THE OXIDATION O F OXALIC ACID I N THE ABSENCE O F OTHER ACIDS BY J. C. WITT
In connection with some work on the various direct and indirect methods for the determination of calcium in cement raw mixtures,l occasion arose to study in some detail the precipitation of calcium as the oxalate, and the titration of oxalic acid and its salts with potassium permanganate. Some phases of the work were suggested by a previous study of oxidation-reduction reaction without the addition of acid. In determining calcium by the permanganate titration method, the calcium is precipitated from an ammoniacal solution by adding an excess of ammonium oxalate. The calcium oxalate is collected, dissolved in an excess of dilute sulphuric acid, and the oxalic acid thus liberated titrated with standard permanganate solution. Another procedure is to precipitate calcium with a standard oxalate solution, and after filtering out the calcium oxalate, acidify the filtrate with sulphuric acid and titrate the oxalic acid present.3 I n either case, the reaction is represented by the equation : 2KMn04 5HzC204 3HzS04-+ . K2SOa 2MnS04 8Hz0 locoz. It is generally believed that this reaction can only take place in the presence of a strong acid. Hydrochloric acid is avoided because of its action on permanganate and nitric acid should not be present because i t is an active oxidizing agent. Consequently sulphuric acid has come into general use. Oxidatiolz of Oxalic Acid by Potassium Permanganate in the Presence of Acetic Acid.-Acetic acid may be substituted for sulphuric acid in the reaction. As a preliminary qualitative test, some ammonium oxalate solution was acidified with acetic
+
+
+
+
+
Philippine Jour. Sci., 15, 107 (1919). (1915); 38,47 (1916); 40, 1027 (1918). Mohr: “Titrirmethode,” p. 239. Kraut: Chem. Centr., 1, 316 (1856).
* Jour. Am. Chem. SOC.,37,2360
436
I
J . C. Witt
acid and the mixture was warmed to about SO" C. A small amount of potassium permanganate was then added and the mixture was stirred vigorously for some seconds. No change in color was observed, and the result was recorded as negative. The beaker was set aside. Some hours later when the beaker was about to be emptied and washed, it was noted that the permanganate color had completely disappeared. Quantitative measurements revealed that satisfactory titration can be made, and that the stoichiometric relations are the same as when sulphuric acid is present. However, the procedure, and the color of the solution when the end-point is reached are somewhat different. It is a common experience that in the presence of sulphuric acid the color of the first portion of permanganate fades slowly and considerable stiring is necessary. After this, the reaction is practically instantaneous, and the titration proceeds as rapidly as the permanganate is added. The temperature of the solution being titrated constantly decreases by radiation, and because a cooler liquid is being mixed with it. When the titration is made in the presence of acetic acid, the temperature must not be allowed to fall-otherwise the velocity of the reaction decreases to such an extent that it is impossible to finish the titration in a reasonable time, if a t all. The longer the end of the titration is delayed, the greater is the decrease in temperature (as the titration is ordinarily carried out) and of course this causes a still greater loss of time, and so on. It is therefore necessary io keep the mixture heated almost to the boiling point during the entire titration. This was done at first by placing the beaker, containing the solution being titrated, on the hot plate several times for reheating. This was not completely satisfactory. Later an electric hot plate about six inches in diameter was placed on the base of the burette stand. The contents of a beaker were first heated to boiling on a gas hot plate and then transferred one by one to the electric hot plate and the titrations made while the liquid continued to boil. When not actually in use, the burette was raised several inches, and the current turned
Oxidation of Oxalic Acid
437
off so that there was no error due to expansion of the solution in the burette. There was no appreciable error from volatilization of oxalic acid. Under these conditions, the titration was almost as rapid as in the presence of sulphuric acid. The electric hot plate was used in the same manner in all the titrations subsequently discussed in this paper. There is a difference in the end-point, however. The permanganate color does not fade directly to colorless but there is an intermediate brown color. As the titration proceeds, this becomes more and more apparent, and longer time and more vigorous stirring is necessary to dissipate it. The end-point is sharp but the color at this stage is not the characteristic pink. It is rather a rose color caused by a blending of the pink of the excess permanganate and the brown of a colloidal solution of manganese compounds. Oxidation in the Absence of Other Acids.-Neutral ammonium oxalate is not oxidized by potassium permanganate, at least under ordinary conditions of concentration, temperature, and the like. However, the fact that the necessary hydrogen ion concentration could be supplied by a weak acid such as acetic suggested that i t should be possible to titrate oxalic acid alone with permanganate, the necessary hydrogen ion concentration being supplied by the oxalic acid itself. It has been shown that oxalic acid alone, may be completely oxidized to water and carbon dioxide by permanganate;I and Neidle and Crombie2 found that ferrous salts may be completely oxidized by’permanganate, without the additions of acid, though the reaction is not suitable for a method of titration because the end-point is masked by colloidal substances. Preliminary tests showed that in starting to titrate oxalic acid alone with permanganate, the behavior is identical with that observed in the presence of sulphuric acid. The color caused by the first addition is slow to fade, but after this the color produced by additional permanganate disappears rapidly. I
Moranski and Stingl: Jour. prakt. Chem., 18, 83 (1878). Jour. Am. Chem. SOC.,38,2607 (1916).
