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The ionic double layer at the ZnOsolution interface
The Ionic Double Layer at the ZnO/Solution Interface I. The Experimental Point of Zero Charge L. BLOK AND P. L. DE BRUYN ~ Department of Metallurgy and Materials Science, Massachusetts Institute of Technology, Cambridge, Massachusetts 02189 Received June 24, 1969; accepted October 21, 1969 An experimental study of the adsorption of potential-determining ions at the zincite (ZnO)/aqueous solution interface is reported. Changes in solubility of the oxide with pH are shown to limit adsorption measurements to the pg range 8-10. The pzc of ZnO, as determined by the intersection of adsorption curves at low ionic strength, is shown to depend on the method of preparation and on the nature of the supporting electrolyte. I. INTRODUCTION
ion inclusion in the soluble oxide when formed by wet precipitation methods will be Quantitative information on the ionic enhanced. This contamination must exerdouble layer at the metal oxide-aqueous solucise some control on the subsequent history tion interface has been obtained almost exof the interface and might confound predicclusively by acid-base potentiometric titrations of surface behavior. tions of oxide suspensions. The initial studies Even though the interaction of ZnO with of Parks and de Bruyn (1) on hematite its aqueous environment might differ in de(FelOn) were continued by Onoda and de tail from that of less soluble oxides, the Bruyn (2) and Atkinson, Posner, and Quirk thermodynamic description of its double (3). Recently, B6rub6 and de Bruyn (4, 5) layer should conform to the general treatemployed this experimental approach to ment developed by B~rub6 and de Bruyn evaluate the ionic double layer on rutile (4) in a recent publication. The tow dif(TiO2); Tadros and Lyklema (6) extended fusivity of Zn ++ and O = in the solid at room this technique to studies of porous SiO~. In temperature suggests that thermodynamic this paper, we discuss results obtained in a equilibrium will be established only between study of the zincite-solution interface. In choosing ZnO for investigation, a new the solid surface (or the first few atomic layers variable is introduced in the experimental of lattice ions) and the surrounding liquid evaluation of the ionic double layer at oxide phase. On the assumption that the chemical surfaces. The higher solubility of this oxide potential (~) of surface lattice ions is not in comparison with either TiO2 or Fe20~ will significantly altered by adsorption from solucomplicate the potentiometric titration be- tion, the change in surface potential (~0) cause now dissolution processes will also may then be written as compete significantly for the added acid or RT &k0 -- ~-- d in (azn++/a~++); base. Furthermore, the probability of foreign [11 RT 1To whom requests for reprints should be 2F d In (ao~/ao=). directed. --
Journal of Colloid and Interface Science, ¥ol. 32, No. 3, March 1970
518
pzc
IONIC DOUBLE LAYEI~ AT ZnO/SOLUTION INTERFACE. I
519
The pK values quoted above were calculated from standard free energy of formation data as listed by Latimer (8). For dilute solutions ( ~ o ~--- constant) in contact.with pure solid phase (~zno = constant), it then follows that
i
d~z~++ = - d ~ o = = - 2 d ~ o s -
z w
-5
Zn(OH) 2
= 2d~H+ = 2d~z~oR+
o
=
IS1
- - d/~zno2 =
= 2/~d~Ez,2+~(om2~+11+++ " In view of Eq. [8], which identifies H + and OH- as p.d. ions, Eq. [1] may be rewritten to give the well-known relation -~5
I 0
2
I 4
6
8,
10
12
d~o = R ~ -T d ln (aH+~ \a~/
p H ~
FIG. 1. Equilibrium solubility diagram of ZnO. Here the ionic activities relate to the aqueous solution phase only and the superscript pzc refers to the point of zero charge. The chemical potentials (and activities) of dissolved lattice ions may be related to those of a number of other ionic species, which collectively are called the potential-determining (p.d.) ions, with the aid of the following dissolution reactions: Zn0(s)
= Zn ++ + O=.
[2]
ZnO(s) + H20 = Zn ++ + 2OH-; pK = 16.4. Zn0(s) + H20 = Zn0H + + OH-; pK = 12.0. ZnO(s) + H20 = ZnO2H- + H+; pK = 17.0. ZnO(s) + H20 = ZnO~= + 2H+; pK = 29.7.
[31
[4] [5] [6]
(2 + n)ZnO(s) + (2 + n)H20 = {Zn2+.(OH)2~+I}~+ + 3OH-. 2 Matijevid,
Couch,
and
Kerker
(7) p r e s e n t e d
evidence for the existence of this polynuclear zinc complex of valence +3 and with a Zn/OH ratio of either 1/1 or 2/1. No thermodynamic information is available on this proposed reaction.
