 200-300 C, depending upon the nature of the polymer, is
attributed to the decarboxylation of the "salicylic acid units" present in the
polymer chain. It has been observed that jp-hydroxybenzoic acid-formaldehyde
and /7-arninobenzoic acid-formaldehyde polymers (Patel 1977) degrades in a single
step when heated in air at a controlled rate.
The phenolic resins are reported to undergo one step random degradation when
heated in air affording low molecular weight compounds (Jackson and Conley
1964). On the basis of these reported observations it is considered that the first
step in the thermal degradation of the copolymers reported in table 4 may be due to
decarboxylation and further degradation in the second step may be a random
degradation reaction affording simpler degradation products.
On the basis of the copolyrner composition of all the copolymers the possible
% weight-loss due to decarboxylation in the first stage of degradation is calculated
for all the copolyrner samples. The values of the calculated % weight-loss are
shown in table 4. The calculated values are found to be comparable with the
observed values of % weight-loss at the end of the first step of degradation. This
r-n-nrme*r
that fhfi first Rtfin in Hp!0rarlfltirm nf the* rrmnlx/tnArc tc
ana men relative proportion u it> UOL pus&ime iu uiaw ttuy
G Hcrusion about the mechanism of degradation reaction of the copolymers.
W J, Cravan W E, Rosenthal A, Ruetman S H, Stephens C W and Weatherbee C
1956 J. Polym. Sci. 20 75
C tiatterjee S K 1970 /. Polym. Sci. Part A-l 8 1299
S K 1971 /. Polym. Sci. Part A-l 9 3225
S K and Gupta U D 1971 J. Polym. Sci. Part A-l
S K and Gupta M D 1974 /. Macro. Sci. Chem. A8 291
S K and Gupta M D 1977 /. Macro. Sci. Chem. All 869
R C, Donaruma L G and Toraic E A 1962 J. Org. Chem. 27 1424
3^>eGeiso R C, Donaruma L G and Tomic E A 1963 Ind. Engg. Chem. Process Des. Dev. 2 43
M M, Temkina V Ya, Kllesnik E S, Yaroshenko G F and Lastovskii R P 1963
2V. Vses. Nauchno Issled, Inst. Khim. Reakt. No. 25 91 ; 1964 Chem. Abstr. 61 777
Shigoru, Inoue and Shigemastu Michiru 1971 Chem. Abstr. 75 505367
Jackson W M and Conley R T 1964 /. Appl Polym. Sci. 8 2873
Shuji and Kashtha Musabiro 1972 Chem. Abstr. 77 76561
T M, Abdullaev S and Akhmedvo K S 1970 Akad. Nauk. Uzb. USSR 27
34 ; 1971 Chem. Abstr. 74 41571a
I*atel H S and Patel S R 1982 /. Macro. Sci. Chem. A17 1383
I*atel R. M 1977 Studies on synthetic resins Ph.D. Thesis, S P University, Vallabh Vidyanagar
I*atel S D, Patel H S and Patel S R 1981 J. Macro. Sci. Chem. A16 1335
It-Ozhanavskaya A A and Zasova A A 1968 Plast Massy Russ 10 56 ; 1969 Chem. Abstr-
7O 2Q530w
Vlasyuk N V and Natanson M E 1976 Ukr. Khim. Zh. (Russ.) 42 473 ; 1976 Chem. Abstr
85 79743 f.
V/aa-dL Cliristopher, Johnj Morgan, Cyvil Alfred, Allen and Richard Raul 1975 Chem. Abstr.
83 142526P
ynatnic distortion of C 4 N f -skeleton in (CH 3 ) 4 NC1
H D BIST*, MAHENDRA PAL, G S RAGHUVANSHI and
V N SARIN
Department of Physics, Indian Institute of Technology, Kanpur 208016, India
MS received 25 March 1982
Abstract. The polarization data and concentration dependence of the Raman
spectra of aqueous solutions of tetramethyl ammonium chloride in the region
2700-3100 cm" 1 are studied. Group theoretically consistent assignments reveal
that the C 4 N "''-skeleton acquires a dynamic distortion to C at) symmetry due to out-
of-phase combinations of the four CH a -stretching oscillations.
Keywords. Raman spectra; tetramethyl ammonium chloride; dynamic distortion.
Introduction
itramethyl ammonium ion [(CH 3 ) 4 N + , henceforth abbreviated as TMA] retains
high symmetry T d ver a wide range of environmental variations. Preliminary
,alitative vibrational assignments, comprising of Raman (R) and infrared (IR)
idies, are available in the literature (Edsall 1937 ; Ebsworth and Sheppard
59 ; Bottger and Geddes 1965 ; Stanley and Tobin 1972 ; Harman et al 1974 ;
mder Ohe 1975 ; Berg 1975 ; Kabisch and Klose 1978 ; Kabisch 1980).
Dwever, the assignments of internal vibrations of TMA are still at a controversial
ige (Ebsworth and Sheppard 1959 ; Stanley and Tobin 1972 ; Berg 1975 ;
ibisch and Klose 1978), especially in the CH 3 -stretching region. Raman and
studies of the polycrystaliine compounds containing TMA have been reported
;ently (Kabisch 1980 ; Kabisch and Klose 1978). Kabisch etal (1978) studied
3 effect of varying concentration of tetramethyl ammonium chloride (TMAC)
D 2 O, on the CH 3 -stretching modes ; the abnormally high intensities of the
nds assigned as combinations and overtones in the CHa-stretching region were
:ributed to the ion associations occurring in concentrated solutions,
[n this paper we report the concentration dependence and polarization data
the Raman spectra of TMAC in aqueous solutions in the region 2700-
OOcnr*. The three strong polarized bands in this region are assigned as
^-symmetric modes due to a dynamic distortion of QN + -skeleton, consistent
th the group theoretical analysis.
2. Experimental
AR-grade TMAC obtained from Fluka was kept under vacuum at 300 K (RT)
to remove any moisture from this very hygroscopic compound. The aqueous
solutions of known concentrations were prepared from the moisture free compound.
The Raman spectra were recorded on a Spex Ramalog spectrophotometer. The
514-5nm beam obtained from a Spectra Physics 165-09 Ar + laser was used to
excite the Raman scattering. The scattered light was focussed into the entrance
slit at the usual 90 geometry. The reported wavenumbers of the sharp and
strong bands are correct to i 1 CUT-'-.
3. Group theoretical consideration
In tetramethyl ammonium compounds the symmetry of the whole TMA group
will be TV if all the four methyl groups in it retain a staggered or an eclipsed
configuration. Hence the distribution of vibrations for the TMA ion for both
the staggered and the eclipsed configurations remains unchanged. The 45
degrees of vibrational freedom f vlb of TMA are distributed on the symmetry
species of point group T d as follows (Berg 1975) :
Out of these, the A l9 E and F 2 species are Raman active and only the.F a species
are IR active.
A correlation of isolated CH 3 -stretching vibrations (C 3e point group) with those
of TMA under T d and C 3a point groups is given in table 1. Under the group T d
in TMA the C-H symmetric stretching (AJ vibrations of four isolated CH 3 groups
produce a totally symmetric species AI and a triply degenerate species F 2 . Like-
wise, the four isolated 'CH 3 -asymmetric stretching modes () under C 3e produce a
doubly degenerate E species and two triply degenerate species F and F 2 under
T A point group as shown in the second column of table 1. As a result under the
point group T a one expects, in the CH 3 stretching region, four Raman active
bands out of which only one should exhibit totally symmetric and polarized (Aj)
character.
If the C 4 N+ group of TMA acquires a C 3u symmetry even dynamically (e.g.,
this could be the case during its v 3 vibration as discussed later and shown in figure 1)
the situation gets changed as illustrated in the third column of table 1. It is
evident that one would expect a maximum of three totally symmetric polarized
(AJ CH 3 -stretching bands along with four doubly degenerate (E) depolarized C-H
stretching modes. This situation will be applicable to the solutions where the
crystalline field effects would be missing.
4. Results and discussion
Table 1. Correlation of the CH 3 -stretching vibrations having point group C w
with corresponding vibrations in TMA having symmetries T d or C 31r
Isolated CH 3
CH 3 -stretching vibr?tions
CEL in TMA units
"otally symmetric
;H 3 - stretching
Mbration,
-A,(R{p),IR)
(R(p),IR)
ElR(dpMR)
Asymmetric CH~-
stretching vibration,
E(R,IR)
-E(R(dp),IRl
E(R(dp),IR)
A 2
E(R(dp),IR)
ie natations regarding species are discussed in the text. The activities in Raman (R) and
Tared (IR) are given for ready reference, p polarized, dp depolarized.
(F 2 )
Figure 1. Diagrammatic representation of the three totally symmetric CH-stretching
motions of CH 3 in tetramethyl ammonium ion. (a) The T d symmetry is retained,
(b) and (c) the symmetry will be dynamically distorted to C sv (cf. column III in
table 1).
re shown in figure 2. In figure 3 the normal Raman spectra for 0-4 M and
OM aqueous solutions in the same region are presented. Table 2 summarises
ie wavenumbers (cmr*), relative peak intensities, half-widths, relative integrated
densities and polarization characteristics of the bands observed in the regions
400 to 1500 cnr* and 2700 to 3100 cnr*. The assignments of the bands observed
i these regions, based on two alternative considerations (i.e., Fermi-resonance
nd anharmonicitv and dvnamic distortion of CiN + -skeleton) are nrooosed in the
3100 2900
Wavenumber (cm~ 1
2700
Figure 2. Raman spectra (514-5 nm exciting laser beam with power on the sample
~ 150mw) of tetramethyl ammonium chloride in aqueous solution of ~ 0-4 M
in quartz cell at right-angle geometry, (a) Parallel and (b) perpendicular polari-
zations.
3100 2900 2700
Wavenumber (crrf 1 )
Figure 3. Raman sneetra nf t*t
Table 2. Wavenumbers (cnr 1 ) of the observed bands and their assignments in
the aqueous Solutions of tetramethyl ammonium chloride in the regions 1400--
1500cm- 1 and 2700-31 00 cnr 1 .
Dilute solution (0-4M)
Concentrated solution (7 M)
Assigaments
1
2
3
4
5 6
26 (dp)
5,20,6)
..
1422 (dp)
(1-2,20,7)
Vi(F a ) v u (jFi)
56 (dp)
8,18,33)
..
1456 (dp)
v,(JS) v,()
85f
1485f
v 15 (F 2 ) v l5 (F 3 )
26 (p)
6,16,21)
2852
(26)
2824 (p)
(4-6,21,30)
2844
(20)
2Vi (F?) 2v lc (Fj,)
78*
2882
(4)
2878*
2878
(0)
Vi B (F 2 ) + v (JS) v 10 (F ? )+v ()
)6*
2912
(6)
2906*
2912
2v () 2v 8 ()
50 (p)
0,15,60)
2941
(11)
2928 (p)
(15,24,114)
2941
(13)
" m+ *>"*>
56 (p)
2,18,63)
2970
(4)
2961 (p)
(13-4,18,76)
2970
(4)
2v 16 (F 2 ) v.Ui)
59 (p)
8,18,119)
2986 (p)
(14,23,102)
nW) 1-.W
*3 (dp)
3,28,100)
3031 (dp)
(12-6,25,100)
v s () v.()
v u (F 9 )
lumnS 1 and 3 indicate tho observed wavenumbers (cnr 1 ). The figures in parentheses under
:h wavenumbor denote the relative peak intensity (arb. units), the half-widths (cnr 1 ) and
relative integrated intensity respectively.
lumns 2 and 4 indicate the calcjlated harmonic values (cnr 1 ) of overtone/combination
ids shown in column 5. The figures in parentheses in these columns represent the differences
the harmonic values (as expected from the assignments in column 5) and the actual observed
venumbers shown in columns 1 and 3.
lumns 5 and 6 represent the assignments based on "Fermi-resonance and anharmonicity "
I "dynamical distortion of the C 4 N" H -skeleton", respectively. The assignments in column 5 are
j reported by Kabisch etal (1978). Here we have used standard notations (see Berg 1975)
numbering the frequencies as Vi 6 (F 2 ), etc. The v a (/4j) in column 6 is the same as Vi(A : )
column 5.
nd dp denote the polarized and depolarized character of bands.
ands are observed in the infrared spectra.
;ction rules for overtones and combinations are
X F. = A* -t- E + F, + F 2 E X E = Ai + A* + E E X F, = F, + F 2 .
drastically.
5. Fermi-resonance and anharmonicity
Our assignments, shown in column 5 of table 2 are analogous to those proposed
by Kabisch etal (1978). The moderate intense band at 2826 cm^ (in 0-4M
solution) shows the largest mechanical anharmonicity and its intensity is about
an order of magnitude smaller than that of the other strong polarized band.
This band has been attributed to an overtone of the v^(F^) mode. Similarly the
other very weak bands at 2878 and 2906 cnr* (in 0-4 M solution), with less
anharmonicities, are assigned to a combination and an overtone band respectively.
The assignments of the bands in the region 1400 to ISOOcnr^ are based on
several unambiguous arguments including the Teller-Redlich product rule (Berg
1975 ; Herzberg 1960). The depolarized band at 3043 cnr* and the polarized
band at 2989 cnr* can safely be assigned to the CH 3 -asymmetric and the CH 3 -
symmetric stretching modes respectively. The polarized band at 2966 cirr 1 - can
also be justified as the allowed A t component of the overtone of v 15 (F 2 ) mode.
The other polarized band at 2930 cnH cannot be assigned as the combination
of the v s (E~) and the v 15 (F 2 ) modes.
In general, due to Fermi-resonance, the resonating band which is nearer to
the fundamental gains more intensity than the band farther from it. Neither the
Fermi-resonance nor the electrical anharmonicity consideration can explain the
observed intensity distribution and its variation with concentration. The Fermi-
resonance combined with the electrical anharmonicity (coupled with mechanical
anharmonicity) may explain the observed abnormal intensity distribution. One
may argue that the band at 2930 cnr^ may have more electrical anharmonicity
(as would be expected on the basis of higher mechanical anharmonicity) than the
band at 2966 cnr*, making it (2930 cnr* band) more intense after Fermi-reso-
nance than the band at 2966 cnr* which is closer to the fundamental at 2989 cnr^.
However, this argument is untenable due to the following reasons :
(i) The fundamental requirement for the Fermi-resonance is that the levels involved
must have the same symmetry (Herzberg 1960). The selection rules show
that the combination of vc() and v 15 (F 2 ) produces levels of symmetry species
F 1 and F z , out of which species F 2 cannot resonate with CH 3 -symmetric mode
which is of symmetry species A^.
(ii) The intensity of the resonating band(s) can never exceed the intensity of the
fundamental band, unless electrical anharmonicity is a dominating factor
which is seldom observed in Raman spectra.
(iii) It is observed that on varying concentration, the intensities of the resonating
bands change drastically whereas the band positions remain almost unchanged.
This is also not observed in Fermi-resonance especially when the positions
of fundamental bands remain unchanged.
Thus the intense and the polarized band at 2930 crrH cannot unambiguously
bynamic distortion of C 4 7\T+ -skeleton 191
an alternative explanation of the three strong polarized and intense bands may be
searched in the dynamic distortion of QN + skeleton.
6. Dynamic distortion of C 4 N + -skeIeton
The totally symmetric CH 3 stretching motions of the four CH 3 groups in TMA ion
could result in three specific configurations consistent with the internal vibrations
of C 4 N + -group as shown in figure 1. In the first configuration (figure la) all the
C-atoms vibrate in phase along C-N bonds. In this case the T a symmetry of the
C 4 N+-skeleton is maintained even dynamically. This situation corresponds to
Vj mode of species A for the skeleton and yields one totally symmetric polarized
CH 3 -stretching vibration.
In the second configuration, three C-atorns vibrate in-phase and one out-of-
phase and vice versa along C-N bonds as shown in figure l(b). In this case the
T d symmetry of C 4 N + -skeleton is distorted to C 3w dynamically corresponding to
^3(^2) mode for the skelton (T d ) and yields under c& one A totally symmetric
polarized CH 3 -stretching mode and another asymmetric depolarized CH 3 -stretching
mode ().
In the third configuration, three C-atoms vibrate in-phase perpendicular to
C-N bonds and one along the C-N bond as shown in figure l(c). In this case, too,
the T d symmetry of C 4 N + -skeleton is distorted to C 3M dynamically and the situation,
corresponding to v 4 CF 2 ) mode for the skeleton (T a ), yielding one totally symmetric
polarized CH 3 stretching mode (A-^ and another aysmrnerric depolarized CH 3
stretching mode (E).
There are two more possible configurations which produce two asymmetric
depolarized CH 3 stretching modes corresponding to the vibrations of the C 4 N + -
skeleton of species E and F v It can, therefore, be concluded that there must be
three totally symmetric polarized C-H stretches due to dynamic distortion of
C 4 N+-skeleton and coupling between the four CH 3 groups. Thus the evidence
for the three expected A^ modes in CH 3 stretching region is unambiguously
established.
The four CH 3 asymmetric stretches may appear at the same frequency giving
thereby a single depolarized band at 3043 cnH (in 0-4M solution). The larger
half-width of this band and shift with concentration may also corroborate the
presence of more than one unresolved bands under the combined envelope.
Acknowledgements
Financial support from Department of Science and Technology (DST), India, for
providing funds to establish a Laser Raman Laboratory is gratefully acknowledged.
l.JLU>J.UJiaU Vi IVi, VJ&JULiLVJV/ J. (UJU JCWIUCIUU U t-l 1 S I T J. JL/iyj. ^/IClll, IO 4iJOJ
Herzberg G 1960 Infrared and Raman spectra of polyatomic molecules (New York : Van
Nostrand)
Kabisch G and Klose M 1978 /. Raman Spectrosb. 7 311
Kabisch G 1980 /. Raman Spectrosc. 9 279
Stanley A A and Tobin M C 1972 Spectrochim. Acta A28 2141
Yonder Ohe W 1975 J. Chem. Phys. 62 3933
i ab initio molecular orbital study of thymine radicals 1
U CHANDRA SINGH and A MURALIKRISHNA RAO*
Solid State and Structural Chemistry Unit, Indian Institute of Science,
Bangalore 560012, India
MS received 18 March 1982
Abstract. Ab Initio calculations at the STO-3G level have been carried out on
the six thymine radicals, \-yl, 6-yl, 5-yl, 1-yl, 4-hydroxyl and the anion. The results
have been compared with those from ESR studies.
Keywords. Thymine radicals ; molecular orbital calculations ; radiation biology.
Introduction
,e of the most well-known radicals in radiation biology is the thymine radical
ossmann etal 1980), which is known to occur in different types (chart 1).
ymine is used as a model compound for analysing radicals generated by radiation
ought ESR spectroscopy. One of the first radicals of thymine to be examined
the 5-yl radical (Pruden et al 1965) and since then a number of papers have
peared on the different radicals of thymine. Various radicals produced in
gle crystals of thymine derivatives by uv and x-radiations have been investi-
:ed recently through ESR spectroscopy by Flossmann et al (1980). The first
empt to study the radicals theoretically was made by Heiberg and Jensen (1977)
10 employed the semi-empirical INDO method, and found that the INDO
;thod does not yield satisfactory agreement with ESR studies. We considered
desirable to perform ab initio calculations on various thymine radicals in view
their obvious importance to radiation biology.
Method of calculation
ilculations were performed at the minimal STO-3G level (Hehre etal 1969).
>r each radical species, one basic set of structural parameters was developed,
hile some of the optimized INDO parameters (Heiberg and Jensen 1977) were
osen for the purpose, most bond angles and some of the bond lengths were
osen with some guidance from the results of the crystallographic investigation
R H
Anion 4-hydroxyl
,.
x-
'
5-yl 6 "V l
of thymine derivatives (Gerdil 1961). In the case of the l-yl radical, bond distances-
such as N1-C6, N1-C2 and C4-O8, were optimized.
Since the UHF calculations mix up the higher spin-component wavefunctions,
spin densities calculated by the UHF method cannot be correlated with experi-
ment well. Hence, all the spin densities were calculated after annihilating the
unwanted higher spin multiplets using the spin-annihilation procedure (Amos
and Snyder 1961 ; Snyder and Amos 1965). In general, two annihilations were
needed to get rid of the contribution of the higher spin-states. ENDOR experi-
ments on the methyl protons in thymine radical in irradiated thymidine is known
to show a large potential barrier to rotation. Hence, the spin properties of the
methyl protons were calculated only for a single conformation and averaging over
the angles was not performed.
3. Results and discussion
Expectation values of the soin operator '(S z } given in table 1 shows that the
ot the spin, densities before -and alter annihilation snows that the annihilation
procedure reduces spin densities strongly.
In table 2, we have listed the calculated TT and total spin densities at C, N and
O atoms for all the radicals studied and in table 3, we list the calculated atomic
spin densities at the hydrogen atoms. The calculated and experimental coupling
constants for the radicals are given in table 4 while in table 5 we have listed the
Mulliken populations.
Table 1. The (S 2 ) values and the relative energies of the radicals."
Radical
,,
.
Relative
energy
kcal/mol
1-yl
0-7478
1-1706
o-o
6-yl
0-7495
0-7859
98-5
5-yt
0-7425
1-0503
114-4
1-yl
0-7654
1-3819
- 3-8
anion
0-7529
0-9036
-47-9
4-hydroxyl
0-7578
1-3304
108-4
"After annihilation (aa); singe determinant (sd).
Table 2. Calculated n and total spin densities at carbon, oxygen and nitrogen
atoms.
Radical
C 4
C 6
(
1
(
-(
J-l
*
]
^3
s o fl
8
1-yl it
Tot.
-0-058
-0-082
0-403
0-451
-0'
-Q-
1.14
157
-0
-0
004
023
0-421
0-449
-0'
o-
016
015
0-158
0-163
0-215
0-222
6-yl
0-01.7
0-026
-0-019
-0-094
o-
I-
861
017
014
033
0-062
0-035
-0'
-o-
003
001
0-118
0-126
-0-002
-0-026
5-yl
-0-102
-0-149
0-659
0-728
-Q-
-Q-
007
034
-0
-0
007
035
0-002
0-004
-o-
-Q-
009
007
-0-015
-0-016
0-396
0-414
1-yl
0-127
0-166
-0-119
-0-195
Q-
o-
434
466
531
606
0-037
0-028
Q-
o-
013
012
0-079
0-086
-0-111
-0-125
anion
0-060
0-054
0-020
-0-009
o-
a-
593
656
002
007
0-032
0-019
Q-
o-
014
015
0-004
0-004
0^270
0-280
4-hydroxyl
0-385
-0-099
o-
447
004
0-082
o-
0.75
-0-119
0-039
Radical H(C5) H(C6) H(N1) H(N3) H(Me) H(O8)
1-yl
006
o-
00
o-
019
6-yl
0-023 -0
060
-o-
004
o-
-0
002
5-yl
Q
017
-Q-
Q-
018
1-yl
-0'
018
-Q-
001
Q-
-0
022
an ion
-Q-
023
-a-
001
o-
Q
4-hydroxyl
-0
019
-Q-
005
-o-
004
-0
004 -0-002
Table 4. Calculated and experimental coupling constants.
Radical a-proton j3-proton jS-protoa
(CH 8 ) (methylene)
[-yl Exp
18-7
ab initio 4-91
15-55
. .
INDO (sd) 6-60
19-35
1-yl -48-0
17-4
-49-1
18-81
. .
-38-1
17-6
-IKDO {aa) 32-8
7-9
w
21-7
40-0
..
14-72
13-91
22-3
31-60
19-2
26-5
l.yl _ 9-7
m i
-15-2
-14-72
. .
-17-99
-12-10
. .
-14-50
- 5-3
*
'- -6-3
,nion 14-3
-18-81
. .
, .
-13-7
- 6-0 .
hydroxyl -r-21-1
-15-54
.. '
"
-12-70
" ' - 5-5 '
Bond
5-yl
1-yl
anion 4-hydroxyl
N1-C6
0-092 0-015 0-002
0-021 -Q'Q05 0-012
C6-C5
0-098 0-006 0-005
0-094 0-067 0-097
C5-C4
0-053 0-005 0-069
0-028 0-097 0-102
C4-N3
0-026 0-036 0-025
0-031 0-005 0-013
N3-C2
0-028 0-034 0-035
0-034 0'049 0'032
C2-N1
0-040 O'OOl 0-053
0-042 0-052 0-033
C2-07
0-139 0-139 0118
0-138 126 0-124
C4-08
0-135 0157 0-136
0-136 0-096 0-013
C5-C9
0-005 0056 0-005
0-089 0-007 0-004
c -MO -
In
|-,o-
>Xi-yi
c anion-<$
'5. X
8-30- X
"4-hydroxy!
$ /
u -50-X-yl
1 1 1
1 1
-0-06 -0-02 + 0-02
atomic spin density of H (C6)
Figure 1. Plot of experimental coupling constant against atomic spin density of
H (C6).
Since the hyperfine splitting constants are proportional t'o spin densities at a
particular hydrogen atom, the calculated hydrogen spin densities were plotted
against the experimental splitting constants. We see from figure 1 that there is. a
linear relation with a slope of 818-5 Gauss per unit spin density. "
The UHF calculations make it possible to correlate the unpaired spin densities
in the 2p 3 atomic orbitals with the unpaired spin densities in the .2^-orbital of the
carbon atoms. There is a linear relationship which justifies the use of McConnell's
(1956) procedure to compare spin densities obtained by the RHF calculations
with exDerimental values obtained from ESR st>ectroscor>v. We shall nr>w brieflv
3-1 l-yl radical
The l-yl radical resulting from the removal of the hydrogen in the Nl position
hydrogen has been observed in single crystals of anhydrous thymine (Dulcic
and Herak 1972). The radical is stable upto 300 K. The optimized values
for the N1-C6, C5-C6 and C4-O8 distances in tuis radical are 1-389 A,
1 375 A and 1 264 A resectively. Unpaired electron densities at the p a orbitals
of Nl, C5 and O8 atoms are 0-421, 0-403 and 0-215 respectively ; the unpaired
electron is therefore localized on these sites. As examination of Mulliken's
^-population analysis shows that there is considerable delocalization of the
^-electrons in the ring. It is interesting that the lone pair electron at the Nl
atom is in the plane of the ring just as in the case of pyridine molecule. The
7r-orbital charges at Nl is 1-0240 while in all other radicals it is around 1-730.
Thus, the l-yl radical could be depicted as follows :
H
This type of behaviour was also observed, in uracil derivatives (Horan and
Snipes 1970 ; Farley and Bernhard 1975 ; Zehner et al 1976). The ESR spec-
trum of l-yl uracil radical shows the unpaired electron to be localized on the
Nl-atom (^ 0-40). The calculated coupling constant for the methyl proton is
15- 6 G which agrees with the experimental value of 18-05G. Our INDO
calculations predict a value of 19-35G.
