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Full text of "The titration of iron by potassium permanganate"

Illinois Institute 

of Technology 

UNIVERSITY LIBRARIES 



AT 332 

-Finkelstein, L. 
The titration of iron by 
potassium permanganate 



For Use fn Librasy Only 



> 



THE TI-TRATION OF IRON BY 
POTASSIUM PERMANGANATE . 

A THESIS 
Presented by 

LEO FIMELSTEIN . ILLINOIS INSTITUTE OF TECHNOLOGY 

PAUL V. GALVIN LIBRARY 
35 WEST 33RO STREET 
To The CHICAGO, IL 60616 

President and Faculty 
■ of 
ARIVIOUR INSTITUTE OP TECHNOLOGY 
FOR THE DEGREE OP 
BACHELOR OF SCIENCE IN CHEMICAL ENGINEERING 
HAVING COMPLETED THE PRESCRIBED COURSE CP STUDY IN 
CHEMICAL ENGII^EERING 
1914. 

APP ROVED :(2^ ^4y?^^!^t^W^?^^-v^^^>^ 

Ty ^Professor of Chemical Engineering. 

APPROVED : ^^y'^^V^-^z^^^^^L-^-^.^^^^^^^ 

'Dean of the Engineering Studies. 






TABLE OF GOIffENTS. 



Introduction ---------- Page 1 

Theoretical ------------ 3 

Experimental ------------ 4 

I. Solubility of Ferrous 

Sulphate ---------- 6 

II. Titration with Sulphuric 

Acid ------------ 6 

III. Titration with Hydrochloric 

Acid ------------ 10 

Conclusion ------------- 11 



msm 



liffRODUCTIOH, 

It is v/ell knovm that ferrous salts cannot 
be oxidized quantitatively by potassium perman- 
ganate in the presence of hydrochloric acid 
without using one of the several "guard solutions" 
recommended for that purpose. The reaction with 
hydrochloric acid and the mechanism of the pro- 
tecting influence of the "guard solution" are 
in general not understood. Ostwald (Foundations 
of Analytical Chemistry) says: "the solution 
must be acid with sulphuric acid but not with 
hydrochloric acid, since permanganate oxidizes 
the latter in the presence of iron salts. It 
is a catalytic reaction 7/hich goes on here, but 
very little is known yet of the laws which reg- 
ulate it" . On account of the desirability of 
titrating an iron solution in the presence of 
hydrochloric acid from the viewpoint of the 
technical analyst, Asso. Professor B. B. Freud 

is engaged in studying the whole problem. As 
a part of this larger problem, I have titrated 



ferrous salts with permanganate in the presence 
of sulphuric acid and hydrochloric acid, with 
wide variations in hydrogen ion and chlorine ion 
concentration. This v/ork is necessary because 
of serious discrepancies in the literature. 
(Qualitatively the disturbing influence of chlor- 
ine ions has long been known, but its exact 
quantitative influence has not been determined. 

The object of this work was to get definite 
and reliable data, and the plan followed was:- 

(1) To titrate ferrous ions of a definite 
concentration in the presence of hydrogen and 
sulphate ions of definite but varying concen- 
trations. The maximum concentration of sulphuric 
acid used was determined by the solubility of 
ferrous sulphate in aqueous sulphuric acid. 

(2) To titrate the same ferrous ion con- 
centration in the presence of hj^drogen and chlor- 
ine ions of the same ion concentrations as were 
employed before. The maximum hydrochloric acid 
concentration used was determined by the ability 
to get a definite end point. 



THE TITRA TION OF IRON BY POT ASSIUId PERivlA NGANATE . 

THEORETICAL. 



If hydrochloric acid is present in the sol- 
ution when ferrous iron is titrated v/ith potass- 
ium permanganate it is knovm that the results 
obtained are inaccurate. The reason for this 
discrepancy is usually attributed to the reaction 
betv/een hydrochloric acid and potassium perman- 
ganate according to the follov/ing equation: 

10 HGltMngOy -> 2MnO -J-SHgO tSGlg 
If this reaction is regular and if discrepancies 
are due merely to liberation of chlorine, then 
the araount of chlorine evolved should be equi- 
valent in oxidizing power to the excess of per- 
manganate. ',7. G. Birch (Ghemical 'Tews, '''01.99) 
measured this excess and found that the amount 
of permanganate was less than that calculated 
from the above equation. This points to another 
relation beside that which liberates chlorine. 

