T H E OXIDATION OF SOLUTIONS O F FERROUS SALTS JAMES R. POUND The School of Mines, Ballarat, Victoria, Australia Received January 8, 1938
The author has made previously some experiments on the oxidation of ferrous sulfate solutions (1); the present work is an extension of these experiments. Solutions of various ferrous salts were exposed to the air a t room temperatures and higher, and the oxidation that occurred was found by titration with 0.1 normal oxidizing agent before and after the exposure. Generally the ferrous salt solutions (5 cc.) were placed in test tubes 2.5 cm. in diameter, and other reagents were added as desired. The quantities of all reagents are expressed in gram-moles per liter. It was necessary to run comparatively small batches at a time, but these may be readily correlated. Selected results with 0.1 molar solutions of ferrous salts are given in tables 1 and 2; table 3 shows results obtained with other than 0.1 molar solutions. I n each series of experiments the “blank”, i.e., the cubic centimeters of N/10 oxidizing agent equivalent to the original solution, is given after the heading of the series in parentheses: e.g., in table l(a) the 5 cc. of original ferrous sulfate solution was equivalent to 3.38 cc. of N/10 permanganate solution (this is an exceptionally low value). Duplicate experiments are not quoted, and many confirmatory tests are omitted. In each series the temperature varied over a few degrees, the average being quoted; but the experiments of each series were, of course, performed under the same conditions. EXPERIMENTS WITH AQUEOUS SOLUTIONS OF FERROUS SALTS
The ferrous sulfate solution of table l(a) was part of a large stock which was originally 0.1 molar in ferrous sulfate and 0.05 molar in sulfuric acid. Thus the presence of considerable ferric salt does not affect the oxidation. The solutions of table 2(a) were made from an exact weight of ferrous sulfate crystals in each test tube. Here the solution without any added acid gave a precipitate within 3 hr.; after 5 days it contained an appreciable brown precipitate, but the liquid waa colorless; the other solutions containing acid were clear and colorless throughout. Solutions of series l(b), l(c), 2(c), and 2(d) were made from ferrous sulfate crystals and were filtered immediately from a slight precipitate; 955
956
JAMES R . POUND
TABLE 1 Results obtained with N / 1 0 solutions of ferrous salts at room temperature The extent of oxidation is expressed in terms of cubic centimeters of N/10 oxidizing agent; the concentrations of the reagents in gram-moles per liter. The oxidation equivalent of the ferrous salt taken (and thus also of complete oxidation) is given in parentheses after the heading of each series
I
-1-
0.02
_(0)
plus CaCOa
Fa06 solution; 7 days a t 7'C.
-_
1.9 0.1 0.05 0.1 0.1 0.1 0.05 0.05 0.1
HBOI
HzS04 HaPOi HsPOd Hap01 HaPOi NHiCl 2.5 NaCl 0.02 NaNOl 0.8 0.5 2.7 0.3 0.5 1.0
1'
0.12 0.W '
(5.34 cc.) OXIDATION
0.07 0.10* 0.33 0.14 2.09' 3.43 2.01 2.77 0.04*
0.14, 2.27
I
1 1
/I
1
0.1 HzSOl 0.8 HzSO, 0.17HCl 0.01 NaN02
0.01 0.01
0.01 0.01
( 9 ) Fe(ClO,)r solution;
0.03 0.05 0.06 , 1.05 ~
14 days at 12'C.
(5.40 ca.)
( i ) Fe(CtHtO+ solution; 1 day a t 9%.
1
HC~H~OI
REAQENT
0.045 HzSOa 0.095 HzSOc 0.15 NaC:HsOz plus CaCOs
( 4 . M cc.)
1
OXIDATION
0.48t 0.13t 0.17t 0.00 0.86t 0.58t
9 57
OXIDATION OF FERROUS SALT SOLUTIONS
TABLE 1-Continued (j) Fe(C&hOh solution; 3 days at 8'C.
HCrHz0i
i
(5.16
00.)
(1) Fe(CrHz0rh solution: 1 and 2 days at 8°C. (4.29 eo.)