[91 _
R T d In ( a o . - ~ F ~ " \ao.-/
The solubility diagram for ZnO in equilibrium with an aqueous solution of variable pH is shown in Fig. 1. The solubility curve is based on the thermodynamic data cited above. The pH-independent concentration of the neutral species Zn(OH)2 was established by Fulton and Swinehart (9). From this diagram the point of minimum solubility may be calculated to fall at pH 9.7. This pH, according to the Parks and de Bruyn criterion (1), should also give the pzc of the oxide. It is important to note that whereas the magnitude of the solubility will depend on the nature of the zinc-bearing solid (ZnO, crystalline or amorphous Zn(OH)2), the point of minimum solubility should be independent of the particular solid modification assumed in equilibrium with the aqueous phase (2). Any ionic species, other than OH-, which complexes zine ions in solution will alter the solubility of ZnO as depicted in Fig. 1 and m u s t therefore also be treated as a potentialdetermining ion. To this important group of anionic species belong the halide ions. The chloride ion, for example, is known to form four ehlorozinc eomplexes, ZnC1+, ZnCl~, Journal of Colloid and Interface Science, Vol. 32, No. 3, March 1970
520
:
BLOK AND DE BRUYN
ZnCla-, and ZnC14=. Tabulated experimental values for the stability constants of these complexes vary over a wide range. The more reliable experimental results at room temperature fall in the range 10-°'~ to 10+0.5 (10). A simple calculation therefore shows that in 10-1 M chloride solutions only the concentration of the species ZnC1+ may approach that of the Zn ++ ion. For chloride ion concentrations lower than 10-1 M the concentration of all zinc-chlorocomplexes is small compared to that of zinc cation. In all our experiments the chloride ion concentrations never exceeded 10-1 M and the pH of the solution was confined to the range pH 8-10; therefore, it is reasonable to neglect the potential-determining role of the chloride ion relative to that of the hydroxyl ion. The inferred pze of ZnO at pH 9.7 may be compared with experimental values. Widely varying values have been claimed by different investigators employing a variety of experimental techniques. Mieroelectrophoretie measurements gave pzc values of p i t 10.3 (11) and 9.3 (12). In unpublished adsorption results Ray of this laboratory suggests a pze at pH 8.7 in NaNO3 and NaCI04 solutions and a pzc at pH 9.2 in NaC1 solutions of concentration less than 0.01 M. Herezynska and Proszyska (13) reported pH 9 4- 0.3 for the pzc of oxidized zinc metal. This wide variation in the observed pzc of ZnO appears to be characteristic in general of all metal oxide systems (12). No coneerted effort has yet been made to find a rationale for this variation. If. EXPERIMENTAL
range. To relate the adsorption of p.d. ions to the surface potential, the following operational definition of pOH was introduced pOH -
Journal of Colloid and Interface Science, Vol. 32, No. 3, March 1970
[10]3
where X± refers to the mean ionic activity coefficient of the sodium salts (NaC1, NaBr, NaI, NaC104, NaNO~) used as supporting electrolytes. The pH of the system is then defined by pH = pKwa~or - pOH.
[111
The potentiometric titration of the ZnO suspension was carried out by adding small amounts, 0.05 to 0.3 ml, of acid or base to the suspension every 3 to 5 min. The pH of the suspension was determined 2 to 3 min after this addition. Stirring of the suspension had no effect on the emf of the glass-calomel electrode assembly at ionic strengths above 0.01 M. A small effect, 1 to 2 mv, was sometimes noticeable at low ionic strengths, 10.3 M and smaller. From these measurements, after correction for solubility effects, socalled fast (2, 4) adsorption isotherms were constructed. The temperature was kept at 25 ° 4- 0.2°C. III. SOLUBILITY CORRECTIONS The addition of a finite amount of acid to a suspension of ZnO will change the concentration of all p.d. species in the solution and will shift the adsorption equilibrium. A materials balance on the solution phase after the addition of AV ml of an N normal acid solution to the suspension, therefore, will show that
PROCEDUI~
A detailed description of the experimental setup and the procedure followed in the potentiometric titration of oxide suspensions has been given in previous publications (i, 2, 4). The emf of the titration cell consisting of a Beckman G.P. glass electrode and a Beckman calomel or Ag/AgCI reference electrode was determined as a function of log concentration of OH- ions because all titrations were performed in the alkaline pI-I
--log X±(samCoE-,
A V × 2( = [ + I [ +
A~r~*, [12]
where I = (V + AV){[H+]y -- [OH-]i} -
v{ [H+]~
-
[OR-ld
[13]
represents the increase in H + concentration 3This is the same definition introduced by B6rub~ and de Bruyn (4) except that these authors used the symbol X~a, which could be misinterpreted to refer to the activity coefficient of the undissociated salt.
IONIC
DOUBLE
70
50 60 40 E
LAYER
AT ZnO/SOLUTION
[
T
i
10-11
-
INTERFACE.
521
I
/
~
-
3O
:i. C~
• 20
o
I0-2M
--
@"--"
e""'-
•
I _ 5
0
! I 10 15 WEIGHT (gr(lms)
~_ 20
I
FIG. 2. A@ as a function of ionic strength and weight of suspended material in pH range 8.5 to 9.0.
in the solution;
can be evaluated. Furthermore, the materials balance as written assumes that all other
I I = (V + AV){2[Zn++]f
+ [ZnOH+]f -- [ZnO2H-]~, - 2[ZnO2=]A - V{2[Zn++], [14] + [ZnOH4]~ -- [ZnO2K-]~ -
-
2[ZnO2=]~}
represents the amount of added acid consumed in changing the solubility of ZnO; and AAF~* = A{AFH+ - AFoH+ 2AFzn++
-
-
2AFzno2=
[15]
+ AFznoH+ -- AFHzno~=} gives the total change in adsorption accompanying the acid addition. In these equations, the square brackets denote concentrations in solution, the subscripts i and f relate to the initial and final concentrations, and A is the surface area of the solid. In writing the materials balance only univalent and divalent zinc hydroxocomplexes are considered; the trivalent complex concentration is expected to be small in comparison. The experimental technique does not allow a distinction to be made among the different adsorbed species and, therefore, only AFt*
chemical reactions which consume or release t t + or OH- ions are suppressed or eliminated. To this end, CO2 was excluded from the experimental system by a continuous flow of prepurified nitrogen gas. The titration technique measures directly the quantity AQ =
AV X N -
I = II +
A~r~*.
