I. IWASAKI AND P. L.
594
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Vol. 62
THE ELECTROCHEMICAL DOUBLE LAYER ON SILVER SULFIDE A T p H 4.7.l I. IN THE ABSENCE OF SPECIFIC ADSORPTION BY I. IWASAKI AND P. L.
DE
BRUYN
Department of Metallurgy, Massachusetts Institute of Technology, Cambridge, Massachusetts Received December 9, 196Y
The electrochemical properties of the silver sulfide solution interface in sodium nitrate and acetate solutions at constant p H were investigated. The zero-point-of-charge of silver sulfide is observed to lie at pAg 10.2 and to be independent of the nature and concentration of the supporting electrolyte. When the surface is positively charged, the adsorption density of the silver ions is greater in acetate than nitrate solutions suggesting that nitrate ions behave more indifferently than acetate ions toward the surface. Differential capacities for silver sulfide in NaN03solutions are in satisfactory agreement with those for silver iodide in NaC104 and mercury in NaF solutions, A theoretical analysis of the silver sulfide system indicates that the observed differential capacity us. potential curves are in quantitative agreement with similar curves calculated by using the modified Gouy-Stern theory of the electrical double layer.
Introduction Several investigator^^-^ have illustrated that a study of the electrical properties of a mineralsolution interface will help in the understanding of the physical-chemical phenomena involved in the froth flotation process. Freyberger and de Bruyn4 were first to apply electrochemical techniques to the study of interactions a t a sulfide mineralsolution interface. As a prototype solid, these authors chose silver sulfide, and they were able to show that in the absence of specific adsorption, the structure of the double layer is in qualitative agreement with the Gouy-Stern model. The differential capacities at fixed ionic strength for Ag2S were found to be in good agreement with similar values obtained by Grahame6 on mercury and Mackor6 on AgI except when the surface has a high positive charge. This paper describes additional information obtained at the thermodynamically reversible silver sulfide-solution interface in the absence of specific adsorption. The experimental phase of this investigation benefited greatly from the careful development work done by Freyberger. Experimental The experimental technique and apparatus used in the present investigation is essentially the same as that of Freyberger and de Bruyn.' The adsorption experiments consisted of adding stepwise, known quantities of a silver salt solution (nitrate or acetate dependin on the supporting electrolyte) to an aqueous suspension o f silver sulfide precipitate a t constant ionic strength and of measuring electrochemically the equilibrium concentration of silver ion in solution. The amount of silver ion adsorbed on silver sulfide was determined by difference. A reversible silver su1fide.electrode and a reference electrode, the glass electrode, comprised the electrical cell used to measure the silver ion concentration in solution. The behavior of the Ag,S electrodes was checked against a glass electrode in a sodium acetate-acetic acid buffer of high silver concentration and also in an alkaline solution of high sulfide concentration. The stepwise addition of silver salt during the titration a t constant ionic strength was continued until the pAg approached a value of 6. Sufficient sodium hydrosulfide was then added to raise the pAg to about 15 and the titration (1) Based on a dissertation submitted by I. Iwasaki in partial fulfillment of the requirements for the degree of Doctor of Science, M.I.T. (2) D. Talmud and N. M. Lubman, Rolloid-Z., 60, 163 (1930). (3) 0. Jyo, Sci. Repts. Research Insls., Tohuku Univ., 8 6 , 259 (1954). (4) W.L. Freyberger and P. L. de Bruyn, THISJOURNAL, 61, 586 (1957). (5) D. C . Grahame, Chem. Reus., 41,441 (1947). (6) E . L. Meckor, Rec. trav. chem., 7 0 , 763 (1951).