3-2 6-yl radical
The 6-yl radical was first analysed in detail by Henriksen and Snipes (1970) in a
single crystal of dihydrothymine exposed to x-rays. Flossmanu et al (1979)
have also identified this radical in thymine irradiated at 230 K and in methyl
had found a value of 28 3 G for the a-proton coupling constant while Flossmarm
et al have found the value to be 48 G. Our calculations predict a value of 49 1 G
which agrees well with that of Flossmanu et al. Further, the experimental value
of the /^-proton coupling constant is 17-4 G which also agrees with the calculated
value of 18-8 G. The Mulliken population analysis also shows that the 7r-elec-
trons are localized on the carbonyl and there is little delocalization in the ring.
3-3 5-yl radical
The 5-yl radical was observed in both thyinine and dihydrothymine and its deri-
vatives. In dihydrothyrnine exposed to x-rays at 300 K, it is found to be present
only as a minor fraction whereas it is predominant under similar conditions in the
parent compound. According to our calculations, this is the most stable of all
the thymine radicals. The spin densities at the C5 and O8 atoms are 0-659e and
0- 396e respectively. Thus, the spin density is distributed between these two atoms.
The calculated splitting constant for the methyl proton is 14-8 G which agrees
well with the experimentally observed value. Our INDO calculations predict a
value of 22-3 G. However, the methylene proton values are substantially lower
than those obtained experimentally. The calculated value is 14 G, the experi-
mental one being 40 G. According to our INDO calculations, the variation of the
methylene group through rotation around the C5-C6 axis cannot increase the
coupling constant substantially. In order to study the effect of the HCH angle
on the coupling constant, calculations were carried out for two different angles.
It was found that the decrease of the HCH angle to 72 only increased the coupling
constant to 19-2G. However, this increases the energy substantially (by about
210 kJ moH). On the other hand, increase in the HCH angle to 137 decreased
the coupling constant to 10-8 G while the energy increased by about 77 kJ moH.
One possible explanation would be the protonation of the carbonyl oxygen as
suggested in the case of uracil derivatives (Zehner et al 1976 ; Flossman et al
1979). We could not carry out calculations due to convergence problem for the
protonated species.
3.4 1-yl radical
This radical, originally proposed by Huttermann (1970), was detected in several
thymine derivatives except dihydrothyrnine. The radical is thermally stable upto
485 K. Our calculations show this radical to be as stable as the l-yl radical.
The calculated splitting constant for the C7 proton is 18 G which agrees well with
the experimental value of 15-3 G. However, the calculated value for the C6
proton is 14-7 G which is somewhat higher than the experimentally observed one.
The unpaired electron density seems to be located mostly on the C9 (0-606) and
C6 (0-466) carbon atoms. The spin densities on the C5, C6 and C9 atoms as
calculated by the INDO method are in agreement with the ab initio calculations.
3-5 Anton radical
Hydrated thymine irradiated with x-rays at 77 K exhibits a doublet ESR spectrum
attributed to the anion radical. This radical is highly sensitive to light and heat
which are the characteristic of ionic radicals. This ar.ion radical has also been
observed in thymidine at 4-2K. Our calculations show that the spin densities
are localized on the C6 (0-656) and O8 (0-280) atoms resulting in the doublet
splitting due to the a-proton. The calculated splitting constant for the C6 proton
is 18-8 G which is close to the experimental value of 14-3G. The calculations
show that this radical is the least stable of all the thymine radicals.
3-6 4-hydroxy radical
In order to explain the doublet spectrum obtained at low temperatures, Henriksen
and Snipes (1970) proposed the 4-hydroxy radical. However, later work showed
the doublet to be due to the anion radical. We have carried out calculations on
the 4-hydroxy radical which can be considered to be the protonated anion. Our
calculations predict a high stability for this radical compared to other radicals.
The unpaired electron density in this radical resides mainly on the C4 (0-429)
and C6 (0-501) atoms. In the anion radical, the spin density on the C4 atom
is very low (0-060) in contrast to this radical. Also, there is a large negative
spin density on the O7 atom ( 0-143) in contrast to the anion radical (0-004).
However, the spin density distribution on the H-atom is remarkably similar.
A comparison of the Mulliken population of ^-electrons shows that, except in the
C4-O7 bond, the populations are quite similar.
Acknowledgements
The authors are thankful to Professor CNR Rao for guidance and helpful
discussions and to the US National Institute of Health for support of this research
(01-136-N).
References
Amos T and Snyder L C 1961 /. Chem. Phys. 41 1773
Dulcic A and Herak J N 1972 /. Chem. Phys. 57 2537
Farley R A and Bernhard W A 1975 Radiat. Res. 61 47
Flossmann W, Zehner H and Westhof E 1979 Int. J. Rod. Biol. 36 249
Elossmatm W, Zehner H and Mulier A 1980 Z. Naturforsch. C35 20
Gerdil R 1961 Acta Crystallogr. 14 333
Hehre W J, Stewart R F and Pople J A 1969 /. Chem. Phys. 51 2657
Henriksen T and Snipes W 1970 /. Chem. Phys. 52 1997 ; see also Radiat. Res. 42 255
Heiberg A B and Jensen H H 1977 Acta Chem. Scand. A31 195
Horan P K and Snipes W 1970 Radiat. Res. 41 24
Huttermann J 1970 Int. J. Rad. Biol. 17 249
McConnell H ]Vt 1956 /. Chem. Phys. 24 764
Pruden B. Snines W and Gordv W 1965 Proc. Nat . Acad. Sci. USA 53 917 anrt rfifW
EXAFS studies of cobalt oxides and oxide glassest
R PARTHASARATHY, RAVI V PRASAD+ P R SARODE and
K J RAO*
Solid State and Structural Chemistry Unit, Indian Institute of Science,
Bangalore 560012, India
MS received 31 December 1981
Abstract. The EXAFS of Co 2 " 1 " has been studied in rare earth cobaltites and in
sulphate and borate glasses. It has been found that the environment of Co 2+ ions
is very similar in these cases. It appears feasible to study local structures in
glasses using probe ion EXAFS.
Keywords. EXAFS ; ionic glasses ; crystals.
1. Introduction
Rare earth cobaltites have been well investigated in recent years and their structures
are known in sufficient detail (Raccah and Goodenough 1967 ; Demazeau et al
1974). An EXAFS study of such model compounds would be expected to furnish
information (e.g., phase-shifts) which could then be used in EXAFS analysis of
glasses containing cobalt. Considering thecomparative paucity of EXAFS data
on crystalline compounds whose structure is known in detail, we felt that an
EXAFS study of these compounds could also provide much insight on the scope
of the technique itself.
In this paper we report results of EXAFS investigations of crystalline LnCoO 3
(where Ln = La, Nd, Dy and Yb) compounds and of some cobalt-containing
glasses belonging to the systems NaaSO^KaSC^-ZnSCVCoSO,! and CoO-B a O 3 .
The ternary sulphate system Na 2 SO 4 -K a SO 4 -ZnSO 4 has been extensively studied
in this laboratory (Sundar and Rao 1980 ; Sundar and Rao 1981). At low concen-
trations, the isovalent Co 2 + ions are known to enter Zn 2+ sites due to similarity
of sizes (Sundar and Rao 1981). At higher concentrations, however, the site
symmetry of the Co 2 ' 1 ' ion is not known. A Co K-edge EXAFS study might be
expected to yield valuable information on both concentration regimes.
t Communication No. 151 from the Solid State and Structural Chemistry Unit.
+ NSTS Scholar from I.I.T., Kanpur.
* Tn ti>Vl/~tm oil l>n-~Bf~n- n /ta.-nr>n. rU^v.l^l Un n ^^.occay4
202 & Parthasarathy et al
2. Experimental
The preparation of the rare earth cobaltites from the cobalticyanide (Demazeau
et al 1974) and of sulphate glasses from component sulphates (Sundar and Rao
1980) have both been reported earlier. Glasses with higher cobalt content were
progressively more unstable with respect to devitrification. The Co-borate
glass was prepared by melting together CoSO 4 and H 3 BO 3 in a graphite crucible.
Co K-absorption edge EXAFS was measured using a bent crystal spectro-
graph and a Carl-Zeiss microdensitometer as described previously (Parthasarathy
etal 1981). Data were available upto about 400 eV beyond the edge. Details
with regard to the procedure for data analysis and Fourier transformation along
with subsequent curve-fitting procedures have been described elsewhere (Cramer
1978 ; Parthasarathy et al 1982).
3. Results and discussion
Figure 1 shows the radial structure functions (RSF'S) obtained for the crystalline
compounds. Figure 2 shows the RSF's obtained for the glasses. Distances corrected
for phase-shifts are listed in table 1.
3-1. Cobalt in rare earth cobaltites LnCoO z
Rare earth cobaltites are known to crystallise in the perovskite structure (Demazeau
et al 1974). The first peak in the RSF at 1- 92 0-03 A may be identified as the
JD
O
tr
-e-
Table 1. Phase Shifts and distances from RSF'S of cobalt oxides and oxide glasses.
No.
Compounds
a* (A)
Distances (A)
1.
LaCoO 3
034
1
91 2-05
2-91 3
35
4'
25
4-
71
2.
NdCo0 3
046
1
92
2-85 3
29
3
83
4-
37
3.
DyCoO 3
020
1
96
2'83 3
26
3
80
4-
34
4.
YbCo0 3
04
1
91 2-38
2-67 2
95
3
35
3-
60
5.
Sulphate Glass 1
o-
0,81
2
02
3'
42
4'
61
5-
73
6.
Sulphate Glass 2
o-
029
1
98
3
71
4
53
5-
43
7.
Sulphate Glass 3
0-
073
2
02
3
31
3
85
4-
57
8.
Borate Glass
o-
065
1
90
3-
49
4-
64
5'
76
Co-O distance. This distance is almost independent of the rare earth ion. Phase
corrections were made using YbCoO 3 as a standard (Cramer 1978). For peaks
other than the first, we felt that use of the curve fitting procedure was in-
advisable since their amplitudes in the RSF were low. Phase corrections, a*,
were calculated using the expression,
* n t n
standard l Fourier transform.
I a* | was then added to the appropriate peak in the RSF'S of the other cobaltites.
The feature at 2-9 A, however, is rather perplexing and we are unable to ascertain
its origin from available crystal data.
3-2. Cobalt in glasses
3-2a Sulphate glasses : The similarity of the RSF'S of these glasses to those
of the crystalline cobaltites suggests that the co-ordination of Co 2+ in these are
similar. This is quite in keeping with the structural model proposed earlier for
these glasses (Sundar and Rao 1980). On the basis of this model, we assign the
peak at 1 -90 A to the Co-O distance. Similarly, the 3-7 A peak may be assigned
to the Co-K separation ; the peak at 5-5 A could possibly describe the Co-Co
or Co-Zn distances.
At low Co 2+ concentration, the sites that Co 2+ ions may occupy are limited to
those of Zn 2+ ions in the glass. With these concentrations, one would expect site
distortions and the asymmetric first peak does in fact reflect the existence of a
range of distortions. Increasing the Co 2 * concentration would result in the
formation of a distinct but similar type of Co 2+ site in the glass and the forma-
tion of characteristic Co 2 + sites is apparent in the symmetric nearest neighbour
peak in figure 2. Glasses 2 and 3 (see figure 2) show a sudden decrease in RSF
Figure 2. RSF'S of cobalt oxide glasses. The numbers indicate the following
molar compositions (I ) ZnSO 4 : Na a SO 4 : K 2 SO 4 : CoSO 4 : : 1 5 : 30 : 30 : 25
(2) 25 : 25 : 25 : 25, (3) 20 : 20 : 20 :40 and (4) CoO : B a 3 : : 20 : 80.
of metallic glasses where dense random-close-packing models are valid (Wong
1980). In glass 1, however, the decrease in RSF amplitude for R > 4A is not
so significant and peaks are discernible to 5-5A- It may be recalled that these
distances in glass 1 are attributed to Co-Zn or to Co-Co. The decreased structure
in the RSF due to increasing Co 2+ concentration, allows for two possible expla-
nations. One is that the scattering amplitude of Co 2 * ions is predominant and
that scatterers further away than the " line-of-sight " Co 2+ scatterers are conse-
quently screened. The RSF in this case will then show only the presence of atoms
in between the absorber Co 2+ and the " line-of-sight " Co 2+ scatterer. The other
explanation could be that the formation of typical Co 2+ sites in glasses containing
higher concentrations of Co 2+ gives rise to an intrinsically less structured RSF
and this may obscure features due to Co 2 + in Zn 2 -'- sites typical of dilute glasses.
3 2b. Borate glass The noise level in the EXAFS of the borate glass is consi-
derably high (note the high background level in the RSF for R< 1-5A). The
nearest neighbour peak at 1 -90 A (the Co-0 distance) (which was merely asymme-
tric in the sulphate glasses) is now resolved into two distinct components. The
sharpness of the first peak at the normal Co~O distance, indicates the presence
nf a smtrlfi striintiira! unit The neak at 2-70 A has a rather larpp> wiHtli smH inrli.
4. Conclusion
In brief, the Co a + site symmetry in ionic and in covalent glasses appears to be very
similar to that in the rare earth cobaltites. This rather interesting feature may
ensure from the characteristic O 2 ~ ligancy of the Co 2+ ions. Further, structural
information from EXAFS (related to the structuredness of the RSF) appears to
be reduced in the presence of a higher concentration of scatterers. The origin of
this effect is, however, not known.
Acknowledgements
The authors are grateful to Professor CNR Rao, who suggested the possibility
of using probe ion EXAFS for structural studies, for his encouragement and for
helpful discussions. Thanks are due to Dr W H Madhusudan for assistance
in the preparation of the cobaltites.
References
Cramer S P 1978 SSRL Report 78/01 Stanford Synchrolron Research Laboratory
(Stanford USA)
Demazeau G, Pouchard M and Hagenmuller P 1974 /. Solid State Cliem. 9 202
Parthasarathy R, Sarode P R and Rao K J 1981 J. Mat. Sci. 16 3222
Parthasarathy R, Rao K J and Rao CNR 1982 /. Phys. C: Solid State Phys., IS 3649
(communicated)
Raccah P M and Goodenough J B 1967 Phys. Rev. 155 932
Sundar H G K and Rao K J 1980 /. Ghent. Soc. Faraday Trans. 776 1617
Sundar H G K and Rao K J 1981 /. Non-Cryst. Solids, (communicated)
Wong J 1980 in Topics in Applied Physics, Vol. 46 Metallic Glasses, (eds) H J Guntherodt and
H Beck Chapter 4 (Berlin : Springer-Verlag)
ectrical conductivity studies in sulphate glasses and the mixed
kali effect t
H O K SUNDAR and K J RAO*
Solid State and Structural Chemistry Unit, Indian Institute of Science,
Bangalore 560012, India
MS received 31 December 1981
Abstract. Electrical conductivities of alkali sulphate-zinc sulphate glasses have
been measured. The variation of conductivity with compositon confirms the
presence of the mixed alkali effect. The origin of mixed alkali effect has been
explained on the basis of structural considerations reported earlier by us.
Keywords. Ionic glasses ; conductivity ; mixed alkali effect.
Introduction
an earlier paper (Rao and Sundar 1980) on conductivity of sulphate glasses,
reported the mixed alkali effect in K 2 SO<rNa 2 SO 4 -ZnSO 4 glasses. We had
ributed the effect to a structural origin using a viable structural model
1 sulphate glasses. Microscopic strain fields develop around the Na+ ions
i it could account for the observed conductivity minimum. We therefore
i.sidered it imperative to examine how other alkali ions behave in mixed alkali
npositions of sulphate glasses. In this paper, the observation of the mixed
;ali effect for two other alkali pairs is reported and the general conductivity
session suggested earlier is found to be adequate to account for the mixed
;ali effect.
Experimental
isses were prepared using Zn.SO 4 7H 2 O, K 2 SO 4 (BDH) and Rb 2 SO 4 or Cs 2 SO 4
erck) samples with a minimum purity of 99% as batch materials. Batches
ighing 5-10 g were melted in platinum crucibles in an electric furnace and
nples suitable for conductivity measurements were obtained as described earlier
arasimham and Rao 1978, Sundar and Rao 1980). The conductivity cell
i the measuring technique have also been discussed earlier (Narasimham et al
lommunication No. 160 from the Solid State and Structural Chemistry Unit.
'o whom all correspondence should be addressed,
207
Keithley elccetromtcr) through the sample for a known applied voltage (5 V).
The current was. passed only for very brief periods (less than 30 sec at a time)
in order to avoid polarization. In the present studies also conductivities were
measured beyond the glass transition and upto temperatures where the samples
either deform under the spring loaded electrodes or crystallised, thus rendering
further measurements inaccurate. All the glasses were annealed at a temperature
10 K below T fl before Snal measurements of a were made.
3. Results and discussion
The conventional plots of logo 1 vs I*" 1 are given for various glasses in figure 1.
For clarity actual points have not been shown in the conductivity plots. The
glasses exhibit an inflection point which corresponds to the glass transition tempe-
rature. Variation of conductivity (d), activation energy (E a ) and the limiting con-
ductivity (cr fl ) are plotted separately for Rb and Cs glasses in figures 2 and 3.
10 -
11 -
11 -
10
11
12
13
10
11
12
13-
10
11
12
13
14
o
's
b
o
o
1-8 2-2 2-6 3-0
10 3 /T
2-0
2-4 2-8
10 3 A
3-2
Figure 1 . Variation of conductivity with temperature.
I 50% ZnSO 4
(a) 45%K a SO 4 5%Rb 2 SO 4
(b) 35% K a SO 4 15% Rb 3 SO 4
(c) 25% K ? SO 4 25% Rb a SO 4
(d) 20% K 3 S0 4 30% Rb a SO 4
(e) 1 5% KoSO 4 35% Rb 2 SO 4
II 60% ZnS0 4
(a) 35%K 2 S0 4 5%Cs a S0 4
(b) 30% K a S0 4 10% Cs a S0 4
(c) 25%K,,SO 4 ]5%Cs 2 S0 4
(d) 20% K 2 S0 4 20% Cs a S0 4
(e) 15%K 2 S0 4 25%CS 2 S0 4
Electrical conductivity studies in sulphate glasses
209
The relevant quantities for the pure K 2 SO 4 -ZnSO 4 glasses were taken from earlier
work (Narasimham et al 1979). The percentages of ZnSO 4 in K-Rb-Zn and
K-Cs-Zn glasses were 50% and 60% respectively since these formed very stable
glasses. The activation energies and log <r a , values for Zn (Rb, K)SO 4 glass
exhibit no variation in the entire composition range. The behaviour of these
quantities in Za(K, Cs)SO 4 glasses exhibited considerable scatter but with an
averagj which had little variation. The conductivity variations in figures 2 and 3
clearly establish the occurrence of the mixed alkali effect in ionic sulphate glasses.
In the structural model of sulphate glasses considered in our earlier study K+
ions are co-ordinated to four sulphate ions. The occupation of K+ ion sites by
smaller Na + ions causes a small degree of inward pull of the surrounding ions
9.0
o
b -10-0
S
11-0
(a)
(b)
0)
>x
0-8
U)
IH ft
I I
0-2 0-6 1-0 0'2 0-6 VO
Rb/Rb*K
0-2 0-6 1-0
Figure 2. Variation of (a) log cr, (b) log a and (c) activation energy due to
mixed alkali effect in 50 ZnSO 4 xK a SO 4 ' (50- x) Rb a SO 4 glasses.
-11-0 -
(b)
0)
1-2
1-0
ti 0-8
(cl
n.? n-6 1-Q
0-2 0-6 1-0
0-2 0.6 1-0
Figure 4. Schematic of possible rearrangement of structure when a K+ ion is
substituted by larger Cs+ iort. Note that diffusion outwards is difficult when the
larger ion is present.
and hence creates a local strain field. This strain field has the effect of locking
up the K+ ions in the neighbourhood which, we assumed, would therefore not
contribute to conductivity. Since the substitution of Na+ ions in general leads
to a decrease in volume, there is an average enhancement of activation barriers.
The contribution of Zn 2 * ions to conductivity was neglected throughout.
Similar considerations may be applied to the other alkali pairs though we do
not have a clear idea of the nature of the co-ordination of GS+ or Rb + ions.
Substitution of K + ions by larger ions may not, in general, affect the activation
energy since the anions can rearrange with no net strain when pushed outwards.
We assume therefore, that the activation energies are unmodified. However, the
K+ ions in the immediate neighbourhood of larger (Rb+ or Cs+) ions are still
locked up as schematically representsd in figure 4 in two dimensions. The fact
that activation energies and the pre-exponentional terms are not much affected
is borne out in figures 2 and 3.
Hence, the conductivity variation is only due to the locking up of a fraction of
dissimilar alkali-ion pairs which are nearest neighbours. We can therefore repre-
sent the total conductivity <7, of a xZnSO A (l-x)lf(Kb, Cs) 2 SO 4 (l-/) K 2 SO]
glass by the expression,
~ + 2 (1- x) (1 -/) A K+ exp
-4(1
(1)
A's and E's represent the respective pre-exponentional (frequency) factor and
activation barriers.
The last term represents the contribution from locfced-up K+ ions whose con"
centration is governed by the joint probability of the presence of K+ and the
other alkali ion as immediate neighbours. Assuming that
RT
in view of figure 2. Thus the value of /corresponding to the minimum con-
ductivity may be obtained by differentiating equation (1) with respect to /,
-4,4(1 -x)
or
The value of/ obtained is in agreement with the conductivity behaviour in
figures 2 and 3. The unique siz^ ratio ol K+ to O 2 " is the factor which, while
causing significant strain contributions when Na+ is present^ does not affect the
activation barriers when larger alkali ions are present. We therefore feel that in
ionic glasses the mixed alkali effect is primarily a structural consequence and
conductivity is dominantly affected by packing considerations, perhaps unlike in
other network and covalently bonded glasses (Isard 1969, Day 1976)
Acknowledgements
The authors are thankful to Professor CNR Rao for his kind encouragement.
We are also thankful to the Department of Science and Technology, Government
of India, for financial assistance.
References
Day D E 1976 /. Non-Cryst. Solids 21 343
Isard J O 1969 J. Non-Cryst. Solids 1 235
Narasimham P S L and Rao K J 1978 7. Non-Cryst. Solids 27 225
Narasimham P S L, Sudha Mahadevan and Rao K J 1979 Proc. Indian Acad. Sci. 88 11
Rao K J and Sundar H G K 1980 Phys. Chem. Glasses 21 216
Sundar H G K and Rao K J 1980 /. Chem. Soc. Faraday Trans. 176 1617
Reactions of co-ordinated ligands : Kinetics and mechanisms in the
harge transfer interaction between dichloro or diaquo triethylene
etrainine Co(III) and ferrocyanide
G ViSALAKSHI and K S VENKATESWARLU*
Reactor Chemistry Section, Cbemcal group, Bhabha Atomic Research Centre,
Bombay 400 OSS, India
MS received 8 November 1980 ; revised 7 April 1982
Abstract. The cherry red-coloured solid state product obtained by the reaction of
cis-a diaquo Co (III) triethylene tetramine with ferrocyanide was examined by
tlwrmogravimetric, infrared and Mossbauer techniques. Its electronic spectrum
in aqueous medium was interpreted to consist of charge transfer IT transition at
440 nm and a ligand field transition around 330 nm. The kinetics of formation of
this 1 : 1 product in solution was also studied.
Keywords. Cis-a diaquo Co (III) triethylene tetramine ; charge transfer IT transi-
tion ; kinetics of formation.
1. Introduction
In recent years there has been increasing interest in intranuclear electron transfer
processes especially for identifying and isolating long-lived dinuclear intermediates.
The studies by Jwo and Hairn (1976) and those of Haim and Sutin (1976) are
typical examples wherein complex fcrrocyanides were employed. If the dinuclear
complex is to be long-lived, the thermal electron transfer reaction should be prefe-
rably absent or be very slow. This paper presents the isolation of a cherry red
dinuclear complex formed in the reaction between cis-a diaquo Co(III) triethylene
tetramine and ferrocyanide, and a study of its behaviour.
2. Experimental
2.1. Preparation
The cherry red complex was prepared as follows : The cis-a dichloro Co(III) trien
chloride was prepared and purified according to the method suggested by Sargeson
and Searls 1967. The cis-a, diaquo complex was prepared by dissolving the
*To whom correspondence should be made.. 213
dichloro complex in water and allowing it to stand for 24 hr for aquation. When
the aquo complex was mixed with aqueous solution of K 4 Fe (CN) 6 a cherry red
colour was immediately formed. This solution showed an electronic absorption
at 440 nm and a pH between 7 and 8. Equal volumes of equimolar solution of
the two reactants were mixed and allowed to stand for an hour at room tempe-
rature with occasional stirring. On adding ethanol and keeping under refrige-
ration, bright cherry red solid separated out. After filtering, the solid was
repeatedly washed with ice-cold ethanol and dried under vacuum in a desiccator.
2.2. Chemical analysis
The solid was dissolved by heating first with nitric acid, then with perchloric acid
and finally with sulphuric acid. The K, Co and Fe content was determined by
standard analytical techniques. K was determined by flame photometric method.
Fe and Co were estimated spectrophotometrically by ortho phenonthroline (Vogel
1961) and nitroso-R-salt (Sandell 1959) methods respectively.
2.3. Instrumental
A Perfcin- Elmer IR spectrophotometer was employed for recording the IR spectra
of the sample using hexachloro butadiene as the mulling agent. A Stan ton TG
balance with a sample size of 400 mg and a heating rate of 6 K min" 1 in static
air was employed. Electronic spectra were recorded with a Hitachi Model
spectrophotometer. A home-made constant acceleration Mossbauer spectrometer
using 3 mci Co-57 (pd) in conjunction with an MCA was used to record the 67 Fe
Mossbauer spectra of the solid.
3. Results and discussion
The results of chemical analysis of the cherry red] solid for Fe, Co and K con-
tent (table 1) show a reasonable fit for the molecular formula K Co(trien)Fe
(CN) 8 . 4H 2 O consistent with the expectation from the *oute followed for its
synthesis.
. The cherry red complex on dissolution in water satisfies Beer's law. The molar
extinction coefficient is 1000 litre mole" 1 cnr 1 and its stoichiometry is found to
Table 1. Results of chemical analysis of cherry red solid.
