T H E ELECTROMETRIC TITRATIOK OF COPPER SALTS BY SODIC11 HYDROXIDE* BY RALPH A L O S Z O B E E B E
In a previous investigation,' the electrometric titration of copper sulphate, chloride and nitrate, using the quinhydrone electrode, was used as a means of estimating the composition of the precipitates formed under various conditions of precipitation. The application of the method to the trichloracetate, acetate, and chlorate is described in this paper. Preparatzon of Solutzons: A solution of sodium hydroxide (0.40 18 normal), free from carbon dioxide, was prepared by the method of Cornog.* This solution was used in all the titrations. Copper trichloracetate (0.4808 normal) was made by the action of trichloracetic acid on a slight excess of basic copper ~ a r b o n a t e . ~ After standing several hours the excess carbonate was removed by filtration. An excess of trichloracetic acid was then added to prevent hydrolysis of the copper salt solution. Allowance for this excess was made by taking into account the position of the first break in the titration curves as is shown in Fig. 1.I Copper chlorate (0.3985 normal) was prepared by adding barium chlorate in slight excess to a solution of copper sulphate, and removing the precipitated barium sulphate by fi1tration.j Copper acetate, 0.4043 normal, was made by dissolving the c.p. salt in water and adding a measured excess of acetic acid to prevent hydrolysis. All the copper solutions were standardized by adding potassium iodide and titrating the liberated iodine against standard thiosulphate solution using starch indicator. A test showed that potassium chlorate liberated no iodine from a potassium iodide solution. The iodimetric method was therefore assumed to be applicable to copper chlorate.
Copper Trichloracetate Two methods of titration were employed. Method A . (Dzrect Tztratzon.) The alkali was added by small increments, allowing apout five minutes after each addition for the solution and precipi-
* Contribution from the
Moore Laboratory of Chemistry, Amherst College. Hopkins and Beebe: J. Phys. Chem., 34, 5 7 0 (1930). J. Cornog: J. Am. Chem. Soc., 43, 2573 (1921). Bateman and Conrad: J. Am. Chem. Soc., 37, 2 5 5 5 (1915). Miiller and Miiller: Z. anal. Chem., 73, 47 (1928). No excess acid was added to the copper chlorate. However, two solutions, one freshly prepared, and the other six months old showed no measurable difference in the titration curve8, although 8ome basic salt had precipitated from the latter.
3678
RALPH ALONZO BEEBE
tate to approach equilibrium, before reading the p0tentiometer.l Approximately fifty minutes was required to complete a titration. In curves I , 2 , 3, and 1 in Fig. I , obtained by this method for solutions 0.4 normal respectively, the moles of of initial concentration 0.01, 0.03, 0.1, sodium hydroxide per mole of copper salt are plotted against the pH of the solutions. Approximately 1.6 moles of alkali are required for complete precipitation in each case. The alkali requirement was very slightly greater
FIG.I Copper Trichloracetate. (Curves 1-4, Method A).-I, 0.01 N.; 2, 0.03 S . ; 3, 0 . 1 N.; 4, 0.4 N.; 5 , Typical end-point by Method B, (Delayed Titration) for 0.4 N. solution.
for the more dilute solutions but the difference was too small to be shown on the graph.? Nethod B. (Delayed Titration.) Since it was suspected that curves 1-4 resulted from the precipitation of a mixture of copper hydroxide aEd the r : 3 basic salt, not in equilibrium, B method of titration was devised t o insure that sufficient time should he given to reach equilibrium conditions as nearly as possible. The alkali was added t o several samples of copper salt solution in quantities just under 1 . 5 moles per mole of copper salt. These samples, of different initial concentrations of copper salt, were allowed t o stand for varying lengths of time, some with constant stirring, some with occasional Even with this time allowance, it was apparent that equilibrium was not completely established because there was a slow drift toward a lower p H indicating the slow disappearance of hydroxyl ions from the solution. * As was shown in the previous paper (loc. cit.), a requirement of 2.0 moles of alkali would indicate the precipitation of the copper in the form of the hydroxide; but 1 . 5 moles would correspond to the precipitation of the 1:3 basic salt, [Cu(CC13COO)2in this case].
3679
ELECTROMETRIC TlTRATION O F COPPER SALTS
shaking, before the titrations were completed electrometrically. The precipitate was bluish white in color after standing and settled readily after shaking. A pH value of 8 was arbitrarily chosen as the end point of the titrations; but, as may be seen in Curve 5 , Fig. I , a deviation of several tenths of a pH above or below this value make no measurable difference in the volume of alkali as read from the curve since the latter is practically vertical throughout the region of pH 6.5 to 8.5. I n Table I are listed the results for the samples which were titrated. The alkali required was decreased when time was allowed to reach equilibrium. All the end-points are closely concordant and in excellent agreement with the theoretical requirement of I .so moles corresponding to the precipitation of the 1:3 basic salt. That the time of standing and amount of stirring had no effect, indicates that all samples had reached equilibrium. TABLE
I
Alkali required to precipitate Copper Ions from Solution (1.50 moles required for I :3 basic salt) Number Salt of used Sample
Kormal-
it?