procedure was repeated several times to establish the adsorption curve. Under the conditions of experimentation, the amount of silver sulfide precipitated during the titration never exceeded 10-6 g., a value which is negligible compared to the amount of precipitate (20 g.) already present in the system. Three different procedures were used in preparing silver sulfide precipitates. Method I consisted of adding to a saturated aqueous solution of purified HzS a 0.6 M AgNOa solution. Oxygen was excluded from the system by bubbling purified Nz through the solutions. During precipitation, HzSwas also passed through the suspension. Most of the AgzS precipitate used in this investigation was prepared by this method. Method I1 involved f i s t the precipitation of silver oxide by mixing equal volumes of 0.6 M AgN03 solution with a 0.8 M NaOH solution. The precipitate was washed wlth conductivity water and then dissolved by addition of NaCN. The silver cyanide complex solution was treated with HzS as in Method I in the presence of 0.1 M NaOH. Method I11 was the silver-ammine method uaed by Freyberger and de Bruyn4 for preparing all the AgZS for their investigation. The precipitates prepared by Methods I and I1 were washed with conductivity water immediately after preparation. All recipitates were then allowed to stand overnight in a 0.2 M%aOH solution saturated with HzS and were subsequently washed with conductivity water by decantation until free of foreign electrolyte. Precipitation and decantation washings were carried out whenever possible in a nitrogen atmosphere. X-Ray analyses of all three types of precipitates indicated no noticeable difference in the crystal structure which was that of p-silver sulfide or acanthite, the stable form at room temperature. Preliminary experiments showed that aging of the precipitates for a t least ten days was necessary to stabilize the AgzS surface and to obtain reproducible results.
Results Effect of Indifferent Electrolyte.-Two series of adsorption isotherms relating silver ion adsorption on silver sulfide to pAg were determined under conditions of constant pH and ionic strength. For each series, three adsorption curves at ionic strengths of and 10-l M were obtained. The adsorption curve at M was first established, then the ionic strength was raised to M by addition of supporting electrolyte at pAg 13, the curve at 10-2 M was then determined and by the same procedure, the curve at 10-1M . In all these experiments the silver sulfide precipitate was prepared by the nitrate method (Method I). The first series of adsorption curves was determined with sodium acetate as supporting electrolyte and the pH of the solution was maintained at 4.7 with an equimolar buffer solution of sodium acetate and acetic acid. The adsorption results are plotted in Fig. 1. I n the second series of tests,
May, 1958
ELECTROCHEMICAL DOUBLE LAYERON SILVERSULFIDE
the ionic strength was fixed with sodium nitrate and the pH was kept a t 4.7 by additions of "03 or NaOH. The experimental curves for this series are shown in Fig. 2. I n both figures the three isotherms a t different ionic strengths intersect in a point located slightly above pAg 10. The pAg value of this point of intersection determines the zero-point-of-charge for silver sulfide in the absence of specific adsorpti~n.~ Zero-point-of-charge.-The zero-point-of-charge (zpc) is one of the fundamental properties of the double layer system and hence it is desirable to locate it as accurately as possible. Furthermore, before the arbitrary adsorption units in Figures 1 and 2 can be replaced by absolute adsorption densities the zpc and specific surface of the silver sulfide must be known. 'To determine the zero-point-of-charge more accurately than the curves in Fig. 1 and 2 permit, the two series of adsorption experiments mentioned previously were repeated but the pAg of the solutions was varied only in a narrow range from pAg 9 to pAg 12. An additional adsorption curve a t ionic strength 10-4 M was determined for both series. The results of these tests are plotted in Fig. 3 and 4. It is apparent from Fig. 3 that all the isotherms cross a t one point, and it may be concluded that in a sodium nitrate solution a t pH 4.7 the zpc lies a t pAg 10.2. For the two curves a t ionic strengths and M the point of inflection is also seen to fall a t this pAg value. The slope of the adsorption curve at this inflection point measures the minimum differential capacity which for a simple double layer system is reached a t the zero-point-of charge,6 thus providing further evidence that pAg 10.2 is the true zpc. I n Fig. 4 all the curves except the one at ionic M again intersect a t pAg 10.2. The strength M in Fig. 4 is not located corisotherm a t rectly, probably due to error in measurement while changing the ionic strength of the titration solution from to M . (An independent measurement (Fig. 1) shows clearly that all curves including the loM2M curve intersect at one point.) As in sodium nitrate solutions, the best value for the zero-point-of-charge in a sodium acetate-acetic acid buffer solution is therefore also pAg 10.2. This value agrees reasonably well with a value of pAg 10 measured by Freyberger and de Bruyn.4 Effect of the Method of Precipitate Preparation. -To investigate whether the method of precipitate preparation has any effect on the adsorption process, experiments were carried out also with silver sulfide prepared by the silver-cyanide (11) and the silver-ammine (111) methods. The supporting electrolyte used in these experiments was a sodium acetate-acetic acid buffer at pH 4.7 and the measurements were made a t ionic strengths 10-3, 10-2 and 10-1 M. The zero-point-of-charge for precipitates I1 and 111 was observed to lie a t pAg 10.2 and pAg 10, respectively. For comparison, the adsorption isotherms determined on the three precipitates at an ionic strength of lo-' M are plotted in Fig. 5. The specific sur-
595
-
IONIC STRENGTH IO-' M. IO-' M. --b IO-' M. -e-
l . . . . . . . . , , ~ SURFACE AREA: 21.OoOcm'lg
6
10 12 14 PAg. Fig. 1.-Adsorption density of potential determining ions on silver sulfide as a function of pAg; supporting electrolyte, sodium acetate at pH 4.7.