Eirch advances the theory that the reaction 
betvireen hydrochloric acid and potassium perman- 
ganate is as follows: 

MngOY^-U HGl-»2 Mn Gl^**- 7H2O4.4GI2 



This would explain the supp-ression of chlorine 
v/hen raanganous salts are added, as Rice (Journal 
Chemical Socfety 1899) has shown that MnCl^ and 
free chlorine when sealed up in a tube, react 
slowly to c^ive MnCl^ . Pickering (Journal Chem- 
ical Society 1879) also has shown that the amount 
of manganic salt produced v;hen MnOp dissolves 
in hydrochloric acid was increased by the pres- 
ence of MnClo • 

The use of manganous salts for this purpose 
was first proposed by G. Zimmerman (Ber. d. Ghem. 
Ges. 1884, 15) . 

EXPERIiBNTAL. 



The ferrous salt used v/as ferrous sulphate 
(FeSO. 'THgO) , Baker's Analyzed Chemicals. Its 
analysis by the manufacturer showed a trace of 
ferric sulphate . Although ferrous ammonium sul- 
phate is much more stable than ferrous sulphate, 
it was thought desirable to use ferrous sulphate, 
as the solution of amraonium sulphate v/ould intro- 
duce NH^ ions and complicate the work. 



The potassium permanganate solution used 
v;as approximately tenth normal. It was standard- 
ized against sodium oxalate obtained from the 
Bureau of Standards. The method was as follows; 
(lI.S.McEride, Journal American Society Vol. 34.) 

In a 400cc, beaker, 0.25 gram of sodium 
oxalate was dissolved in 200cc . hot water (80- 
90°G.) and lOcc . (1:1) sulphuric acid added. The 
solution was titrated at once. The permanganate 
was added slowly, not mor~- rapidly than 10 to 15 
cc. per minutG, and the last 0.5 to l.Occ. were 
added dropwise, allowing each drop to he discolor- 
ized before the next was introduced. The excess 
of permanganate used to cause an end point color 
was estimated by matching the color in another 
beaker containing the same amount of acid and hot 

water . 

The sulphuric acid used was chemically pure 
sulphuric acid, sp.gr. 1.83, manufactured by the 
Grasselli Chemical Go. The hydrochloric acid was 
chemically pure hjrdrochloric acid, sp.gr. 1.20 
manufactured by the same firm. 



I. Solubility of Ferrous Sulphate 
in Sulphiiric Acid . 

It was necessary to determine the solubility 
of ferrous sulphate in aqueous sulphuric acid 
solution in order to titrate as concentrated sul- 
phuric acid solution as possible. One gram of 
ferrous sulphate ivas dissolved in increasing 
amounts of v/ater and concentrated sulphuric acid 
run in from a burette until precipitation of 
anliydrous ferrous sulphate occured. 

TABLa I . 

Water . Sulphuric a ci d to pre cipi tate. 

5.0----------3.50 

10.0 ---------- 9.15 

15.0 --------- 15.05 

20.0 --------- 24.15 

25.0 --------- 50.00 

30.0 --------- 63.50 

40.0--------- 200. (no precipitation) 

These values were plotted and from the curve 
it was decided to start titrating v/lth a maximum 
concentration of 160cc. sulphuric acid and 40cc . 
water . 

II. Ti tr ation with Sulphuric Acid ._ 

One gram of ferrous sulphate vv-as used for 



each titration. The procedure was as follows: 
The ferrous sulphate was weighed into a 400cc. 
beaker and the water added from, pipettes. When 
the salt was dissolved, sulphuric acid, sp.gr. 
1,83, was added from pipettes. In every case 
the total bulk of solution was 200cc . The sol- ' 
ution was cooled in a bath of running water to 
20"g. It was then titrated with Xtenth noraal 
permanganate . 