Time of Standing
Treatment
hours 20 hours 20 hours 20 hours 20 hours I Z days 20 hours 8 days 2 0 hours 8 days 20 hours 8 days
constant stirring occasional shaking constant stirring occasional shaking
Moles Total moles NaOH
XaOH added
initially I 2
Trichlor- 0 . 4 8 acetate 0 . 4 8
3
0.40 0.40
4
)’
5 6
’.
0.40
”
0.40
i
”
0.10
8
”
0 .IO
9
’’
0.075
IO
”
0.07j
II
”
0.05
I2
”
0.05
20
,3
,
,
f’
1,
If
,I
)’
1
>>
,,
,, ,, ,,
,, I,
I
,461 I ,461 I ,311 1.381
I . 502
1.345 1.480 1,446
1,495 I
1.453 1.453 1.455 1.460
1.497 1.492 1.503 1.500
I
I.
chlorate o 10
20
hours
21
’’
o 40
4 days
22
”
9 days 20
”
o 40 o IO o IO
2j
’*
0.10
20
26
”
20
2 j
”
o oj o.oj
23 24
”
20
20
hours hours hours hours hours
1,
,, ,, ,I
constant stirring occasional shaking
,,
1,
I,
,,
,
1.333 1.434 1.384 1.333 1.475 1.190 1.166 1.536
505
,497 1.197
I
I . 500
I 20
,493
I . 461
,494
,498 (Av.)
1.586 1.593 1.590 1.633 1.631 1.636 1.682 1.698
3680
RALPH ALONZO BEEBE
Copper Chlorate Titrations were carried out by methods A and B as described in the case of the trichloracetate. With the chlorate the E.M.F. readings showed almost no drift, and as a result, the horizontal portions of the curves 1-3 obtained by method A shown In Fig. 2 were more easily reproducible than were the corresponding portions of the trichloracetate curves. Curves 1-3 have breaks in the region between 1.5 and 2.0 moles of alkali. I t is worthy of note that the initial concentration of the copper chlorate has a
FIG.2 Copper Chlorate (Direct Titration). I ,
0.01
S . ;2 .
0.10 S . ;3,
0.4 N.
marked effect on the position of the break, a larger amount of alkali being required for the more dilute sol~i~ions.% . 1:3 basic copper chlorate has been reported by Sabatier.' I t is probable that the precipitate obtained under the conditions described above is a mixture of this 1:3 basic salt and copper hydroxide. An attempt to get more definite evidence for the presence of the I :3 salt by using Method 13 was not successful. Less alkali was required for this delayed titration than for the direct titration by method *A. For example, the 0.4 normal solution required approximately 1 . 7 moles by method A (see Fig. 2 ) , and 1.59 moles by method B (see Table I). However, the amounts required for different dilutions were not in agreement and even in the case of the most concentrated solution, considerable more than 1.5 moles was used. Hence no conclusion can be drawn from these daia about a definite basic salt. Sabatier: Compt. rend., 125, 104 f1897).
ELECTROMETRIC TITRATION O F COPPER SALTS
3681
Copper Acetate The occurrence of an inflection a t about 1.75 moles in Curve A, Fig. 3, obtained by method A for 0.4 normal copper acetate indicates that the precipitate is even more basic in nature than was the one from a copper chlorate solution of the same concentration. The gradual slope of the inflection was evidently due to the partia.1 decomposition of the basic acetate a t first precipitated. Britton has cited a similar case in the precipitation of magnesium by alkali.' Titrations by method B showed that there wa,s a demonstrable though rather small diminution in the amount' of alkali required if compared to method A. Slow Titration: Since methods A and B gave unsatisfactory information about the composition of the precipitate, a third method of titration was tried. To equal samples of the copper acetate solution (0.4K) were added quantities of alkali varying from about 0.1mole up to 2.0 moles per mole of copper salt. From these samples, kept at 30'C for many days, portions were removed from time to time to determine their pII. Care was taken to shake well before removal of a given portion to insure no change in the relative amounts of solution and precipitate. This method gave the same results as would an extremely slow titration in which several days or weeks were allowed for establishing equilibrium after each addition of the alkali. It was suspected that, by this rather tedious type of titration, some inflection might be found corresponding to the precipitation of a basic copper acetate. The rather unexpected results of the slow t,itration, though difficult t o interpret, seemed sufficiently striking to be worthy of reporting. They are shown in Fig. 3. The experimental work necessary for the data of Fig. 3 was repeated three separate times, so there can be no question about the existence of the phenomena involved. Compared to curve A obtained by direct titration, it is seen that there has been a decrease in pH all along the curve after st,anding one day (Curve I ) . Up to about ~ . 2 moles 5 the amount of this decrease in the first day is relatively small and is negligible after that time. Beyond 1.5 moles the decrease is great. Moreover, a maximum occurs in the curve just above I . j moles producing the anomalous effect that a solution to which more base has been added is actually more acidic. This maximum gradually flattens out with time as shown in Curve 4 taken after 49 days. Readings taken after I 18 days (not shown on the chart) differed very little from Curve 5 . I t was not practicable to carry the experiment beyond that point. The final sharp inflection of the curves corresponds quantitatively to the precipitation of copper hydroxide.2 Similar results were obtained from copper chlorate and copper sulphate by this method of extremely slow titration. In both cases up to the 1.5 moles point the pH values changed little upon standing, but changed greatly between 1.5 and 2 . 0 moles of alkali. In this region, the solutions were at first H. T. S. Britton: "Hydrogen Ions," p. zjj. * A new method by Shiptalskii, Katzen and Klyachko, [ J . Russ. Phys.-Chem. SOC., 61, 1497 i I g Z g ) ] . , for the estimation of acetic acid in neutral and basic copper acetates has been based on the precipitation of copper oxide from hot solutions of copper acetate by sodium hydroxide. From hot solutions the dehydration would occur much more rapidly. 1
3682
RALPH ALONZO BEEBE
strongly basic, but changed to a pH of 5-7 when allowed time. In the case of the chlorate a sharp maximum occurred just above the 1.j moles mark. With the sulphate, the curve suddenly changed direction at the I.j moles point becoming almost horizontal. Although the significance of these curves is not obvious, the presence of the maxima is apparently connected in some way with the dehydration of the precipitated copper hydroxide or basic salt which could be detected by the blackening of the precipitate. There was no blackening of the precipitates in
3 (1-4-Slow
FIG.