1 OXIDATIOI
REAQENT
1.26t 1.37t 1.25t 0.95t 0.48t 0.39t 0.05
(0) (0 ) 0.01 0.01 0.01 0.01 0.01
(k) (CHEO0)zFe solution; 1 day at 12'C. (4.59 cc.) REAQENT
OXIDATIO>
0.1 HI SO^ 0.04 HsPOi 0.1 HsPO4 0.32 HsPO4 0.07 (?) ( C H Z C O O H ) ~ 0 . 1 (1) (CHzC0OH)z 0.05 (CH&OONa)z 0.11 (CH&OONa)z 0.12 NaCzHaOz
0.69t 0.04 0.67t 0.40t 0.27 0.44t 0,241 0.727 0.77t 0.86t
IXIDATIOh (1 DAY)
REAQENT
0.1 0.0075 0.015 0.030 0.12 0.21 0.37 0.0089 0.027 0.054 0.18 0.44
( m ) FeSO, and F e d ),)I solution: 2 and 4 days at 18'C.
HBO,
0.05 0.10 0.05 0.05 0.05 0.05 0.05 0.05 0.50 0.05
HnCd).
0 0 0
Citric acid
0
0 0 0.001 0.005 0.01 0 0 0 0.005
0 . 8 HsPGI
0.005
0
0 0.0004 0,002 0.004
2 days in
the light
2 days in the ..
laboratory
0.13 0.01 (-0.25) (- 1.08)
0.77t 0.94 0.97* 0.90'
0.38 0.25 0.26 0.21
0.62 0.53 0.37
0.12t
(5.8
,
I 1laboratory days in the I
1 ~
(-0.10) (-0.04) (-0.34) (-0.70) (-0.06)
0.73
(REDUCTION IS OWEN I N PARENTEEBES WITH MINUS SIQNS)
OXIDATION
REAGENT8
0.38 0.01 0.39 0.43 0.43 0.55
H&O4 NaCzHiOz NaCzHsOz NaCzHsOz NaCzHaOz NaCzHsOz NaCtHaOz HCzHsOz HCzHsOz HCzHsOz HCzHaOz HCzHaOz
)XIDATION ( 2 DAYS)
(-0.16)
1
(-0.14)
1
(-0.14)
I
1.19
0.10
1
4 days in
the dark
0.08 0.00 0.07 (-0.01) (-0.10) 0.03
958
JAMES R . POUND
TABLE 1-Concluded (n) Fa904 solution; 16 days at 17‘C. in the dnrk (5.21
00.)
OXIDATION
Sulfuric .wid
0.05 0.05 0.05 0.05 0.05 0.05 0.05
Citric wid
0 0.004
0.014 0.025 0.1 0.2 0.5 0.0
Sulfuric acid
O d i c acid
0.21 0.23 0.295 0.24 0.34 0.43 0.69
0.05 0.05 0.05 0.05 0.05 0.05 0.05
0 0.004 0.016 0.032 0.1 0.2 0.5
0.21 0.27 0.29 0.41 0.38t o.29t 0.42t
0.55 0.55 0.55
0.0 0.032 0.1
0.16 0.21 0.84$
0.55 0.55 0.55
0.1
0.16 0.17 0.23
0.05
, 0.1 &PO&
4.13
0.026
In these experiments the concentration of the ferrous salt was lowered, generally halved, by addition of the reagent solution, t Precipitates were formed here soon after the start of the experiment. $ A trace of precipitate was formed before the end of the experiment. This value is certainly high.
more precipitate formed in these solutions when they were kept at room temperature. In the first experiment of table l(b) the pH fell from 3.9 to 3.0, and in the first experiment of table I(c) the pH fell from 3.5 to 3.3. The first experiments of table l(k) and table l(1) had p H values of 5.0 and 2.5, respectively. I n series 2(c) after 6 days the solution with no added acid was yellow and contained considerable yellow-brown precipitate; the next solutions with increasing acid were decreasingly straw-colored; and the last solution was colorless. The first solution in 2(c) changed during the 6 days from pH 3.9 to 2.6, Le., from 0.00005 M to 0.001 M sulfuric acid; this increase in acidity proceeded a t a diminishing rate, i.e., the acidity developed most rapidly at the start. The first solutions in 2(d) and 2(i) had p H values of 3.5 and 2.1, respectively. In series 2(d) the solutions with less than 0.005 molar acid gave precipitates; here the more alkali was used to reduce the acid concentration, the greater was the amount of precipitate, but the oxidation was not increased proportionately. The alkali probably precipitated a proportional amount of basic ferric salt early in the experiment, and thus the total oxidation would be the same as that occurring in the original “neutral” solution. When the ferrous salt solutions were in contact with solids like calcite, zinc oxide, and magnesia, which would continuously reduce the acid concentration, the solutions were rapidly oxidized.