Element Weight (%)
Observed Calculated
K 7-4 7-4
Fe 10-5 10-6
ui wuuuuucuu^ vajuauuu cum
results support the above-mentioned molecular formula for the complex arrived
at from chemical analysis. Its stability constant was found to be 4-5 x 10 s litre
mole- 1 . To ascertain whether the colour of the complex is vested in the cationic
or anionic moiety of the solid, its solution was passed through the Na form of
a cationic exchanger, which did not retain the colour. The moiety can either be
neutral or anionic in nature. A similar treatment with the chloride form of
an anionic exchanger resulted in the retention of the colour on the exchanger
proving that the moiety responsible for the colour is anionic in nature.
The TG analysis of the cherry red complex was performed with the object of
probing the thermal stability of the solid complex as well as to determine if all
the four molecules of watsr are bound equally strongly from its dehydration beha-
viour. The observed TG run is given in figure 1. The thermogram shows a very
slow.and gradual loss in weight in the temperature interval 313-873 K with a super-
posed very sharp loss in weight around 480 K. The overall loss in weight between
31 3-873 K accounts for loss of one molecule of trien and four molecules of H a O.
However, if one calculates the loss in weight represented by the steep loss alone,
Weight of sample : 400mgs
Rate of heating : 6C/min
static air
Container : platinum make
313
773 873
Temp. (K)
water molecules in 3 to 1 ratio. Considering that ferrocyanide was used as the
starting material and in the light of the TG and ion exchange results mentioned
above, the following structural units are proposed for the cherry red complex.
K+ [(H 2 0) (trien) - Co - NC - Fe (CN) 5 )]- 3H 2 O.
This does not disturb the ferrocyanide unit and involves a single linear CN
bridge, which is well documented as far as cyanides are concerned.
This type of structural unit gets independent support from IR results. The IR
spectrum shows an absorption at 2040cm" 1 with a clear shoulder at 2100cm" 1 .
The former is attributed to stretching of terminal CN groups whereas the latter
comes from that of bridging C==N groups (Nafcamoto 1963). A band at 585 cnr 1
attributable to Fe-C stretching in ferrocyanide was also observed in the IR
spectrum of the cherry red solid.
It is reported (Nakamoto 1963) that in solids both co-ordinated as well as lattice
water molecules give characteristic IR absorption in the range 3550-3200 cm- 1
(due to antisymmetric and symmetric O-H stretching) and also in the range
1630-1600 cm- 1 (due to H-O-<H bending). On the other hand, the co-ordinated
H 2 O molecules give an additional absorption in the region 880-650 cnr 1 due to
rocking mode of the water molecule. In the IR spectrum of our complex at room
temperature, there are absorptions in the neighbourhood of 3000cm- 1 and
1600 cnr 1 in addition to several sharp ones in the range 600-850 cm,- 1 . It is diffi-
cult to unequivocally discern the absorption arising out of rocking modes of H a O
in this group of peaks. The spectrum in this 600-850 cnr 1 region remains
unaffected after heating the sample at 473 K, i.e., after the steep loss of 3H 2 O
molecules, as indicated by the TG scan. It implies that the lor.e co-ordinated
water molecule is lost after 473 K.
The foregoing results enable us to propose the following structure for the cherry
red complex.
NH-
CN
CN
HN-
HoN"
NC
-CN
NC'
OH 2
CN
3H 2
In order to investigate the origin of colour of this cherry red complex, the
electronic spectra of solution of the complex in water with different concentra-
tions were taken. The spectra are shown in figure 2. Two well-defined maxima
at 330 and 440 nm were observed. Vogler and Kunkely (1975) have reported the
electronic spectra of dinuclear complex [(CN) 5 Co (III) KCFe (II) (CN) 6 ] 8 - as
having two maxima at 326 and 385 nm. They assigned the 326 nm band to
Co(III)ligand flsld separation, and the 385 nm band to intervalence charge transfer
process involving the two cations. On the same lines, in our case, the 440 nm
1-0
0-8
0-2 -
300
400 500
Wavelength (nm)
600
Figure 2. Absorption spectra of solutions obtained by the dissolution of different
amounts of cherry red solid in water.
band is assigned to charge transfer involving the interaction of the two metals
of the anionic species. The 330 nm band is assigned to the Co(III) ligand field.
According to the qualitative potential energy diagram constructed by Vogler and
Kunkely (1975), the 440 nm band observed here represents a spin allowed
in tervalence transi ti on .
It would be worthwhile and interesting to probe the extent of electron transfer
in these charge transfer complexes. A bulk magnetic susceptibility measurement
would not be a very sensitive tool to do this since such a transfer is very small
in practice and the macroscopic measurement could not detect it. A microscopic
technique which looks more closely at either cobalt (III) or iron (II) site is
desirable. At least for iron site, it was felt that Mossbauer might throw some
light on the situation.
It is well-known that the two extreme situations in which the 6th d electron
of Fe either stays on it or goes over to Co completely far away from it can be
easily distinguished by Mossbauer spectroscopy from different isomer shifts
observed for Fe(II) and Fe(III) cases. With this in mir.d, we recorded the 67 Fe
Mossbauer spectra of the cherry red complex and to our surprise observed that
The first step in the formation of the cherry red complex is the constitution of
an ion-pair.
KI
[Co(trien) (H 2 O) 2 ] 3 + + [J
[(trien) (H 2 O) 2 Co. . .Fe [CNe]- 1 (1)
The- rate of formation of this ion-pair is very fast, so that KL compared to
KI can be neglected in practice. If this ion-pair gives the final cherry red product
through a uaimolecular reaction, the next step can be logically visualised as
follows :
K*
[(trien) (H 2 O 2 ) Co ... FeCCN),]- 1 ^
KL Z
[(trien) (H 2 O) Co NC. Fe(CN) 5 ]- 1 + H 2 O (2)
cherry red complex
The slowest step in this process is the removal of water molecule and hence
it is the rate-determining step. The overall reaction rate depends on the con*
centration of the ion-pair, which in turn depends on the concentration of the
starting ions, viz.,
[Co(III) trien] 8 * and [Fe(CN) 6 ] 4 "
The rate expression may then be written as
dc/dt = K[Co (trien) (H 2 O) 2 ] 3 + [Fe (CN) 6 ]*- (3)
whers K
Data pertaining to kinetic studies namely values for CD at 440 nm vs time
plots presented in figure 3 show that in diaquo complex, it is difficult to resolve
the semilog plots. In order to investigate if the two *i /2 's can be resolved, the
starting diaquo complex of cobalt was replaced by the corresponding dichloro
complex of cobalt. The resulting kinetic data plotted in figure 3 also show that
for dichloro Co (III) trien complex the semilog plots are resolvable into two linear
components having distinct f 1/2 , values. This clearly shows that in dichloro
complex its aquation tates place prior to the interaction with ferrocyanide.
A linear plot of dc/dt (figure 4) against the product of concentrations of the
reactants indicates an overall second order kinetics. The reaction rate and t l/z
for the system [9 x 10- 4 MCo(III) trien complex and 9 x 10~ 4 M ferrocyanide]
were measured at temperatures 301, 308, 318 and 333 K to estimate E a , the acti-
vation energy of the reaction. A linear plot of log K vs \\T was obtained. The
activation energy for the interaction of c?s-a diaquo Co(III) trien complex with
ferrocyanide was found to be 28-8. kJ.
Ferrocyanide concentration = 6x10 M
80 120
Time (min)
7-5 12-5
Figure 3. Plot of log (OD^, OD t ) vs. time for cis-a diaquo/dichloro cobalt(III)
trien complex and ferrocyanide system. Curve (1) diaquo cobalt(Ul) complex and
ferrocyanide system. Curve (2) dichloro cobalt(lll) complex and ferrocyanide
system.
Cobalt complex
concn. 5x10 M
4 8 12
(Cobalt complex] [Ferrocyanide] x 10 7 M
Cobalt complex _ 4
concn, 3 x 10 M
024
[Diaquo cobalt (III) Trien complex]
[Ferrocyanide] x 10 M
Figure 4. Plot of dc\dtvs. [Diaquo Co(Ill) trien complex] [Ferrocyanide].
calculations of N-methyl substituted acrylamides
G RAMANA RAO
Department of Physics, University College, Kakatiya University, Vidyaranyapuri,
Warangal 506 009, India g.
MS received 11 February 1982
Abstract. CNDO calculation is made for N-methyl acrylamide (both in as and
trans configurations) and N,N-dimethyl acrylamide. The charges, bond orders and
dipole moments are discussed and compared with those of acrylamide. The trans
form of N-methyl acrylamide is found to be more stable than cis isomer by 4-5
kcal/mole.
Keywords. CNDO calculation ; acrylamides ; dipole moments ; bond orders,
Introduction
is wellknown that in amide molecules, no single valence bond structure is
isistent with all their 'properties. This is due to the de-localization of the
bonyl Tr-electrons and lone pair electrons of nitrogen, resulting in the partial
ible bond character of the C-N bond (Morris and Orville Thomas 1961 ;
nnon 1955, 1956 ; Venkata Chalapathi and Venkata Ramiah 1968, 1971).
N-methyl substituted acrylamides, in addition to nitrogen lone pair, ethylenic
lectrons also take part in such conjugation. The effect of such conjugation
5 been studied in the case of acrylamide (AC) by CNDO method (Ramana
and Venkata Ramiah 1975). The aim of the present paper is to make similar
dies on N-methyl acrylamide (NMAC) and N,N-dimethyl acrylamide (DMAC)
1 to examine the effect of the above conjugation as well as the effect of N-methy-
!on in AC on charge distributions, bond orders, dipole moments and barriers
internal rotation about the C-N bond.
Method of calculation
e CNDo/2 calculations (Pople and Beveridge 1971) of AC, NMAC (both
and trans forms) and DMAC are accomplished by programme QCPE 142
fNDO (Dobosh). The structure of the molecules is presented in figure 1.
e structure parameters used are y (C = C) = 1 -337 A, y (C = O) = 1-243 A,
C-N) = 1-315 A, y(C-C) = l-47A, y (N-H) = 1-02A, y (N-C) = 1-47 A,
Figure 1. Structure of acrylamide (J?j = R$ H), N-methylacrylamide (Ri CH 3 ,
/J 2 =H) and N,N-dimethylacrylamide (R^ = R 2 = CH 3 ) 6 is the angle the dipole-
moment makes with C-N bond and $ is the angle of rotation of amino group
from the rest of the molecular plane.
are assumed to be 120 each. The atomic orbitals used in the calculations are
Is for hydrogen, 25, 2p for carbon, nitrogen and oxygen.
3. Results and discussion
3 1 Charges, bond orders and dipole moments
The total charge (rr + or), net pi-charge and mobile bond orders of NMAC
and DMAC are presented in table 1. The results of AC are given for comparison.
The net total charges on oxygen, functional carbon and nitrogen in AC are
0-371e, 4 0-358e and 0-241e respectively. On successive methylation of
AC, the net total charges on oxygen and 'functional carbon atomsr emain almost
unaltered whereas that on nitrogen atom shows a considerable decrease in magni-
tude as can be seen from table 1. The total charge in the amide system (i.e.,
peptide bond) in AC, NMAC and DMAC is 0-254e, 0-212e and 0-179e
respectively. Thus the net total charge in the peptide bond is negative and the
magnitude of the charge decreases with successive methylation of AC. This
result is in agreement with that reported for primary, secondary and tertiary
amides (Van et al 1970).
In AC, NMAC and DMAC the pi-bonding between nitrogen lone pair and
carbonyl group results in de-localization of nitrogen lone pair electrons. The
pi-orbitals of ethylenic bond also take part in such a conjugation. The net pi-charge
on nitrogen atom is 0-185e in AC. The corresponding charge in NMAC and
DMAC on nitrogen atom is 0-218e and 0-249e respectively. Thus we observe
that the de4ocalization of nitrogen lone pair increases with successive methylation
of AC. However, the sigma charge (this is the total charge minus net pi-charge)
on nitrogen atom remained almost constant.
The vinyl substitution results in a small amount of pi-charge of about 0-036e
being transferred from the ethylenic bond to carbonyl bond in AC which remains
unaltered on successive methylation of AC. On the other hand, the carbonyl
bond receives quite a large amount of pi-charge from nitrogen lone pair, the
amount of charge being 0-185e, 0-218e and 0-245e in AC, NMAC and DMAC
llia rntrnor/vn I/M-IA -noif e^hai-na -n1a\ra a /^/-n-inont
o o o
o
$
fl ^ cO
o*>
^
?) OO CO OO
Y P* C4 C4
oo
CS
U 6 o o
6
% 1 S
1
1
U o o o
o
s
Ooo in in
o o o
II oo oo oo
in
O
oo
^666
o
^T" v*i m r^
j^jv Vj \J s5
**o
o
S *? ^ 7
^
VI
u I I '
i
&
in oo oO
0\
"o
^. OO r 1 T-H
^1*
-a
r. ^
n
'~ N *
rn
O OOO
U rl n r^
Os
* i
O T-1 "'
* 9 -j-
rt
rt
1
tn
to
3
c?T ^ m '- 1
Tj-
43
u ~ ^
a
^
ON
1
S
V-<
g T T T
T
a
i
"3
"o
T-l r-l CO
M
o
1
13 "*
^J' ON oo
1
1
VH
13
o
o
o
1
S ^ S
^^ *n in ^*^
U f*"' f*" 1 rr>
S
3
fl
o
IS
P 1 f? oo
O rn to <r,
s
8
s
1 1 I
i
VI
4-*
TJ
'
s
o
*
<^N
gj
=0
.s
V-i
^ -2
c!
*o
tf
.3
*w^ *^*^
J< 1
s
>.
w -2
o
M 1 t"J
<=!
"WJ
Y-f
s
1 1
s
"
O
f^J
cC
Cv
i
o
(D CQ ^
S
1
1st
^S "S ^
5 5 fL
(ounp-Nft.
S
ja
U
#
those of the related amide molecules. As such, the vinyl substitution does
not have much effect on the amide system. This agrees with the result already
obtained on the basis of normal coordinates analyses (Ramana Rao and Venkata
Ramiah 1981).
The mobile bond orders of these molecules are also shown in table 1. The
C-N and C-C bond orders have considerable double bond character whereas C = O
and C = C have reduced double bond character. This is because the carbonyl
bond receives w-charge both from nitrogen lone pair and ethylenic ?r-bond. The
bond orders of C-N and C = O are close to the values reported for formamide
(Pople and Segal 1966). Further, the bond orders of three molecules consi-
dered are almost the same.
The calculated dipole moments for AC, NMAC (trans), NMAC (cis) and DMAC
are 3-81, 3-80, 3-71 and 3-63 debyes respectively with corresponding directions
(as shown in figure 1) as 54, 55 24', 55 33 and 55 10'. These values are close
to one another and are also comparable to those of formamide and acetamide
which are 3-71 and 3-82 debyes respectively (Yan etal 1970).
3.2 Barrier to internal rotation and cis-trans isomerism
By assuming the transition state to correspond to < = 90 (see figure 1) in accor-
dance with the findings in the case of alkyl substituted amides (Yan et a! 1970),
the barrier to internal rotation about C-N bond of AC, NMAC and DMAC have
been calculated to be 20, 17-5 and 11 kcal/mole respectively. The experimental
value of barrier height in DMAC is 16-8 kcal/mole (Rogers and Woodbrey 1962).
The barrier heights in formamide, N-methyl formamide and N,N-dimethyl
formamide are calculated to be 20-3, 18-5 and 15-03 kcal/mole respectively by
CNDO/2 method (Yan et al 1970). The corresponding experimental values are
19-2, 18-05 and 21 kcal/mole (Stewart and Siddal 1970). One can see that on
successive methyl substitution the calculated barriers have decreased in acrylamides
and also formamides. Further, the calculated values of barriers in N,N-dimethyl-
amides are lower than the experimental values.
The charge density in the lone pair orbital of nitrogen in the ground state
(<f> = o) is l-815e, l-782e and l-762e in AC, NMAC and DMAC. The correspond-
ing quantities in the transition state (fi =90) are l-966e, l-935e and l-904e.
That is the lone pair charge is getting localized in the transition state as it increases
to a value close to 2e from a value around 1 8e in the ground state. This shows
that the de-localization of nitrogen lone pair is responsible for the barrier to a
large extent. The same result was obtained in the case of thioamides (Srinivas
Rao and Venkata Ramiah 1976).
NMAC can exist in both cis and trans forms. Our calculations show that the
trans form of NMAC is more stable than cis form of NMAC by 4- 5 kcal/mole.
4. Conclusions
(ii) ine deiocauzation 01 nitrogen lone pair increases witn successive metny-
lation of AC whereas the sigma charge remains almost unaltered.
(iii) The nitrogen lone pair charge plays an important role in determining the
properties of AC, NMAC and DMAC as the pi-charge received by carbonyl
bond from nitrogen lone pair is several times greater than that received from
the ethylenic bond.
(iv) The mobile bond orders of C = O, C N, C C and C = C remain
almost unchanged on successive methylation in AC.
(v) The dipole moments of AC, NMAC and DMAC are of the same magnitude
as for formamide and acetamide.
(vi) The calculated barrier heights for AC, NMAC and DMAC are found to
decrease on success ive methylation of AC as is the case with formamide
series. The trans form of NMAC is found to be more stable than the cis
form of NMAC by 4-5kcal/mole.
Acknowledgements
The author is thankful to Prof. K Venkata Ramiah, Member, Union Public
Service Commission, Government of India, and to Dr L Srinivas Rao, Reader
in Physics, Osmania University, for helpful suggestions.
References
Cannon C G 1955 Mikrochlm. Acta 2 555
Cannon C G 1956 J. Chem. Phys. 24 491
Dobosh Paul A Quantum chemistry programme exchange, Indiana University, Bloomington,
Indiana
Morris R E and Orville Thomas W J 1961 /. Mol. Spectrosc. 6 572
Pople J A and Beveridge D L 1971 Approximate molecular orbital theory (New York : McGraw-
Hill)
Pople J A and Segal G A 1966 J. Chem. Phys. 44 3289
Ramana Rao G and Venkata Ramiah K 1975 Indian J. Pure AppL Phys. 13 569
Ramana Rao G and Venkata Ramiah K 1981 Indian J. Pure AppL Phys. 19 232
Rogers M T and Woodbrey J C 1962 /. Phys. Chem. 66 540
Srinivas Rao L and Venkata Ramiah K 1976 Pramana 6 244
Stewart W E and Siddal T H 1970 Chem. Rev. 70 517
Venkata Chalapathi V and Venkata Ramiah K 1968 /. Mol. Spectrosc. 26 444
Venkata Ramiah K and Venkata Chlapathi V 1971 Curr. Sci. 40 365
Yan S E, Momany F A, Hoffman R and Sheraga H A 1970 /. Chem. Phys. 74 420
Proc. tndian Acad. Sci. (Chem. Scij, Vol. 91, lumber 3, June 1982, pp. 229-229
Printed in India.
Red edge excitation and proton association in the excited
state of acridine
P GANGOLA,* N B JOSHI and 1> D PANT
Department of Physics, DSB College, Kumaun University, Naini Tal 263 002, India
MS received 14 July 1981 ; revised 25 January 1982
Abstract. A comprehensive study of acridine spectra with variation of pH, wave"
length of excitation, deuteration of the solvent, etc., has been made. The excited
state protonation of acridine is found extra-ordinarily excitation wavelength sensitive
n^ar the red edge of the first absorption band. The proton association takes place
very fast (K PT ~ 10 10 sec" 1 ) on excitation at the red edge of the first absorption band
(REE) and acridinium emission is observed while it is slow on short wavelength
excitation (SWE). The reaction rate slows down at lower temperature which is
indicated by a delay in the initiation of the effect by ~8 nm on REE. The acri-
dinium type emission with REE at 80 K shows that proton tunnelling is the chief
mechanism of proton transfer. The quantum yields are also found wavelength
dependent. Contrary to previous observations acridinium ion also shows a REE
shift at 80 K.
Keywords. Red edge excitation ; proton transfer ; acridine.
1. Introduction
Despite the well-accepted generalisation that in large molecules luminescence
occurs only from the vibrationally equilibrated lowest electronic excited state,
three major violations have been observed in some aromatic molecules in solu-
tions on excitation near the red edge of the first absorption band. The first of
these relates to change in polarisation of emission (Weber 1960 ; Weber and
Shinitzky 1970 ; Valeur and Weber 1978), the second to a slight change in wave-
length (Chen 1967 ; Fletcher 1968) and the third to a change in chemical reactivity
of the excited state (Gangola et al 1977, 1979 ; Pande et al 1980 ; Shah et al
1980). In molecules like acridine which undergo a large change in pK value in
the excited state, the rate of proton transfer changes on red edge excitation (REE).
Weller (1957) who studied acridine emission with only the short wavelength exci*
tation (SWB) did not observe any proton -abstraction in alcoholic solution and
of acridine in neutral, alkaline and acidic solutions at low and room temperatures.
2. Experimental
In view of weak emission on RBB, careful attention was given to the elimination
of the stray light, impurity emission extraneous or due to ionised species of the
solute and solvent Raman scattering. Spex fluorolog model 1902 has two sepa-
rate double monochromators each for excitation and emission and the stray
light problem was not present.
Cbromatographed (courtesy CDRI, Lucknow) acridine was prepared from
Fluka acridine base material and was similar to those obtained from Eastman
Kodak and CIBA Research Centre, Bombay. Methanol solutions (10~ 4 M) were
prepared at pH ~ 9. Absorption, emission and excitation spectra were taken
with Spex fluorolog. The spurious effect produced by Raman scattering at
extreme RBB was ^0-8%, hence was neglected. Acridinium absorption was
absent even in non-alkaline solutions taken with 10cm cell in Beckmann DK-2A
spectrophotometer. The excitation spectra were taken with a more dilute solu-
tion (~ 10~ 5 M) by monitoring at 420 nm and 480 nm emission peaks respectively
of SWE and RBB. The excitation spectrum of the solvent by monitoring at
480 nm (due to Raman scattering) was also obtained separately to make allowance
for distortion in the actual excitation curve.
All experiments reported here are performed with the chromatographed sample
mostly in methanol. If a different solvent is used in any particular experiment,
it is mentioned in the text. Methanol (BDH) was used after distillation. CH 3 OD
(isotopic purity > 95%) obtained from BARC, Bombay, was used as such. The
alkalinity of the solution was varied by adding NaOH (GR grade) and H 2 SO 4
(Anala R) and the pH of the solution was measured by a Phillips pH meter.
3. Results and discussion
In acidic methanol solution (pH ~ 1) only the absorption and emission spectra
of acridinium are observed. The emission spectrum at 295 K is a little diffuse,
however, at 80 K a well-structured spectrum is observed with a temperature blue
shift of ~ 400cm- 1 . The lowest excited states of acridine are nn* (Landner and
Becker 1963) ; *L a located at 384 nm (e max - 3 x 10 3 ) slightly lower than the
^ which is at 358 nm (e max - 10 x 10 a ). The ^and^j states separate out in
acridinium, the *L U (e max - 3-6 x 10 3 ) band is very much red shifted while 1 L J
(^max 17 x 10 3 ) is not displaced but intensified (figure 1). The emission and
absorption spectra of acridinium agree well with those given by earlier workers
(Landner and Becker 1963 ; Mataga et al 1956, 1957 ; Kellmann 1977 ; Wbitten
and Lee 1971). The emission spectrum at 295 K is definitely independent of
excitation wavelength. However, at 80 K there is minor red shift ~ 250cm.- 1
on excitation by /l ex = 446 nm without change in the structure (figure 2). While
the absence of red edge effect at 295 K is in keeping with the observation of
300
380
Wavelength (nm)
460
Figure 1. Absorption spectrum (1) and excitation spectrum (2) of acridinium in
acidic methanol (pH /% 1) at 295 K.
Fletcher (1068), the mino* red shift in acridinium emission at 80 KL was not
observed earlier. The magnitude of the shift and the shortest wavelength of
excitation which initiates the shift appear to be pH dependent. With pH ~
the magnitude of the shift A* increases to 500cm- 1 compared with 250cm.- 1 at
pH ~ 1. The effect is discernible with /l ex =440nm at pH~0 whereas it
starts with A ex = 446 nm at pH -^ 1. The emission spectra of acridinium (pH ~ 0)
at 80 K for different wavelengths of excitation are given in figure 3.
At pH ~ 7 absorption at longer wavelengths is absent indicating the absence
of acridinium ions in the ground state and the blue fluorescence of acridine is
observed. The absorption and emission spectra of acridine also agree with those
given by other workers. The fluorescence bands become sharper at 80 K with a
very small temperature blue shift of ~ 100 cnr 1 . In this solution there is appa*
rently no wavelength shift on RBB at 295 K as well as at 80 K, an observation
440 520 600
Wavelength (nm)
680
Figure 2. Emission spectra of acridinium in acidic methauol (pH ~ 1) with
d)350nm at 295 K, (2) 350 nm at 80 K and (3) 446 nm at 80 K.
10 -
400
480 560
Wavelength (nm)
640
Figure 3.
with % n
Emission spectra of acridinium in acidic methanol (pH
(1) 350 nm, (2) 430 nm, ' (3) 440 nra and (4) 446 nm.
0) at 8C K
50 -
30
o
0)
CC
10
300 340 380
Wavelength ( nm)
420
Figure 4. Excitation and absorption spectra of acridine in alkaline methanol
(pH ~ 9) at 295 K ; excitation spectra with emission monitored at (1) 420 nm and
(2) 480 nm and (3) absorption spectrum.
E 30
.3
O
20
'w
c
jjj
i i
10
400 480
Wavelength (nm)
560
Figure 5. Emission spectra of acridine in methanol (pH
(1) 350 nm, (2) 440 nm and (3) 4lOnm.
7) at 295 K with
Around pH ~ 9 in alkaline solutions in methanol or ethanol or ethanol *
methanol (1 : 1), although only acridine exists in the ground state, the fluorescence
emission now shows edge excitation effect definitely. On adding alkali there is
not the slightest change in absorption or SWB emission spectrum of acridine but
on REE acridinium type emission is observed. REE effect in the pH range 7-5
to 9- 5 was investigated and acridinium emission was studied. The effect was
best observed in pH range 8 to 9 after which again the relative intensity of acrji
Table 1. Excitation wavelength dependence of emission bands of acridine and
acridmiura at 295 K and 80 K.