TABLS II. 

No. cc HgO cc HgSO^ NorKality. cc XN KMnO^ 



1 


40 


160 


27.0 


37.50 


2 


40 


160 


27.0 


37.80 


3 


40 


160 


27,0 


37.55 


4 


40 


160 


27.0 


37.50 


5 


45 


155 


26.2 


37.65 


6 


45 


155 


26.2 


37.30 


7 


45 


155 


26.2 


37.27 


8 


45 


155 


26.2 


37.30 


9 


50 


150 


25.3 


37.32 


10 


50 


150 


25.3 


37.46 


11 


50 


150 


25.3 


37.45 


12 


50 


150 


25.3 


37.43 


13 


55 


14 5 


24.5 


37.05 


14 


55 


145 


24.5 


37.08 


15 


55 


145 


24.5 


37.05 


16 


60 


140 


23.6 


37.35 


17 


60 


140 


23.6 


37.30 


18 


60 


140 


23.6 


37.30 


19 


65 


135 


22.8 


37.30 


20 


65 


135 


22.8 


37.30 


21 


65 


135 


22.8 


37.30 


22 


70 


130 


21.9 


37.18 



8 



TABLE II GOIfT'D 



No. 


ccHpO 


cc H2SO4 


Normality 


cc XN KMnO^ 


23 


70 


130 


21.9 


37.35 


24 


70 


130 


21.9 


37.36 


25 


70 


130 


21.9 


37.38 


26 


75 


125 


21.1 


37.30(:?.) 


27 


75 


125 


21.1 


36.80 


28 


75 


125 


21. L 


36.75 


29 


75 


125 


21.1 


36.73 


30 


75 


125 


21.1 


36.72 


31 


80 


120 


20.3 


36.80 


32 


80 


120 


20.3 


36.65 


33 


80 


120 


20.3 


36.80 


34 


80 


120 


20.3 


36.81 


35 


85 


115 


19.4 


36.40 


36 


85 


115 


19.4 


36.43 


37 


85 


115 


19.4 


36.42 


38 


90 


110 


18.6 


36.50 


39 


90 


110 


18.6 


36.50 


40 


90 


110 


18.6 


36.50 


41 


95 


105 


17.75 


36.64 


42 


95 


105 


17.75 


36.70 


43 


95 


105 


17.75 


36.70 


44 


100 


100 


16.9 


36.82 


45 


100 


100 


16.9 


36.85 


46 


100 


100 


16.9 


36.85 


47 


100 


100 


16.9 


36.88 


48 


105 


95 


16.5 


37.10 


49 


105 


95 


16.5 


36.93 


50 


105 


95 


16.5 


36.94 


51 


105 


95 


16.5 


36.92 


52 


110 


90 


15.2 


37.00 


53 


110 


90 


15.2 


36.98 


54 


110 


90 


15.2 


36.97 


55 


115 


85 


14.36 


37.00 


56 


115 


85 


14.36 


36.95 


57 


115 


85 


14.36 


36.95 


58 


120 


80 


13.5 


36.96 


59 


120 


80 


13.5 


36.94 


60 


120 


80 


13.5 


36.98 


61 


125 


75 


13.7 


36.95 



9 



TABLS 11 CONT'D. 









-^ 




No. 