TABLE 2 Results obtained with N/10 solutions of jerrous salts (a) F a , salution: 5 days at 28%. H90r
(f) FeCL solution; 4 days at 26%.
(5.2 cc.)
0.45t 0.30 0.31 0.29
(0)
0.05 0.1 0.2
-
0.1 1.9 3.2 4.2 5.0
(b)F a r solution; 8 days at 27.C. (3.13 CD.) -
(g)
OXIDATION
____ 6.0
-
FeCL solution; 7 days at 25'C.
1 ____ HC1
-
0.28 0.13 0.10 __
1
O.ooOo5 0.005 0.01 0.02 0.05
o.l
(O")
0.54
(O")
i (6.34 00.)
/
(0.1)
-
( e ) FeClr solution; 4 days at 25'C.
HC1
-_-
RXAQXNT
__
0.1 0.32 1.9 0.1
0.1
0.1
1.99 3.22. 4.06.
(4.88 w.)
-AGENT
OXIDATION
BEAQlENT
1 OXIDATION 0.23 1.03 0.16 0.53* 0.16* 0.56
0.06 HsBOa 1.8 MgCl.2 2.3 NaCl 5.6 NHiCl
1
Platinum black
1
4.00
(i) Fe(N0i)r solution; 3 days at 26°C. (5.17 cc.) (4.92
REAQENT
00.)
-___._.
0.11 2.7 Hap04 0.8 HISO,
I
0.1 1.9 0.1 0.1 0.1 0.1
0.01 HCl 0.01 Hd?O4 0.014 HsBOs 0.0001 NaOH 0.0003 NaOH 0.0006 NaOH 0.001 NaOH
0.11
(h) FeClt solution; 8 days at WC. (4.91 00.)
HCI
-
-
0.10 0.30 0.33 H a 0 8 0.14 0.2q 0.05 NarB40, 0.35 HCiHaOI 0.07 0.10 1.7 HCzHrOt 0.15 N a G H ~ O ~ 1.08t o.wt plus CaCOa plus ZnO 1.42t plus MgO 3.60t plus BaCOs 1.31t
1.1 0.1
0.52
(d) F&OI solution; 3 days at 28'C.
1
,
_-
0. 00015 0.005 0.00015 0.00015 0.00015 0.00015 0.00015 0.00015 0.00015
(4.92 a?.)
1 OXIDATION
1
4.00
0.02 HaPo, 0.04 HaPo, 0.1 HIPOI 0.2 H I P O ~ plus CaCOa 0.03 NaCzHaOz 0.12 NaCsHaOz
OXIDATION
0.12 o.nt 0.36t 0.87t 1.19 0.6t 0.12
2.1t * In these experiments the concentration of the ferrous salt was lowered, generally halved, by addition of the reagent solution. t In these cases a precipitate formed during the experiment. 959
960
JAMES R . POUND
TABLE 3 Results obtained with ferrous salt solutions of various concentrations (a) Various ferrous aalt solutions; 14 day8 at 12'C.
!
BOLUTION
ADDEDREAOENT
1
OXIDATION
M/10 F e ( S O a ) z . I ., . . . . , , , . . . . . , , , . , . . . ,
None 0.2MHNOs
0.00 0.03
M/10 Fe(C104)z.,, , , . , , . . , , , . . . , . , , , . , .
None 0.2 MHNOs
0.04 0.07
M/4 Fe(NO& , . . . . . . , . . , , . , . , . . . . , . . . .
None 0.2 MHNOs
0.22 0.39
., . . , . ,..........
None 0.2 MHNOs
0.14
M/4 Fe(CIO,)Z.,
, , , ,
. .,
M/2 Fe(SOs)z.. , , , . , . . . . . .
,
,
{ . { . { { {
Sone 0.2MHSOs
..........
0.21
~
0.88 1.27
(b) Ferrous phosphate solutions; 18 days at 11'C. SOLUTIONE
(1
M/12 Fe(HzP04)z
1
M/5 Fe(HzP04)z, . . . , . . . . , , . , . , . . , , . . . .