295 K
Acridiite (in alkaline methanol)
SWE (JL. = 350) 400, 420, 440, 465, 398, 418, 444, 468,
50Q 506, 544
REE (A, x = 420) 448, 476, 500, 535 444, 472, 500, 535
Acridinium (in acidic methanol}
SWE U OI = 350) 452, 480, 500 444, 470, 500, 535
REE (/l ex = 446) no change, structure 482, 515, 555
improves
350 nm excitation (SWE) usual fluorescence of acridine is observed. On excitation
with longer wavelengths around 390 nm, however, some subtle changes in the
spectrum start appearing, finally leading to acridinium type emission on excitation
by 415 nm. There is no further change in the structure of the spectrum on
excitation with 420 nm (REE) except for an additional red shift of 100 cmr 1
The subtle changes appearing on excitation by A ex ~ 390 nm may be due to a
transient state of H + in the solvation cage of acridine or it may be due to an
admixture of acridine and acridinium spectra. The intensity ratio of 475 nm
and 445 nm emission bands goes on increasing by excitation with longer wave-
lengths and finally beyond 415 nm excitation only the acridinium type spectrum
appears. The REE spectrum of acridine at 295 K, however, has a much better
resolved structure than the acridinium spectrum at the same temperature. In
fact, it has the same spectral structure as the acridinium spectrum at 80 K observed
on SWB with A^ ~ 472 nm and vibronic spacing Av e ~ 1270 cnr 1 .
The edge effect has also been investigated at 80 K in alkaline methanol +
ethanol (1 : 1) glass. This solvent is chosen because it forms a clear and unc racked
glass at 80 K, however, in alkaline methanol also similar results were obtained.
At 80 K, the vibronic structure of the emission spectrum is well resolved &v,
~ 1270cm" 1 and the edge effect can be studied carefully. Figure 7 shows the spectra
obtained with alkaline methanol + ethanol (pH ~ 9) glass at 80 K. The usual
acridine spectrum is observed with SWB. Perceptible change starts appearing
with 403 nm excitation which is seen not as a change in the emission wavelength,
but as a relative change in the intensities of the bands. On excitation with 415 nm
acridinium type emission but slightly blue shifted from SWB acridinium emission
is observed. Finally, with REE (420 nm) purely acridinium emission both in
wavelengths of the bands and their intensity distribution is obtained. Although
the complete change from acridine to acridinium spectrum is observed on REE
at 420 nm both at 295 K and 80 K, the intensity change starts appearing at a
shorter wavelength of excitation at 295 K. This is evident from the curve 4
360
440 520
Wavelength -(nm)
600
Figure 6. Emission spectra of acridine in alkaline methanol (pH ~ 9) at 295 K
with A iS :. (1) 350 nm, (2) 395 nm ; (3) 400 nm, (4) 405 nm and (5) 420 nm.
70 -
360 440 520
Wavelength (nm)
60O
295 K. The pH dependent wavelength effect in acridinium at 80 K may also
be due to the proton transfer rate being wavelength dependent. However, for
acridinium the heterogeneity of the emitters might cause similar effect. On the
other hand, it is also known that the presence of other H+ or OH~ ions greatly
affects the hydrogen bonding or ion-pair formation (Zundal 1976).
Deuterium effect at 190 K was observed in CH 3 OD alkaline solution. With
SWB the spectra are identical in wavelength and intensity distribution in both
CH 3 OH and CH 3 OD solutions and the emission is due to acridine. Intensity
changes on REB at 190 K appeared to be slower in CH 3 OI> solution. However,
some uncertainty existed in the pH value of CH 3 OD which did not allow a careful
assessment of deuterium effect. At 295 K, however, the two spectra were identical
on REE. We believe that the rate of proton transfer becomes faster at 295 K
and the solvent isotope effect is discernible only at low temperature.
The acridinium excitation and absorption spectra are given in figure 1. The
absorption and excitation spectra agree almost quantitatively in short wavelength
region, however, in the long wavelength transition the quantum yield increases
by about 34%. The excitation and absorption spectra for acridine in alkaline
methan'ol solution are given in figure 4. Here again there is an increase of 34%
in the quantum yield value for long wavelength transition. On the long wave-
length part (REE), however, the excitation curve 1 monitored at 420 nrn (acridine
emission peak) goes below, while the curve 2 monitored at 480 nm (acridinium
emission peak) passes above the absorption curve. Normally one should have
expscted both the excitation curves to lie above the absorption curve as the
quantum yield for long wavelength transition, i.e., X L, is higher. The fall in quantum
yield as seen by curve 1 clearjy indicates that acridine is being converted into
acridinium in the excited state. The quantum yields ($4) for aoridine and acridi"
nium (as measured in aqueous, solutions on A ex = 350 nm) are 0-38 and 0-55
respectively, i.e. ^ is higher by 44-7% than j> A . Calculated on this basis, if
acridinium is being formed in the excited state by REE, curve 2 sliould reveal a
further increase of 44-7% over the initial 34%, thus a total increase of 93-9%
should be observed. Quantitative measurements from enlarged curves show that
the total increase in revealed by the excitation curve monitored at 480nmis
~ 93% neap the long wavelength region, although error in these experiments is
large (~ 20%). Thus we believe that these results though only semiquantitative
can be explained if acridin?um is formed in the excited state. The absence of red
edge effect at pH ~ 7 and its. observation at pH ~ 9 also rules out the presence of
acridinium in the ground state to be the origin of the effect.
The nitrogen of aza-anthracene becomes highly basic on excitation. However,
Weller on excitation with 365 nm did not observe any change in acridine emission
although an excited state protonation was expected. Weller ascribed the failure
of proton association to slow rate of reaction (K PT ~ 10 G sec- 1 ) compared to the
lifetime of acridine molecule (16ns). O.ir results obtained on REE, however,
.suggest that the protonation reaction increases around 400 nm, and around
410 nm the reaction is so fast (K PT ~ 10 10 sec- 1 ) that only acridinium type of
emission is observed.
are present in the ground state we expect the ratio of acridinium and
emissions to be
where (A ex ) denote the extinction coefficient at the wavelength of excitation, Ctte
concentration and & 5 the quantum yield.
Substituting the numeric?! values at pH = 9 we obtain a theoretical value of
4 x 1(H, whereas the actual observed value is 4 x 1Q- 2 . Thus the presence of
acrddtnium ions in the ground state is ruled out. An exciplex of acridine with
OH also does not explain the red edge effect as in more alkaline solution the
red edge emission diminishes. The structured emission may not be expected in
an exciplex. Further, the effect is observed in frozen glass which cannot be due
to the formation of exciplex. A ground state complex of acridine with OH~ is
unlikely and is not observed in absorption or in excitation. As has already
been mentioned, the non-alkaline solution shows a different type of red edge
effect in as much as there is no change in the wavelength of emission but the
structure of the bands improves. This sharpening of the structure is different
from what has been observed in solid solutions or glasses due to heterogeneity
of the emitting centres (Rudik and Pikulik 1971). However, in the solution state
also we may assume various configurations of solute-solvent complexes and as
suggested by Itoh and Azumi (1975) due to selective absorption at the red edge
sharpening of the bands may result. However, the peculiarity about REE emis-
sion in alkaline solution at 295 K is that it resembles the low temperature acri-
dinium emission in acidic solution. The REE emission at low temperature is
also almost similar except that the room temperature REE emission is shifted to
the red by a small amount ~ 170cm." 1 , It is also interesting to note that the
acridinium type REE emission itself shifts slightly on increasing the wavelength
of excitation. This is apart from the subtle intensity changes which start appearing
around 390 nm excitation. The temperature dependence of the reaction rate
is an additional factor for our assumption of proton transfer in the excited state.
Recent observations on proton reactions in the excited state have revealed that
several types of hydrogen bonded complexes can occur in the excited state in the
same solvent. They can be weakly or strongly hydrogen bonded, ion-pairs or
completely protonated complexes. If it is assumed that K PT is wavelength depen-
dent at the red edge, many subtle changes occurring in the emission spectra can
be qualitatively understood on this basis.
The protonation reactions in the excited state are generally diffusion controlled
and/or there is proton tunnelling. While the diffusion-controlled reactions are
viscosity dependent, the tunnelling rates are only temperature dependent. In any
protonation reaction both the processes may be present simultaneously or exclu-
sively any one may be present. The fact that the red edge effect is seen in frozen
glass shows that the tunnelling effect is the predominant mode of proton transfer.
Even the inconclusive isotope effect points towards the tunnelling effect. If it is
assumed that the temperature dependence of the effect is indicated by an initiation
JL * " *
a tunnelling process of proton transfer between the solvent oxygen atom and
the nitrogen atom of the acridine.
The acridine molecule forms hydrogen bonded complexes, as shown below
with solvent molecules
H
O
CH 3
The hydrogen bonded proton can occupy two positions one near the oxygen
of the alcohol and the other at nitrogen of acridine. In the latter case it behaves
like a protonated molecule and/or an ion-pair complex. Due to the polarization
of hydrogen bond by the nature of the medium and other factors many other
stages may be possible. For our discussion we assume two potential minima
for the proton in the excited state, the one with proton near N having a lower
energy. According to Lowdin (1965) the rate constant for tunnelling of a particle
of mass m through a double minimum potential barrier can be expressed as
K =
where v is the hit frequency, i.e., the number of times per second that the particle
falls incident on the barrier, F , a a are respectively the height and width of the
potential barrier. K is the fraction of energy measured from the top of the barrier.
At low temperatures most of the protons will be at the potential minimum while
at higher temperatures there will be quite a large fraction of molecules in the
higher energy region. Therefore at 80 K we may assume K = 1 and the rate of
tunnelling becomes slower assuming that there are no other changes at low tempe-
ratures in the configurational diagram. We believe that the effect of OH" ions
on REE may also be due to changes in symmetries of the potential minima
(Zundal 1976). This change may also account for the sharpening of the bands
at room temperature because as the asymmetry increases the low frequency
modes vanish. A tentative explanation for the change in proton transfer rate
viscosity dependent, or based, on tunnelling can be had by suggesting that during
the reaction process either promoting or non-promoting modes are active at
REE. We suggest that after vibrational relaxation from the Franck-Condon state
on SWB in a time ~ 10~ 12 sec the excited molecule has enough time ~ 10~ 8 sec
to be equilibriated with the vibrationless excited Si state. If cage relaxation,
orientational relaxation and other geometrical distortional relaxations are not
completed within this time, a slight difference in energy may still be revealed in
SWE acridine type of emission is observed, which changes to acridinium type on
REE near (0, 0) band. It is, thus, not a simple shift as is observed in some of
the edge excitation red shift experiments reported by earlier workers. In recent
years many photochemical reactions, quantum yields, lifetimes and other non-
radiative processes (Avouris et al 1977 ;. Rice 1974 ; Freed 1976) have been
found to be dependent on the wavelength of excitation and certain broad features
for these phenomena seem to be explicable on the basis of theories of non-radiative
processes. However, the enormous difference in the proton transfer reactivity
of the molecule relaxing to vibrationless excited singlet state and an energetically
similar molecule prepared on absorption of the corresponding radiation frequency,
is, however, not easily understood. Bsrhaps this situation is analogous to wave-
length dependent non-radiative processes in aromatic molecules, e.g., the third
channel in benzene (Jacon et al 1977).
Acknowledgement
Authors are grateful to the University Grants Commission, New Delhi, for financial
assistance.
References
Avouris P, Gelbart W M and EUSayed M A 1977 Chem. Rev. 77 793
Chen R F 1967 Anal Biochem. 19 374
Fletcher N 1968 /. Phys. Chem, 72 742
Freed K F 1976 Top. Appl. Phys. 15 (ed.) F K Fong p. 168
Gangola P, Joshi N B and Pant D D 1977 Chem. Phys. Lett. 51 144
Gangola, P, Joshi N B and Pant D D 1979 Chem. Phys. Lett. 60 329
Itoh K and Azumi T 1975 J. Chem. Phys. 62 3431
Jacon M, Lardeaux C, Lopez-Delgado R and Tramer A 1977 Chem. Phys. 24 145
Kellmann A 1977 /. Phys. Chem. 81 1195
Landner S J and Becker R S 1963 /. Phys. Chem. 67 2481
Lowdin O P 1965 Adv. Quantum Chem. 2 213
Mataga N, Kaifu Y and Koizumi M 1956 Bull. Chem. Soc. Jpn. 29 373
Mataga N, Kaifu Y and Koizumi M 1957 Bull. Chem. Soc. Jpn. 30 368
Pande U, Joshi N B and Pant D D 1980 Chem. Phys. Lett. 72 209
Rice S A 1974 Excited states 2 (ed.) E C Lim (New York : Academic Press) p. 232
Rudik K I and Pikulik L G 1971 Opt. Spectrosc. 30 147
Shah J, Joshi N B and Pant D D 1980 Curr. Sci. 49 609
Shapiro S L and Winn K R 1980 /. Chem. Phys. 73 5958
Valeur B and Weber G 1978 J. Chem. Phys. 69 2393
Weber G 1960 Biochem. J. 76 335
Weber G and Shinitzky M 1970 Proc. Natl. Acad. Set. (USA.) 65 823
Weller A 1957 Z. Electrochem. 61 956
Whitten D G and Lee Y J 1971 /. Am. Chem. Soc. 93 961
Zundal G 1976 The hydrogen bondVL (eds.) P Schuster, G Zundal and C Sandorfy (Amsterdam;
North Holland Pub. Co.) p. 283
Micellar catalysed chlorination of acetophenone by chloramine-T
V RAGHUNATHAN, P S RAGHAVAN, K VAfDYANATHAN
and V S SRINfVASAN*
Department of Chemistry, Ramakrishna Mission, Vivekananda College,
Madras 600 004, India
MS received 21 November 1981 ; revised 31 March 1982
Abstract. The chlorinPtion of acetophenone by chloramine-T CAT has been
catalysed by anionic micelle, sodium lauryl sulphate (NaLS). Though the order
in CAT is one, the order in acetophenone is fractional at lower concentration and
becomes zero at higher concentration in the presence and absence of NaLS. This
is probably due to the change in rate-determining step. At lower acetophenone
concentration, the decomposition of enol-chlorinating species complex is rate-deter-
mining whereas at higher concentration, the formation of chlorinating species is
rate determining. The graph of k z versus detergent concentration is sigmoidal and the
positive co-operativity versus log[>] graph is 1-11, indicating possible interaction
between micelle and substrate.
Keywords. Chlorination ; acetophenanc; anionic micelle ; NaLS effect ; positive
co-operativity ; chloramine-T.
1. Introduction
Chlorination of acetophenone using chloramine-T (CAT) has been investigated
by Balasubramanian and Thiagarajan (1976) in HC1O 4 medium and the reac-
tion is an acid catalysed one. The present work is the result of our earlier finding
that the surfactant, sodium laurylsulphate (NaLS), catalyses the chlorination of
phenols (Rengarajan et al 1980) and amines (Raghavan et al 1980) by CAT.
The present paper attempts to study the niicellar influence on the chlorination of
acetophenone by CAT.
The experiments have been so patterned such that the nature of the transi-
tion state and its stabilization in the ttv.cellar phase can be understood and a
formal comparison can be attempted with enzyme catalysed reactions. Due
to the limited solubility of acetophenone in water, all the reactions were carried
out in 20% HOAc-80% H 2 O (v/v) at 30 C.
2. Results and discussion
The kinetics ofchlorination of acetophenone by CAT has been investigated in the
binary solvent mixture of acetic acid-water in the presence of 0-20 M HC1O 4 .
2*1 Dependence of rate on [CAT] and [Acetophenone}
The dependence of rate on [CAT] has been determined by varying the initial con-
centration of CAT at a given concentration of acetophenone, NaLS, HC1O 4 and
solvent composition. The disappearance of CAT follows first order knetics both
in the presence and absence of NaLS (table 1). The reaction rate depends on
acetophenone concentration at lower levels. The order with respect to ketone
is fractional but at a higher concentration of ketone, the reaction exhibits zero
i order kinetics (table 1). Similar trends have been observed in the chlorination
of acetophenone (Balasubramanian and Thiagarajan 1976) and phenol
(Balasubramanian and Thiagarajan 1975) by CAT.
2.2. Dependence of rate on detergent concentration
There is a gradual increase in rate with increasing concentration of anionic
micelle, NaLS and this is marked above the initial micelle concentration 0-010M
(figure 1). This can be traced to a possible hydrophobic interaction of the phenyl
ring with hydrocarbon core of the micelle while the side chain prefers the Stern
layer of the micelle. The chlorinating species may be oriented such that the
attack on the substrate becomes more facile. Similar catalysis by NaLS has been
observed in the chlorination of anilines (Raghavan et al 1980) and phenols
(Rengarajan et al 1980) by CAT.
Table 1. Dependence of rate on [CAT] and [Acetophenone].
[HClOi] = 0-20 M; [NaLS] = -8-0 x ]Q- 3 M; 20% HOAc-80% H 2 O (v/v) 30 C
[CAT] MX JQ 8 [Acetophenone]
M X 10 a
fcj x IQ^ec" 1
1-50
4-0
5-5
2-0
4-0
5-8
4-0
4-0
5-7
2-Q
Q-8Q
2-6
2-Q
1-00
2-8
2-Q
1-50
4-2
2-Q
2-0
5-5
2-Q
3-Q
5-6
2-0
4-0
5-8
2-Q
5-0
5-6
10 20 40 60
10 3 [NoLS).
80
100
Figure 1. Dependence of rate on detergent concentration.
2.3 Temperature influence
The temperature influence over the reaction rate has been studied in the range
of 30 and 50 Cj and from the plots of log k z versus 1/T, the activation energies
have been calculated. The enthalpies and entropies of activation thus evaluated
are summarised in table 2.
The lower activation energy observed in the micellar phase is in keeping with
the rate enhancement noted in the presence of micelle.
2.4 Solvent influence
The rate of chlorination of acetophenone is affected by change in polarity of
the solvent and the rate increases with increasing percentage of acetic acid
(table 3).
3. Mechanism of chlorination of acetophenone in NaLS
As the order in acetophenone is fractional which then approaches zero at higher
concentration, this reaction probably proceeds via the formation of a complex
between the enol of acetophenone and CAT, exhibiting Michaelis-Menton type
of kinetics. The above observations can be explained by the following reaction
sequence ;
O OH
O
CAT
OH
K,
CgHs C
chlorinating species
C fl H 5 C = CH 2 + chlorinating species v^ complex
complex - products
WJiere A, ana .iv are composite ui me ^uubuuus jv ls j\ z , j\$ cum ,j. /-
concentration of acetophenone, the decomposition of the complex seems to be
rate-determining, exhibiting fractional order dependence on acetophenone. At
higher concentration of acetophenone, the formation of active species from CAT
seems to be rate-determining, exhibiting zero order dependence on keton^. It
is further evidenced by the insensitivity of the rate to structural variation in
acetophenone J both ^-methyl and p-methoxy react with the rate comparable to
the specific rate of chlorination of acetophenone itself at 0-040 M.
Table 2. Temperature dependence and therniodyitamic parameters.
[Acetophenone] = Q 04 M ; [CAT] - 002 M ; [HCIO 4 ] = 2 M ; 20% HOAc-80% H 3 O (v/v)
Temperature C
(NaLS)
/q x 10 s Sec- 1
M x 10 3
30
4-5
30
8-0
5-8
40
tl-8
40
8-0
12-2
50
25
50
8-0
29
(a)
(b)
E a k.cal/molo
14-2.
13-2
Aff* kcal/mole
13-6
12-6
A 5* kcal/mole
-Q-033
-0-036
(a) in the absence of NaLS; (b) in the presence of NaLS.
Table 3. Dependence of rate on solvent composition.
[Acetophenone] =0 -04 M; [CAT] = 0-002 M; [HC1O 4 ] =Q-G2M; [NaLSj =0'008 M;
Temp. 30 C.
Solvent composition
%HOAc-%H 2 O
20-80 5-8
30-70 7-1
40-60 8-5
4. Theoretical treatment of micellar catalysis
4.1. Positive co-operativity
On the basis of the mathematical model proposed by Bruice et al (1968)
nD + S ^ D,,S
U M
KQ >. products
where n is the number of detergent molecules (D), D n S is catalytic micelle, K D
is the dissociation constant of micelle, k m and k are the rate constants in the
micellar and aqueous phases respectively, the rate expression obtained is
= log [D] log Kom
From a graph of
lo g r~ bs k ~] versus log [D], (figure 2) the slope, ,
\_K-m "-oljsj
obtained is 1-11 which is considered as index of co-operativity (Piszkiewicz 1977).
As the n value is greater than 1, it is referred t.o-as. positive c o-operativity which
implies the stimulation of the interaction of additional substrate molecules, by the
interaction of the first molecule with nrcelle : K D value obtained in this case is
2-1 x 10~ 2 M.
4-2. Binding model
By a scheme similar to the one considered above and taking into account the
CMC of the detergents, the rate expression derived is,
C D is the d;tergftnt concentration and N is the aggregation number of the deter-
gent. From a graph of 1/(fc fl . - fc oba ) against l/(C fi - CMC), the K/N value
obtained is 15-8 M (figure 2).
5. Experimental
All liquid organic compounds were distilled using glass apparatus, rejecting head
and tail fractions. HCIO 4 (E Merck) was standardised, after dilution, against
carbonate-free sodium hydroxide solution. HaLS is purified by washing with
anhvdroiK ether anrf th^n r^nrvsta lit sine thrw: tr ffvi* tim^s with Q<; / rtKa-n rt i 1-411
0-6 -
20 40 60
1/pD-CMC]
-2-0
-1-5
log[D]
-1-0
Figure 2. Plots for Piszkiewicz and binding models.
The reaction rate has been followed by estimating unreacted CAT iodometri-
cally at various intervals of time. The rate constants were evaluated by the
numerical method using appropriate integrated rate expression. The specific
rates are reproducible to within 4%. The stoichiometry of the acetophenone-
CAT reaction is 1:1. The product of the reaction has been identified as
phenacyl chloride by carrying out the analysis on a preparative scale.
'Acknowledgement
One of the authors (PSR) thanks the CSIR, New Delhi, for the grant of Senior
Research Fellowship.
References
Balasubramattian V and Thiagarajan V 1976 Aust. J. Chem. 29 1449
Balasubramanian V and Thiagarajan V 1975 Int. J. Chem. Kinet. 7 605
Bruice T C, Katzhendler J and Fedor L R 1968 /. Am. Chem. Soc. 90 1333
Piszkiewicz D 1977 /. Am. Chem. Soc. 99 1550
Raghavan P S, Srinivasan V S and Venkatasubramanian N 1980 Indian J. Chem. 19A 322
Rengarajan K, Vaidyanathan K, Srinivasan V S and Venkatasubramanian N 1980 Indian J.
Chem. 19A 79
Reactions of Moles with mercury (II) salts
AVIJIT BANERJI* and MANJUSHA SARKAR (nee CHAUDHURI)
Department of Pure Chemistry, University College of Science, Calcutta 700009*
India
MS received 29 August 1980 ; revised 16 February 1981
Abstract. A number of mercurated indoies were prepared. Their spectroscopic
data (uv, iR, ^-NMR and 13 C-NMR) are reported. Our results contradict certain
data appearing in previous reports on similar compounds. When heated in acetic
acid merourated compounds decomposed, although the products formed were not
identical with those obtained earlier from the same substrates by the action of
th?llium(III) acetate in acetic acid. Reaction of l-methylindolc-3-mercuriacetate
with styrene in the presence of lithium tetrachloropalladate furnished a product
which was tentatively assigned the structure 9-methyl-l,3-diphenyl-l, 2,
3, 4-tetrahydrocarbazole.
Keywords. Indoles ; mercuration ; organomercurials.
1. Introduction
In the course of our work on the reactions of heterocycles with metal salts ;we
investigated the reactions of indole and substituted indoles with thallium(HI)
acetate (Banerji and Ray 1978). A variety of oxidation products were obtained
whose structures depended on the substituents at the 2- and 3-positions of indole.
The results could be explained by assuming initial thallation of the indole to
a 3-thallated indolenine, which reacted in situ to yield the products. IB no case
were we able to isolate the organ o-thallium derivatives.
In contrast it was reported earlier (Mingoia 1930; Ramachandran and Witkop
1964; Yudin et al 1971) that mercury(II) salts react with indoles to give fairly
stable isolable organo-mercori derivatives. The complete characterisation of only
a few of these has teen described. Also, in the above-mentioned reports, there
exists a great deal of confusion regarding the structures of these compounds as
well as their spectroscopic data. Hence, we decided to re-investigate various
aspects of the indole-HgX a (X - -OAc, -Cl) reaction with special reference
to the spectroscopic properties of the products. The results of this investigation
are reported in the present paper.
* To whom correspondence should be made.
247
a-u room lemperaiure vjs.iimacn.ar.aran HI,U YVIIK.OP JL^VHJ. J.JD.C
properties of the products obtained by using (i) one molar, (ii) two molar and
(iii) excess of mercuric acetate, were very similar to those reported by Rama-
chandran and Witkop who characterised these as ths " mono ", " di- " and " tri-
acetoxymercuri " derivatives respectively. The elemental analyses, however,
showed that none of these products was pure. Products from (i) and (ii) con-
tained less nitrogen than the formulation suggested by Ramachandran and Witkop.
Hence it appears that mercuration occurred to a greater extent than the " mono"
and " di-acetoxymercuri " stage. It was not possible to accurately check mercu-
ration at these stages. As all the products were extremely insoluble in nature
they could not be purified by crystallisation.
l-Mbthylindole-3-mercuriacetate (I), l-methylmdole-3-mercurichloride (II) and
l,3-dimethylindole-3-mercuriacetate(III) wereobtained following the method of
Yudin et al (1971) while 3-methylmdole-2-mercuriacetate (IV) was also prepared
by the method of Ramachandran and Witkop. These compounds could be
obtained in pure state and gave the expected elemental (C, H, Nj analyses.
The uv spectra of all the compounds were measured in ethanol and the data
arc given in table \. Our results indicate that for compounds (I-IV) the absorp-
tion maxima suffered only a small but significant hypsochromic shift whereas the
extinction coefficients were virtually unchanged. The mercury substituent has
two vacant />-orbitals and exerts a -R conjugative effect on the aromatic nucleus
which causes the observed hypsochromic shift in the uv spectra. Similar obser-
vations for phenyl- thienyl-. and furyl-mercury derivatives were noted earlier
(Leandri and Tundo 1954). Ramachandran and Witkop (1964) had reported the
uv spectral data of what they formulated as " mono- ", " di/' and " tri-acetoxy-
mcrcuri " indoles in about 2 -5% acetic acid. These workers reported marked
changes, in the portion of absorption maxima as well as very large enhancements
in extinction coefficients when compared with the starting materials. Our results
contradict these reports. The absorption maxima of the "mono-" (^S H 218,
272, 277-5, 288 nm), and " di^cetoxymercuri " (^ H 218, 272, 277-5, 288 nm)
compounds showed the same small hypsochromic shift compared to indole
(>S H 216, 266, 276, 287 nm) as observed for the other compounds. Moreover,
no .perceptible shifts of the absorption maxima of these two compounds, or in
fact of any of the other mercurated products, were observed when the medium
was changed from ethanol to 2-5% acetic acid. The extinction coefficients for
the two products obtair.ed from indole calculated on the basis of mono- and di-
acetoxymercuration are considerably less than thos.e expected. This is in confor-
mity with the elemental analyses which showed the occurrence of mercuration
beyond the desired stage. - - -
We report here the IR spectra of the mercurated indoles for the first time.