8
-
IONIC STRENGTH M. IBLM. IO-' M.
+
SURFACE AREA: 20,000 c d l g
10 12 14 PAg. Fig. 2.-Adsorption density of potential determining ions on silver sulfide as a function of pAg; supporting electrolyte, sodium nitrate a t pH 4.7.
6
8
$0
3x 1
4 1
1
.g 8 2
-a
-
8
IONIC STRENGTH M.
.3
13
e3
8
-.+- I Q a -+-
M.
IO-'
M.
--b
M.
I
9
10 11 12 PAg. Fig. 3.-Adsorption density of potential determining ions on silver sulfide as a function of pAg; supporting electrolyte, sodium nitrate a t pH 4.7.
faces of the precipitates were calculated by assuming the zpc. to lie a t pAg 10.2 and that the minimum M has differential capacity a t ionic strength the same value (6 p farads per cm.z) as that for mercury in a 10-3 M sodium fluoride solution.6 With this assumption the adsorption density may be expressed as pcoulombs per cm.2. In Fig. 5 the isotherm determined by Freyberger and de Bruyn at the same ionic strength is also included.
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I. IWASAKI AND P. L.
DE
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Vol. 62
facial free energy. (y) for the silver sulfide-solution system at constant pH and ionic strength may be expressed by the differential equation d-y = u s d E
IONIC STRENGTH 10- M. IO-' M. * IO-' M. -&IO-' M.
-..o--
-=-
10
9
0'
12
11
PA& Fig. 4.-Adsorption density of potential determining ions on silver sulfide as a function of pAg; supporting electrolyte, sodium acetate a t p H 4.7.
(1)
where us is the surface charge of potential determining ions and E is the e.m.f. of the reversible cell used in the titration experiments. Equation 1 resembles the well-known Lippmann equation for an ideal polarized electrode or interface. The surface charge on AgzS is related to the adsorption density (I?) of the potential determining ions by the relation4 =
F(rAp+
- 2rS-)
(2)
This definition of us is significant only if the concentration of the supporting electrolyte is large compared to those of the silver and sulfur-bearing ions. I n Fig. 6 the change (y yo) in interfacial free energy where yo refers to the value a t the zpc. is plotted against pAg. The curves were obtained by graphical integration of the adsorption isotherms for AgzS in NaN03 solutions. For positively charged surfaces the ( y - yo) values shown in Fig. 6 are approximately equal to two-thirds of the corresponding values for Precipitate I in sodium acetate solutions and about one-half of the corresponding values quoted by Freyberger and de Bruyn4 for Precipitate I11 in sodium acetate solutions. Further differentiation of equation 1 leads t o a definition of another thermodynamic property of the system, namely, the differential capacity (CT)of the double layer a t a reversibly charged interface
-
METHOD OF SUPPORTING PREPARATION ELECTROLYTE AgN03+HzS NaNOa AgN03+HrS NaAc --r- Ag(CN);+ H,S No Ac ---c Ag(NH3);+H2S NaAc FROM FREYEEAGER I , , AND DE ERUYN
--P-
----
I
6
8
10
12
14
PA^. Fig. 5.-Adsorption density i f potential determining ions on silver sulfide a8 a function of VAE. Effect of the method of precipitate preparation and of tkk supporting electrolyte is shown; ionic strength, 0.1 M .