ADDED REAOENT
0.16MHaPO4 3.0 M HsP04 0.4 MHaPOi 3.0 M HaPo,
~
I
,
I
1 I
~
OXIDATION
2.69t 3.87 7.75 6.90
1
M/2 Fe(HzPO,)z.,. , . . , . , , , . . . . . , . . . . . . (0)
1.0 M HsPO, 3.0 M HaPOi
I
1
20.53t 12.75
M / 2 Fe(N0s)z solution OXIDATION
ADDEDREAQENT
0.2 M HSOs 0.6 MHNO, 0.1 M Ha904 0.45 M His04 0.32 M HaPo, 1.5 MHaPO, 2.7 MHaPO, 0.03 M HsBOa 0.06 M HaBOa 0.33 M HaBOs 0.16 M (CH&OOH)z 0.35 M HCIH~OI
7 days at 11%
3 days at 2s'C.
0.25 0.90 1.66
0.59 2.06
4 days at 30%
1.69 2.21
0.59 3.57 3.94
6.28 8.78 23.44 (complete) 0.96
0.24
t In these c u e s a precipitate formed during the experiment.
1.14 1.23 1.27
OXIDATION OF FERROUS SALT SOLUTIONS
961
All the yellow to brown precipitates formed would seem to be basic ferric salts, but in 2(g) with magnesia the precipitate was greenish, and in series I(c), l(k), 2(d), and 3(b), where phosphoric acid was used, the precipitate was white (ferric phosphate?). The M/2 ferrous nitrate solution would contain about 0.04 molar nitric acid and a little barium nitrate; the other ferrous nitrate solutions, made by dilution, would contain proportional amounts of these substances. The ferrous perchlorate solution, made from iron and perchloric acid, contained a little of the free acid,-about 0.05 molar acid in the 0.25 molar solution of ferrous perchlorate and 0.02 molar acid in the 0.1 molar solution of the perchlorate. The ferrous phosphate solution was prepared from iron and phosphoric acid; unless there is an excess of acid present a white solid (ferrous phosphate?) separates during the preparation. The 0.5 normal ferrous phosphate solution would contain 0.5 molar Fe(H2P04)2 and about 1.0 molar free (excess) H8POa, while the 0.2 and 0.08 normal ferrous phosphate solutions would contain proportional amounts of these constituents. These solutions form precipitates on standing, especially the two dilute solutions; addition of more phosphoric acid is necessary to prevent this precipitation. It is evident that ezact formulation of the constituents of these solutions requires a study of the pH, etc. Ferrous acetate solutions,were prepared from ferrous sulfate and barium acetate solutions, the latter being in slight excess. This excess of barium acetate, 0.005 gram-mole per liter, is allowed for in tables l(i) and l ( j ) . These ferrous acetate solutions oxidized very rapidly at room temperatures, considerable brown precipitate and scum forming within 24 hr.; this is undoubtedly due to the weakness of acetic acid. It is reported that ferrous bicarbonate solutions are oxidized similarly with the greatest readiness. The ferrous acetate solution (l(1)) had a n initial p H of 2.5, which changed to 2.1 in 2 days. Ferrous succinate solution was prepared from ferrous sulfate solution and barium succinate, but the reaction was incomplete, and the solution (l(k)) contained 0.02 mole of sulfate to 0.08 mole of succinate radical. This solution oxidized rapidly, forming much dark brown precipitate and scum; the initial pH was about 5 . I n the two experiments in l(k) where free succinic acid was added, it did not dissolve completely, and the concentrations quoted are approximate. Only the ferrous chloride solutions were titrated with potassium dichromate solution. The solutions of the other ferrous salts were titrated with potassium permanganate solution; but the solutions containing oxalic and citric acids, those of l(m) and l(n), were titrated with ceric sulfate solution.
