180
LAWRENCE li. SCHARFSTEIN A N D CECILV. KING
Vol. 58
THE ADSORPTION OF COPPER SULFATE ON COPPER AND SILVER1 BY LAWREXCE R. SCHARFSTEIN AND CECILV. KING Department of Chemistry, New York Univeisily, New York, N . Y . Received J u l y 6 , 195s
The amount of copper(I1) sulfate removed from solutions in water and methanol, on percolation through columns of finely divided copper and silver, has been measured. The nature of the adsorption isotherms obtained is discussed. At low concentrations in aqueous Aolution, the adsorption on c,opperfalls off rapidly, as if a null point would be found a t about M . The adsorption in all cases is equivalent to less than a monatomic layer, but is greater than necessary to charge the electrical double layer, if this is considered in the simple Helmholtz sense.
Introduction The precise measurement of adsorption of metal salts on metals is difficult, because so little is adsorbed. It is necessary t o bring a very large surface area into equilibrium with a small amount of solution if the latter is to decrease appreciably in concentration. Very dilute solutions are sometimes used if a suitable analytical method is available, since the per cent. change in concentrationisgreater. However, it is advantageous to cover a wide concentration range in order to find the type of adsorption isotherm obeyed, and whether an upper limit is approached. The use of radioactive tracers for such measurements looks inviting, but presents additional problems. Metals exchange with their own ions to an extent equivalent to many atomic layers, and it is not certain how much of this is due to statistical fluctuations at the surface, how much to local cell action, and how much to internal diffusion. No one has attempted to distinguish between adsorption and exchange in such cases. Adsorption studies have been limited to the case of less noble ions on more noble metals, or on passive metals. I n order to obtain more information about adsorption on metal surfaces by a simple, direct experimental method, it was decided t o make measurements of the retention of copper sulfate from solution, on the surface of copper and silver. T o obtain the advantage of large surface area and small solution volume, the method of percolation through a column of finely divided metal was chosen. If equilibrium were attained rapidly as the liquid seeps through the (initially dry) column, the first effluent should be free of the metal salt, and increase to the initial concentration should be very sharp. This is not quite true in practice, but the total amount of retention can be found by analyzing a few ml. of effluent rather than the much larger volume which would be required to bathe the entire metal sample. The main disadvantage of this method is that it is time-consuming; the metal must he washcd thoroughly and dried in the column for each esperimeat. With corrodable metals air must be excluded. Experimental The Metals.-Electrolytic copper, in the form of a granular powder, was employed.2 The manufacturer stated t h a t it was prepared by ejection of molten metal through a fine nozzle, with rapid cooling in an inert fluid. The pmticlt% were approximately spherical, and varied in diameter from (1) From the Ph.D. thesis suhlnitted by Lawrence R . Sclwrfstein in the Graduate Sc1.001 of New York Univcmity, Rlay, 1053. (2) From Bolinont Sineltine and Refining Works, Inc.
0.003 to 0.1 mm., with over 5070 between 0.01 and 0.08 mm. Measurement of 20 particles from each of 7 samplings gave an estimated area of 180 cm.2/g. Spectroscopic analysisashowed 0.005% Fe, 0.003y0 Ni, 0.002% each of Ag and Pb, 0.001% each of Zn and Sn, with smaller amounts of Mg and Si. The copper was washed with alcohol and ether, and in the columns it was washed with acidified copper sulfate solution to remove exchangeable base metals, until a negative test for iron was obtained. The silver was a sample previously prepared and used in this Laboratory.416 It consisted of uniform crystals about 9 X 9 X 14 p , and had an estimated area of 500 cm.e/g. The Columns.-Most of the work reported with copper was done with a column about 3 cm. in diameter, filled with 536 g. of copper, to a height of 41 cm. The bottom of the glass tube was melted down to a short, thick-walled capillary tube, and the copper was retained by a 14-mg. plug of glass wool. Side arms at the top of the tube permitted passage of nitrogen and hydrogen. A similar smaller column was used for preliminary experiments. The silver column was 1 cm. in diameter and contained 96 g. of silver. Operation of the Columns.