Illinois Institute
of Technology
UNIVERSITY LIBRARIES
AT 332
-Finkelstein, L.
The titration of iron by
potassium permanganate
For Use fn Librasy Only
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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 TITRATION OF IRON BY POTASSIUId PERivlANGANATE .
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 acid to precipitate.
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. Titration 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 ._ Titration with Hydrochloric 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 .
CONCLUSION.
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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