Copper Acetate. .S-Direct 3. 2 1 days; 4, 49 days.
titration.
titration).
I, I
day;
2,
4 days;
Curve A On standing, holyever, those precipitates from the solution to which the greater amounts of alkali had been added gradually darkened, and as time went on the precipitates from samples to which less alkali had been added began to turn dark. The maximum in curves 1-4 gradually moved to the left with time; it was observed that the blackening of the precipitates roughly kept pace with the maximum, precipitates from samples represented by points to the right of the maxima being darkened and those by points to the left, remaining blue in color. The appearance of the maxima near the I . j moles point suggests the presence of the 1:3 basic copper acetate. So far as the author is aware, no such salt has been reported although basic acetates of composition z : I , I : I , and I :2 have been described.’ General Discussion From the results of this paper and the previous paper by Hopkins and Beebe,*it is apparent that under favorable conditions a definite I :3 basic salt 1 2
Iof, Kobrin and Klyachko: Zhur. Prikladnoi Khim., 3, 336 (1930). loc. cit.
ELECTROMETRIC TITRATION O F COPPER SALTS
3683
is precipitated from solutions of copper sulphate, chloride, nitrate, and trichloracetate by addition of alkali, and that less stable basic salts, probably of I :3 composition, are likewise precipitated from solutions of copper chlorate and acetate, although the evidence for the composition of the latter is not to be considered conclusive. Certain generalizations can be drawn from the data at hand. All the basic salts are unstable in presence of excess alkali, probably reacting to form copper hydroxide. The order of this instability varies with the different salts, basic copper acetat’e being the least stable. The latter is so unstable that even in the direct titration it reacts rapidly with hydroxyl ions at relatively low concentrations with the result that the inflection in the titration curve is gradual and almost two moles of alkali, corresponding to the precipitation of copper hydroxide, are required to make the solution alkaline. (See Curve A, Fig. 3). Basic copper chloride, on the other hand, is so stable that it is necessary to reach a relatively high concentration of hydroxyl ions before appreciable rapid decomposition of the basic salt occurs, with the result that a sharp inflection is found a t the I . j moles point. The tendency toward decomposition of the basic salt is apparently increased by dilution. This is most noticeable in the case of the nitrate and chlorate and is easily explained by the increased tendency toward hydrolysis of the basic salt in more dilute solution. A comparison of the direct titration curves for copper nitrate and copper chlorate, shows that) the basic nitrate is the more stable of t,he two since for solutions of the same initial concentration, say 0.1 normal, more alkali is required in the case of the chlorate. On standing in the presence of excess copper salt solution, copper hydroxide tends to react to form the I :3 basic salt,, as shown experimentally by the results of method B in this paper and of the so-called “effect of prolonged stirring,” in the previous paper. Here it is found that cupric chloride and sulphate react most readily with copper hydroxide to form the basic salt, and cupric acetate least readily. By comparison of the relative tendencies of the copper salts to react with copper hydroxide and of the basic salts to be hydrolyzed or to react with excess hydroxyl ions, the basic salts of copper which have been studied may be listed in the following order of stability: chloride > sulphate > trichloracetate > nitrate > chlorate > acetate.l
summary Electrometric titration curves are shown for copper trichloracetate, acetate, and chlorate with sodium hydroxide. 2. Evidence is presented for the formation of a definite I :3 basic copper trichloracetate, Cu(CC13COO)2.3C u ( O H ) ~ . x H ~not O previously reported. 3. The order of stability of the basic salts of copper thus far studied is chloride > sulphate > trichloracetate > nitrate > chlorate > acetate. The author’s thanks are due to Howard Jones and Stewart Seass who checked a part of the experimental work reported in this paper. I.
The position of the sulphate relative to the chloride is rather uncertain owing to the complicating tendency to form a I : 2 basic copper sulphate.