-The solutions dripped onto the metal from a Teflon plug buret, which could be adjusted very easily to the desired rate of delivery. The flow rate was ordinarily one or a few drops per minute; the solution spread over and down through the metal without flooding at any point. About 33 ml. of solution was required to wet the copper in the large column; the first effluent appeared after 2 to 6 hours. Ordinarily 3 to 25 ml. was collected, in portions, for analysis; the first effluent was copper-free, and the rise to full strength took place in 3 to 5 ml. of solution. Before uSe in each experiment the copper, after thorough washing with water, was dried in a stream of nitrogen passed in a t the capillary tip, while the column was warmed in a tub: furnace. Hydrogen was then passed in for 2 hours at 350 , and the metal was cooled in nitrogen. With nitrogen stsillflowing, the buret was attached. All solutions were prcviously deaerated, and nitrogen was passed through the top of the buret, through the top of the column, and around the tip of the column until effluent appeared. The silver was heated in hydrogen to 150", which would reduce sulfide if present, and it was not considered necessary to protect with nitrogen during runs. Analysis of Solutions.-Most of the analyses were done with a Lumetron photoelectric colorimeter, some with a colorimeter which used a smaller solution cell. The effluents were collected in weighing bottles and diluted with water and reagents as required, by weight. Volume concentrations were obtained by estimating densities as closely as possible. The copper ammonia complex was used for concentrations above 10-3 ill. For more dilute solutions, a copper test proposed by Clarke and Jones, and further described by Kolthoff was developed into a quantitative proccdure, as follows. To the solution to be analyzed were addcd one drop of 0.5 M sulfuric acid, 0.2 ml. of 0.1% dimethylglyoxime in ethanol and one ml. of saturated potassium periodate solution; the whole was diluted (by weight) to approximately 10 ml. The mixture could now be kept until a convenient time for analysis, since no color appears in acid solution. On adding 5 mg. of solid sodium bicarbonate a red-violet color develops (absorption maximum a t 525 mp), reaches full intensity in about, 3 minutes, and starts to fade several (3) Conrtesy of t h e Chemistry Dermrtment, Columbia University. ( 4 ) C. V. King and F. S. Lang, J . Electrochem. Soc.. 99,295 (1952). (5) C. V. Kin. and R. K. Schochot, THISJOURNAL, 67, 895 (1953). (6) I. M. Kolthoff, J . Am. C h e n . Soc., 61, 2422 (1930).
Feb., 1954
181
ADSORPTION O F C O P P E R SULFATE ON COPPER AR'D SILVER
minutes later. With a suitable filter in the colorimeter, Beer's law was shown to be obeyed in t,he range 4 X 10-6 t o 1.2 X IO-* f i f coppcr sulfalc. Addition of nictlianol up to 20yo by volume did not affect the absorption. I t was shown that solutions analyzed with t,hc aninionin complex could be diluted and analyzed with dinic lhylglyoxinie with the same results, within the usual prccision of about 3%. Since water from copper stills may contain as much as 5X M copper ion, water with a much lower copper content was obtained and used for the most dilute solutions. While part of the copper ion in the column effluent has been reduced to the cuprous state, this need not be considered in the analysis, since it is reoxidiaed during preparation of the solutions for the colorimeter. Preliminary Experiments.-The first runs were made, in the smaller column, with (analytical grade) copper sulfate dissolved in pure water; the pH varies in the range 4 to 6, depending on the concentration.' While the amounts of copper ion retained seemed reasonable, the pH of the effluent remains lower than the original after equilibrium has been established in the column. For example, 4.68 X 10-8 M copper sulfate went into the column a t pH 5.1; about 2 ml. of effluent was deficient in copper ion, and after another 16 ml. was collected the pH was still 3.6 to 3.7. All the older work indicates that cuprous oxide will be formed on the copper under these conditions; an excellent treatment of the system is given by Pourbaix.* The question of stability of cuprous oxide, however, depends mainly on a knowledge of the equilibrium constant of the reaction C u + + Cu 2Cu+; I< = acu+2/acu++ The accepted value of I P O has been measured with acidified solutions. More recently El Wakkadll has found the value 10-6 for unacidified solut,ions. El Waltlcad measured the pH of equilibrated solutions, and found that the amounts of cuprous and hydroxyl ions could not exceed the solubility product constant. Because of the possibility of cuprous oxide formation, it was decided to employ acidified solutions in further runs. I t was found that if the pH were adjusted to any value from 3.5 to 2.5 by addition of sulfuric acid, there was no change on passing through the column, and the amount of copper retained was independent of pH. When sulfuric acid, pH 3.5, was passed through, the effluent contained no more copper than the water used. During the experiments, the results were checked repeatedly by returning to two reference concentrations, and 6.7 X 10-3 M. I t was found that namely, 2.1 X if the column were left open to the air, or in contact with aircontaining solutions, new measurements at, these concentrations showed an increased adsorption of a few per cent. I n the course of two years the adsorption capacity of each of the two columns was nearly doubled. This is attributed to the repeated formation and removal of an oxide film; it has been shown by Rhodinlz that oxidation of a very smooth copper surface, followed by reduction in hydregen, leads to an increase in area, as measured by nitrogen adsorption.