[16]
The adsorption Ar~* can then be evaluated if the solubility change / / is known either by thermodynamic calculation or by direct measurement. Now the quantity H is independent of the amount of solid in the suspension and depends only on the volume of the solution and the solubility difference between the final and the initial equilibrium state. The adsorption density, being proportional to the surface area, will depend on the amount of solid present at a fixed solution volume. A linear dependence of AQon the mass of the solid phase at cons{ant solution volume is, therefore, expected. A series of experiments was carried out by suspending up to 20 gm of ZnO in 300 ml of solution to determine AQ as a function of pH at three ionic strengths, 10-3, 10-2 and Journal of Colloid and Interface Science, Vol. 32, No. 3, March 1970
522
BLOK AND DE BRUYN 9o
t
I
I I Salt
70
I
60
2
& ~o 0
I
li0" M
I
\\
50-
1
I _~
No
\
,,\ --
4 0 --
M
--
30--
E 20-10--
8.o
8.5
9.0
9.~
lo.o
pH Fro. 3. Experimentally determined solubility curves for ZnO precipitate. The dashed curves are calculated from thermodynamic data. 10-I M. A typical result of such a measurement is shown in Fig. 2 for the pH range 8.5-9.0. The intercept on the ordinate axis determines the amount of H + ions consumed in the dissolution reaction for that pH interval and the slope of the straight lines measures the adsorption per gram of ZnO at the chosen ionic strength. From a series of plots such as those shown in Fig. 2, experimental solubility curves at fixed ionic strength were calculated. The results of such calculations are reproduced in Fig. 3, which gives the amount of H + ions consumed per liter of solution as a function of pH. Since the experimental determinations yield only solubility differences, the solubility of ZnO has been set arbitrarily equal to zero at the point of minimum solubility (pit = 9.7). Figure 3 also includes solubility curves at zero and 10-I M concentrations of uni-univalent supporting electrolyte as calculated from Fig. 1. To establish the theoretical solubility curve for 0.1 M solutions, the concentrations of the various p.d. ions were calculated from expressions similar to Eq. [10] by assuming Journal of Colloid and Interface Science, Vol, 32, No. 3, March 197 0
values of 0.114 and 0.44 for - l o g h± of univalent and divalent p.d. ions, respectively. The solubility values calculated from the thermodynamic data of Fig. 1 are seen to be smaller than the values measured in this investigation except at relatively low pH values. This discrepancy can be ascribed to a failure to include the polynuelear, triply charged, zinc hydroxocomplex in the thermodynamic analysis. However, since the pK for reaction [(7)] is not known, it is not possible to evaluate the true significance of this oversight. We rather believe that the deviation can be accounted for by the presence of small amounts of finely divided amorphous Zn(OH)~ or some other less stable form of ZnO. These experiments clearly demonstrate that the operational solubility values cannot be assumed known from available thermodynamic data. The experimental measurements themselves are not too accurate as 4 T h e e x p e r i m e n t a l v a l u e ( a p p r o x i m a t e d to t h e s e c o n d decimal place for m o s t u n i v a l e n t s o d i u m s a l t s (14)).
IONIC i ~'uEE -1
NoNO3 o 0.00021,,t ,~, 0.0007 M 0 0.002 M • 0.005 M • 0.016 M
:2..
DOUBLE
I
LAYER
1
'
AT
ZnO/SOLUTION
I "~ ~"
INTERFACE.
523
I
of at least 1 m s. Below pH 8.5, the accuracy of the adsorption measurements decreases rapidly because the solubility of the oxide increases more rapidly with decreasing pH NO NO 3 o 0.0015 M eel.
o<
I
8.0
8.5
I
I
9.0
9.5
pH
-4 -
FIG. 4. F a s t a d s o r p t i o n i s o t h e r m s (adsorption d e n s i t y in m i e r o c o u l o m b s / c m ~ vs. pH) on ZnO I I I in NaNO3 solutions (see T a b l e I).
shown by the scatter of points in Fig. 2. It has been established that to obtMn a 10 % accuracy in the adsorption density in the pH range 8.5-10, the amount of precipitate in 1 ml of solution must have a surface area
-
Precipitare
I
II
III
IV
V
VI
VII
/
<
J / ,~
°/Tj °/
" i.-.-o -I°'~
8.0
jo/
I
8.5
1
[
9.0
9.5
p H ~
FIG. 5. F a s t adsorption isotherms on ZnO V I in NaNO~ solutions (see T a b l e I).
TABLE DEPENDENCE
/ °
I
O F PZC ON M E T H O D
OF PREPARATION Specific Surface
Method of preparation
(m~/gm)
R e a c t 1 l i t e r of a 1 M ZnCI~ s o l u t i o n w i t h excess 2 M N a O H at boiling point. Reflux for 10 days. C o n t a m i n a t e d b y silica because of dissolution of glass. Same r e a g e n t c o n c e n t r a t i o n as for I, b u t r e a c t i o n carried out at room t e m p e r a t u r e . P r e c i p i t a t e refluxed, t h e n washed, ZnC12 added to lower supern a t a n t p H to 7. Reflux for 10 days. E x p e r i m e n t a tion a f t e r 1.5 years storage. R e a c t 1 M ZnCl~ w i t h excess 8 M K O H at room t e m p e r a t u r e . Age at 100°C in p o l y e t h y l e n e cont a i n e r in w a t e r b a t h for 10 days. E x p e r i m e n t a tion a f t e r 1.5 years storage. Room t e m p e r a t u r e p r e c i p i t a t i o n of 1 M ZnCI~ w i t h excess 8 M KOH. F i n a l pH 10.5. Reflux for 20 days. R e a c t a t boiling p o i n t cone. ZnC12 s o l u t i o n (450 gm ZnC12 in 400 ml water) w i t h 8 M KOII. Viol e n t reaction. Washed, refluxed for 2 days. R e a c t solution of 450 gm Zn(NO3)2 in 1.5 liters of w a t e r w i t h one-half a m o u n t of 10 N N a O t t required for complete p r e c i p i t a t i o n a t boiling point. Aged for 7 days a t 100°C. R e a c t 1 M Zn(NO3)~ w i t h 10 N N a O I t in presence of 0.003 a t o m % In. Reflux for 4 days. Precipit a t e fired a t 1250°C to homogenize i n d i u m distribution.
pzc in N a salt Solutions NO~-
13.1
C104-
C1-
Br-
Not determined
13.1
8.8
11.5
9.5 10.0
16.5
8.8
8.5
8.6
--
8.8
<8.0
--
8.25 - -
9.5
--
6.5 ~
13.5
0.5 ~'
I-
9.1
9.2
8.7
8.7
--
Low specific surface due to conc. ZnC12 s o l u t i o n used. Low specific surface due to c a l c i n a t i o n of p r e c i p i t a t e . Journal of Colloid and Interface Science, Vol. 32, No. 3, March 1970
524
BLOK AND DE BRUYN
-31
a
L :~
.o.o3
--2
o A [] •
--1
c~
I
/
than does the adsorption. Potentiometric titrations were thus limited to the relatively narrow p H range 10-8.5.