The important bands and their assignments are listed in table 2. The spectra
were recorded in Kujol mull using polyethylene discs. The C-Hg stretching bands
appeared at 400-415 cm" 1 for all the compounds. When the IR spectra were
recorded in KBr pellets, additional bands; appeared at 335-340 and 505-550 cnr 1
. - - . .. . . ...
spectrum of l,3-dimethyl-2-acetoxymercuri-ir.dole is given in figure 1, with
the more important assignments. The different impure mercurated products
("mono-", "di-" and " tri-acetoxymercuri '') from indole also showed C-Hg
stretching at 400-415 cnr 1 .
NMR investigation of mercurHndole was also carried out (table 3). Some
!H-NMR data had been earlier reported by Ramachandran and Witkop (1964)
using trifluoroacetic acid (TFA) as the solvent. They made certain structural
assignments on the basis of these observations; in particular the " di-acetoxymcr-
curi" derivative was assigned the structure 2,3-diacetoxymercuri-ir.dole. It seems
that no reliance can be placed on these data for the following reasons: Firstly
in TFA medium the indoles will certanly exist as 3-protonated indolenine species
rather than free species; Secondly in the strong acid medium used demercu'-ation
would easily occur so that the reported values may not refer to the mercurated
compounds at all.
"We found that for the l-methyl-3-acetoxymercuri-indole, the C-2 proton and
N-methyl group were slightly de-shielded with respect to 1-methyl indole. There
were also changes in the signals for the benzenoid protons. In the parent hetero-
cycle all these protons appeared as a complex multiplet centred around 57-2,
whereas in the mercuri-derivative one of thsse moved downfleld to appear as a
double doublet at 5 7- 50 (J Q = 6-l Hz, J m = l5Hz). This signal was presu-
mably due to the C-4 proton, which is shielded by the 3-acetoxymercuri grouping.
The 20 MHz 18 C-NMR of I methyl-3-acetoxymercuri-indole was recorded in
DMSo-^/ 6 . Tentative assignments, based on single-frequency off-resonar.ce
decoupling (SFORD) multiplicities and chemical-shift theory are given in figure 2.
18 C-NMR spectra of the other compounds could not be recorded on account
of their poor solubility in the more common deuterated solvents.
Wavelength
7 8 9
0-disubstituted c-Hg
benzene stretching
3500
1800
1400 1000
WdvenumberCcnf 1 )
600 300
Figure 1. IR spectrum of l,3-dimethyl-2-acetoxy-mercuri-:ndole. recorded in
23-47
Figure 2. 20 MHz 13 C- NMR Spectral data (CDCl a ) (in ppm) of l-methyl-3-
acetoxy mercuri-Jndole.
Table 1. uv absorption maxima of mercurated indoles (in 95 % ethanol),
Compound (in EtOH)
c)innm
2-Acetoxymercuri-skatole
Skatole*
l-Methyl-3-acetoxymercuri-indole
l-Methyliadale-3-mercuu-chloride
1-Methyliadole*
226(4-32), 290(3-90), 296(3-90)
222(4-50), 275(3-73), 282(3-78), 2-90(3-69)
222(4-58), 272-5 (3-90), 284(3-89), 295(3-80)
222(4-58), 270 (sh) (3 -76), 283(3-77), 295(3-67)
219 (4-54), 275 (3-77), 282(3-78), 293(3-66)
1 ,3-Dimethyl-2-ac<3taxy-mercui:i-indole 228 (4 43), 292 5 (3 99)
1,3-Dimethyl-indole* 225 (4- 50), 248 (3-72), 278 (3 68)
ladole* 216(4-54), 266(3-76), 276(3-76), 287(3-68)
* Houlihan (1972).
Table 2. IR absorption bands of mercurated indoles (in nujol mull recorded in
polythene discs).
v (in cm" 1 )
fVminOinnfl
\^\J LJL*.\J V- U14U
Acetate*
o-Disubsti-
tuted Hg-O
Hg-C
>C= 3 O -O-C-
benzene
stretching
2- A" Jtoxymercuri-skatole
1580 1010
725 680
405
l-Methyl-3-acetoxyinercuri-
iadole
1600 1005
730 690
400
1 , 3-Dimethyl-2-acetoxy-
mercuri-indole
1600 1010
750 700
415
l-]VEethylindole-3-mercuri-
chloride
... ...
750
410
Table 3. 80 MHz
spectral data of mercurated indoJes.
Compound
Signal at 5 (ppm)
Solvent
-OCOCH 3 Aromatic protons Other protons
Mona-acetoxy-mercuri- DMSO-de 1 87
6-70-7-80
-CH 3 merged with
skatole
DMSO peak
l-Methylindole-3- CDC1 3 2-05
6-95(lH,br,s)
-N-CH 3
mercur i-acetate (3 H , s)
7-08-7-30
3-76(3H,s)
(3H,m)
7-50(lH,dd,
/ =6-lHz,
/ m =l-5Hz)
l,3-Dimethylindole-2- CDC1 3 2-01
6-94-7-37
-N-CH 3
mercuri-acetate (3H,s)
(4H,m)
3-37(3H,s)
C-CH 3
2-21 (3H,s)
l-Methylindole-3- DMSO-rf 6
6-94-7-30
-N-CH 3
mercuri-chloride
(3H,m);
7-38(lH,dd,
3-74(3H,s)
/ =8-8 Hz,
jr.-l-SHz;
7-65(lH,dd,
/ a =7-7Hz,
J m = 2-6Hz)
As mentioned earlier, the reaction of ir.doles with TTA in acetic acid gave
oxidation products. It was of interest to find out whether the mercuri-jndoles
would also be converted to similar products under similar or more vigorous
reaction conditions. It was found that the mercurated compounds remained
largely unaffected when subjected to reaction conditions similar to those used for
the TTA reaction. When more vigorous conditions were used, none of the pro-
ducts previously obtained with TTA was formed though some, demercuration to
the original compounds occurred. Most of the material underwent decompo-
sition to intractable products. In the case of mono-acetoxymercurHndole,
oxindole was formed in low yield on heating with glacial acetic acid for 4 hours.
2.1. Reaction of \-methylindole-3-mercuri~acetate with styrene
The palladium salt-catalysed olefto arylation reaction provides a very convenient
route to a wide variety of olefinic compounds (Hecfc 1968). We attempted the
of unstable products were formed, of which only one of the major compounds
could be isolated in a reasonably pure state by preparative TLC. None of the
others could be isolated by TLC as they underwent decomposition. The isolated
product, which exhibited a typical indolic uv spectrum W H 225, 277-278
(sh) nm), was tentatively assigned the structure 9-methyH,3-diphenyH,2,3,4-tetra-
hydrocarbazole (3) primarily on the basis of its ^-NMR spectrum. A variety
of possibilities exist for alternative combinations with styrene, of which only that
shown was compatible with the ^-NMR (see experimental for chemical shifts)
(figure 3).
The formation of the product (3) can be mechanistically rationalised as shown
in scheme 1.
3. Experimental
All milting points are u ".corrected and were determined on an electrically heated
Kofhr Block molting point apparatus. The uv spectra were measured on a Varian
634S spectrophotometer in 95% aldehyde free ethanol. The IR spectra were
'J (n (d)
(e)
(f)
& 2-81' (2H, d, .J= 4 Hz)
8 3-24 (3H, K.S)
63-75 ' (IH , m)
&B-83-7-27 (14 H, br. ,
unresolved signal)
S1-20 (2H,m)
Figure 3. 80MHz 1 H-NMR spectral data of 9-methyl-l, 3-diphenyl-il,2,3,4-tetra-
hydrocarbazole.
Scheme -1*
PdCl
PhCh=CH-
.Ph
recorded in polythene disc on a Beckman IR-20 spectrophotometer. The 1 H- and
i 3 C-NMR spectra were recorded on Varian Associates CFT-20 NMR
spectrometer.
3.1. 2 -Acetoxymercuri-skatole
Mercuric acetate (1 59 g, 5 mmol) in absolute ethanol (15 ml) was added in
portions to a magnetically stirred solution of skatole (0-66 g, 5 mmol) in absolute
ethanol (10 ml) whereby a pale yellow dispersion was formed. The mixture was
stirred fcr 2 hr and left overnight. The pale yellow precipitate was filtered and
washed thoroughly with absolute ethanol (yield M g, 57-9%); m.p. 180 (d).
Found C 35-5%, H 3-1%, N 3-6%; C n H 13 O 2 NHg requires C 35-7%, H 3-2%
H 3-6%.
3.2. 1 Methyl-3-acetoxymercuri-indole and l^-dimethyl-l-acetoxymercuri-indole
1-Methylindole and 1,3-dimethylindole were prepared according to Hear.ey and
Steven (1973).
To a stirred solution of the methylated indole (5 mmol) in absolute ethanol
(10 ml), mercuric acetate (1-59 g, 5 mmol) in absolute ethanol (15 ml) was added
and left overnight. The white precipitate was filtered off, washed and dried.
3.2a. l-Methyl4-acetoxymercuri-irtdole. Yield l-7g, 89-4%; m.p. 187 (d).
Found C 34-2%, H 2-2%, N 3-5%; C 11 Et 1 O 2 NHg requires C 33-9%, H 2-3%,
N3-6%.
3.2b. l^-Dimethyl'l-acetoxymercuri-indole. Yield 1-1 g, 55-0%; m.p. 144-6.
Found'C 35-6%, H-3-2%, N 3-7%, C 12 H 13 O 2 NHg requires .C 35-7%, H 3-2%,
N 3-4%.
3.3. 1 -Methylindole-3 -mer cur ichlo ride
1-Methylindole (0-66g, 5 mmol) and sodium acetate (l-6g, 20 mmol) were
dissolved in anhydrous methanol (10 ml), and mercuric chloride (1-36 g, 5 mmp])
in methanol (15ml) was added. A white precipitate appeared instantar.eoufy,
which was filtered, washed and dried (yield 1-5 g, 83-3%), m.p. 170(d) (Yudin
et al 1971, 170-3). Found C 35-6%, H 1-9%, N 3-7%; C 9 H 8 ClKHg requires
C 35-6%; H, 1-9%, N 3-8%.
3.4. Reaction of l-methylindole-3~mercuriacetate and styrene in presence of
To a mixture of l-methylindole-3-mercuriacetate (0-97g, 2- 5 mmol) and styrene
(0-6 ml, > 5 mmol); 0-1 (M) Li 2 PdCl 4 (50 ml, 5 mmol) in methanol was added
dropwise over a period of 1 hr and then stirred for 48 hr in an atmosphere of
nitrogen. The mixture was filtered, the solvent removed and the residue was
extracted with methylene chloride. The residue from the methylene chloride
extract was rmrifled bv column chromatosraphy and PTLC. 9-meth.yl-l,3-diDhenvl-
80 MHz NMR spectral data of (3) (in CDC/ 3 ): (a) 82-81 (2H, rf,/ =
(b) 53-24 (3H, broad s); (c) 53-75 (1H, m); (d) 54-44 (1H t, / =
(e)56-83~7-27 (4H, broad unresolved signal), (f)51-20 (2H, m).
3 . 5. Mono -acetoxymer cur i-indole
To a stiired solution of indole (0-59g, 5 mmol) in absolute ethanol (10ml);
mercuric acetate (1-59 g, 5 mmol) in absolute ethanol (15ml) was added portion-
wise. The reaction mixture was allowed to stand overnight, filtered, the residue
washed with ethanol and dried (yield l-2g, 66-6% on basis of mono -mere ura-
tion); m.p. 266 (d) [Ramachandran and Witkop 1964, 270 (d)]. Found
N2-2%, C 10 H 9 2 NHg requires N3-7%.
3.6. Di^acetoxymercuri-indole
Di-acetoxymercuri indole was prepared similarly using two moles of mercuric
acetate (3-17 g, 10 mmol). The product was isolated in a similar manner (yield
2-7g, 98-1% on basis of di-mercuration); m.p. 205 (d) [Ramachandran and
Witkop 1964, 205 (d)]. Found Nl-9%; C 12 HnO 4 NHg 2 requires N 2-2%.
3 . 7. Tri-acetoxyme-curi-indote
Indole was mercurated with excess of mercuric acetate (6-36 g, 20 mmol) following
the same procedure. A precipitate rapidly formed and then redissolved. The
alcohol was removed under reduced pressure. The residue was triturated with
water, filtered and then washed successively with water, ethanol and ether
(yield 0-5 g, 10-7% on basis oftri-mercuration); m.p. 268-9 (d) [Ramachandran
and Witkop 1964, 270 (d)]. Found N 1-8%, C 14 H 13 O 6 NHg 8 requires N 1-6%.
Acknowledgement
One of the authors (MS) is grateful to CSIR, New Delhi, for financial support.
References
Banerji A and Ray R 1978 Indian J. Chem. B16 422
Heck R F 1968 J. Am. Chem. Soc. 90 5518
Heanly H and Steven V L 1973 /. Chem. Soc. Perkin Trans. 1 499
Leandri G and Tundo A 1964 /. Chem. Soc. p. 3377
Mingoia Q 1930 Gazz. Chim. Ital. 60 509
Ramachandran L K and Witkop B 1964 Biochemistry 3 1603
Yudin L G, Kost A N and Pavlyuchenko A I 1971 Khim. Geterotsikl. Soedin 7 1517
Houlihan W J (ed.) 1972 The chemistry of heterocyclic compounds Part 1 (New York ;
Wiley Interscience) p. 21
Quantitative structure activity relationships Part V. Release and
uptake of norepinephrine in murine heart by phenethylaimnes
VIJAY GOMBAR
Department of Phamaceutical Sciences, Panjab University, Chandigarh 160014,
India.
MS received 24 September 1981
Abstract. Quantitative structure activity studies have been carried out on a series
of hydroxyphenethylamines. The calculated Fujita-Ban group contributions indi-
cate that the m and jp-hydroxyphenethylamines have high affinity for uptake and
efflux of radioactive norepinephrine. The highly negative contribution of the
OH group at the second ortho position indicates that the derivatives withhydroxyl
groups at both the ortho positions should have no activity or extremely low activity
as inhibitors of uptake and as releasing agents.
Keywords, Fujita-Ban calculations ; norepinephrine ; phene'thylamines.
1. Introduction
The in vivo effects of various compounds on uptake and release of radioactive
norepinephrine have been studied in detail. The results of a series of investi-
gations by Axelrod eta! (1961, 1962), Hertting etal (1961, 1962), Daly etal (1966)
and Creveling et al (1966, 1967, 1968, 1974) include structure-activity relationships
concerning inhibition of norepinephrine uptake at plasma membrane and its
displacement from the storage sites. The different behaviours of ortho-hydroxy-
phenethylamines and ortho-hydroxyphenethanolamines towards inhibition of
uptake of radioactive norepinephrine (Rotman etal 1975) provide some insight
into the structural requirements of uptake sites.
The present paper embodies quantitative structure-activity studies concerning
the effects of hydroxyl groups from various positions in phenethylamine on in vivo
inhibition of uptake of [ 8 H] norepinephrine into, and its release from, murine heart,
2. Method and data set
For phenethylamine(I) and its hydroxy derivatives Rotman et al (1975) have assayed
inhibition of uptake and release of radioactive norepinephrine by measurement of
Part I ; Singh etal (1980), Part II ; Gombar et al (1951), Part III ; Jain and Gombar (1981),'
Part IV ; Gombar and Wadhwa (1982).
?55
me amount 01 i~i-j-jLu J -norepineparine. jooin inese acuvmes nave oeen
expressed in terms of ED 60 (mol/kg).
In the present work the observed activity, however, refers to log ED 50 where
ED SO is taken in mol/kg units. The present quantitative structure-activity studies
are carried out in the light of the Fujita-Ban de novo model (Fujita and Ban 1971)
because the purpose of this study is to quantify the effect of hydroxyl group at
different positions in the aromatic ring of phenethylamine. The group contri-
butions have been claculated as per Kubinyi's algorithm of converting the Fujita-
Ban matrix into normal equations matrix followed by simultaneous equations'
solution by any of the standard methods (Kubinyi 1977). A computer program
FUJKUB, fully incorporating this algorithm was developed by the author and used
for the present work. All data were processed on the DECSYSTEM-2050 at the
Regional Computer Centre (North), Chandigarh, India.
3. Results
The structures, observed activities and Fujita-Ban matrix for phenethylamines
exhibting inhibition for uptake of norepinephrine are given in table 1. This infor-
mation about phenethylamines eliciting release of norepinephrine is collected in
in table 2. Both these sample sets are statistically unbiased as the observed activity
is found to vary about two log units which corresponds to 100-fold difference in
activity. Tables 3 and 4 respectively contain the normal equations matrices corres-
ponding to the Fujita-Ban matrices in tables 1 and 2. In tables 3 and 4 // represents
contribution by the unsubstituted arbitrary reference (phenethylamine) and Ax
represents the contribution by the group X. The group contributions calculated
in the two cases (figure 1) lead to the following correlations.
- log ED 50 (uptake) = 5-80-0-40 [2-OH] + 0-36 [3-OH]
+ 0-27 [4-OH] + 003 [5-OH] -1-10 [6-OH]
n =15 ; r = 0-806 ; j = 0-416 (1)
_ logED 5 o (release) = 3-87 - 0-39 [2-OH] + 0-36 [3-OH]
+ 0-50 [4-OH) + 0-42 (5-OH)
M= ^14; r==0 -755; 5 = 0-468 (2)
It is observed that these correlations are significant almost at 95 % level.
4. Discussion
4-1. Inhibition of uptake of norepinephrine
It is evident from the contributions of hydroxyl group at different positions
(figure la) that phenethylamines with -OH substituent at m or/and p-positions
have relatively high' affinity as in vivo inhibitors of the uptake 'of norepinephrine.
It is obvious from the minus sign of the group contribution of 2-OH that the
presence of an orthohydroxy group in phenethylamine would reduce the inhi-
bition affinity. In (1) the negative sign and fairly high magnitude of the] contri-
bution by 6-OH group further reveals that phenethylamines with hydroxy groups
occupying both the ortho positions must be extremely weak inhibitors. This is
in accord with the results reported by Rotman et al (1975). In, most of the
phenethylamines bearing two ortho hydroxy groups they did not observe any
inhibition even for concentrations as high as 100^mol/kg.
Table 1. Structures, Fujita-Ban matrix and activity of phenethylamines investi-
gated for in vivo Ithibition of a [H] norepijtephrine uptake into murine heart.
CH 2 CH 2 NH 2
Compound R
No.
Independent substituents Activity
2-OH 3-OH 4-OH 5-OH 6-OH Obsd.
Calcd.
1
H
5-42
5-80
2
2-OH
1 4-86
5-40
3
3-OH
1 6-46
6-16
4
4-OH
1 6-16
6-07
5
2,3-(OH) 2
1 1 5-93
5-76
6
2,4-(OH) 2
1 1 5-77
5-66
7
2,5-(OH),
1 1 5-63
5-40
8
3,4-(OH) a
1 1 6-49
6-43
9
3,5-(OH) a
1 1 6-21
6-17
10
2,3,4-(OH),
111 6-20
6-03
11
2,3,5-(OH; 8
1 1 1 5-92
5-76
12
2,3,6-(OH) 3
11 1 4-66
4-66
13
2,4,5-(OH) 3
1 1 1 6-14
5-66
14
3,4,5-(OH) 3
11 1 ' 6-31
6-43
Table 2. Structures, Fujita-Ban. matrix and activity of phenethylamines investi-
gated for in vivo release of 1 3 H] rtorepinephriae from murine heart.
Compi
Nc
Ii dependent
-blind T?
Substitueats Activity
JUUU IV
2-OH 3-OH
4~OH 5-OH Obsd. Calcd.
1
H
3-29 3-87
2
2-OH 1
3-22 3-47
3
3-OH 1
4-64 4-23
4
4-OH
1 4-42 4-37
5
2,3-(OH) a 1 1
4-23 3-83
6
2,4-(OH) a 1
1 4-19 3-97
7
2,5(OH) a 1
1 3-70 3-89
8
3,4-(OH), 1
1 4-72 4-73
9
3,5-(OH) 2 1
1 4-94 4-65
10
2,3,4-(OH) 3 1 1
1 4-13 4-33
11
2,3,5-(OH) 8 1 1
1 4-20 4-25
12
2,4,5~(OH) S 1
1 1 5-15 4-39
13
3,4,5-(OH) 3 1
1 1 5-00 5-15
14
2,3,4 > 5~(OH) 4 1 1
1 1 4-10 4-75
Table 3. Normal equations matrix for norepinephrine uptake inhibition Studies.
15 9 9 7 6 1 to
87-39
995441 VOH
50-35
959441 A 8 _ OH
= 53-43
744730 VOH
42-31
644360 AS-OH
35-45
111001 A 8 ^ H
4-66
Quantitative structure activity relationships
Table 4. Normal equations matrix for norepincphrine release studies.
259
14 8 8 7 6
.0
59-92
88444
A 2
32-91
84844
VoH
=
35-95
74473
VoH
31-71
64436
Vj
27-09
* See table 3.
CH 2 CH 2 NH 2
(0-36)
CH 2 CH 2 NH 2
(0-27)
V.
(o)
Figure 1. Fujita-Ban group contribution of -OH group at different positions i
phenethylamine as a. inhibitors of uptake, b. releasors of norepinephrine.
in
4-2. Release of norepinephrine
The pre-requisite for release of norepinephrine is its active uptake into the
neuronal terminal though transport of amine into the cell after binding to the
uptake site is an important factor. Therefore, the compounds with poor affinity
for uptake are expected to be poor releasors of norepinephrine as well. This is
quite evident from (2). The contribution of 2-OH group is once again negative
(figure Ib). This amounts to say that phenethylamines with -OH group at
ortho position are poor releasing agents. These phenethylamines have already
been shown to have poor affinity for uptake ( 4- 1). It can be further seen that
2,3,6-trihydroxyphenethylamine which has both the ortho positions occupied by
-OH groups, has extremely low affinity for uptake and should be expected to be a
still poorer releasing agent. Indeed, this compound does not figure in table 2
as Rotman etal (1975) observed that ED 5n for this compound was verv hieh
Graveling C R, Daly J W and Witkop B 1968 /. Med. Chem. 11 595
Creveling C R, Luudstrom J, McNeal E T, Tice L and Daly J W 1974 Mol. Pharmacol.
Daly J W, Creveling C R and Witkop B 1966 /. Med. Chem. 9 273
Fujita T and Ban T 1971 /. Med. Chem. 14 148
Gombar V, Kapoor V K and Singh H 1982 Arzneim. Forsch. 32 7
Gombar V and Wadhwa L Arzneim. Forsch. (in press)
Hertting G, Axelrod J and Patrick R W 1961 Biochem. Pharmacol. 8 246
Hertting G, Axelrod J and Patrick R W 1962 Br. J. Pharmacol. Chemother. 18 161
Jain D V S and Gombar V 1981 Int. J. Quantum Chem. 20 419
Kubinyi H 1977 Arzneim. Forsch. 27 750
Rotman A, Lundstrom J, McNeal E, Daly J and Creveling C R 1975 /. Med. Chem. 18
138
Singh H, Gombar V and Jain D V S 1980 Proc. Indian Acad. Sci. (Chem. Sci.) 89 77
Magnetic susceptibility studies of Mn^M^O (M = Zn, Mg, Fe)
C E DESHPANDE, p p fiAKARE, M N S MURTHY,
N Y VASANTHACHARYA* and P GANGULY*
National Chemical Laboratory, Pune 411 008, India
*Solid State and Structural Chemistry Unit, Indian Institute of Science,
Bangalore 560012, India
MS received 21 April 1982
Abstract. Magnetic susceptibility studies of Mni-^M^O (M = Zn,
Af = Mg, x ^0-12; M = Fe, x<0-4) in the range 77 to 300K are reported.
The methods of preparation of Mn^MgO systems preclude the presence of trivalent
ions. The Mu^Fe/) system shows anomalous behaviour around x =0-2 0-3.
The results are discussed in terms of competition between the nearest neighbour
and the next-near-neighbour interactions, dilution effects and cooperative effects
of FeO 6 octahedra.
Keywords. MnO ; FeO ; magnetic susceptibility.
1. Introduction
Magnetic susceptibility of MnO and its solid solutions with other ions such as Mg,
Zn and Fe have been extensively studied by several workers (Millar 1928 ; Foex
1948; Seino et at 1973; Jagadeesh and Seera 1980; Evrard 1971; Hope et of).
Hope et al have investigated the system (Mn^Fe^O (y < 1-0) which always
has a slight excess of oxygen. Murthy and coworkers (Deshapande and Murthy
1981; Deshapande et al 1978) have recently reported the preparation of MnO
stabilized by small amounts of Zn and also Mn^.Fe^O (x < 0-4)which are very
resistant to oxidation. Thus, the Mnj^Fe.0 system has no excess oxygen for
x< 0-4. In this paper, we report the results of our studies of the magnetic
susceptibilities of Mn^Af.O (M = Mg, Fe and Zn) in the 77-300 K range.
2. Experimental
The Mui^MjO systems were prepared according to the procedure of Murthy
and coworfeers (Deshapande and Murthy 1981; Deshapande et al 1978). X-ray
lattice parameters were determined using a Phillips PW 1050 diffractometer.
measured using HgCo(SGN) 4 for calibration purposes. The field used was 3000
gauss.
3. Results and discussion
3-1. M-ray diffraction studies
Crystal structures of these solid solutions are all of the rocfcsalt type. The
cubic unit cell parameters of these solid solutions are given in table 1. All the
samples show a linear decrease in the unit cell dimensions with increasing x as
expected of Vegard's law behaviour (figure 1).
3.2. Magnetic susceptibility studies
Magnetic susceptibility studies of these compounds were carried out below 300 K.
All the samples showed a maximum in the susceptibility at a temperature r max
which could be associated with the antiferromagnetic ordering temperature T Nf
The ;Q I vs T plots of these compounds are shown in figures 2-4. The magnetic
susceptibility of the Ma^Z^O and Mn a _,Mg,O samples were normalised for one
g atom of Mn. Such a normalisation was not carried out for the Mnj.^Fe.O
compounds. The values of /H eff and calculated from the slopes of these lines
above 200 K from a least squares fit are given in table 1. Because of the broad
Table 1. X-ray parameters and magnetic properties of Mni_ B M a O compounds.