438
J . C. Witt
However, considerably before the equivalent amount has been added, the solution becomes turbid, and then becomes clear on stirring. As the titration proceeds the turbidity appears with smaller additions of permanganate, and more stirring is necessary to dispel it. Finally a point is reached at which the solution does not become clear on continued stirring. It remains turbid when boiled for 20-30 seconds. At the end of this time the colloid coagulates and a slight permanent precipitate is formed. A further addition of permanganate immediately produces a dense chocolate-colored precipitate. The formation of the first permanent precipitate is taken as the end-point of the titration. To avoid confusion in the discussion that follows, the true end-point, obtained by titrating with permanganate in the presence of sulphuric acid and completely oxidizing the oxalic acid, is termed the “acid” end-point, and the end-point just described, obtained in the absence of any acid other than oxalic is termed the “non-acid” end-point. The reaction was studied further to determine whether or not the non-acid end-point could be obtained with sufficient regularity and accuracy to find out the mechanism of the reaction, and obtain a formula for calculating the acid end-point from the non-acid end-point. Solutions of potassium permanganate, oxalic acid and calcium chloride (each, 0.2 N) equivalent t o each other were prepared. The calcium chloride solution was prepared by dissolving Iceland spar in a small amount of hydrochloric acid and diluting to volume. The concentration of this solution was such that the calcium in 1 cc combined with the oxalate ion in 1 cc of the oxalic acid solution, and 1 cc of the potassium permanganate solution was required to oxidize the calcium oxalate thus formed, when dissolved in dilute sulphuric acid. Portions of 5, 10, 15, 25, 35, 45 and 50 cc, respectively, of the oxalic acid were placed in beakers and diluted t o approximately 100 cc. Each portion was then titrated at the boiling point, with permanganates until the non-acid endpoint was reached. The results are shown in Table I (column
c> *
Oxidation of Oxalic Acid
439
TABLE I The Titration of Oxalic Acid with Potassium Permanganate in the Absence of Sulphuric Acid . . Cc 0.2 N oxalic acid taken
Cc 0 . 2 N c c 0.2 N KMn04 reKMn04 required for acid quired for non' 2nd-point (thee- acid end-point retical)
0
vo
5 10 16 25 35 45 50
10 15 25 35 45
v 1
0
i
3.31 6.58 8.85 16.43 22.76 29.34
0.662 0.65s 0.657 0.657 0.650 0.652
-K or, V;
3.13 6.25 0.38 15.62 21.87 28.13 31.25
0.18 0.33 0.47 0.81 0.s0 1.21 1.42
5.06 10.06 1.5.06 25.12 34. SO 44.86 40.95
0.06 0.06 0.06 0.12 -0.20 -0.14 -0.05
The non-acid titrations bear an approximately constant ratio to the acid titrations (column D) . Referring to the equation for the oxidation of oxalic acid in the presence of sulphuric acid, it will be noted that for every 5 molecules of oxalic acid oxidized, three molecules of sulphuric acid are required. That is, five of the eight equivalents, of acid accounted for in the equation, or 0.625 of the total amount, is oxidized. Multiplying the values in column B by this factor, we find that the results differ only slightly, and in a regular manner, from the results obtained by titration (column F). Dividing the sum of the terms in column F by the sum of the terms in column A, we obtain the quantity 0.029 which we shall
J . C. Witt
440
call k . Now we may derive a formula for the value of the non-acid titration as follows: Vl=O.625 Vo kO I n which, Vo= cc of potassium permanganate necessary completely to oxidize a given volume of oxalic acid solution in the presence of excess sulphuric acid (acid end-point) V1= cc of potassium permanganate necessary t o produce the non-acid end-point. 