962
JAMES R. POUND EXPERIMENTS WITH OTHER SOLUTIONS
A . Ferrous chloride in alcohol Anhydrous ferrous chloride was made; it was insoluble in ether, but freely soluble in alcohol. The solution in absolute alcohol (99 per cent by weight) turned brown rapidly in the air, but the addition of a little concentrated hydrochloric acid (aqueous solution) made the alcoholic solution a clear yellow color. The stock alcoholic solution was 0.5 molar in hydrochloric acid and 0.4 molar in ferrous chloride; the more dilute solutions were made by diluting this stock solution with alcohol. I n these solutions the ferrous salt was determined by diluting with water, distilling off the alcohol in a stream of carbonic acid gas, cooling, diluting, and titrating with potassium dichromate solution. The 0.4 molar solution of ferrous chloride in alcohol oxidized to the extent of 15 per cent in a couple of hours, and of 36 per cent in 1 day a t room temperatures. With 0.4 M , 0.16 M , and 0.06 M solutions in alcohol the oxidation in 6 days a t 11°C. was 90 per cent complete, and in the presence of extra 2 molar hydrochloric acid (aqueous) it was quite complete. Other experiments indicated that in 89 per cent (by weight) alcohol the ferrous chloride oxidized somewhat more slowly than in the 99 per cent alcohol. Ferrous chloride thus undergoes oxidation much more rapidly in alcoholic solution than in aqueous solution. Addition of hydrochloric acid hastens the oxidation in both solvents.
B. Ferrous sulfate in concentra&d sulfuric acid Ferrous sulfate crystals were shaken with concentrated sulfuric acid for 2 days, and the solution was filtered through a porous fritted-glass plate. The resulting slightly turbid solution of ferrous sulfate in about 85 per cent (by weight) sulfuric acid was kept in a current of dry air for 4 days a t 28OC. There were crystals in each sample a t the end of that time, and it was therefore doubtful whether the ferrous sulfate was ever all dissolved. The original 5 cc. of solution (8.66 9.) required 1.36 cc. of N / 1 0 permanganate; under the conditions stated the oxidation was equivalent to 0.24 cc. of permanganate. With the addition of phosphoric acid (90 per cent acid) to give a concentration of 0.65 molar, the oxidation was equivalent to 0.48 cc. of permanganate. Oxidation in this solvent is therefore proportionately greater than in water, and even here phosphoric acid accelerates the oxidation. C. The oxidation of aqueous ferrous sulfate solutions in the presence of oxalic and citric acids T o prevent the oxidation of ferrous salt solutions the addition of citric acid has been proposed, but citric acid and similar acids have the power of
OXID.4TION O F
FERROUS
963
SALT SOLUTIONS
reducing ferric salts, especially in sunlight; moreover, the simple estimation of ferrous salts by permanganate and dichromate breaks down in the presence of these organic acids. A drop of ferric salt solution mixed with one of oxalic acid, both M / 2 0 , after standing a few minutes in ordinary laboratory light, gave a blue color with potassium ferricyanide, while under similar conditions oxalic acid only slowly reduced potassium ferricyanide. With potassium permanganate solution the titration of ferrous salt in the presence of oxalic acid, even in the presence of phosphoric acid in the cold, was out of the question. With potassium dichromate solution, using potassium ferricyanide as outside indicator, the ferrous salt could be estimated in the cold with moderate accuracy, say 1 per cent, if the indications of the outside drop were observed a t once, Le., within half a minute; in a longer time the blue color appeared, since the oxalic acid then reduced the ferric salt, and light would accelerate this reduction. Thus with ferrous sulfate solution (0.0587 M ) and oxalic acid (0.0837 M ) , in the presence of sulfuric acid (except where phosphoric acid was used), the following results were obtained: FeSO, SOLCFXON cc.
H&Lh
BOLUFXON
cc.
~
1
Haor (16 M )
i
KMnOd TITRATION Cold
~
Hot
cc.
0
~
20
(11.8)*
I
16.24 27.61
~~
cc.
20
20
O 0
i
I
9.25 9.25 9.35
*Fugitive.
Though cold oxalic acid solution was fairly rapidly oxidized by permanganate and ceric sulfate and only slowly by dichromate, yet the mixture of ferrous salt and oxalic acid was oxidized fairly rapidly by the dichromate, or the ferrous salt could only be approximately estimated thus in the presence of oxalic acid. Similarly, potassium dichromate oxidized ferrous salt and citric acid simult*aneously,and there was no direct relation between the amount of iron oxidized and the amount of citric acid oxidized. As potassium dichromate oxidized citric acid solution very slowly in the cold, this simultaneous oxidation is an example of induced reaction, or the oxidation of ferrous salt catalyzed the oxidation of citric and oxalic acids.