Compound x
a, A
(0-003)
B
9
(K)
T
-*max
(K)
Mni-aZn^O O'OOl
4-449
5-63
533
127
0-005
..
5-77
568
125
o-oio
..
5-98
629
123
0-050
4-443
6-14
667
119
Mn a _.Mg/) 0-001
4-450
5-82
573
120
0-005
4-451
5-72
547
119
0-010
4-449
5-73
568
118
0-050
4-439
5-44
460
116
0-10
4-430
5-44
440
114
0-12
4-426
5-47
435
110
0-15
4-419
5-45
445
100
tfai_,Fe ffl O 0-1
4-435
5-27
358
126
0-2
4-425
4-94
210
138
0-3
4-412
5-20
238
130
Magnetic susceptibility studies o
263
4.45
Mn 1-x Zn x O
M ni _ x Fe x O
A Mn 1 _ x Mg x O
4-43 -
4.41 -
.0-1 .0-2 0-3 0-4
x
Figure 1. Variation of unit cell parameter with increasing x for the Mn, .M a O
system.
Mn. Zriv
1 X ^
220
Q
""^ O i^oOO
200
"~ ooooo 00 0-050
A
180
220
;>
OQ oOO
200
, ^Vx O
\
9^30X00 00000
"c
4 0-010
2 180
o>
T X 220
_
OA _o
' 200
o O
180
QjscDOOOOO 000
U'UUO
210
\
- \ ^ooO-
Qnm cOXiOO r\_r\r\* *
200
Mn Mg r
1 f. X Q
180
o
~ o O o X=0-12
160
^ OqpOO
190
r A
T o
o
o
170
o0 X=0-10
-i
O O ( v.oO O
150
4
r- ' o ' '
180
1 160
- o
o ^ n X=0-05
I 000
~ S 190
o
X
_ o X = 0-01
170
- *A o
190
^
ooo X=0-005
O O Q
170
- f o
I o
190
- o
o o X = 0-OO1
17O
- - ooo
I ' I I
100 200
Temp.(K)
300
Figure 3. . X' 1 vs T plots
200
100
Mrv x Fe x O
= 0-10
o
oo
oCb
100 200 300
nature 01 tne maximum me /j flff ana t/ values inus caicuiaiea may not
the true high temperature values. The trends are, however, likely to be real.
Mn.O exhibits a broad susceptibility maximum which is slightly above the order-
ing temperature T N ; 2# itself is characterized by a sharp in flection point in the
susceptibility. In the rock-salt structure both the nearest neighbour cation-cation
interaction / and the next -near-neighbour cation-anion-cati on interaction J nnn
ar possible. In 3d transition metal monoxides, J Mn is the dominant interaction
and this leads to magnetic ordering of the second kind in the fee structure
(Goodenough 1963). The broad maximum in the susceptibility is probably
due to competing interactions and frustrations inherent in the rock-salt .structure
which cannot sustain a magnetic order in which all interactions are antiferro-
magnetic. At T N rhombohedral distortion takes place which reduces the unit
cell length thus strengthening antiferromagnetic interactions and weakening ferro-
magnetic interactions between nearest neighbours. This exchange striction which
is a cooperative process is likely to be damped strongly in the presence of non-
magnetic impurities.
3.3. Mrti-iAfg, and Mrt^Zn,
In these compounds the sharpness of the inflection point decreases with increas-
ing Mg or Zn concentration (figures 2 and 3). The maxima broaden consider-
ably and !T max shifts to lower temperatures with increasing x. The substitution of
Zn or Mg for Mn probably affects the cooperative nature of the exchange stric-
tion process so that long-range ordering is affected. The magnetic structure is
broken into clusters with short-range interactions. This is the usual dilution
effect.
3.4. Mn^Fe^O
Substitution of Fe for Mn leads to a decrease in 6. In FeO, 0(T N ~ 1 which
is indicative of a / = whereas a OfT N ~ 5 in MnOis indicative of J nn ^= 0. The
and ;u e f t values of these solid solutions are given in table 1. The values and
TN(~ r max ) are in reasonably good agreement with earlier results (Evrard 1971;
Hope et al). There is however an anomaly around x = 0-3 with respect to the
3"mM and (table 1). This sample also shows a broad minimum in. the inverse
susceptibility curve (figure 4). The anomalous behaviour could be representa-
tive of the solid solutions Mn^FeaO as distinct from the (Mr^-a-Fe^O^O (y < 1)
studied by earlier workers (Evrard 1961 ; Hope et al). A Mossbauer study
of our samples has shown the absence of Fe 3+ ions. The broad nature of the
susceptibility around T N for the x = 0-3 samples seem to show the presence of
iron-rich and Mn-rich clusters. Hope et al have reported the results of the
quadrupole splitting. A# a at tne ^e nucleus in (Fd^Mn.)/) systems. Their
results show that for small values of x, A# a > * s small reflecting essentially the sym-
metrical environment of MnO 6 octahedra. With increasing x, A-# a increases
rapidly showing the formation of distorted FeO 6 octahedra. There seems to be
some evidence of a change in the slope of the A-# a v s x plot around x = 0-2-0 -3
and the A^a values are very close to the value found in 'FcO' in this range.
Acknowledgement
Tiie authors are thankful to Professor CNR Rao for continued interest in
the problem and for valuable discussions.
References
Deshapaude C E and Murthy M N S 1981 Bull. Mat. Sd. 3 261
Deshapande C E, Pant L M and Murthy M N S 1978 Indian J. Chem. A16 251
Evrard O 1971 Rev. Chim. Miner. 8 63
Foex M 1948 C.R. Acad. Sd. 227 193
Goodenaugh J B 1963 Magnetism and the chemical bond (New York : John Wiley)
Hope D A O, Cheethara A K and Long G J (Preprint)
Jagadeesh M S and Seera M S 1980 Phys. Rev. B21 2897
Millar R W 1928 /. Am. Chem. Soc. 50 1875
Seiao M, Miyahara S and Nore Y 1973 Phys. Lett. A44 35
> , pp. zo/-z/7.
Printed m India.
Liquid structure of vanadium tetrachloride from
neutron diffraction study
R V GOPALA RAO* ar.d B M SATPATHY
Physical Chemistry Section, Jadavpur University, Calcutta 700032, India
MS received 12 June 1981 ; revised 24 March 1982
Abstract. Assuming the separation of the inter-molecular scattering function into
the radial and angular parts and using Egelstaff et a/'s orientational model for
tetrachlorides, the structure of liquid vanadium tetrachloride has been studied. It
has been observed that such a separation is approximate for this liquid and the
introduction of a third correction term is required to account for the molecular
structure function. The chlorine-chlorine partial structure and effective angle-
averaged intermolecular chlorine-chlorine potential in the liquid has been evaluated.
Without taking the third correction term, introduced to generate theoretically the
molecular structure function, the centre structure function has been obtained in an
approximate way from the experimentally observed molecular structure function
and from it the centre radial distribution function, centre direct correlation function
and the angle-averaged vanadium-vanadium effective potential has been evaluated.
Keywords. Liquid structure ; vanadium tetrachloride ; intramolecular function ;
intermolecular scattering function ; orientational correlation ; partial structure!
1. Introduction
Tetrachloride liquids possessing a high degree of symmetry in three dimensions
offer themselves as an interesting group of two-component systems for neutron
and x-ray diffraction study. Thus, attempts have been made by various workers
to account for the molecular structure function (Egelstaff et al 1971 ; Narten 1976 ;
van Tricht 1977a; Granada et al 1979) in this direction in the recent past. In
the vanadium tetrachloride molecule, the central vanadium atom makes negligible
contribution to the diffraction pattern, as the scattering amplitude of vanadium "is
-0-41 fm (1 fm = 10~ 13 cm)in comparison to the chlorine atom scattering ampli-
tude of 9-58 fm, and this makes the study simpler while one attempts to obtain
partial distributions of constituent atoms.
The existence of orientational correlation in molecular liquids have been esta-
blished by several workers (Egelstaff et al 1971; Powles 1973; Sandier et al
1974; Gibson and I>ore 1979; Murad et al 1979). Gibson and Dore (1979)
have concluded that in VC1 4 , strong correlation between neighbouring molecules
exist due to non-sphericity of the molecule.
In this paper, we use the Apollo model of Egelstaff et al (1971) as a basis for
bond of one molecule is in line with that of the other molecule and the chlorine
atom of one lies in the hollow formed by the three off axis chlorines of the second.
If it is assumed that the orientation of one molecule to the other is statistically
independent of the.'r relative separation, one can then uncouple the radial and
angular correlations of the two molecules. Though, this is orly an approxima-
tion, several workers (Page and Powles 1971; Suzuki and Egelstaff 1974; Gopala
Rao and Joardar 1979, 1980) have assumed this separation in their work and
in fact su^h a separation in the case of methane (Murad et al 1979) works very
well as they put it. In general, the results are good, when the distance between
the molecular centres is large. However, at short distances between the mole-
cules, where the correlations are large the interlocking of two molecules makes
it probably difficult for rotations to occur independent of translations and thus
the uncoupling of radial and angular correlations is not exact. In the first part
of this work, the motto is tj generate theoretically the molecular structure func-
tion, S m (0). The centre structure function, S C (Q) has been computed with
hard sphere model and the approximation of free rotation is accounted by taking
a third term which we call as a decoupling correction term. The idea of such
an additional term in the expression for S m (>) was also given earlier by EgelstafF
et al (1971) and Weis and LeVesque (1976). The molecular structure function
thus obtained is compared with the experimental results of Gibson and Dore (1979).
In the neutron diffraction study of VC1 4 molecule, the contribution of the central
vanadium atom scattering to the total intensity being negligibly small, the diffrac-
tion pattern is effectively from hollow tetrahedral chloride units. Thus it behaves
li.fo a homonuclear system which enables to extract out the chlorine-chlorine
partial pair distribution function, g c i-ci( r )> from the neutron data (Gibson and
Dore 1979). From this, we evaluate the chlorine-chlorine partial structure,
SX-GI (Q)> tne intermolecular chlorine-chlorine direct correlation function,
Cci-ciO 4 )* which is analogous to the function, C a7 (r) in the reference interaction
sita model (RISM) theory (Lowden and Chandler 1973) and also the intermolecular
chlorine-chlorine angle-averaged potential function ^ C i-ci 00> through a method
given by Gopala Rao and Joardar (1978a).
In the second part of the paper, to get an approximate idea of the potential
governing the molecular centres, we assume that to a first approximation equa-
tion (3) holds and from experimental S m (Q), we deduce the S c (Q) function. It
may be pointed out that in the present molecule, the vanadium atom lying at the
centre of mass of the system, the S c (Q} function represents the vanadium-vanadium
partial structure, S v -v (&) From -Vv (fi) We evaluate the centre radial distribution
function, g v -v (0 tne centre direct correlation function, C v _ v (r), and the angle-
averaged intermolecular potential function (or the vanadium-vanadium potential)
2. Theory
The molecular structure function, S m (g), may be written as
scattering lengths of the ith an d./th nuclei, which are separated by a distance r.
The angular bracket denotes an ensemble average and the summation extends. over
all pairs of nuclei in the system. The contributions to the summation arising from
atoms within the same molecule may be separated out of the molecular structure
function so that
S (0 =/!() + A (Q), (2)
where / t (Q) is the molecular form factor and D M (Q) corresponds to the inter-
molecular contributions. At large Q values, D m (fi) becomes negligible and the
observed diffraction pattern is characteristic of a single molecule. At small
Q "alues, the D m (Q) function is more prominent and gives information about
the liquid structure.
The term, D m (Q) can be approximately separated into a molecular centre struc-
ture function term, S (Q) and an orientation dependent form factor, / 2 (fi), such
that
S m (fi) = A (0 + /. (0 [S, (0 - 1], ' (3)
However, as discussed earlier, this is only a rough approximation though several
workers (Page and Powles 1971; Suzuki and Egelstaff 1974; Gopala Rao and
Joardar 1979, 1980) have found it to be good to explain certain liquid struc-
tures. Particularly in a molecule like VC1 4 , where the non-sphericity is large, the
above relation (3) fails to account for the molecular structure. The addition of a
third term, / 3 (0, is required to give a proper description of the molecular
structure function, thus
S* (0 = /! (0 + A (0 IS. (0 - 1] + A (0 (4)
where we call/ 3 (0 as the decoupling correction term and this gives a measure
of the orientational correlation between the molecules in the system.
For tetrahedral VC1 4 molecule, the molecular form factor, f v (0, is given by
/i (0 = @v + 4& C i)- 2 [*v + 4 ^ci + 86 V Wo
(firy-ci) + 12 *W (2'ci-ci)] '(5)
where b v and b c \ are the coherent scattering lengths for the vanadium and chlorine
atoms respectively, / (fir) is a spherical Bessel function of zeroth order and r v -ci
and r cl _ cl are the appropriate distances between the vanadium and chlorine nuclei.
For the orientation dependent term,/ 2 (0, we use Egelstaff *s- Apollo model foi
tetrachloride liquids (Egelstaff et al 1971), where / 2 (0 is given by
/a (0 = (&v
+ 6 C1 ./ (fir ci-a) + WCL I (Qrtt (6)
where
irl2
\ _ r rt fr\ v cin ff\ cin ft H ft
We assume S C (Q} to be that given by Percus-Yevick approximation for hard
sphere potential.
The term/ 3 (0, used to make up the discrepancy of equation (3) in accounting
for the observed S m (0, is taken to be in the following form :
f (ri\ - Sin {A! (go ~ 0} m
h(Q} A.(fi,-fi) ' C)
where ^ and ^ a are two parameters, ^ being angular and A a radial in nature.
This function is a short range function and becomes almost zero at large Q values
and also when Q -> 0. Thus in the long wave limit ie as Q -* 0, / 3 (0 gives a
limiting value of 0-01 and for large Q, / 3 (0 becomes vanishingly small. Though
/s (0 term cannot be derived rigourously from theory, it can be formulated intui-
tively. Thus out of the two parameters one stands for the radial part and the
other for the angular part of the correlations.
At this juncture it may be pointed out that several other mathematical forms
were tried for the term/ 3 (0 and these arc mentioned below.
(i) /a (0 = (4 n r } i/ 2 e*P [ - (0 - ft) 8 /**] . (8)
Here r is a parameter and Q is the point where the decoupling correction is
the highest.
Oi) /a (0 = exp ( - Afi) sin Afi, (9)
where A is a parameter.
(iii) A (0 = W exp ( - A<2)> (10)
where A and rc are parameters. Equations (8)- (10) were used with several varia-
tions and were found to be inferior to equation (7) in accounting for 5^ (0.
The total pair-distribution function, g (r), in two-component system like VC1 4
is the sum of three partial components corresponding to the intramolecular, g M (/)
and intermolecular, g L (r) components :
g(r)-&,(r) + &(r). (11)
But since the terms depending on the position of the vanadium atom, due to its
very small scattering length, are negligible, the intermolecular pair distribution
function, L (r), is approximated by:
'). (12)
and is obtained by the Fourier transformation of D m (0, such that
^max
gci-ci W ~ 1 = ^~ \QD m (Q}M(Q) sin Qr dQ (13)
where p ci is the chlorine number density and M(Q) is a Lore h: modification
function introduced into the- transformation in order to reduce termination
effects and is given by
From equation (10), the chlorine-chlorine partial structure, Sci-ci (0
the Fourier transform of g c i-ci ( r )> can b e found out to be
l - (14)
From Sci-ci(0> the effective chlorine-chlorine potential can be obtained (Gopala
Rao and Joardar 1978a).
As already stated, assuming equation (3) holds good as a first approximatior
we deduce S e (0 from the experimental S m (0, and subsequently from S e (0,
other quantities liks g v _ v (/), C v _ v (r) and $ v -v 00 car > be obtained (Gopa]a Rao
and Joardar 1978a).
3. Results and discussion
The molecular structure of VC1 4 molecule has been investigated by Morino and
Uehara (1966) using gas phase electron diffraction. The study has shown that the
tetrahedral symmetry of the molecule is not disturbed by vibronic interactions.
Even in the condensed liquid state the spherical symmetry of the molecule may be
assumed to exist still to a good approximation and therefore we choose to calcu-
late the centre structure factor, S e (0 through a simple Percus Yevick approxi-
mation for the hard sphere potential with a hard sphere diameter, a = 5 -50 A.
In this connection, it may be pointed out that S a (0 generated with a Sutherland's
potential as a perturbation over the hard sphere in the random phase approxi-
mation has been found not to improve the results over the simple hard sphere.
This choice of a is based on the molecular dynamics calculation on VC1 4 liquid
with a Lennard- Jones potential by Murad and Gubbins (1980), who get the
v-v( r ) peafc at r = 6-30 A. To get<r from this, we make use of the approxi-
mate relation cr LJ = 2 1/a <r H . s . This value of 5-50 A compares well with the
(rvalue of 5-57 A for a similar molecule like TiCl 4 , given by van Tricht (1977b).
The (0 and f z (0 were calculated through equations (5) and (6) respectively.
These are plotted in figure 1. The molecular parameters used are taken from
Gibson and Dore (1979) and are tabulated below along with other parameters.
Temperature; T = 294 K
Molecular number density, p = 0-00565 A 3
Hard sphere diameter, a = 5 -50 A
r v _ cl = 2- 14 A, r c ,_ cl = 3-49 A
deuced by Fourier transforming the real space distribution, d L (r\ data (where
JL (r) = 4w/ C! [g c i-ci W "" !]) obtained by Gibson and Dore (1979).
In Sgarc 2, we give / 3 (Q) as obtained from equation (7). The parameters ^
and Aa were chosen so as to give the best possible fit of S m (0 as given by
equation (4).
Now, as mentioned earlier, the total in termolecular scattering is effectively from
chlorine atoms and thus the chlorine-chlorine partial structure, Sci-ci (0 1S obtai-
ned from D m (Q) through equation (11). This S C i- C i(Q) function and the effective
in termolecular chlorine pair potential, <ci-ciO") h as been obtained and presented
in figures 4 and 5 respectively. In figure 6, the C i-ci( r ) a nd Cei-ciO") functions
are also shown.
0-8
o
o
12
Figure 1. The intramolecular form factor, f^ (Q) and the intermolecular form
factor, / a (Q).
Figures. The molecular structure function, S m (Q).
2-0
o
o
o
CO
4
Q(A H )
Figure 4. The chlorine-chlorine partial structure functions Sci-ci(2)-
Figure 7 gives the centre structure, Sv-v (2) evaluated through the approximate
equation (3) obtained from the experimental S m (Q). Figure 8 gives the inter-
molecular potential < C i_ cl (r) and figure 9 the g v -vO") nd C v _ v (r) functions, all
obtained from SV-v(fi)-
Figure 2 gives an idea of the contribution of the / 3 ()) correction term to the
total molecular structure function. At the principal peafc for S m (0, this
contribution is 0-22 out of 0-56, i.e., about 39%. Unlike, in' me thaj\e (Murad
et al 1979), the large centra] vanadium atom in VC1 4 makes the chlorine atoms
more protruding, thereby increasing the non-sphericity and thus the orientations
of the molecules in the Apollo model are liJaely to be strong, which , inhibits the
free rotation of the molecules with respect to each other independent of trans-
0-05 -
0-03 -
0-01 -
-0-01 I 1
11
Figures. The intermolecular chlorine-chlorine potential function, # C i-ciO).
--2-0
10
Figured. The chlorine-chlorine direct correlation function, C7 cl _ P1 (r) and the
chlorine-chlorine pair distribution function, C1 _ C1 <. '
.The addition -of such a correction term gives a molecular structure function
ich is given in figure 3 and compared with experimental S m (Q) data The S (Q)
obtained compares reasonably wall with the experimental one, except giving a
tellite peafc with a height of 0- 18- at Q = 1-3 A' 1 . The experimental result,
oe pointed onit&at other tetrachlorides lifce SiCl 4 , and TiCl 4 show such
pronounced satellite peaks corresponding to this point (van Tricht 1977b).
2-0
o
>
i
>
to
1-0
10
Figure 7. The vanadium-vanadium partial structure functions, 5 V _ V (0.
0-04 -
u
-1-0
Figure 9. The vanadium-vanadium direct correlation function, C v _ v (r), and the
vanadium-vanadium pair distribution function Y _ V 00-
It is seen from figure 5 that at the nearest neighbour distance, i.e., 3-95 A, for
the chlorine atoms, the potential has a positive value of 1 52 x 10~ 2 eV. It may be
pointed out here that liquid chlorine also gives such a positive minimum (Gopala
Rao and Joardar 1978b) in its potential function. However, the second neighbour
rests at a negative potential well of depth 0-66 x 10~ 2 eV at a distance of 6-65 A.
The potential dies very fast and thus the chlorine-chlorine interactions are weak
beyond the second neighbour distance.
The centre structure, 8, (Q), shown in figure 7, also has a subsidiary peak at a
Q value 1-5 A" 1 , like the calculated molecular structure function. The oscilla-
tions in it almost vanish beyond Q = 6-0 A" 1 . The intermolecular potential func-
tion, jj v _ v (r) in figure 8 gives an equilibirum distance between the molecular centres
of 6- 55 A, and the potential depth at this point is 2-84 x 1Q- 2 eV.
The study shows that the present orientational model can explain the mole-
cular structure of liquid VC1 4 , if we introduce a correction factor for the sepa-
ration of the radial and angular part in the intermolecular scattering term. This
term plays a significant role in explaining the molecular structure for tetrahedral
molecules where the sphericity has been lost to a great extent by the large size
of the central and peripheral atoms. However, for molecules like CH 4 and CC1 4
which are more spherical and where positional correlations dominate over orien-
tational correlations (Murad et al 1979; Gopala Rao and Murty 1974, 1976)
the role of this term will be diminished. Here it is pertinent to point out that
Murad et al (19 79) have found the free rotation approximation is superior to the
Reference Interaction Site Model (RISM) theory in explaining th,e structure
of tile tetra^ecjral liquid methane molecule,
Acknowledgements
The authors wish to thank Dr J C Dore of the University of Kent at Canter-
bury, UK, for providing them with the experimental results and other related
papers on his study in liquid VC1 4 . They also acknowledge the financial assistance
received from the Department of Science and Technology (DST), Government
of India, for this work. One of the authors (BMS) is thankful to DST for the
award of a fellowship. The authors also thank the UGC for the financial
assistance during the early part of the work.
References
Egelstaff P A, Page D I and Powles J G 1971 Mol. Phys. 20 881
Gibson I P and Dore J C 1979 Mol. Phys. 37 1218
Gopala Rao R V and Joardar R N 1978a Phys. Lett. A67 71
Gopala Rao R V and Joardar R N 1978b Unpublished work
Gopala Rao R V and Joardar R N 1979 J. Phys. C28 4129
Gopala Rao R V and Joardar R N 1980 Mol. Phys. 40 309
Gopala Rao R V and Murty A K 1974 /. Phys. C7 2249
Gopala Rao R V and Murty A K 1976 Indian J. Phys. 50 1
Granada J G, Stanton G W, Clarke J H and Dore J C 1979 Mol. Phys. 37 1297
Lowden L J and Chandler D 1973 /. Chem. Phys. 59 6587
Morino Y and Uehara H 1966 /. Chem. Phys. 45 4543
Murad S, Evans D J, Gubbins K E, Streett W B and Tildesley D J 1979 Mol. Phys. 37
725
Murad S and Gubbins K E 1980 Mol. Phys. 39 271
Narten A H 1976 /. Chem. Phys. 65 573
Page D I and Powles J G 1971 Mol. Phys. 21 901
Powles J G 1973 Adv. Phys. 22 1
Sandier S I, Das Gupta A and Steele W A 1974 J. Chem. Phys. 61 1326
Suzuki K and Egelstaff P A 1974 Can. J. Phys. 52 241
van Tricht J B 1977a /. Chem. Phys. 66 85
van Tricht J B 1977b Ph.D. Thesis, Neutron diffraction study of carbontetrachloride, silicium-
tetrachloride, titanium tetrachloride and tintetrachloride in the liquid state at 295 K.
Delft University
Weis J J and Levesque D 1976 Phys. Rev. 13 450
Periodic precipitation of cobalt(II) oxinate in agar gel : factors
influencing the flocculation
N KANNIAH, S AMBROSE, F D GNANAM and P RAMASAMY*
A C College of Technology, Perarignar Anna University of Technology,
Madras 600025, India
MS received 2 February 1982 ; revised 18 May 1982
Abstract. The influence of concentration, ageing and pH of the gel medium on
the periodic precipitation of cobalt oxinate in agar gel is reported. The results are
explained on the basis of Shinohara's revised flocculation theory. The flocculation
value (F) increases with increase in the gel concentration whereas it decreases with
increase in ageing and pH of the gel. The raise in temperature of the gel increases
the solubility of the sparingly soluble substance and hence the F value. The
effect of additives on the periodic precipitation of cobalt(II) oxinate is reported.
Keywords. Periodic precipitation ; cobalt oxinate ; flocculation value ; gel concen-
tration ; ageing.
1. Introduction
Many compounds that form insoluble precipitates in a counter diffusion system
exhibit the Liesegang (1896) phenomenon in which, a series of concentric rings
are produced rather than a continuous precipitate. Few quantitative data are
available despite over 800 publications on this subject (Stem 1967). The -formation
of Liesegang rings has been influenced by various factors like concentration of the
reactants, concentration, ageing, temperature and pH of the gel medium and the
amount of additives present.
We have recently published the experimental conditions for obtaining the
Liesegang rings of cobalt(II) oxinate in agar and the influence of the concen-
tration of the reactants on the periodic precipitation (Kanniah et al 1981). In
this paper the influence of the concentration, ageing, pH and temperature
of the gel and the effect of additives on the formation of Liesegang rings of
cobalt oxinate are discussed in detail.
2. Theory
The/periodic precipitation of cobalt oxinate has been explained on the basis of
the revised coagulation theory of Shinohara (1970). As the outer electrolyte
* To whom correspondence should be made.