I n general, the different isotherms agree quite closely except that below pAg 8 the curve established with Precipitate I11 shows a more rapid increase in silver ion adsorption density than the other curves. This same difference was observed when comparing the isotherms a t other ionic strengths. The isotherm on Precipitate 111, however, is seen to be in good agreement with that determined by Freyberger and de Bruyn. To compare the role of different electrolytes, the adsorption curve for 10-1 M NaN03 (see Fig. 2) is also included in Fig. 5. The shape of the adsorption curve and the magnitude of the adsorption density when the surface is negatively charged are seen to be nearly independent of the type of indifferent electrolyte and the method of preparation of the silver sulfide. On positively charged surfaces both the nature of the supporting electrolyte and the method of preparation of the precipitate are seen to affect the course of the adsorption curve. For precipitate I the adsorption curve is steeper in sodium acetate solutions than in the presence of sodium nitrate. This difference has been observed to become larger with increasing ionic strength. Discussion Thermodynamical Considerations.-Freyberger and de Bruyn4 showed that the change in inter-
(5)
=
all chemical potentials (p) constant except p ~ ~ + + ( p s ' ) and paoivent
except p~4+(ps') and psolvent
A measure of the magnitude of CT for silver sulfide is obtained by differentiating the experimental adsorption curve. The absolute value of the differential capacity may be determined from the @pAg curve by assuming a value of 6 microfarads per cm.2 for the differential capacity at the zeropoint-of-charge in a solution of ionic strength M . This value has been obtained experimentally by Grahame6 for mercury in contact with a M NaF solution. I n making this assumption it is implied that in the vicinity of the zpc, the magnitude of the differential capacity is independent of the nature of the solid phase provided specific adsorption is absent. The differential capacities for AgzS prepared by three different chemical methods in solutions of low2and sodium acctate at ionic strengths of lo-' M are plotted against pAg in Fig. 7. For comparison, CT-pAg curves for AgzS prepared by . the nitrate method in NaN03 solutions of ionic strength 10-8, lo-' M are also included. The exceptionally high capacities a t low pAg values for Precipitate I11 are in agreement with the results previously obtained by Freyberger and
ELECTROCHEMICAL DOUBLE LAYERON SILVERSULFIDE
May, 1958
de Bruyn.4 These anomalous results might be explained by specific adsorption of the acetate ion, although the lower CT values registered by Precipitate I in the same solutions and the failure to detect a measurable shift in the zero-point-of-charge in the more concentrated acetate solutions argue against such an interpretation at least in the vicinity of the zpc. It is possible that the surface of a Ag2S precipitate prepared by the silver-ammine method is initially more reactive than the other Ag2S precipitates. The observed higher silver adsorption densities and the large differential capacities may be the result of a slow aging process of the more reactive precipitate. When the solid is positively charged, the differential capacity of Precipitate I is always slightly higher in sodium acetate than in sodium nitrate solutions. This observation suggests that the acetate ions can get closer to the surface than nitrate ions. Similar observations have been made on the behavior of various indifferent anions in the silver iodide-solution system. Lijklema' noted that for a positively charged AgI surface in contact with KF, KNOa or NaC104 solutions of the same concentration and pAg, the differential capacity is highest in KN03 and lowest in KF solutions. He ascribes these results to differences in the degree of hydration of the anions, the nitrate ion is assumed to be the least hydrated of the three ionic species. For the silver sulfide system, the pAg may be related to the reversible cell potential by the relation E
- EO = -2.303
y(pAg
- PA$)
6
8
10
12
597
14
16
PA& Fig. 6.-Interfacial free energy lowering a t the silver sulfide-solution interface as a function of p A g ; supporting electrolyte, sodium nitrate a t pH 4.7.