964
JAMES R . POUND
Ferrous salts, however, could be titrated by ceric sulfate solution in the presence of oxalic and citric acids without appreciable error. I n these titrations diphenylamine, dissolved in sulfuric acid, was used as indicator. With diphenylamine the titration must be taken to a definite purple color; too rapid a titration, too great a dilution, insufficient acid (both sulfuric and phosphoric acid), and the formation of a precipitate of ferric phosphate, all made for high results, while an excess of sulfuric acid made for a low result and a less sensitive one. Considerable care was needed to obtain consistent results with this indicator, and undoubtedly the recently introduced indicators would be better. The best results were secured in a solution which was from 0.2 to 1 or 2 molar in both sulfuric and phosphoric acid. The following results were obtained: CITRIC ACID
cc.
5 5 5 5
10 10 10
cc
.
Ce(S0r)r USED
IN
TITRATION
ce.
0 5
20
1 g. 0 10
10 cc. of M/10
2.07 2.04* 2.08*t 2.05 4.08 4.07 4.07
* This end point faded in the sunlight.
t A rapid titration gave higher results. Using the ceric sulfate titration, experiments on the oxidation of ferrous sulfate solutions in the air in the presence of oxalic and citric acids, and in the presence and absence of light, were carried out. In series l(m) 10-cc. samples of 0.05 molar ferrous sulfate solution, which was also 0.025 molar in ferric sulfate, were used. The experiments “ in the light” were done indirectly in sunlight in the daytime and in electric light a t night; those “in the laboratory” were done in diffused light in the daytime only; those “in the dark” were kept in a dark cupboard. I n series l(n) freshly made 0.1 molar ferrous sulfate solution was used, and the oxidation was allowed to proceed a t room temperatures, the mean temperature being 17’C., and in the dark; 5 cc. of the ferrous sulfate solution was used in each test. The solutions containing appreciable quantities of oxalic and citric acids, over 0.05 molar, became yellow more or less rapidly, and those containing much oxalic acid soon gave a precipitate of yellow ferrous oxalate, FeC204.2Hz0; this separation of insoluble ferrous salt decreased the extent of oxidation (see the three ex-
OXIDATION OF FERROUS SALT SOLUTIONS
96 5
amples markedt in table l(n)); in the fourth example (marked:) in the more concentrated sulfuric acid solution the precipitate only appeared towards the end of the experiment and the oxidation of the ferrous salt was not retarded. SUMMARY
1. Phosphoric acid accelerates the oxidation of ferrous salts. When the concentration of phosphoric acid is high, the rate of oxidation slackens a little. 2. Hydrochloric acid increases the oxidation of ferrous salts. At room temperature the increase in oxidation is slight up to 1 molar hydrochloric acid, but thereafter it rises at an increasing rate; a t 25°C. the influence of hydrochloric acid is more marked. 3. Sulfuric acid does not increase the rate of oxidation of ferrous salts, except that a t concentrations above 18 molar the rate of oxidation is increased slightly. 4. Neutral solutions of ferrous salts are oxidized more rapidly than acid solutions (Le., than moderately acid solutions). At 25°C. the rate of oxidation decreases from that of a neutral solution until a concentration of sulfuric acid of about 1 molar is reached, or a concentration of hydrochloric acid of about 0.5 molar. At room temperatures the rate of oxidation decreases until the acid concentration is about 0.1 molar. 5. The ferrous salts of weak acids are oxidized more rapidly than those of strong acids; e.g., there is rapid oxidation with ferrous acetate solutions, with ferrous chloride solutions buffered with borax or sodium acetate, with ferrous salt solutions containing calcite and other substances which would remove free acid. 5a. Similarly, 0.1 molar sulfuric acid prevents the rapid oxidation of ferrous acetate solutions a t room temperature; the addition of acetic acid to ferrous chloride solutions containing 0.1 molar hydrochloric acid has no effect on the oxidation; boric acid behaves similarly. Sulfuric acid, and even phosphoric acid, reduces the rate of oxidation of ferrous succinate solutions. 6. Phosphoric acid may increase the rate of oxidation of ferrous salts partly because it is a relatively weak acid, but its main effect seems more specific. Phosphorous acid and hypophosphorous acid also increase the rate of oxidation. The oxidation of ferrous phosphate solutions is particularly rapid. 7. At room temperatures small concentrations of nitric acid, up to 0.5 molar, affect the rate of oxidation just as sulfuric or hydrochloric acid does, i.e., very slightly. Above this concentration the nitric acid oxidizes the ferrous salt rapidly. Nitric and hydrochloric acids together in small