279
formed as a sol at the contact plane of the two reactants. This boundary known
as sol-front advances spreading the sol region. As more and more outer electro-
lyte diffuses, the ionic concentration reaches a characteristic value F which triggers
tha flocculation of the sol. F is expressed by the equation
where C 10 is the concentration of cobalt nitrate
Afc = */> and ^ = 1 , (3)
*
p is known as the spacing coefficient. x n+1 and # are the positions of the
(n + l)th and nth rings from the gel boundary, k is called the front constant
which is estimated using the Adair's equation. The concentration of the super-
saturated solution of the product (C^) formed just before the formation of sol
is given by
exp (-*)
The flocculation value JP is calculated as
F=C 3o + r. (5)
3. Experimental
3.1. jE^fec* of gel concentration
2-178 grams of analytical .grade oxine were dissolved in minimum amount of
2N acetic acid. This solution was mixed with hot agar agar solution and the
final volume was made up to 300 ml. The pH of the solution was adjusted to
4-25. Thus 1-0% agar agar solution impregnated with 0-05 mole/lit oxine was
prepared at 4-25 pH. 50 ml of this solution was poured into a corning tube of
20mm diameter and allowed to set. After 3 hr 10ml of 1-031 mole /I cobalt
nitrate was taken over the gel. To study the effect of gel concentration of the
periodic precipitation of cobalt oxinate, the gel concentration was varied from
0-4% to 2-0%. In all the experiments the concentration of the inner electrolyte
and that of the outer electrolyte was kept as 0-05 and 1-031 mole/1 respectively.
3.2. Effect of ageing of the gel
1% agar agar gel impregnated with 0-05 mole/1 oxine was prepared as before.
1-031 mole/1 cobalt nitrate was taken over the set gel at different time. The
Figure 1. The Liesegang rings of cobalt oxinate in agar gel in presence of additives
potassium sodium tartarate, potassium thiocyanate and potassium chloride.
The periodic precipitation of cobalt oxinate was normally carried out at room
temperature (30 C). The influence of the temperature was studied by loeeping
the gel at different temperatures. The glass tubes containing the hot gel solution
with oxine were immersed in a thermostat toept at the required temperature.
The temperature of gel was varied from 30 to 42 C. In all these experiments
the concentrations of the gel, oxine and cobalt nitrate was 1%, 0-05 mole/1 and
1-031 mole/1 respectively.
3 .4. The effect ofpH of the gel medium
Oxine was dissolved in minimum amount of 2N acetic acid and the hot agar agar
solution was mixed. The pH of this solution was adjusted with aqueous
ammonia to 4-25. Similarly the pH of the gel solution with oxine wa's varied
from 3-7 to 4-7. Oxine is precipitated, when the pH is above 4-7. When the
pH was below 3-7 gel set was not observed. The pH of the solution was
measured at 60 C. The gel solutions were allowed to set at room temperature
(30 C) and cobalt nitrate solution was taken over the set gel after 3 hr. The
concentrations of gel, oxine and cobalt nitrate were the same as in the previous
experiment.
3 . 5. Influence of additives
To study the effect of additives on the periodic precipitation of cobalt oxinate,
suitable additives were taken along with oxine in the gel medium. The amount
of additive was varied from 0-001 to 0-026 mole/1. When the gel containing
oxine and additive was set, cobalt nitrate was taken as the outer electrolyte. The
concentrations of the gel, inner electrolyte and out^r electrolyte were kept as 1 0%,
0-05 and 1-031 mole/1 respectively. Potassium sodium tartarate, potassium
chloride and potassium thiocyanate are the additives taken along with oxine in
the gel.
In all the experiments, sharp brown coloured disc-like precipitate rings demar-
cated by clear void spaces were obtained within a week (figure 1). The interspac-
ing between successive rings increases with the number of ring (n) from the gel
boundary. At the lower rings, small crystals of cobalt oxinate were observed.
The distance measurements were made with cathetometer. The IBM 1130 com-
puter was used for the calculation of flocculation values using Shinohara's
coagulation theory.
4. Results and discussion
4.1. Effect of concentration of the gel medium
Matalon and Packter (1955) have established that the gel has a great influence
on the periodic precipitation of insoluble salts and that the gel interacts apprec-i
ciably with the precipitated substance. They have modified Wagner's relation
(1950) and derived the following equation:
_ 7?
the interaction between the gel and the substance precipitated. S = a constant
which is a measure of super-saturation of sparingly soluble substance J m = a
constant with integral value J and Ci = outer electrolyte concentration."
From ,the above equation it is very clear that the spacing coefficient (/>) will
increase with increase in the gel concentration. As the concentration of gel
increases the interaction of the gel with, the precipitated substance increases.
Hence 'the sol of the sparingly soluble substance is well protected. The stability
of the sol increases. Therefore the amount of the outer electrolyte required to'
flocculate the sol will be more. Thus; the flocculation value must increase with
increase in the gel concentration (figure 2) leading to an increase in the spacing
coefficient (p). '
4.2. Effect of ageing of the gel .
Daring ageing of the gel the micellae and the intervening capillary spaces become
coarser. This will decrease the solubility of cobalt oxinate' leading to rapid
precipitation. Hence the flocculation value (figure 3) decreases with increase in
the ageing of the gel. ,
4.3. Effect of temperature
The increase in temperature of gel leads to a progressive diminution of the total
volume of the micellae (Clayton 1932). Hence the pore size increases. This in
turn leads to an increase in the diffusion coefficient. Moreover the solubility of
7.75-
o
\J
O
x
U.
7-50 -
7-25 -
h)
Figures. Effect of ageing on the flocculation value (F). C i0 = 1-031 mole/1;
C 20 = 0-05 mole/1., C g = l%.
the sparingly soluble product increases with increase in temperature. As the
solubility increases the precipitation taJoes a longer time and hence the rings are
formed at greater distances leading to an increase in the spacing coefficient. As
the solubility increases, the concentration of the supersaturated solution (C^)
formed just before flocculation increases leading to a higher flocculation value
(F). Hence increase in temperature of the gel medium increases the flocculatjon
value (table 1).
4.4. Effect of pH of the gel medium
The pH of the gel medium plays a predominant role in the periodic precipitation of
the sparingly soluble salts (Varma 1953). Cobalt oxinate is soluble in acids. As
the pH of the gel medium increases, the solubility decreases leading to a decrease
in the spacing coefficient and .flocculation value (table 2). The periodic precipi-
tation of cobalt oxinate is observed only in the narrow range of pH 4-3 to 4-0.
4.5. Effect of additives
The characteristic features of the periodic precipitation of cobalt oxinate are
very much influenced by the presence of additives. The effect of impurities on
the periodic precipitation of calcite has been studied by Bugazh and Fraknoy
(1961), Gnanam et al (1980) and Krishnan et al (1981). When potassium sodium
tartarate, potassium thiocyanate and potassium chloride are used as additives, the
flocculation value increases with increase in the concentration of the impurity.
TMc reveals that solubility of the sparingly soluble product increases with increase .
C\0
mole/1
C 20 Temperature p
mole/1 C
Fx 1Q
mole/I
1-031
0-05
30
1-049
7-038
1-031
0-05
34
1-058
7-380
1-031
0-05
38
1-064
7-426-
1-031
0-05
42
1-082
7-897
0-859
0-05
30
1-051
7-035
0-859
0-05
34
1-061
7-126
0-859
0-05
38
1-065
7-212
0-859
0-05
42
1-085
7-658
0-687
0-05
30
1-053
6-997
0-687
0-05
34
1-065
6-983
0-687
0-05
38
1-071
7-093
0-687
0-05
42
1-090
7-425
Q'515
0-05
30
1-059
7-067
0-515
0-05
34
1-071
6-881
0-515
0-05
38
1-075
6-926
0-515
0-05
42
1-097
7-227
Table 2. Effect
of pH on
flocculation
value.
PH
CIQ
mole/1
Cj
mole/1
P
FxlO a
mole/1
4-05
1-031
0-05
1-101
8-412
4-20
1 -031
0-05
1-092
8-167
impurities the flocculation value is the lowest in the case of nntassi nm
o
O
X
LL
KC\
2-5 5-0 7-5
C a x10 3 (mol lit" 1 )
10-0
Figure 4. Effect of additives on the flocculation value (F). C i0 ==1-031 mole/1;
C ao =0-05 mole/1, C g = 1 % ; C a = concentration of additives.
When excess cobalt nitrate diffuses into a gel impregnated with oxire, a posi-
tively charged sol of cobalt oxinate is formed due to the adsorption of excess
Co 2+ ion. In that case the counter-ion, plays an important role in flocculatin g
the sol. Among the counter-ions (tartarate, thiocyanate and chloride), the tri-
valent tartarate will be more effective in flocculating the sol. Hence the floccih
lation value should be the lowest for a particular concentration of the tartarate.
This can also be accounted by the lyotropic order of anions (McBain 1950).
5. Conclusion
The results thus conclusively prove that the concentration, ageing, pH and
temperature of the gel have pronounced influence on the flocculation of cobalt
oxinate. The flocculation values of different anions are in the lyotropic order.
References
Bugazh V A and Fraknoy V 1961 Kolloidn. Zh. 181 10
Clayton W 1932 Colloid aspects of food chemistry and technology (London : J A Churchill
Gnanam F D, Krishnan S, Ramasamy P and Laddha G S 1980 /. Colloid. Interface Sci. 73
193
Kanniah N, Gnanam F D and Ramasamy P 1981 /. Colloid. Interface Sci. 80 377
Krishnan S, Gnanam F D, Ramasamy P and Laddha G S 1981 Krist. Tech. 10
Liesegang R 1896 Naturwiss. Wochenschr. 11 353
Matalon R and Pacter A 1955 /. Colloid. Sci. 10 46
McBain J W 1950 Colloid science (Eoston : P C Heath and Company)
Shouji Shinohara 1970 J. Phys. Soc. Jpn. 29 1073
Stern K H 1967 Bibliography of Liesegang rings (USA : National Standards)
VnllnMti 7h 133 32
Bis (^-cyclopentadienyl/indenyl) N-aryl dithiocarbamato chloro
oxotungsten(VI) complexes
G S SODHI, H S SANGARI and N K KAUSHIK*
Department of Chemistry, University of Delhi, Delhi 110007, India
MS received 16 October 1981 ; revised 24 December 1981
Abstract. Some bis(ij 5 -cyclopentadienyl) N-aryl-dithiocarbamato chloro oxotungsten
(VI) complexes of the type (C 5 H s ) 2 WO(S a CNHR)Cl (R =o-, m-, ^-tolyl and Ph)
have been prepared by the reaction of stoichiometric amounts of bis 0>j B -cyclo-
pentadienyl) oxotungsten(VI) dichloride with sodium salts of dithiocarbamic acids
in refluxing tetrahydrofuran. The corresponding indenyl complexes of the type
(C 9 H 7 ) 2 WO (S 2 CNHR) Cl were similarly synthesised by refluxing equimolar quantities
of bis (V-indenyl) oxotungsten(VI) dichloride and sodium dithiocarbamate in tetra-
hydrofuran. Infrared spectral studies demonstrate that in these complexes dithio-
carbamate ligands are bidentate. Electronic spectra, magnetic susceptibility and
elemental analysis have also been carried out for the complexes.
Keywords. Cyclopentadienyl ; indenyl ; oxotungsten(VI) ; sodium dithiocarbamates.
1. Introduction .
In earlier papers, we reported cyclopentadienyl (Sangari etal 1980) and indenyl
(Kaushik etal 1980) derivatives of oxomolybdenum(VI) complexes. Our interest
in the investigation of bonding mode of various dithiocarbamate ligands
(Sangari etal 1981) prompted us to synthesise and characterise a number of
organo oxotungsten(VI) dithiocarbamates.
2. Experimental
2-1. Materials
Sodium dithiocarbamates were prepared by standard method (Klopping etal
1,951). Bis (j? 5 -cyclopentadienyl) and bis 0? 5 -indenyl) oxotungsten(VI) dichloride
were prepared by the interaction of tungsten(VI) oxytetrachloride with sodium
cyclopentadienyl and sodium indenyl respectively (Anand etal 1968).
*. TO whom correspondence should be made,
spectrophotometer. The electronic spectra were recorded on a Perkin-Elmer
4000 A instrument in the 450-700 nm range. Magnetic susceptibilities at room
temperature were measured on a standard Gouy's balance. Mercury tetrathio-
cyanatocobaltate (II) was used as calibrant.
2-3. Preparation of the complexes
The complexes were obtained as insoluble products after refluxing bis (7? 5 -cyclo-
pentadienyl/indenyl) oxotungsten(VI) dichloride with stoichiometric amounts of
sodium dithiocarbamate for 14-18 hr. The yellowish-brown to brown complexes
were filtered, washed successively with tetrahydrofuran, acetone and water and
finally recrystallised from petroleum ether.
3. Results and discussion
Bis (i? 5 -cyclopentadienyl/indenyl) oxotungsten(VI) dichloride reacts with sodium
dithiocarbamates in equimolar quantities according to the following general
equations :
(C 5 H 5 ) 2 WOCl 2 + Na (S a CNHR) - (C 5 H 5 ) 2 WO (S 2 CNHR) Cl + NaCl
(C 9 H 7 ) a WOCl 2 + Na (S 2 CNHR) -* (CgE^WO (S 2 CNHR) Cl + NaCl
(where R = o-, w-, p-tolyl and Ph).
The method used to prepare all the oxotungsten(VI) complexes gave materials
of good purity as evidenced by satisfactory elemental analysis (table 1) and IR
studies (table 2).
. All the oxotungsten(VI) complexes are soluble in hot petroleum ether but
insoluble in most organic solvents. They are quite stable in air. Their decompo-
sition and melting points are very high, i.e., above 300 C. They are yellowish
brown to brown in colour. Magnetic susceptibility value at room temperature
shows that all compounds are diamagnetic.
3 1 . Infrared spectra
The main interest in the preparation of these complexes is the attachment of the
dithiocarbmate ligand. If the dithiocarbamate ligand is bidentate, a single band
at "- 1000 crrH is found (Bonati et al 1967) which is due to two equivalent C-S
stretching vibrations. In unidentate dithiocarbamate, as in Ru(NO) (S 2 CNEt 2 ) 3
(Domenicano et al 1966) a doublet arises at ~ 1005 cnH and 983 cnH which is
due to two non-equivalent C-S stretching vibrations. Thus this a reliable criterion
for determining the bonding mode of the dithiocarbamate ligand. All the prepared
complexes possess one medium intensity band at ~ 1000 cnH. This indicates
the presence of four-membered ring system in these complexes and also supports
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Organo oxoiungsten(VjT) dithiocarbamates 293
16 bidentate nature of the dithiocarbamate ligand. The band at ^ 950 cnr^
idicates the presence of W=O group in the complexes (Cousins etal 1964).
A number of skeletal vibrations appear around the 1000 cnH region due to the
:esence of cyclopentadineyl or indenyl groups. However these are quite weak
> compared to the v(CmS) or v(W=O) stretching frequencies. Hence it is
)ssible to distinguish these latter stretching frequencies even in the presence of
:eletal vibrations due to organic ring systems. Further, the doublet observed
, the unidentate behaviour of the dithiocarbamate moiety is a splitting of the
ind appearing at ~ 1000 cm- 1 . Such splitting can be easily marked, even in
ie presence of other weak bands in the region. Thus IR spectral studies are a
(liable criterion for determining the. bonding mode of the dithiocarbmate ligand.
While assigning the co-ordination number to the cyclopentadienyl compounds,
ie latter group is assumed to occupy a single co-ordination site (Coutts etal 1970,
)74, 1975). Thus the compounds of the type 0? 5 -C 5 H 5 ) a Ti (SaCNRa) Cl [R =
kyl] (Kaushik etal 1978) and (r/ 5 -C 5 H 5 ) a Ti (S 2 CNHR) Cl [R = aryl] (Kaushik
al 1979) were assigned a co-ordination number five. Recently a co-ordination
amber six was assigned to the oxotungsten(VI) complexes of the type 0? 5 -C 5 H 5 ) 2
^O (SaCNRa) Cl (Kaushik etal 1981). On this basis, the complexes reported
\ this paper may be assumed to be hexacoordinate.
All the complexes reported in this paper possess one band at ~ 1500 crrr* which
;veals the thioureide band. As the frequency of this band lies between that
-N (1350-1250 cnr*) and C=N (1690-1640 cur*), this suggests that this bond
Dssesses some double bond character.
The C-H stretching frequencies at ^ 2990-3000 cm" 1 in the complexes reported
.dicate indenyl and cyclopentadienyl group.
The band at ~ 365 cnr 1 indicates the presence of W-C1 group in the complexes,
hile that at ^ 350 cnr 1 is assigned to the tungsten-sulphur stretching frequency
Jradley et al 1969).
2. Electronic spectra
he electronic spectra of the complexes recorded in nujol, exhibited a single band
i the range 24600-24200 cnH region which may be assigned to the charge transfer
and (Dunn et al 1962) in accord with the electronic configuration (n-1) dns of
ingsten in each case.
cknowledgement
ine of the authors (GSS) wishes to thank the CSIR, New Delhi for, providing a
>llowship.
Coutts R S P and Wailes P C 1974 Aust, J. Chem. 27 2483
Courts R S P and Wailes P C 1975 /. Organomet. Chem. 84 47
Damenicano A, Vacigo A, Zambonelli L, Loader P L and Venanzi L M 1966 Chem. Commun.
476
Dunn T M, Nyholm R S and Yamad S 1962 /. Chem. Soc. 1564
Kaushik N K, Bhushan B and Chhatwal G R 1978 Transition Met. Chem. 3 215
Kaushik N K, Bhushan B and Chhatwal G R 1979 Z. Naturforsch. 34b 949
Kaushik N K, Singh R P, Sangari H S and Sodhi G S 1980 Synth. React. Inorg. Met-org.
Chem. 10 617
Kaushik N K, Sodhi G S, Sangari H S and Chhatwal G R 1981 J. Chin. Chem. Soc. Taipei
(in press) 28 231
Kldpping H L and Vander Kerk G J M 1951 Red. Trav. Chim. 70 917
Sangari H S, Sodhi G S, Kaushik N K and Singh R P 1980 Z. Naturforsch. B35 1254
Sangari H S, Sodhi G S, Kaushik N K and Singh R P 1981 /. Inorg. Nucl. Chem. 43 1080
A theoretical study on the specific interaction of
hexafluorobenzeae with benzene and p-xylene
D V S JAIN* and F S NANDEL
Dspartment of Chemistry, Paujab University, Chandigarh 160 014, India
MS received 4 February 1982 ; revised IS June 1982
Abstract. A CNDO/2 study has been carried out for C 8 F 4 + CaHeand C 6 F 6 +
composites and individual molecules. The favoured configurations of the adducts
have been decided on the basis of energy calculations of various geometries. For
the CsFe+CfiHa adduct the lowest energy corresponds to the configuration in which
the molecular planes are parallel to each other with a twist angle of 30. For the
C 8 F 6 + C 8 Hi adduct lowest energy corresponds to a geometry in which the two
molecular planes are inclined by a small angle with the angle of twist between the
molecular planes being 30. It is shown that the complexes are not of the charge
transfer type.
Keywords. CNDO study ; benzene ; hexafluorobcnzene ; p-xylene ; charge transfer.
1. Introduction
The formation of complexes between hexafluorobenzene and alkyl benzenes is
quite well-known (Patrick and Prosser 1960 ; Duncan et al 1966 ; Boeyens and
Herbstein 1966). However, the nature of the specific interaction has not been
well understood. It was earlier suggested (Patrick and Prosser 1960 ; Frnby
etal 1966) that this has the nature of electron transfer interaction with benzene
acting as donor and hexafluorobenzene as an acceptor. Thermodynamic studies
on alkyl benzenes + hexafluorobenzene showed (Fenby and Scott 1967 ; Duncan
etal 1966) that the excess enthalpy became more negative with the substitution
of aikyl group in benzene. This was taken as an evidence for the charge transfer
type of interaction because alkyl substitution also increased the electron donating
power of the alkyl benzenes. However, no charge transfer bands were observed
spectroscopically. As a matter of fact more recent (Gaw and Swinton 1968)
thermodynamic studies and calculations based on quadrupole interactions (Broth
and Swinton 1974) indicate that these interactions arise principally from electro-
static forces. Depolarised Rayleigh light scattering study (Brown etal 1978)
gives indirect evidence for the existence of C 6 H 8 + C 6 F 6 adduct with their planes
parallel to each other even in the liquid solutions. In the absence of any unequi -
vocal experimental evidence we report the CNDo/2 calculations (Pople and
295
296
D V S Jain and F S Nandel
Beveridge 1970) on the energies and charge distributions of the benzene (p-xylene)
and hexafluorobenzene molecules separately and also for the composite mole-
cules in which the distance between the centres of mass of C 6 H 6 (CBHio) and C 6 F 6
and their relative orientations have been varied.
2. Geometries
The following relative orientations (figure 1) of benzene (/7-xylene) with respect
to hexafluorobenzene in the complex have been considered:
(i) Those in which the benzene (p-xylene) and hexafluorobenzene molecules
were in the same plane as shown in figures la and Ib.
(ii) Those in which the six-fold axis of benzene (p-xylene) lies in the molecular
plane of hexafluorobenzene and one of the fluorine atom points towards the centre
of the benzene molecule (figure lc). For this configuration the HFB molecular
plane has been rotated around the two-fold axis passing through the flourine
atom pointing towards benzene ring by to 30.
(iii) Those in which the molecules are placed one above the other and the
wist angle between the molecular planes has been varied from to 30 (figure Id)
(iv) Those in which the molecular planes are inclined to each other (figure le).
For symmetry consideration this was considered only for the hexafluorobenzene/
p-xylene complex.
The energy of the molecular systems hexafluorobenzene and benzene (jp-xylene)
has been calculated as a function of the distance between the interacting molecules
for all the above configurations. The molecular geometries of benzene (p-xylene)
and hexafluorobenzene in the complex are assumed to be those of the isolated
molecules (jR cc =1-397 A, jR CH = 1 10 A, R CP 1 = 37 A, and R cc ' = 1 52 A).
The intermolecular interaction energy (A) is defined by the difference :
AE = Energy of composite molecule the sum of the energies of isolated
molecules.
3. Results and discussion
The results obtained for the various configurations of different classes are summa-
rised in tables 1 and 2. It is found that the configurations falling in categories
(i) and (ii) have the least stabilization energy. The intermolecular distances for
these configurations were also varied but only the distances corresponding to the
lowest energy is recorded for these geometries. Configurations with 0=30
were found to be the most stable from category (iii) for benzene + hexafluoro-
benzene complex. A plot of AE vs. Rfot this is shown in figure 2 which shows a
Table 1. CNDO stabilization energy for the banzene-hexafluorobenzene adduct
as a function of distance (R) and geomtry.
Geometry 0/degree R\A - A X 10 3 /a.u
I (a) ... 7-30 0-576
(b) ... 5-50 0-400
II 3-00 5-400
15 3-00 5-372
30 3-00 0-123
m 3-00 0-654
15 3-00 5-911 -
Table 2. Stabilization energy A^xlOs f 1 ' p-xylene-hexafluoro benzene system
at various tilt angles (<) between the molecular planes [twist angle (6) being 30].
J/A
2-60
2-315
2-70
2-860
3-215
2-80
2-396 2-843
3-168
3-046
2-90
2-555
2-803
3-625
3-00
1-910 2-163
2-317
3-10
1-759
-5-00
-6-QO-
lU
<J
- 7-00
3-0
-R/A
4-0
Figure 2. CNDO stabilization energies for benzene -hexafluorobenzene as a
function of distance between the two molecules.
minimum at R 2-65 A. Crystallograpbic studies (Dahl 1972, 1975) on
some of the alkyl benzenes and hexafluorobenzene adducts also indicate that these
Interaction of hexafluorobenzette
299
for HFB-/>-xylene complex, the most stable geometry is that in which the mole-
cular planes of both the molecules are twisted by 30 and are inclined by a small
angle $ (figure le). It was found that the maximum stabilization takes place at
<f> ~ 6 (table 2), This is consistent with the crystallographic studies
(Dahl 1975) in which the inclination angle of 5 '4 is reported. ,
The charges on various atoms in the isolated molecules and in the composite
molecules for stabilized geometries are given in figure 3, It can be seen that there
is very little intermolecular charge transfer. There is only intramolecular adjust-
ment of charges on atoms on each molecule. It is also apparent from the HOMO
and LUMO of the molecules (figure 4) that the charge transfer from benzene
(p-xylene) to hexafluorobenzene is not energetically favoured.
F- 0.158
-0.004 .0-004
/H
[168 /
~/f\ r\r\
0158
F^ie 0-004
-0-004
C fiW
-n.nii. N
0-030
0-016 -0-02
016 .Q 02 Q,
(c)
0-20
0-10-
if
-0 60
C 6 F 6
Figure 4. Eaorgies of HOMO and LUMO of benzene (p-xylene) and liexafluorobenzene
and their adducts.
We have carried out the correlation of the molecular orbitals of the composite
molecule with the molecular orbitals of the individual hexafluorobenzene and
benzene (p-xylene) molecules. It is evident from figure 4 that the interaction
of highest occupied molecular orbitals play insignificant role in stabilization.
Actually this leads to slight destabilization. It was noted for both the complexes
that the stability arises due to the interaction of some lower molecular orbitals.
The theoretical calculations agree with the crystallographic studies regarding
the configuration of the complexes. The calculated intermolecular distances are
small. This, however, is a general weakness of CNDO method. Our calculations
also explain the results of depolarized Rayleigh scattering study of benzene-f-
hexafluorobenzene system.
References
ahl T 1975 Acta Chem. Scandi. A 28 170
uncan W A, Sheridan J Pand Swinton F L 1966 Trans. Faraday Soc. 62 1082, 1090
uncan W A and Swinton F L 1966 /. Phys. Chem. 70 2417
snby D V, Mclure I A and Scott R L 1966 /. Phys. Chem. 70 602
jnby D V and Scott R L 1967 /. Phys. Chem. 71 4103
aw W J and Swinton F L 1968 Trans. Faraday Soc. 64 2023
itrick C R and Prosser G S 1960 Nature (London) 187 1021
>ple J A and Beveridge D L 1970 Approximate molecular orbital theory (New York :
[cGraw Hill)
Conformations! behaviour and vibrational spectra of
3-methyl 2-butanethiol
S K NANDY and G S KASTHA*
Department of Physics, Jadavpur University, Calcutta 700032, India
* Optics Department, Indian Association for the Cultivation of Science, Jadavpur,
Calcutta 700 032, India
MS received 5 March 1982 ; revised 27 July 1982
Abstract. The Raman spectra of 3-methyl 2-butanethiol in the temperature range
120 C to + 60 C have been recorded together with its liquid phase infrared
spectrum at room temperature. The spectral analysis shows that the molecule of
the compound exists in the liquid state, in three different rotameric configurations
A, S and C of which the form A is the stablest. Besides, a tentative assignment
of the observed vibrational frequencies arising from the rotameric forms has been
presented.
Keywords. Raman spactra; infrared spectra; temperature dependence of Raman
band intensit/es; rotational conformers; energy differences; vibrational experiment.
1. Introduction
Rotational isoraeriam in substituted alkanes has been studied both experimen-
tally and theoretically. It is now fairly well understood how the different rota-
meric properties, such as the number of stable rotational conformers, their stabi.
lities and energy differences in these molecules change with the nature, position
and number of substituents specially, when the substituents are halogen atoms.