0 = cc of oxalic acid titrated. Now the oxalic acid and the permanganate were taken as equivalent. Therefore we may substitute Vo for 0. Substituting the value of k and collecting we have, Vi=O.654 Vo. Considering the coefficient of VOas a constant, we may write V I = K VO. Substituting the successive values of V1 and dividing by V,, we obtain the values given in column G, designated as Vi, The values of Vi-Vo are shown in column H. The average numerical value of the terms in this column is 0.099 cc. That is, we can calculate the amount of oxalic acid present from the value of the non-acid titration within practically the experimental error. Also on this basis we may consider that k and K are constants. However k must be determined for each series of titrations as will be shown later. The results indicate that the value 0.625 does not vary. Consequently the reaction between oxalic acid and permanganate may be represented by the equation : 2 KMn04 8H2C2O4+ K2Cz04 2MnC204 8H20+loco2 The oxalic acid is reduced by the permanganate and as long as there is a sufficient concentration of hydrogen ion, present, the manganese compounds resulting from the reduction of the permanganate are dissolved by the excess of oxalic acid. The hydrogen ion concentration decreases during the titration until a point is reached at which the manganese com-
+
+
+
+
441
Oxidation of Oxalic Acid
pounds are no longer dissolved and further addition of potassium permanganate results in the formation of a precipitate. At the end of the titration, the solution is no longer acid-in fact, in most cases it shows a faintly alkaline reaction. In Figure 1 curve A shows the volume of potassium permanganate that is required completely 'to oxidize oxalic acid ; curve C, the theoretical amount necessary to produce the non-acid end-point ; and curve B, the non-acid titration value obtained. The area between B and C is determined by the constant k . Eflect of Certain Variables on the Non-acid End-point.--In making a series of titrations it. l is necessary to keep constant 1 cc.OXALlC k i n I as many factors as possible 20 30 40 if concordant results are to be Fig. 1 obtained. Results are affected by the temperature, the total volume of the liquid titrated, the presence of other electrolytes and probably by other factors. TABLEI1 Effect of Total Volume Cc 0 . 2 N oxalic acid
Water added
50 50 50
0 50
150
Cc 0 . 2 N potassium permanganate t o obtain nonacid end-point -
31.83 32.67
32.80
I n obtaining the results shown inTabIe I, the temperature was kept nearly constant since the titration was made a t the boiling point-though the concentration of the mixture changed during the titration owing to evaporation and to the addition of permanganate, which of course affected the boiling point.
J.
442
Cc 0 . 2 N oxalic acid
!
c. Witt
I.
25 25 25 25 25
Cc potassium permanganate to give non-acid end-point
Other acid added (after neutralizing with sodium hydroxide) I
None 5 cc HzS0.j 5 cc HCl 1 cc HC1 5 cc CH3COOH
16.34 15.15 15.80 16.26
Sodium acetate practically prevents the reaction. Large amounts of sodium sulphate and of sodium chloride decrease the titration values. That is, before the amount of potassium permanganate necessary to produce the non-acid end-point under normal conditions is added, a colloid is precipitated by the high concentration of electrolytes present. When sodium chloride equivalent to only 1 cc of concentrated hydrochloric acid is present, the effect is negligible. The Addition of S d p h w i c Acid to Complete the Reaction.-
Oxidation of Oxalic Acid
Cc 0 2 oxalic acid
Sulphuric acid present
25 25 25
Large excess None 15 cc of approx. 0.1 N 25 cc of approx. 0.1 N 50 cc of approx. 0.1 N
443
cc potassium -
25 25
permanganate
25.00 16.25
19.77 23.87 25.00
J . C. Witt
444
case, water was added to make the total volume titrated 100 cc. Between 25 and 50 cubic centimeters of the sulphuric acid was needed, or approximately 4 to 7 times as much as the theoretical. With 50 cc, the liquid, when the end-point was reached, was not a clear pink b u t slightly brownish and turbid, resembling the non-acid end-point. The equation for the oxidation of oxalic acid in the presence of sulphuric acid is more properly written :
+
+