1
/ .
00003M 0002M 0013M 0035M
IV. ADSORPTION RESULTS
0
S L
i
<
I 8.0
8.5
I
I
9.0
9.5
pH
(u E
-3
I
u ~:L - 2 >~
L Z~
b
I
I
NaCl o A u --1 • -- • •
/
0.0002 M 000061'4 0,002 b4 0.0055 M 0.016M 0.105 M
S 8.0
8.5
9.0
9,5
pH
I
E u
~.-2 _
1
c
I
A
NaBr
/ / ~ "
0,0004 M 0.0018M 0.006 M 0.016 M
--
5 d3 I-
2 < 8.0
8.5
9D
9.5
pH
-31
I
I
~L
z~ U o •
--1
d
t
oooo .
/
I
0.0015 M 0.005 M o.o15 M 0.04M
~
_
E 8
-
~ 8D
I
I
1
8,5
9.0
9.5
-
pH
F r o . 6. F a s t
adsorption
isotherms
on ZnO V in
various uni-univMeat electrolyte solutions. (a) NaNO~ , (b) NaC1, (c) NaBr, (d) NaI. Journal of Collc4dandlnterface ~cience, Vol. 32, No. 3, ~Iarch 1970
The ZnO used in the experimental adsorption measurements was prepared b y precipitation from aqueous ZnCl2 or Zn(NO~)~ solutions with N a O H or KOH. X - R a y diffraction studies verified the crystalline zincire structure of the precipitate, and electron micrographs revealed the crystals to be needle shaped with a length of the order of 1 # and a diameter of a few hundred angstroms. The specific surface was obtained from B E T krypton adsorption measurements. Solubility corrections were made as outlined in the previous section. A strong dependence of the pzc of the precipitate on the method of preparation is demonstrated b y the experimental study. This dependency is illustrated in Fig. 4 and 5, which give the fast adsorption isotherms for two precipitates in NaNO3 solutions at 25°C. The precipitate, ZnO III, used in obtaining the adsorption results of Fig. 4, was prepared b y precipitation from I M ZnCI~ solutions with excess 8 M K O H 5 at room temperature. I t was subsequently aged for 10 days at 100°C in a polyethylene container in a water bath. The adsorption measurements were made 1.5 years after the preparation. The pzc, as determined by the intersection of the low ionic strength curves, is seen to lie above p H 9.5. Precipitate VI, which was used in obtaining the adsorption results of Fig. 5, was prepared b y allowing a concentrated solution of ZnC12 to react with an amount of 10 N N a O H just sufficient to precipitate one-half of the available zinc as ZnO. The reaction was carried out at the solution boiling point and 5 The precipitation was carried out with KOH because one might expect less foreign cation incorporation in ZnO when the radius of the foreign cation is larger than that of the lattice cation. The radii of IC+, Na +, and Zn++ are 1.33, 0.98, and 0.69 A, respectively.
IONIC DOUBLE LAYER AT ZnO/SOLUTION INTERFACE. I the precipitate was aged for 7 days at 100°C. A pzc below pH 8 is extrapolated for this preparation. Altogether, seven different precipitates were used. The pertinent information on the preparation and the observed pzc in NaNO3 solutions is summarized in Table I. In contrast to the large effects of the preparation method on the location of the pzc, the influence of different univalent anions of the supporting electrolyte on this parameter was relatively small. Figure 6 illustrates the adsorption curves obtained on precipitate V with NaNO~, NaC1, NaBr, and NaI. The pzc varies between pH 8.9 and pH 8.5, with ZnO in NaNO~ solutions showing the lowest value for the pze. V. SUMMARY AND CONCLUSIONS An experimental study of the adsorption of p.d. ions at the zincite-aqueous solution interface was made in the pH interval 8.5-10. The measurements were limited to this narrow pH range because of the relatively high solubility (10 -s moles per liter and higher) of the oxide. In order to account for changes in solubility in the potentiometric titration of the suspension, it was necessary to make direct measurements of the solubility differences with pH. Although it is not expected that the point of minimum solubility will differ much from that calculated from available thermodynamic data, the actual solubility of the precipitates deviates substantially from the thermodynamic calculations (Fig. 3). The pzc of the zinc oxide precipitates used in this study was established by the intersection of fast adsorption isotherms at low ionic strengths. In agreement with observations of previous investigators (11-13), the pzc was observed to vary over a wide pH range (about 2 p i t units). It is now possible to correlate this variation with the particular conditions used in the preparation of the oxide from aqueous solutions. A pze value approaching that of the point of minimum solubility (pH 9.7) is indicated for precipi-
525
tares made by reacting ZnCI2 solutions with a large excess of base. If the precipitation is carried out with less than the stoichiometric amount of base, the pzc drops to a value as low as pH 8. A low pzc (below pH 9) is also found when precipitation occurs from a concentrated zinc chloride or nitrate solution. These observations suggest that the incorporation of foreign anionic species in the solid during precipitation may exercise an important control on the surface behavior of this oxide. Further consideration will be given this conclusion in a second paper. In contrast to results from adsorption studies with TiO2 (~), the pzc of ZnO precipitates is also influenced by the nature of the supporting electrolyte. The precipitates consistently showed a lower pze in NaNOs solutions than in NaG1 solutions. InsufiL cient experiments were made with other electrolytes to establish any further trends. The magnitude of the observed shift is always less than that observed with precipitates prepared by different methods. A small shift in pzc with increasing concentration (especially above 0.05 M) of supporting electrolyte is also observed. This effect cannot be ascribed simply to specific adsorption of the anion because the observed shift is in a direction (increasing pH) opposite to that expected. ACKNOWLEDGMENTS The authors gratefully acknowledgethe counsel and advice offeredby ProfessorJ. Th. G. Overbeek and also the helpful discussions with Dr. Y. G. B~rub6. This investigation was made possible by the liberal financial support provided by the U. S. Army Research Office, Durham, North Carolina. REFERENCES 1. I:}ARKS, G. A., AND DE BRUYN, P. L., or. Phys.