(4)
I n equation 4,EOand pAgo represent the cell e.m.f. and pAg, respectively, at the zero-point-of-charge; the quantity ( E - Eo)measures a relative potential difference. For the ideal polarized electrode, EOis known as the potential of the electrocapillary maximum. I n Fig. 8 the differential capacity of Ag2S prepared by the nitrate method and in contact with NaN03 solutions is plotted against the potential difference relative to the potential a t the zero-point-of-charge as defined by equation 4. Similar plots for silver iodide in NaC104 solutions6 and mercury in NaF solutionsE are included for comparison. The shape of the capacity curves for silver iodide and silver sulfide is remarkably similar at all ionic strengths; only at ionic strengths below 10-1 M do the curves for mercury resemble those of the two solids. However, the difference between the mercury curve and the Ag2S or AgI curve on positive polarization should not be stressed too much because at potentials more positive than 400 mv. an increase in capacity is also observed for mercury, The same characteristic steep rise in capacity on positive polarization shown by Ag2S and AgI is noted for the Hg/NaF system at temperatures greater than 250° and for methanol as solvent at room temperature.'" The dotted curve shown in Fig. 8 is for 0.001 N NaF in conductivity water at (7) J. Lijklema, Doctor's Dissertation, University of Utreoht, 1957. ( 8 ) D. C. Grahame, J. Am. Chem. Soc., 76, 4819 (1954). See also O.N.R. Tech. Report, No. 14, 1954, by same author. (9) D. C. Grahame, J. Am. Chem. SOC.. 79, 2093 (1957). (10) D. C. Grahame, 2. Elektrochem., 69, 740 (1955).
6 8 10 12 14 6 8 10 12 14 6 8 10 12 14 PAg. Fig. 7.-DifTerential capacity of the double layer on silver sulfide a8 a function of pAg: - - -, precipitate prepared by silver amine method in sodium acetate solution; - - -, precipitate prepared by nitrate and cyanide method in sodium acetate solution; ---, precipitate prepared by nitrate method in sodium nitrate solution.
-
85O.I1 At relatively high negative potentials, the differential capacity for silver sulfide is seen to approach a value of about 15 microfarads/cm.2. This levelling off in the vicinity of 15 microfarads/ cma2is also shown by AgI and Hg and was observed recently for liquid gallium12 thereby suggesting that the nature of the solid exercises no control on the magnitude of the capacity at high negative potentials. The steeper rise in the capacity at increasing positive potentials compared to increasing negative potentials for Ag2S and also AgI suggests that the anion moves closer to the surface than the cation which is always strongly hydrated. Consequently, to achieve the same potential jump across the double layer when the Ag2S surface is positively charged compared to a negatively charged surface, (11) The information needed for constructing this curve kindly supplied by Dr. D. C. Grahame. (12) Private communication from Dr. D. C. Grahame, Amherst College, Amherst, Mass.
598
I. IWASAKI AND P. L. I I
I
"O"OF
8
0
Eu
I
0
I , I
.
I . ' .
s g o w 3 8 + + I I +
8
9
Ap2S/NoN0,
8I
Potential relative to zero-point-of-charge (mv.). Fig. &-Differential capacities of the double layer on silver iodide, mercury and silver sulfide at room temperature.
l'
'
'
'
'
"
DE
Surface charge, pcoul./cm.z. Fig. 9.-Differential capacity, C, of the inner region of the electrical double layer as a function of surface charge on AgI, Ag2S and Hg a t room temperature.
01.
,
,6 8
.
,
. I .
10 12 14
6
. 8
.
.
J
10 12 14
P&. Fig. 10.-Differential capacity of double layer on silver sulfide in NaNOa solutions as a function of pAg. Dotted curves by calculation, other curves by differentiation of adsorption isotherms.
silver ions must adsorb in greater amounts than sulfide ions (see Fig. 5). Quantitative Interpretation of Capacity Curves.The behavior of the total differential capacity versus relative potential curves may be explained by the theory of the electrical double layer. Grahame8J0.18has shown that the experimental capacity curves for mercury in contact with aqueous NaF solutions may be reproduced by calculations based on this non-thermodynamic theory. This author obtained excellent agreement between observation and theory by applying in the absence of (13) D. C. Grahame, 2. EEekfroehem., 69, 773 (1955).