However, this is not so if the substiiuent is a group of atoms like the thiol group.
Experimental data on rotational isomerism in alkane thiols and theoretical compu-
tations (Freeman 1974) are meagre and far from adequate. Nevertheless, it has
been possible to interpret the experimental data by assuming that the rotational
conformers in mercaptoalkanes arise mainly due to rotations about the skeletal
C-C bonds and that the thiol group remains oriented in a fixed configuration
except in the rare case of ethyl mercaptan (Smith et al 1968 ; Wilson 1972).
But the task of ascertaining how the number of stable rotational isomers and their
stabilities depend on the position of the substituent thiol group requires the
acquisition of more experimental data in differently-substituted alkanethiols.
Accordingly detailed Raman spectroscopic investigations on the vibrational spec-
trum of 3-methyl 2-butanethiol in the temperature range -120 C to + 60 C
and the IR spectrum of the same compound in the liquid phase has been studied.
*To whom correspondence should be made,
303
alkanethiols including the two very similar molecules of 2-methyl-l-propanethiol
(Ozaki et al 1975) and 2-butanethiol (McCulJough et al 1958). These experi-
mental data together with their discussion form the subject-matter of this paper,
2. Experimental
3-methyl 2-butanethiol from M/s. Schuardt (Germany) was distilled under reduced
pressure and its Raman spectrum in the liquid state was obtained both photo-
graphically and with a 200 mW 4880 A radiation of argon ion laser source of a
Gary 82 and Spex C laser Raman spectrophotometers. The Raman spectrum in
the solid state, the polarisation character of the Raman lines and the temperature
dependence of the intensities of some of the Raman lines in the range ~ 120 C
to 60 C were studied with the same spectrophotometers. The m spectrum was
recorded in a Perkin Elmer model 21 spectrophotometer with rock salt optics.
3. Results
The Raman and IR frequencies with estimated relative intensities in different phases
are given in table 1. The polarisation character of the Raman lines are also shown
in the table including the probable assignments of the observed frequencies in
terms of the modes of vibration in different rotameric forms of the molecule.
The variation in the intensities of Raman lines due to C-S stretching mode of
vibration at three temperatures is shown in figure 3.
4. Discussions
4.1. Rotameric forms and their stabilities
If the CH 8 groups are considered rigid and the SH group given a fixed orientation
the molecule of 3-methyl 2-butanethiol will have only one central C-C axis of
rotation. The three rotational conformers arising due to orientation about this
bond are shown in figure 1 and are indicated as A, B and C. It may be noted
that while in form C, the thiol group is in the trans-position with respect to the
hydrogen atom, in forms A and 5 they are gauche with respect to each other.
The configuration of these rotamers is very similar to that obtained in 2-butane-
thiol in which there is a H-atom in place of one of the two CH 3 groups in the
second carbon atom. These forms shown in figure 2 have energies, according
to McCulJough et al (1958), in the order E C >E B >E A . From a comparison of these
three rotameric forms with those of 3-methyl 2-butanethiol and considerations of
the nonbonded interactions in the various groups in the different conformers of
the 'two molecules the energies of the three conformers of 3-methyJ 2-butanethiol
are found to be E B >E a > E A in the free state. The three rotamers will have
approximately the same dipole moment and Ibwering of energy in the liquid phase
is not expected to change the relative energy differences significantly. In other
IR bands (cm" 1 )
liquid (thin film)
Raman
shifts (cm" 1 )
- Assignment
Rolamer
Liquid
Glassy mass
( 120 C)
109 (3)
C-C torsion
133 (1)
C-C torsion
225 (4) D
absent
C-C-S defoimation
BIC
320 (3) D
320 (2)
C-G-S deformation
A
358 (5) P
absent
C-C-C deformation
B
425 (5) P
425 (2)
. C-C-C deformation
BIC
483 (6) P
483 (3)
C-C-C deformation
BIC
513 (4) P
513 (5)
C-C-C deformation
A
623 (5) P
623 (8)
C-S stretch
A
650 (10) P
650(6)
C-S stretch
C
680 (8) P
680 (5)
C-S stretch
B
780 (w)
786 (6) D
786 (10)
CSH angle deformation
A
870 (w)
872 (3) P
872 (3)
CSH angle deformation
B]C
900 (w)
915 (5) P
915 (4)
CH 3 rock
B
960 (w)
960 (3) D
960 (3)
CH 3 rock
C
990 (3) D
990 (4)
C-C stretch/ CH a rock,
A
1015 (m)
1015(1) P
absent
C-C stretch
BIC
1030 (2) D
1030 (4)
C-C stretch
A
1080 (m)
1080 (3) P
1080 (2)
C-C stretch/CHg rock
BIC
1110 (w)
11 15 (3) D
1115 (4)
CH 3 Rock
A
1150 (m)
1152 (3) D
1152 (5)
C-C stretch
A
1188(3) P
1188(5)
C-C stretch
A
1235 (m)
1237 (5) P
1237 (4)
CH deformation
B(C
1260 (2) D
1260 (4)
CH deformation
A
1290 (3) P
1290 (3)
CH deformation
BIC
1330 (m)
1326(3) D
1326(5)
CH wagg
A
1343 (2) D
1343 (2)
CH wagg
BIC
1360 (ssh)
1370 (s)
1368 (1) P
1368 (2)
(CH 3 ) bend sym.
1390 (ssh)
1388(2) P
1388(2)
1455 (vs)
1454 (8) D
1454 (8)
(CH 3 ) def. asym.
1472(7) D
1472 (8)
2570 (m)
2569 (9) P
2569 (9)
(S-H) stretch
2880 (s)
2860 (10) P
2860 (10)
(CH) g of CH 3 stretch
2930 (ssh)
2910 (10) P
2910 (10)
(CH) of CH stretch
2960 (vs)
2958 (5) D
2958 (9)
(CH) as of CH 3 stretch
P, polarised ; D, depolarised ; s, strong ; m, medium ; w, weak ; v, very ; sh, shoulder.
It is seen from table 1 that there are three polarised Raman bands at 623, 650
and 680 cm" 1 in the spectrum of 3-methyl 2-butanethiol. They correspond to the
three frequenc-es 620, 659 and 684 cnr 1 in 2-butanethiol which have been assigned
to the C-S stretching vibrations in the three retainers of the molecule by McCullough
et al (1958). The former three Raman bands, by analogy, represent the v (C-S)
frequencies in the three rotamers of the present molecule. It is seen from figure 3
that the relative intensities of the three bands vary with change of temperature
and this change is the largest for the band at 623 cm" 1 . The intensity of this
band increases appreciably with lowering of temperature and therefore, it is
attributed to the most stable rotamer (A) of the molecule. The observed variation
in the intensifies of the Raman bands 650 and 680 cm" 1 suggest their origin to
forms C and B respsctively. From plots of the variation of log /eW-feso and
log ^623/^eso against reciprocal of absolute temperature (figures 4, 5) the energy diffe-
rences A# are obtained as 0-25 and 0-49 kcal/mol. It is seen that form C is
more stable than form B by 250 cal/mol. This reasonably confirms the existence
of three rotanaeric forms as assumed in the very beginning.
4.2. Assignment of the vibrational frequencies
4.2a. Group vibrations: The molecule of 3-methyl 2-butanethiol with 18 atoms
will have 48 modes of vibration and 48 vibration frequencies. These may be
classified roughly in terms of vibrations of the methyl groups, the C-H group,
the CSH group and skeletal modes. To each of the three methyl groups, there
belongs three C~H stretchings, three CH deformations, two CH 3 rocking and
one H 3 C-C torsional modes of vibrations. The CH group will give rise to on
CH 3
CH
H
(B)
Gz(SH-H)
Figure 1. Three rotational conformers in 3-methyl 2-butanethiol.
CH
Figure 2. Three possible isomers in 2-butanethiol.
600
Figure 3. Variations in the intensities of Raman lines due to C-S stretching modes
of vibration of 3-methyl, 2-butanethiol. (a) 333 K, fb)260-5K, (c) 213 K.
CH stretching and two CH deformation modes while there will be one S-H
stretching vibration, one CSH angle deformation and one CS torsional mode for
the SH group. The assignment of most of these modes is straightforward and is
not given here. However, it is difficult to assign the vibrations arising from the
torsional modes. Further, difficulties are experienced in separating the CH 3
rocking modes from those arising from C-C skeletal stretching vibrations. These
are considered in the next section.
4.2b. Skeletal vibrations : The skeletal of 3-methyl 2-butanethiol molecules gives
rise to 12 vibrational frequencies in each rotamer and they may be broadly classi-
fied as C-C torsion (1), C-C-S deformations (2), C-C-C deformations (3), C-C
stretching (4), and C-S stretching (1). All these vibrations are sensitive to the
configuration of the rotameys. Some of their assignments are discussed below.
The two low frequency Raman bands 109 and 133cm" 1 observed in the Raman
spectrum of the liquid at room* temperature are believed to arise from torsional
1.9 Or
1,7:
Figure 4. Plot of log wWw vs - ] / r -
0,10
1,90
60
I.80
30 40 50
(T xj V
Figures. Plot of log / ( I>/'IM vs. 1/2".
C-C-S deformation vibration. In the present molecule, the two frequencies 225
and 320cm" 1 most probably represent this mode of vibration 1 . Since with
lowering of temperature, the former vanishes and there is little change of intensity
in the latter, the frequency 225 cmr 1 corresponds to the least stable of the forms
B and C while the latter represents on& of the two d (C-C-S) modes due to form A.
Accordingly the polarised Raman bands 358, 425, 483 and 513 cm" 1 observed in
the liquid phase spectrum of 3-methyl 2-butanethiol are assigned to this mode of
vibration. It may be noted that these bands strikingly correspond to the frequencies
377, 412, 453 and 517cm- 1 assigned to <5(C-C-C) mode in 2-butanethiol by
McCullough et al (1958). The first of these bands vanishes at low temperature
and should be attributed to the least stable form, the other two bands whose
intensities decrease appreciably on cooling should correspond to the forms B or C-
The intensity of the band 513 cm" 1 , on the other hand, slightly increases at low
temperature and thus is attributed to form A. Though not all the possible twelve
frequencies due to <5(C-C-C) modes in the retainers have been recorded, the
presence of Raman bands whose intensities vary differently with lowering of
temperature confirms the presence of at least two rotamers. From a plot of log
^sia/^ss against 1/T (figure 6) the energy difference between the conformers A and
B or C or both, is obtained as 0-34kcal/mol, which is roughly the average of
the energy difference values between (i) forms A and B, and (ii) forms A and C,
obtained from the temperature dependence of Raman bands due to C-S stretching
modes of vibration.
The frequencies due to C-S stretching vibrations have already been discussed
and those due to C-C skeletal stretching are now considered. As with the
5 (C-C-C) modes, in this case also we should expect twelve C-C stretching vibra-
tions appropriate to the three rotamers. From the data obtained from published
literature, the Raman bands in the frequency region 900-1200 cnr 1 are believed to
arise from C-C stretching modes. However, where there are methyl groups in
the molecule, the two CH 3 -rocking modes appear respectively in the region 850-
1000cm.- 1 and at about 1100 cnr 1 , which makes reliable assignment of the v (C-C)
frequencies difficult. The frequencies 915, 960, 990, 1015, 1030, 1080, 1115.
1152 and 1183cm- 1 observed in the vibration spectra of 3-methyl 2-butanethiol
certainly represent the two CH a -rocking modes and C-C stretching modes of
T.95r
1.85
1.75
1-65
55
65
310 SB Nandy and G S Kastha
vibrations in the three rotamers. Of these 1152 and 1183cm" 1 definitely belong
to v (C-C) mode and since their intensity increases when the temperature is lowered
they are associated with the rotamer A. From a comparison with the v (C-C)
frequencies observed in 1,2, ethanedithiol 1, 3 propanedithiol and 2-mercapto-
ethanol (Hayashi etal 1965 ; Nandy etal 1973a, b ;Som et al 1975) where there
are no complications arising from CH 3 rocking modes of vibration, the Raman
bands 1015 and 1030 cnr 1 are assigned to the v (C-C) vibrations, the former
belonging to the less stable forms B or C and the latter to form A. In view of
the CH 3 rocking frequencies proposed for 2-methyl 1-propanethiol (Ozaki et a!
1975; Scott etal 1958) and 2-butanethiol (McCullough etal 1958), the bands
915, 960 and 1115 cnr 1 may be reasonably assigned to this mode in the three
rotamers as shown in table 1. The two Raman bands 990 and 1080 cm" 1
may arise from either R(CH 3 ) or v (C-C) modes but their assignment is not
certain.
Some comments on the C-S-H deformation frequencies of 3-methyl 2-bulanethiol
are in order. In different alkanethiols the frequencies corresponding to these
modes have variously been put in the frequency interval 775 to 900 cm" 1 . For
example Ozaki et al (1975) has assigned the Raman band at 774 cm" 1 in 2 methyl
1-propanethiol to 5 (C-S-H) modes while Torgrimsen and Klaeboe (1970) has
proposed for this mode two Raman frequencies 778 and 814cm." 1 in
1 -propane thiol. In 1-2 ethanedithiol (Hayashi etal 1965), the two modes are at 800
and 890 cm" 1 and McCullough et al (1958) have attributed the frequency 863 cm" 1 to
the d (C-S-H) mode in 2-butanethiol. Following these observations the two
Raman frequencies 786 and 872cm" 1 observed with the present molecule are
assigned to the CSH angle deformation mode. Since the intensity of 786cm" 1
Raman band increases at low temperature it certainly originates from the most
stable A rotamer.
Acknowledgements
Thanks are due to Prof. D A Long of Bradford University and Prof. W J Orville
Thomas of Salford University, England, for their help in recording the laser Raman
spectra of the compound.
References
Freeman S K 1974 Application of laser Raman spectroscopy (New York : John Wiley) pp. 203-205
Hayashi M, Shiro Y, Oshiraa T and Murata H 1965 Bull Chem. Soc. Jpn. 38 1734
McCullough J P, Finke H L, Scott D W, Pennington R E, Cross M E, Messerly J F and
Waddington G 1958 /. Am. Chem. Soc. 80 4786
Nandy S K, Mukherjee D K, Roy S B and. Kastha G S 1973a Indian J. Phys. 47 528
Nandv S K, Mukheriee D K, Roy S B and Kastha G S 1973b Can. J. Chem. 51 1139
Determination of ionisation constants of nitrobenzidines
S ARAVAMUTHAN, C KALIDAS and C S VENKATACHALAM *
Department of Chemistry, Indian Institute of Technology, Madras 600036, India
MS received 5 August 1981 ; revised 17 June 1982
Abstract. The ionisation constants of 2-nitrobenzidine, 2,2'-dinitrobenzidine and
2,3'-dinitrobenzidine were determined spectrophotometrically in 33-3% (w/w)
methanol. The low pK value obtained for 2, 3'-dinitrobenzidine when compared
to. that of 2-nitro. and 2,2'-dinitrobenzidines is explained on the basis of electron
withdrawing nature of the nitrogroup and intramolecular hydrogen bonding.
Keywords. Ionisation constant ; spectrophotometry ; nitrobenzidines.
1. Introduction
Nitrobenzidines particularly, 2-nitrobenzidine (2-NB) and 1, 2'-dinitrobenzidine
(2, 2'-DNB) have considerable importance in the preparation of azo-dyes (Knhn
1959) and polycylic cinnoline derivatives (Braith Waite etal 1958). It is well-
known that organic electrode processes are pH dependent (Kolthoff and Lingane
1952) and in order to study the variation of polarographic half-wave potential
(i/ 2 ) of an electro-active material (depolariser) with the pH of the medium, the
determination of ionisation constant (pK) of the depolariser becomes necessary.
The importance of pK determination towards the elucidation of electro-reduction
of organic compounds is well-known (Holubek and Volke 1962; Laviron 1962).
Hence, in this paper we describe a spectrophotometric method for the determi-
nation of ionisation constants of 2-NB, 2,2'-DNB and 2, S'-dinitrobenzidine (2,3'-
DNB) in methanol-water mixtures. It may be pointed out that pK values for
these nitrobenzidines have not been reported earlier.
2. Experimental
2-Nitrobenzidine (4,4'-diamino-2-nitro biphenyl, Kovar 1964), 2, 2'-dinilrc-
benzidine (4, 4'-diamino-2, 2'-dinitro biphenyl, Porai etal 1945) and 2,3'-dinitro.
benzidine (4,4'-diamino-2,3'-dinitro biphenyl, Lefevre and Turner 1926) were
prepared according to the literature procedure. Their purity was checked by
S Aravamuthcm, C Kalidas and C S V&ikatachatam
infrared spsotra and mass spactrometry. Methanol was purified by the standard
procedure before use.
The ionisation constants of 2-NB, 2, 2'-DNB and 2, S'-DNB were determined by
a spsctrophotometric method in 33-3% (w/w) which corresponds to 40% (v/v)
msthanol-water mixtures (Chattanathan 1971). Solutions of different pH values
were prepared in 33-3% (w/w) methanol-water mixture according to the proce-
dure given by Bates etal (1963). Solutions of lower pH values were prepared
from psrchloric acid of different molalities. The pH of various solutions were
measured using a KNICK. precision pH meter (accuracy + 0-01 pH unit). The
ultraviolet absorption data were obtained with a Carl-Zeiss (ZFM4) spectrophoto-
meter for a definite concentration of each of the nitrobenzidines at different pH
values. In the case of 2, S'-DNB, the pK determination was carried out in
psrchloric acid of different molalities since it was found that no appreciable
change in the absorption was noticed for 2, S'-DNB in the pH range 0-4-2-3.
3. Results and discussion
The ioaisation constant (pK BH +) of a base (B) for the equilibrium BH + ^ B + H +
can be calculated from the relation (Bates et al 1963),
= log (C DHf /C B ) - log C H+ , (1)
where the apparent ionisalion constants, pK' BH f tends to pK BH -f. as the molar
concentration of the acid in the solvent tends to zero. In the present work, the
ionisation ratios (C BH +/C B ) of various n it robenzidines were measured spectrophoto-
metrically. Using these and the equilibrium concentration of H+ ions, the appa-
rant ionisation constants (pK' BH f) for 2-NB, 2,2'-DNB and 2,3 '-DNB were calcu-
lated from equation (1). pK BHl - for the nitrobenzidines were obtained by plotting
pK' against the acid concentration and extrapolating the linear plot to infinite
dilution. A typical plot is shown in figure 1. The spectral data are presented in
1.6
0.8
0.8
'1.6
table 1. The pK BH -t- values obtained from such a plot for 2-NB, 2,2'-DNB and
2, S'-DNB at the wavelength where the ionisation ratios (C BH +/C B ) between the
protonated and unprotonated nitrobenzidines exhibit a large difference (/I =
400nm in all the cases) were 1-76 0-02, 1-34 + 0-04 and -2-10 0-04,
respectively. It may be pointed out that pK^n- values calculated at other
wavelengths corresponding to 1 =380 and 390 nm for 2-NB and 2, 2'-DNB and
A = 410 and 430 nm for 2, S'-DNB were found to be independent of the wave-
length within the limits of experimental error.
Table 1. Absorbance (D) and molar extinction coefficients () for nitrobenzidines
at the wavelength (A = 400nm) at various pH values in 33-3% (w/w) methanol-
water mixtures. Temperature 30 0-1 C.
Substance pH D e C BH+ IC B pK'
2-nitrobenzidine
(6-0 x 10-* M)
0-45
0-120
200
4-92
1-14
0-S3
0-155
258
...
...
1-32
0-205
342
2-22
1-67
1-45
0-275
458
1-34
1-58
I'll
0-335
558
0-90
1-66
1-93
0-380
633
0-66
1-75
2-11
0-490
817
...
...
2 , 2'-dinitrobenzidine
(3-0 x 10~ 4 M)
2, 3'-dinitrobenzidine
(2-0 x 10~ 4 M)
0-17
0-11
367
9-83
1-16
0-28
0-13
433
7-13
1-14
0-45
0-14
467
6-22
1-25
0-65
0-17
550
4-65
1-32
1-04
0-28
917
1-88
1-32
1-22
0-35
1167
1-17
1-29
1-45
0-44
1467
0-67
1-27
1-92
0-56
1867
0-27
1-36
2-26
0-63
21CO
0-12
1-34
-0-61
0-31
1550
3-81
-0-03
-0-57
0-42
2100
1-85
-0-30
0-53
0-53
2650
1-03
0-52
-0-47
0-62
3100
0-75-
0-60
-0-41
0-70
3500
0-40
0-81
-0-35
0-79
3950
0-20
-1-04
The low pK BK !- values obtained for nitrobenzidines when compared tobenzidine
(Albert and Serpant 1971) may be attributed to the electron withdrawing tendency
of the nitro group. The low pK BH h values of 2, 2'-DNB in comparison with 2-NB
indicates that 2,2'-DNB is less basic than 2-NB due to the presence of additional
nitro group in the molecule. The negative value of pK DH ^- obtained for 2, S'-DNB
(-2-10 0-04) may be explained on the basis of the fact that the presence of
one of the nitro-groups either in 3 or 3' position would decrease the basic
character of the compound due to intramolecular hydrogen bonding with the
amino group in 4 or 4' position of the biphenyl ring.
References
Albert A and Serjeant E P 1971 The determination of ionisation constants (London : Chapman
and Hall) p. 58
Bates R G, Maya Paabo and Robinson R A 1963 /. Phys. Chem. 67 1933
Braith Waite R S W, Holt P F and Hughes A N 1958 /. Chem. Soc. 4073
Chattanathan N 1971 Studies on indicator acidity function in non-aqueous media, Ph.D. Thesis,
Indian Institute of Technology, Madras
Holubek J and Volke J 1962 Coll. Czech. Chem. Commun. 27 680
Koithoff I M and Lingane J J 1952 Polarography (New York : Interscience) Vol. 2 p. 624
Kovar V 1964 Chem. Abst. 60 454h
Kuhn H 1959 Angew. Chem. 71 93
Laviroa E 1962 Bull. Soc. Chim. France 418
Lefevre R J W and Turner E E 1926 /. Chem. Soc. 1759
Porai B A, Koshits and Salyamon G S 1945 Chem. Abst. 39 45998
Equilibria in the system containing chloride and
sulphates of potassium and magnesium
V R K S SUSARLA* and K SESHADRI
Central Salt and Marine Chemicals Research Institute, Bhavnagar 364002, India
MS received 13 April 1981 ; revised 27 October 1981
Abstract. Reciprocal salt-pair system 2KC1 + MgSO 4 ^K 2 SO 4 + MgCl a has been
studied at 35 C to. eliminate the discrepancies reported by different workers and
for correlating the experimental data with natural evaporation of brine (without
NaCl) so as to recover potash salts.
Keywords. Schoenite ; leonite ; kainite ; kieserite.
1. Introduction
The various processes employed for the production of potassium salts from sea
water concentrates depend upon the equilibria existing in the reciprocal salt-pair
system 2KC1-H MgSO 4 %> K^SO 4 + MgCl 2 at the desired temperature. Van't
Hoff (1918) studied the system at 25 C and 83 C and pointed out the existence
of leonite and kainite at 25 C, but the boundaries of these salts were not clear.
Autenrieth (1954) studied both stable and metastable equilibria to establish the
transition from schoenite to leonite. Further work was carried out by D'Ans
(1933) and Campbell (1934) at other temperatures such as at 0C, 55 C and
100 C. Campbell did not report the existence of kainite at 100 C, but Van't
Hoff pointed out its existence at 83 C. When the above data were used to
interpret the extraction process of potassium and magnesium salts from brine
and bittern, considerable variations were observed.
In tropical countries like India, bittern gets evaporated between 30 to 45 C
and therefore an understanding of the above system is necessary. To establish
the various equilibria the above reciprocal salt-pair system has been studied at
35 C.
2. Experimental methods
For studying the reciprocal salt-pair system at 35 C, a constant temperature
water bath rn.iintain.ed at the above temperature was used. For establishing the
time for various equilibrium stages, the contents of the system were allowed to stir
* To wham correspondence should be made. 315
of the two successive results only further progress was made in studying the system.
The common methods of analysis such as magnesium by EDTA, chloride by
Mohr's method, potassium by tetraphenylborate, sulphate by gravimetric esti-
mation of BaSO 4 were used. The EEL flamephotometer was used to check the
analysis of potassium by the above method.
In the preliminary stages the equilibria in the various connected binary systems
were obtained. The compositions of various binary solutions of the salts
potassium chloride, magnesium chloride hexahydrate, magnesium sulphate hepta-
hydrate and potassium sulphate were noted. A saturated solution of one salt was
then added to a second solid salt having a common ion to get the invariant point
as represented by the composition of the solution where the solid in contact with
the solution showed the presence of both the salts. In the more complicated
system, the third salt was added to a system where the solution was in equili-
brium with two salts. This gave the various invariant points on the diagram with
respect to the solid phase with which three salts were in equilibrium. The various
fields or boundaries were established for the double salts 'Schoenite, leonite,
kainite, kieserite etc.
In all the above cases the solid phase compositions have been obtained by
chemical methods of analysis and were confirmed by using x-ray diffraction and
differential thermal methods of analysis (DTA). This enabled checking and establish-
ing the nature of various equilibria in the system and information about the nature
of the solid phase at each stage.
3. Results and discussion
The above system, i.e., 2KC1 + MgSO 4 ^ K 2 SO 4 + MgCl 2 at 35 C formed a
reciprocal salt-pair. The analytical results obtained are given in table 1 and
diagrammatically shown in figures 1 and 2 by Lowenherz's and Janecke's method of
projection. In figure 1 the horizontal axis was used to represent 2KC1 (or K 2 C1 2 )
and MgSO 4 and the vertical ordinate to represent K 2 SO 4 and MgCl 2 and the
results were expressed in moles/1000 moles of water. In Janecke's method of
projection (figure 2) the compositions of the solutions were expressed in mole
per cent of the dissolved salts. The composition of various solutions occurring
at different invaraint points is given in alphabetical order of names.
3.1. Critical points consisting of one or two .salts with a common ion
In all the points A, B, C andD where KC1, K 2 SO 4 , MgSO 4 . 7H 2 OandMgCl 2 .
6H 2 O are in equilibrium, the results fairly agree with previous data (Autenrieth
1954). The same is true for the points E, F, G, H, I, J, K and L where the
solutions were saturated with respect to the salts having a common, ion.
3.2. Critical points representing solutions in equilibrium with three salts
As mentioned above, all the points have been obtained by starting the experi-
ments with a solution saturated with respect to two salts. Starting from point
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