CC HgO 


CC HgSO^ 


Normality 


CC XN KMn04 


62 


125 


75 


12.7 


36.95 


63 


125 


75 


12.7 


36.80 


64 


130 


70 


11.8 


37.00 


65 


130 


70 


11.8 


37.00 


66 


130 


70 


11.8 


37.00 


67 


13 5 


65 


11.0 


36.98 


68 


135 


65 


11.0 


36.97 


69 


135 


65 


11.0 


36.98 


70 


140 


60 


10.13 


36.95 


71 


140 


60 


10.13 


36.95 


72 


140 


60 


10.13 


36.95 


73 


145 


55 


9.2 


36.90 


74 


14 5 


55 


9.2 


36.92 


75 


145 


55 


9.2 


36.93 


76 


150 


50 


8.45 


3 6.90 


77 


150 


50 


8.45 


36.90 


78 


150 


50 


8.45 


36.90 


79 


160 


40 


6.75 


36.95 


80 


160 


40 


6.75 


36.95 


81 


160 


40 


6.75 


36.95 


82 


170 


30 


5.07 


36,92 


83 


170 


30 


5.07 


36.94 


84 


170 


30 


5.07 


36.92 


85 


180 


20 


3.38 


36.83 


86 


180 


20 


3.38 


36.80 


87 


180 


20 


3.38 


36.85 


88 


185 


15 


2.53 


37.05 


89 


185 


15 


2.53 


36.95 


90 


185 


15 


2.53 


36.95 


91 


185 


15 


2.53 


36.90 


92 


190 


10 


1.69 


36.80 


93 


190 


10 


1.69 


36.82 


94 


190 


10 


1.69 


36.70 


95 


190 


10 


1.69 


36.80 


96 


195 


5 


0.845 


36.80 


97 


195 


5 


0.845 


36.80 


98 


195 


5 


0.845 


36.80 



10 



The strength of permanganate used was as follows ; - 

Titration !Io. 1 to 5 incl. .0981 normal. 
" " 6 " 30 " .0979 " 

" " 31 " 98 " .0980 " 

In titrating the higher acid concentrations 

it was found that if the titration v/as carried on 

slowly the solutions turned brown thus obscuring 

the end points. If titrated rapidly the end 

point is fairly sharp. 

Ill ._ Titrat i on with Hydrochl oric Acid. 
One gram of ferrous sulphate was used for 
each titration. The water and acid were added in 
the same manner as in the titration v\?ith sul- 
phuric acid. The solutions were cooled to 20°C. 



bef. 


ore titra 


ting. 


TABLS 


III. 


ITo. 


cc HpO 


cc HGl 


Normality 


99 


199 


1 




0.164 


100 


199 


1 




0.164 


101 


198 


2 




0.328 


lOii 


198 


2 




0.328 


103 


197 


3 




0.492 


104 


197 


3 




0.492 


105 


197 


3 




0.492 


106 


195 


5 




0.820 


107 


195 


5 




0.820 



cc XN KMnO^ 



37 


.50 


37 


.53 


37 


.60 


37, 


.63 


37, 


.69 


37, 


.60 


37. 


,67 


37. 


,65 


37. 


,63 



11 



TAELS III CONT'D. 



No. 


cc HgO 


cc HGl 


Normality 


cc XN KMn04 


108 


190 


10 


1.64 


37.58 


109 


190 


10 


1.64 


37.56 


110 


180 


20 


3.28 


37.60 


111 


180 


20 


3.28 


37.40 


112 


175 


25 


4.10 


37.60 


113 


175 


25 


4.10 


37.55 


114 


170 


30 


4.92 


37.40 


115 


170 


30 


4.92 


37.50 


116 


160 


40 


6.56 


no definite 
end point. 



When the concentration of the acid increased 
beyond half normal the yellow color of the sol- 
ution obscured the end point to such an extent 
that the results could not be taken very accur- 
ately. No odor of chlorine was given off during 
the titrations . 

CONCLUS ION. 

1. It is definitely determined that var- 
iations in the concentrations of hydrogen ions 
from O.ON to 15. N in the presence of 30/, ions 
of the same concentration does not influence 
the amount of permanganate required. 

2. A concentration of hydrogen ions great- 



12 



er than 15. N in the presence of SO^ ions of the 
same concentration influences the amount of 
permanganate required by as much as 3fo. 

3. In case of a concentration of hydro- 
gen and chlorine ions greater than 0.5N the 
end point cann ;t be accurately determined. 

4. For hydrogen ion concentrations below 
0.5N in the presence of chlorine ions of the 
same concentration, ferrous salts can be titrated 
with an error of not more than one half of one 
percent . 



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