Vol. G2
specific adsorption the Gouy-Chapman Theory of the diffuse double layer to a physical model which divides the liquid layer into a charge-free inner region in contact with the mercury surface and a diffuse layer. The electrical analog of this model is seen to be two condensers in series for which it follows that
where C, and cd denote, respectively, the differential capacity of the inner or Stern region and the diffuse layer. The thickness of the inner region is determined by the closest distance of approach of the electrical centers of the counter ions. Both c d and C, are non-thermodynamic quantities. The magnitude of C d is given by the diffuse double layer theory, cd =
t
BRUYN
+
19.46~[(u~)~ 137.36~]'/2microfarads/cm.* at 25" (6)
where x is the valence of the ions of a symmetrical electrolyte, uB is the surface charge in microcoulombs per cm.2 and c is the concentration of the electrolyte in moles per liter far from the interface. There is no simple relationship defining C, in terms of fundamental variables; Grahame postulated that its magnitude should depend on the surface charge but not on the electrolyte concentration. By using this postulate and equation G it is then possible to calculate CT from equation 5. The same analysis was applied to silver sulfide in contact with aqueous solutions of NaN03. Figure 9 shows the variation of C, with 8 for this system as calculated from the experimental results at ionic strength 10-l M . Similar plots for Hgs and Ag17 are also included in Fig. 9. The C,-as curve for silver iodide is also derived from adsorption data at an ionic strength of 10-1 M while the curve for mercury is based on direct capacity measurements in a 0.916 M NaF solution. The horizontal scale for mercury has been condensed to show the whole curve. In Fig. 10 the calculated CT versus pAg curves for silver sulfide at ionic strengths of and M are compared with similar curves obtained by direct differentiation of the adsorption isotherms. The agreement is quite satisfactory and may be cited as proof that the theory of the diffuse double layer and the Grahame postulate regarding the dependence of C, on CTS are applicable to the completely reversible silver sulfide-solution interface in aqueous solutions of NaN03 at p H 4.7. Figure 9 reveals some interesting features. The resemblance between the Hg and AgzS curves is quite striking. Both mercury and silver sulfide show a capacity minimum at a negative surface charge, the minimum of about 11.5 microfarads/ cm.2 lies at - 1.2 microcoulombs/cm.2 for AgzS and a minimum of 17.2 microfarads/cm.2 is observed at - 12 microcoulombs/cm.2 for mercury.1° The authors believe that silver sulfide is the first heteroionic solid for which the existence of this capacity minimum has been demonstrated. The AgI curves do not show such minima within the range of surface charge accessible to experimentation. Liquid gallium12 also shows a minimum at a lower value
May, 1958
ORGANIC CATIONEXCHANGE PROPERTIES OF CALCIUM MONTMORILLONITE
of US than - 12 microcoulombs/cm.2. It would appear therefore, that for different solids the minimum C, value occurs a t different us values. The Cs curves are of interest because they reflect the interaction between the solvent molecules and the surface charge in the inner region. Although a quantitative explanation of these curves is not possible as yet, Grahame9~io presented qualitative evidence t o suggest that the increase in C, on both sides of the capacity minimum for Hg is perhaps largely due to a decrease in the thickness of the inner region. This decrease in thickness is postulated to result from the compression of the solvent molecules which are probably completely or almost completely oriented, by the electrical field emanating from the surface. I n view of the demonstrated similarity in the Cs versus us curves for Hg and AgzS this proposed qualitative theory by Grahame may also be applied t o AgzS. The similarity between the Hg and AgzS curves shown in Fig. 9 becomes even more marked when Cs versus us curves for the mercury system a t higher temperatures9 (>25O) and also for mercury in contact with methanolic solutions of NaF,1° are
599
compared with the Ag2S curve. The plateau-Iike region of slowly increasing capacity at small positive surface charge shown by mercury in Fig. 9 is seen to disappear in the other Hg curves. The existence of an “ice-like layer” of oriented solvents which excludes anions has been postulated by Grahameg to explain the behavior on low anodic polarization of the C, versus US curves a t 0 and 25”. The disappearance of the “hump” a t higher temperatures is ascribed to the “melting” of the “ice layer.” It would seem that this ‘‘ice layer” which is manifested by the characteristic humps in the CT versus US and C, versus US curves for mercury, is completely absent on AgzSand AgI. Acknowledgments.-The authors are grateful to the Atomic Energy Commission for financial support in conducting this investigation. To Dr. J. Th. G. Overbeek of the University of Utrecht, Dr. A. M. Gaudin and Dr. C. Wagner of the Massachusetts Institute of Technology, Dr. D. C. Grahame, of Amherst College, and Dr. W. L. Freyberger of the New Jersey Zinc Company, the authors are indebted for their helpful suggestions and interest shown in this investigation.