+
Results The measurements on copper reported below were obtained with the column in almost daily use and filled with deaerated copper sulfate solution when idle. In this way good checks were found on returning to the reference solutions. Table I gives results a t such concentrations that reduction t o copper(1) ion is negligible. The methanol solutions were made by dissoving CuSOc.5H20in anhydrous methanol, and were not acidified. At lower concentrations reduction to cuprous ion becomes appreciable, and after equilibrium is established in the column, the total copper concentration of the effluent is greater than that of the (7) C. '8. Davies, J . Chem. Sac., 140, 2093 (1938). (8) h4. J. N. Porirbaix, "Thermodynamics of Dilute Aqueous Solution-," Edward Arnold and Co., London, 1949, Chap. VI. (9) F. Fenwiok, J . A m . Chem. Sac., 48, 860 (1926). (IO) E. Heinerth, Z . Elektrochem., 37, 61 (1931). (11) S. E. S. El Wakkad, J . Chem. Sac., 3563 (1950). (12) T. N. Rhodin, J . A m . Chem. Sac., 72. 4313 (1950).
TABLE I AUSOI~PTICN O F COPPER SULFATE ON COPPER T w i p . 20-27"; 536 g. copper; p H 3.0 in first 5 solutions, 3.5 in otlitw ccuso4
M
\Vnt.er
x 162
No. of expt.
4.01 2.10 1.73 1.30 1.12 0.672 .421
3
7
2 3 2 2
(1
2.30 1.42 1 .oo 1 .oo
108
8.2 6.2 6.1 5.4 5.5 5.1 4.4 3.4 8.5 8.0 5.4 6.3"
2 3 2 4
.290 Methanol
Moles ads./g.
cu x
1 2
Av. dev.,
% 4.5 3.3 3.7 3.4 4.4 0.8 2.5 5.7 4.3 3.8
...
4.3
At 0" with ice jacket around column.
initial solution. The analytical method was not good enough to fix the final concentration reliably, and this had to be calculated. Measurements in more dilute solutions are given in Table 11; under "moles ads.," min. and Max. refer to the amounts of adsorbed copper ion based on calculations of the final concentration employing the two values of K mentioned above. As will be seen later the difference in values is unimportant; in fact, if no correction were made, the conclusions would be no different. TABLE I1 ADSORPTIONFROM AQUEOUSSOLUTION Temp. 20-27 ", 536 g. copper, pH 3.5 CCUBO~, AI
No. of expt.
7 . 5 2 x 10-4 1.30 10-4 7.71 x 10-6 3.41 x 10-b 1.45 10-6 7 . 3 1 X 10-6 5.25 X 10-6
3 2 2 3 2 2 3
x
x
Moles ad&/& c u x 10'0 Min. Max.
180 48 33
17 5.2 2.5 I .3
190 52 36 19 6.2 3.0 1.6
Av. dev.,
% 4.3 10 7.3 2.1 12 5.3 7.2
Adsorption on Silver.-From aqueous solutions the adsorption on silver was similar, in moles per gram, to the adsorption on copper. Much more copper salt was retained from methanol, and most of the experiments were done with solutions of CuS04.5H20 in methanol. Addition of the' same amount of sulfuric acid that was used in the aqueous solutions had little or no effect, so the solutions reported were not acidified. It was found that in successive runs the retention of copper decreased; for example, from 2.27 X loF8M solution the adsorption per g. silver in four runs was 44,37, 31, 30 X lo-* mole. A small amount of copper was so firmly retained that it could not be washed out with water or methanol; it was displaced by passing silver nitrate solution through the column, and about g. atom was recovered in this way in each of three runs. Thereafter, the column was washed with water, with silver nitrate and again with water before drying and heating in hydrogen; agreement was then found in successive runs.