Chem. 66, 967 (1962). 2. ONODA, G. Y., JR., AND DE BRUYN, P. L.,
Surface Sci. 4, 48 (1966). 3. ATKINSON,R. J., POSNEI~, A. M., ANDQVlRI<, J. P., J. Phys. Chem. 71,550 (1967). 4. BI~RUB]~, Y. G., AND DE BRUYN, P . L., or.
Colloid Interfac. Sci. 27, 305 (1968). Journal of Colloid and Interface Science, Vol. 32, No. 3, March 1970
526
BLOK AND DE BlZUYN
5. B]~RUB]~, Y. G., AND DE Bt~UYN, P. L., Or. Colloid Interfac. Sei. 28, 92 (1968). 6. TADROS, TH. F., AND LYKLEMA, J., J. Electroanal. Chem. 17, 267 (1968). 7. MATIJE¥I¢, E., COUCH, J. P., AND KERKER, M., J. Phys. Chem. 66, 111 (1962). 8. LATIMER,W. M., "Oxidation PotentiMs." 2nd ed. Prentice Hall, Englewood Cliffs, New Jersey (1956). 9. FULTON, J. W., AND SWINEHART, D. F., Or. Am. Chem. Soc. 76, 863 (1954). 10. SILL~N, G. L., AND MARTELL, A. E., "Stability
Journal of Colloid and Interface Science,
Vol. 32, No. 3, March 1970
11. 12. 13.
14.
Constants of Metal-Ion Complexes," Chem. Soc. (London) Spec. Publ. 17 (1964). MATTSON, S., AND PUGH, A. J., Soil Sci. 38, 229 (1934). PAm~s, G. A., Chem. Rev. 65, 177 (1965). HERczYNsKA, E., AND PROSZYSKA, K., Polish Academy of Sciences Institute of Nuclear Research, Warsaw, Report No. 372/V, 1 (1962). ROBINSON, R. A., AND STOKES, R. I-I., "Electrolyte Solutions," Butterworths, London, (1965).
ion inclusion in the soluble oxide when formed by wet precipitation methods will be Quantitative information on the ionic enhanced. This contamination must exerdouble layer at the metal oxide-aqueous solucise some control on the subsequent history tion interface has been obtained almost exof the interface and might confound predicclusively by acid-base potentiometric titrations of surface behavior. tions of oxide suspensions. The initial studies Even though the interaction of ZnO with of Parks and de Bruyn (1) on hematite its aqueous environment might differ in de(FelOn) were continued by Onoda and de tail from that of less soluble oxides, the Bruyn (2) and Atkinson, Posner, and Quirk thermodynamic description of its double (3). Recently, B6rub6 and de Bruyn (4, 5) layer should conform to the general treatemployed this experimental approach to ment developed by B~rub6 and de Bruyn evaluate the ionic double layer on rutile (4) in a recent publication. The tow dif(TiO2); Tadros and Lyklema (6) extended fusivity of Zn ++ and O = in the solid at room this technique to studies of porous SiO~. In temperature suggests that thermodynamic this paper, we discuss results obtained in a equilibrium will be established only between study of the zincite-solution interface. In choosing ZnO for investigation, a new the solid surface (or the first few atomic layers variable is introduced in the experimental of lattice ions) and the surrounding liquid evaluation of the ionic double layer at oxide phase. On the assumption that the chemical surfaces. The higher solubility of this oxide potential (~) of surface lattice ions is not in comparison with either TiO2 or Fe20~ will significantly altered by adsorption from solucomplicate the potentiometric titration be- tion, the change in surface potential (~0) cause now dissolution processes will also may then be written as compete significantly for the added acid or RT &k0 -- ~-- d in (azn++/a~++); base. Furthermore, the probability of foreign [11 RT 1To whom requests for reprints should be 2F d In (ao~/ao=). directed. --
Journal of Colloid and Interface Science, ¥ol. 32, No. 3, March 1970
518
pzc
IONIC DOUBLE LAYEI~ AT ZnO/SOLUTION INTERFACE. I
519
The pK values quoted above were calculated from standard free energy of formation data as listed by Latimer (8). For dilute solutions ( ~ o ~--- constant) in contact.with pure solid phase (~zno = constant), it then follows that
i
d~z~++ = - d ~ o = = - 2 d ~ o s -
z w
-5
Zn(OH) 2
= 2d~H+ = 2d~z~oR+
o
=
IS1
- - d/~zno2 =
= 2/~d~Ez,2+~(om2~+11+++ " In view of Eq. [8], which identifies H + and OH- as p.d. ions, Eq. [1] may be rewritten to give the well-known relation -~5
I 0
2
I 4
6
8,
10
12
d~o = R ~ -T d ln (aH+~ \a~/
p H ~
FIG. 1. Equilibrium solubility diagram of ZnO. Here the ionic activities relate to the aqueous solution phase only and the superscript pzc refers to the point of zero charge. The chemical potentials (and activities) of dissolved lattice ions may be related to those of a number of other ionic species, which collectively are called the potential-determining (p.d.) ions, with the aid of the following dissolution reactions: Zn0(s)
= Zn ++ + O=.
[2]
ZnO(s) + H20 = Zn ++ + 2OH-; pK = 16.4. Zn0(s) + H20 = Zn0H + + OH-; pK = 12.0. ZnO(s) + H20 = ZnO2H- + H+; pK = 17.0. ZnO(s) + H20 = ZnO~= + 2H+; pK = 29.7.