ORGANIC CATION EXCHANGE PROPERTIES OF CALCIUM M0NTMORILLONITE1j2 BY W. H. SLABAUGH AND F. KUPKA Department of Chemistry, Oregon State College, Corvallis, Oregon Received December 1 1 , 1967
Selectivity coefficients for the exchange of diethylamine, n-butylamine, di-n-butylamine and n-octylamine ions with calcium montmorillonite were determined at 25 and 45’, and from these data isosteric heab of exchange, free energies and entropies were calculated. A mechanism for the exchange process is proposed in which the length of hydrocarbon chain and nature of the amine group are considered.
Introduction The cation-exchange properties of montmorillonites and the change in physical properties brought about by the replacement of inorganic ions with amine ions have been studied in considerable detail by a number of workers. However, determination of exchange equilibria on homoionic montmorillonites at various temperatures has been confined mainly to the exchange of inorganic cati o n ~ . ~Study , ~ of the exchange equilibrium constant with organic cations has been limited to isothermal conditions.6 I n this study, a specially prepared calcium montmorillonite was treated with perchlorates of primary and secondary aliphatic amines. The selectivity constants a t different equivalent fraction of calcium montmorillonite for different temperatures were obtained and, when possible, thermodynamic (1) Taken in part from the Doctoral Thesis of F. Kupka, August, 1957, Oregon State College. (2) This work was performed under a fellowship sponsored by the Baroid Division of the National Lead Company. (3) G. L. Gaines, Jr., and H. C. Thomas, J . Chem. Phys., 23, 2322 (1965). (4) J. A. Faucher, Jr., and H. C. Thomas, ibid., !42, 258 (1954). ( 5 ) W. H. Slabaugh, THISJOURNAL, 68, 162 (1954).
functions were calculated.6 The isosteric heats of exchange and free energy for the reaction at 0.30 equivalent fraction of calcium montmorillonite have been calculated. Experimental Raw bentonite from Clay Spur, Wyoming, was converted to a 2% suspension, allowed to stand three days and decanted from the residue of quartz, large particles and other dense material. The suspended material was diluted to 0.3% solids and passed through three columns containing the. calcium form of a sulfonic acid type organic exchange resin. Each column converted more than 98% of the montmorillonite to the calcium form. The resulting suspension was concentrated by decantation after standing for several weeks and the excess salts were then removed by dialysis. The calcium montmorillonite resulting from this treatment contained an excess of only 1.1 mole % of calcium ion and had a base exchange capacity of 102 milliequivalents per 100 g. of solid. The amine perchlorates were prepared from the amines by neutralization of a water solution to a pH of 5.8 to 6.0.’ Secondary amines were obtained from K and K Laboratories, and the n-octylamine was redistilled Armeen 8D from the Armour Chemical Division of the Armour Company. The concentration of these solutions was determined by conventional means. Primary standard disodium dihydro(6) W. J. Argersinger, A. W. Davidson and 0. D. Bonner, Trans. Kans. Acad. Sci., 63, 404 (1950). (7) W. H. Slabaugh and V. E. Cates, AnaE. Chem., 27, 151-(1955).