2KMn04 5HtC204 xH2S04+ K2S04 2MnS04 8H20 loco2 +(x-3) H2S04 Theoretical Considerations.-When oxalic acid or one of its salts, is titrated with potassium permanganate, there is a question of equilibrium between what may be termed oxidation and acid agencies. The effect of the latter may be divided into solution and precipitation. On the basis of this, the various phases of the reaction that have been studied will be discussed, including (a) the presence of excess sulphuric acid, (b) a limited amount of sulphuric acid, (c) oxalic acid alone (d) the effect of electrolytes, volume, temperature, etc. Excess Sulphuric Acid.-When an excess of this strong acid is present, the manganese compounds formed are dissolved at once, producing a clear nearly colorless solution, and the first drop of permanganate added after the reaction is complete is easily seen. Because of the high hydrogen ion concentration, the reaction is almost independent of temperature, concentration, and the presence of other electrolytes (so long as the latter do not have any oxidizing or reducing power). Acetic Acid.-Even with a large excess of acetic acid, the hydrogen ion concentration is not high, and as the titration proceeds, the concentration is diminishing both by neutralizations and on account of the formation of acetates and the consequent common ion effect. The result is that while the titration can be made and the same values obtained as with sulphuric acid, the temperature must be kept at or near the boiling point, and a longer time is required.
+
+
+
Oxidation of Oxalic Acid
445
Limited Amount of Sulphuric Acid.-The titration starts in the normal manner but the solution of the manganese compound becomes slower and slower as the hydrogen ion concentration diminishes, and a colloid is formed. When this first appears it may be dissolved by continued stirring. Long before the hydrogen ion concentration reaches zero, the precipitating power of the ions present in the mixture (principally the sulphate ions) is sufficient to throw down manganese compounds. Further addition of permanganate produces more of the precipitate according to the so-called Volhardt reaction. Various investigators disagree as to the composition of this precipitate. It is probably a hydrous oxide of manganese, and may vary in composition, with the condition under which the precipitation is made. Indeed Orloffl says there are thirty-three possible intermediate oxides or combinations of oxides in the reduction of Mnz07 to MnO. E#& of Electrolytes, Volume, and Temperatures.-When conditions are favorable, the titration of oxalic acid alone with potassium permanganate produces no precipitation until all the oxalic acid has either been oxidized or neutralized. However, changes in volume or temperature may cause a premature precipitation. When neutral electrolytes are added, the precipitating power is increased without increasing the hydrogen ion concentration-consequently the equilibrium is upset and a precipitate forms. We are now in a position to discuss the constant K , or why slightly more than the theoretical amount of permanganate is necessary to produce the non-acid end-point. The theoretical amount is required to remove the last of the oxalic acid from the solution, and then a slight excess is required to produce the colloid, which is later precipitated either by heating or by adding still more permanganate. Further, it has been shown that oxidation-reduction reactions in the absence of acids are not instantaneous, and that an excess of one of the reacting substances may be required to drive the reaction to completion. Zeit. phys. Chem., 80, 230 (1912).
J . C. Witt
446
Summary Oxalic acid may be titrated with potassium permanganate without the addition of any other acid. The oxalic acid serves as both reducing agent and as acid, and there are two distinct reactions although they proceed simultaneously. Of 8 equivalents of oxalic acid present, (a)5 are oxidized by permanganate, and ( b ) the other 3 equivalents combine with the products of (a). When this point has been reached the further addition of permanganate solution produces a colloidal manganese compound which is precipitated on boiling for a few seconds. The appearance of turbidity followed by a slight permanent precipitate marks the end-point. The end-point is affected by (a)the temperature at which the titration is made, (b) the concentration and (c) the presence of electrolytes in solution. When the solution of an oxalate is titrated in the presence of sulphuric acid, the colloid precipitating power of the sulphate ion in a sense acts in opposition to the tendency of the hydrogen ion to keep the products of the reaction in solution, The former is sufficiently powerful to produce a precipitate in the presence of a fairly high concentration of the hydrogen ion. For this reason, the quantity of sulphuric acid actually required is much greater than the theoretical. Chicago, IlJ.