[31
[4] [5] [6]
(2 + n)ZnO(s) + (2 + n)H20 = {Zn2+.(OH)2~+I}~+ + 3OH-. 2 Matijevid,
Couch,
and
Kerker
(7) p r e s e n t e d
evidence for the existence of this polynuclear zinc complex of valence +3 and with a Zn/OH ratio of either 1/1 or 2/1. No thermodynamic information is available on this proposed reaction.
[91 _
R T d In ( a o . - ~ F ~ " \ao.-/
The solubility diagram for ZnO in equilibrium with an aqueous solution of variable pH is shown in Fig. 1. The solubility curve is based on the thermodynamic data cited above. The pH-independent concentration of the neutral species Zn(OH)2 was established by Fulton and Swinehart (9). From this diagram the point of minimum solubility may be calculated to fall at pH 9.7. This pH, according to the Parks and de Bruyn criterion (1), should also give the pzc of the oxide. It is important to note that whereas the magnitude of the solubility will depend on the nature of the zinc-bearing solid (ZnO, crystalline or amorphous Zn(OH)2), the point of minimum solubility should be independent of the particular solid modification assumed in equilibrium with the aqueous phase (2). Any ionic species, other than OH-, which complexes zine ions in solution will alter the solubility of ZnO as depicted in Fig. 1 and m u s t therefore also be treated as a potentialdetermining ion. To this important group of anionic species belong the halide ions. The chloride ion, for example, is known to form four ehlorozinc eomplexes, ZnC1+, ZnCl~, Journal of Colloid and Interface Science, Vol. 32, No. 3, March 1970
520
:
BLOK AND DE BRUYN
ZnCla-, and ZnC14=. Tabulated experimental values for the stability constants of these complexes vary over a wide range. The more reliable experimental results at room temperature fall in the range 10-°'~ to 10+0.5 (10). A simple calculation therefore shows that in 10-1 M chloride solutions only the concentration of the species ZnC1+ may approach that of the Zn ++ ion. For chloride ion concentrations lower than 10-1 M the concentration of all zinc-chlorocomplexes is small compared to that of zinc cation. In all our experiments the chloride ion concentrations never exceeded 10-1 M and the pH of the solution was confined to the range pH 8-10; therefore, it is reasonable to neglect the potential-determining role of the chloride ion relative to that of the hydroxyl ion. The inferred pze of ZnO at pH 9.7 may be compared with experimental values. Widely varying values have been claimed by different investigators employing a variety of experimental techniques. Mieroelectrophoretie measurements gave pzc values of p i t 10.3 (11) and 9.3 (12). In unpublished adsorption results Ray of this laboratory suggests a pze at pH 8.7 in NaNO3 and NaCI04 solutions and a pzc at pH 9.2 in NaC1 solutions of concentration less than 0.01 M. Herezynska and Proszyska (13) reported pH 9 4- 0.3 for the pzc of oxidized zinc metal. This wide variation in the observed pzc of ZnO appears to be characteristic in general of all metal oxide systems (12). No coneerted effort has yet been made to find a rationale for this variation. If. EXPERIMENTAL
range. To relate the adsorption of p.d. ions to the surface potential, the following operational definition of pOH was introduced pOH -
Journal of Colloid and Interface Science, Vol. 32, No. 3, March 1970
[10]3
where X± refers to the mean ionic activity coefficient of the sodium salts (NaC1, NaBr, NaI, NaC104, NaNO~) used as supporting electrolytes. The pH of the system is then defined by pH = pKwa~or - pOH.
[111
The potentiometric titration of the ZnO suspension was carried out by adding small amounts, 0.05 to 0.3 ml, of acid or base to the suspension every 3 to 5 min. The pH of the suspension was determined 2 to 3 min after this addition. Stirring of the suspension had no effect on the emf of the glass-calomel electrode assembly at ionic strengths above 0.01 M. A small effect, 1 to 2 mv, was sometimes noticeable at low ionic strengths, 10.3 M and smaller. From these measurements, after correction for solubility effects, socalled fast (2, 4) adsorption isotherms were constructed. The temperature was kept at 25 ° 4- 0.2°C. III. SOLUBILITY CORRECTIONS The addition of a finite amount of acid to a suspension of ZnO will change the concentration of all p.d. species in the solution and will shift the adsorption equilibrium. A materials balance on the solution phase after the addition of AV ml of an N normal acid solution to the suspension, therefore, will show that
PROCEDUI~
A detailed description of the experimental setup and the procedure followed in the potentiometric titration of oxide suspensions has been given in previous publications (i, 2, 4). The emf of the titration cell consisting of a Beckman G.P. glass electrode and a Beckman calomel or Ag/AgCI reference electrode was determined as a function of log concentration of OH- ions because all titrations were performed in the alkaline pI-I
--log X±(samCoE-,
A V × 2( = [ + I [ +
A~r~*, [12]
where I = (V + AV){[H+]y -- [OH-]i} -
v{ [H+]~
-
[OR-ld
[13]
represents the increase in H + concentration 3This is the same definition introduced by B6rub~ and de Bruyn (4) except that these authors used the symbol X~a, which could be misinterpreted to refer to the activity coefficient of the undissociated salt.
IONIC
DOUBLE
70
50 60 40 E
LAYER
AT ZnO/SOLUTION
[
T
i
10-11
-
INTERFACE.