LAWRENCE R. SCHARFSTEIN AND CECILV. KING
182
Vol. 58
Table 111 summarizes these experiments. More dilute solutions were not used because the amount of copper-free effluent would be so large; a t the lowest concentration over 50 ml. had to be collected to reach the initial concentration.
0 gives about, 7.8 X IOw8 g. atoms copper ion per g. copper as a limiting value. The true surface area of the copper cannot be estimated accurately, but an approximation can be made. The roughness factor probably lies between 3 and 5 . Some of the surface must have been inTABLE I11 accessible to solution; on dismantling the columns, the pellets were found to be partly sintered or ADSORPTION OF COPPER SULFATE AND METHANOL cemented together. A reasonable estimate of the SOLUTIONS effectfive area should be from 500 to 750 cm.2/g. Temp. 20-27 96 g. silver, two runs each concentration. This would give 1.0 to 1.6 X g. atoms/cm.2 Moles ads./g. Av. dev. CCUSO~, M Ag X 108 % as the maximum adsorption. 1 . 9 9 x 10-2 69 6.1 The amount of adsorption is then always far less 1.07 X 59 8.6 than the equivalent of a monolayer, which would be 4.80 x 10-3 53 9.4 near 25 X lq-"Jg. atoms/cm.2 (based on the reticu2 . 2 9 X 10-8 43 2.8 lar density of (100) planes in the copper crystal). 8 . 9 2 x 10-4 36 8.5 On the other hand, the amount of adsorbed cop6.33 X 31 13 per(I1) ion necessary t o charge the electrical double 1.17 X 63" 5.0 layer, if considered in the simple Helmholtz sense, 3.07 X 50b 8.0 would be only 1 X g. atoms/cm.2 per 0.1 1.69 X 4.0b ... volt. The potential difference a t the interface probably does not exceed 0.2 volt. a At 0". From aqueous solution, pH 3.5. The results of the adsorption from methanol on Discussion silver (Table 111) do not lie on a straight line when The data of Tables I and 111show fair agreement plotted as reciprocals, but a rough extrapolation to with Freundlich isotherms; Le., the points fall l / c = 0 leads to a saturation value of about 8 X reasonably close to straight lines on log-log plots. lo-' mole per g. silver. Taking the area as 500 The data of Table I1 deviate widely from such a cm.2/g., the maximum adsorption is 16 X 10-lO plot. It can be seen in Table I1 that the amount of mole/cm.2, a value not much smaller than was adsorption falls off rapidly a t low concentrations ; found for silver salts on silver from aqueous soluactually, the volume of copper-deficient effluent tions. Rudberg and v. EulerL4found little difference in went through a maximum a t 3.4 X M , and the decrease a t lower concentrations was not the amount of silver nitrate adsorbed on silver from entirely due to the increased total copper content of water and from 96% alcohol. There is no great difthe equilibrium solution. Possibly copper(1) ion ference in the amount of copper sulfate adsorbed on is adsorbed to a smaller extent than copper(II), but copper from water or from methanol, as seen in it is probable that no adsorption at all would occur Table I. The adsorption on silver is some 12 to 15 a t approximately M total copper ion. This times as great from methanol, and a t least part of situation has been found in the silver-silver ion the copper is very firmly bound, as evidenced by the system a t M,13and is in accord with the theory difficulty in removing it. It was thought possible that some exchange took place, but silver ion could of electrode potentials. If the data of Table I are plotted as reciprocal of not be detected in the effluent solution. The few experiments a t 0" do not indicate a much amount adsorbed us. reciprocal of concentration (Langmuir isotherm) , there is again reasonable greater temperature effect than the experimental agreement with a straight line, but no assurance error. Rudberg and v. Euler found little effect of that results at still higher concentrations would temperature with silver nitrate on gold, but King not deviate from this line. Extrapolation to l / c = and Schochets found that much larger amounts might be adsorbed on silver a t 0". (13) M. Proskurnin and A. Frumkin, 2. physik. Chem., 8 1 5 5 , O,
2D (1931).
(14) E. 0.Rudberg and H. v. Euler, 2. Physa'k, 19, 275 (1923).
b