966
JAMES R. POUND
concentrations (less than 0.2 molar) do not deet the oxidation. The rate of oxidation of ferrous nitrate solutions is similar to that of solutions of the chloride and the sulfate. 8. Oxides of nitrogen (from acidified sodium nitrite) rapidly accelerate the oxidation of ferrous salts. Platinum black also catalyzes the oxidation (1). 9. Neutral salts make little difference in the rate of oxidation of ferrous salt solutions. The presence of a considerable quantity of chloride increases somewhat the rate of oxidation of ferrous chloride solution. 10. No relation can be seen between the rate of oxidation and the solubility of oxygen in the various solutions; the latter must be secondary to the other factors involved. Similarly, the formation of a precipitate (generally of basic salt) in the solutions is due to low acid concentration and thus accompanies high rate of oxidation, although it does not seem to cause the latter directly. 11. The rate of oxidation (see table 3) is proportional to the square of the concentration of the ferrous salt. In table 3(a) the rate of oxidation of the M/4 solutions is six times that of the M/10 solutions, and that of the M/2 solution is four times that of the M/4 solutions (see also reference 1). 12. The addition of acid, up to a concentration of 0.2 molar, increases the rate of oxidation of the M/2 and M/4 solutions. Hydrochloric and sulfuric acids have the same effect as nitric acid; hence the experiments with the first two are not reported. 13. Oxidation of the ferrous phosphate solutions is very rapid; in 16 days at room temperatures it is about 75 per cent complete with the solutions containing the least phosphoric acid. Addition of phosphoric acid decreases the rate of oxidation for the concentrated ferrous phosphate solutions, but increases it for the dilute solutions (M/12). I n these experiments (table 3(b)) the rate of oxidation is not proportional to the square of the ferrous salt concentration, but the exact “free acid” concentrations in these solutions are difficult to determine. 14. Addition of sulfuric, hydrochloric, or nitric acid, up to 0.2 molar, decreases the rate of oxidation of the M/2 ferrous phosphate solution by 20 per cent, corresponding to 16.5 instead of 20.53 cc. of permanganate (compare 5a and 21). 15. With the addition of 2.7 molar phosphoric acid the oxidation of M/2 ferrous nitrate solution was completed within 4 days a t 30°C.(compare table 2 ( e ) ) . 16. Boric, succinic, and acetic acids decrease slightly the rate of oxidation of M/2 ferrous nitrate solution a t 30°C.; such a decrease is not notieeable in tables 1 and 2: 17. I n the dark, ferric salt solution is not reduced by citric acid and is only slightly reduced by oxalic acid.
OXIDATION OF FERROUS SALT SOLUTIONS
967
18. I n the light, citric acid and oxalic acid reduce ferric salt continuously; the reduction is proportional to the concentration of these acids. In equimolar solutions the reduction by citric acid is greater than by oxalic acid. 19. Addition of sulfuric acid prevents the reduction by citric acid. Phosphoric acid accelerates the oxidation of the ferrous salt even in the presence of citric acid. 20. In the absence of light, oxalic and citric acids increase the rate of oxidation of ferrous salt solutions. The effect is slight a t concentrations less than 0.05 molar, but greater at higher concentrations. 21. This effect of oxalic and citric acids is diminished by increasing the sulfuric acid concentration (see 5 and 5a). 22. I n general, then, oxalic and citric acids behave 1t9 the other organic acids (weak acids), but the formation of insoluble ferrous oxalate and especially the influence of light introduce further complications with these two acids. CONCLUSIONS
No simple or single explanation will account for the facts summarized above. I n very dilute acid solutions some relation between the hydrogenion concentration and the rate of oxidation waits to be unravelled, but in concentrated acid solutions and in non-aqueous solutions the hydrogen ions may play a minor part in the oxidation. High rates of oxidation are undoubtedly caused by ( a ) specific catalysts, such as phosphoric acid, hydrochloric acid (in high concentrations), nitric acid (direct oxidizer), oxides of nitrogen, and surface catalysts such as platinum black, and (b) low acidity. A solution of a ferrous salt undergoes the minimum oxidation in the presence of a moderate amount of a strong non-oxidizing acid, like sulfuric and hydrochloric acids; a concentration of such an acid of about 0.2 normal is suitable, or of sulfuric acid even up to 6.0 normal. REFERENCE (1)
POUND:J. Soc. Chem. Ind.
66, 327T (1936).