521
I
/
~
-
3O
:i. C~
• 20
o
I0-2M
--
@"--"
e""'-
•
I _ 5
0
! I 10 15 WEIGHT (gr(lms)
~_ 20
I
FIG. 2. A@ as a function of ionic strength and weight of suspended material in pH range 8.5 to 9.0.
in the solution;
can be evaluated. Furthermore, the materials balance as written assumes that all other
I I = (V + AV){2[Zn++]f
+ [ZnOH+]f -- [ZnO2H-]~, - 2[ZnO2=]A - V{2[Zn++], [14] + [ZnOH4]~ -- [ZnO2K-]~ -
-
2[ZnO2=]~}
represents the amount of added acid consumed in changing the solubility of ZnO; and AAF~* = A{AFH+ - AFoH+ 2AFzn++
-
-
2AFzno2=
[15]
+ AFznoH+ -- AFHzno~=} gives the total change in adsorption accompanying the acid addition. In these equations, the square brackets denote concentrations in solution, the subscripts i and f relate to the initial and final concentrations, and A is the surface area of the solid. In writing the materials balance only univalent and divalent zinc hydroxocomplexes are considered; the trivalent complex concentration is expected to be small in comparison. The experimental technique does not allow a distinction to be made among the different adsorbed species and, therefore, only AFt*
chemical reactions which consume or release t t + or OH- ions are suppressed or eliminated. To this end, CO2 was excluded from the experimental system by a continuous flow of prepurified nitrogen gas. The titration technique measures directly the quantity AQ =
AV X N -
I = II +
A~r~*.
[16]
The adsorption Ar~* can then be evaluated if the solubility change / / is known either by thermodynamic calculation or by direct measurement. Now the quantity H is independent of the amount of solid in the suspension and depends only on the volume of the solution and the solubility difference between the final and the initial equilibrium state. The adsorption density, being proportional to the surface area, will depend on the amount of solid present at a fixed solution volume. A linear dependence of AQon the mass of the solid phase at cons{ant solution volume is, therefore, expected. A series of experiments was carried out by suspending up to 20 gm of ZnO in 300 ml of solution to determine AQ as a function of pH at three ionic strengths, 10-3, 10-2 and Journal of Colloid and Interface Science, Vol. 32, No. 3, March 1970
522
BLOK AND DE BRUYN 9o
t
I
I I Salt
70
I
60
2
& ~o 0
I
li0" M
I
\\
50-
1
I _~
No
\
,,\ --
4 0 --
M
--
30--
E 20-10--
8.o
8.5
9.0
9.~
lo.o
pH Fro. 3. Experimentally determined solubility curves for ZnO precipitate. The dashed curves are calculated from thermodynamic data. 10-I M. A typical result of such a measurement is shown in Fig. 2 for the pH range 8.5-9.0. The intercept on the ordinate axis determines the amount of H + ions consumed in the dissolution reaction for that pH interval and the slope of the straight lines measures the adsorption per gram of ZnO at the chosen ionic strength. From a series of plots such as those shown in Fig. 2, experimental solubility curves at fixed ionic strength were calculated. The results of such calculations are reproduced in Fig. 3, which gives the amount of H + ions consumed per liter of solution as a function of pH. Since the experimental determinations yield only solubility differences, the solubility of ZnO has been set arbitrarily equal to zero at the point of minimum solubility (pit = 9.7). Figure 3 also includes solubility curves at zero and 10-I M concentrations of uni-univalent supporting electrolyte as calculated from Fig. 1. To establish the theoretical solubility curve for 0.1 M solutions, the concentrations of the various p.d. ions were calculated from expressions similar to Eq. [10] by assuming Journal of Colloid and Interface Science, Vol, 32, No. 3, March 197 0
values of 0.114 and 0.44 for - l o g h± of univalent and divalent p.d. ions, respectively. The solubility values calculated from the thermodynamic data of Fig. 1 are seen to be smaller than the values measured in this investigation except at relatively low pH values. This discrepancy can be ascribed to a failure to include the polynuelear, triply charged, zinc hydroxocomplex in the thermodynamic analysis. However, since the pK for reaction [(7)] is not known, it is not possible to evaluate the true significance of this oversight. We rather believe that the deviation can be accounted for by the presence of small amounts of finely divided amorphous Zn(OH)~ or some other less stable form of ZnO. These experiments clearly demonstrate that the operational solubility values cannot be assumed known from available thermodynamic data. The experimental measurements themselves are not too accurate as 4 T h e e x p e r i m e n t a l v a l u e ( a p p r o x i m a t e d to t h e s e c o n d decimal place for m o s t u n i v a l e n t s o d i u m s a l t s (14)).
IONIC i ~'uEE -1
NoNO3 o 0.00021,,t ,~, 0.0007 M 0 0.002 M • 0.005 M • 0.016 M
:2..
DOUBLE
I
LAYER
1
'
AT
ZnO/SOLUTION
I "~ ~"
INTERFACE.
523
I
of at least 1 m s. Below pH 8.5, the accuracy of the adsorption measurements decreases rapidly because the solubility of the oxide increases more rapidly with decreasing pH NO NO 3 o 0.0015 M eel.
o<
I
8.0
8.5
I
I
9.0
9.5
pH
-4 -
FIG. 4. F a s t a d s o r p t i o n i s o t h e r m s (adsorption d e n s i t y in m i e r o c o u l o m b s / c m ~ vs. pH) on ZnO I I I in NaNO3 solutions (see T a b l e I).
shown by the scatter of points in Fig. 2. It has been established that to obtMn a 10 % accuracy in the adsorption density in the pH range 8.5-10, the amount of precipitate in 1 ml of solution must have a surface area
-
Precipitare
I
II
III
IV
V
VI
VII
/
<
J / ,~
°/Tj °/
" i.-.-o -I°'~
8.0
jo/
I
8.5
1
[
9.0
9.5
p H ~
FIG. 5. F a s t adsorption isotherms on ZnO V I in NaNO~ solutions (see T a b l e I).
TABLE DEPENDENCE
/ °
I
O F PZC ON M E T H O D
OF PREPARATION Specific Surface
Method of preparation
(m~/gm)
R e a c t 1 l i t e r of a 1 M ZnCI~ s o l u t i o n w i t h excess 2 M N a O H at boiling point. Reflux for 10 days. C o n t a m i n a t e d b y silica because of dissolution of glass. Same r e a g e n t c o n c e n t r a t i o n as for I, b u t r e a c t i o n carried out at room t e m p e r a t u r e . P r e c i p i t a t e refluxed, t h e n washed, ZnC12 added to lower supern a t a n t p H to 7. Reflux for 10 days. E x p e r i m e n t a tion a f t e r 1.5 years storage. R e a c t 1 M ZnCl~ w i t h excess 8 M K O H at room t e m p e r a t u r e . Age at 100°C in p o l y e t h y l e n e cont a i n e r in w a t e r b a t h for 10 days. E x p e r i m e n t a tion a f t e r 1.5 years storage. Room t e m p e r a t u r e p r e c i p i t a t i o n of 1 M ZnCI~ w i t h excess 8 M KOH. F i n a l pH 10.5. Reflux for 20 days. R e a c t a t boiling p o i n t cone. ZnC12 s o l u t i o n (450 gm ZnC12 in 400 ml water) w i t h 8 M KOII. Viol e n t reaction. Washed, refluxed for 2 days. R e a c t solution of 450 gm Zn(NO3)2 in 1.5 liters of w a t e r w i t h one-half a m o u n t of 10 N N a O t t required for complete p r e c i p i t a t i o n a t boiling point. Aged for 7 days a t 100°C. R e a c t 1 M Zn(NO3)~ w i t h 10 N N a O I t in presence of 0.003 a t o m % In. Reflux for 4 days. Precipit a t e fired a t 1250°C to homogenize i n d i u m distribution.
pzc in N a salt Solutions NO~-
13.1
C104-
C1-
Br-
Not determined
13.1
8.8
11.5
9.5 10.0
16.5
8.8
8.5
8.6
--
8.8
<8.0
--
8.25 - -
9.5
--
6.5 ~
13.5
0.5 ~'
I-
9.1
9.2
8.7
8.7
--
Low specific surface due to conc. ZnC12 s o l u t i o n used. Low specific surface due to c a l c i n a t i o n of p r e c i p i t a t e . Journal of Colloid and Interface Science, Vol. 32, No. 3, March 1970
524
BLOK AND DE BRUYN
-31
a
L :~
.o.o3
--2
o A [] •
--1
c~
I
/
than does the adsorption. Potentiometric titrations were thus limited to the relatively narrow p H range 10-8.5.
1
/ .
00003M 0002M 0013M 0035M
IV. ADSORPTION RESULTS
0
S L
i
<
I 8.0
8.5
I
I
9.0
9.5
pH
(u E
-3
I
u ~:L - 2 >~
L Z~
b
I
I
NaCl o A u --1 • -- • •
/
0.0002 M 000061'4 0,002 b4 0.0055 M 0.016M 0.105 M
S 8.0
8.5
9.0
9,5
pH
I
E u
~.-2 _
1
c
I
A
NaBr
/ / ~ "
0,0004 M 0.0018M 0.006 M 0.016 M
--
5 d3 I-
2 < 8.0
8.5
9D
9.5
pH
-31
I
I
~L
z~ U o •
--1
d
t
oooo .
/
I
0.0015 M 0.005 M o.o15 M 0.04M
~
_
E 8
-
~ 8D
I
I
1
8,5
9.0
9.5
-
pH
F r o . 6. F a s t
adsorption
isotherms
on ZnO V in
various uni-univMeat electrolyte solutions. (a) NaNO~ , (b) NaC1, (c) NaBr, (d) NaI. Journal of Collc4dandlnterface ~cience, Vol. 32, No. 3, ~Iarch 1970
The ZnO used in the experimental adsorption measurements was prepared b y precipitation from aqueous ZnCl2 or Zn(NO~)~ solutions with N a O H or KOH. X - R a y diffraction studies verified the crystalline zincire structure of the precipitate, and electron micrographs revealed the crystals to be needle shaped with a length of the order of 1 # and a diameter of a few hundred angstroms. The specific surface was obtained from B E T krypton adsorption measurements. Solubility corrections were made as outlined in the previous section. A strong dependence of the pzc of the precipitate on the method of preparation is demonstrated b y the experimental study. This dependency is illustrated in Fig. 4 and 5, which give the fast adsorption isotherms for two precipitates in NaNO3 solutions at 25°C. The precipitate, ZnO III, used in obtaining the adsorption results of Fig. 4, was prepared b y precipitation from I M ZnCI~ solutions with excess 8 M K O H 5 at room temperature. I t was subsequently aged for 10 days at 100°C in a polyethylene container in a water bath. The adsorption measurements were made 1.5 years after the preparation. The pzc, as determined by the intersection of the low ionic strength curves, is seen to lie above p H 9.5. Precipitate VI, which was used in obtaining the adsorption results of Fig. 5, was prepared b y allowing a concentrated solution of ZnC12 to react with an amount of 10 N N a O H just sufficient to precipitate one-half of the available zinc as ZnO. The reaction was carried out at the solution boiling point and 5 The precipitation was carried out with KOH because one might expect less foreign cation incorporation in ZnO when the radius of the foreign cation is larger than that of the lattice cation. The radii of IC+, Na +, and Zn++ are 1.33, 0.98, and 0.69 A, respectively.
IONIC DOUBLE LAYER AT ZnO/SOLUTION INTERFACE. I the precipitate was aged for 7 days at 100°C. A pzc below pH 8 is extrapolated for this preparation. Altogether, seven different precipitates were used. The pertinent information on the preparation and the observed pzc in NaNO3 solutions is summarized in Table I. In contrast to the large effects of the preparation method on the location of the pzc, the influence of different univalent anions of the supporting electrolyte on this parameter was relatively small. Figure 6 illustrates the adsorption curves obtained on precipitate V with NaNO~, NaC1, NaBr, and NaI. The pzc varies between pH 8.9 and pH 8.5, with ZnO in NaNO~ solutions showing the lowest value for the pze. V. SUMMARY AND CONCLUSIONS An experimental study of the adsorption of p.d. ions at the zincite-aqueous solution interface was made in the pH interval 8.5-10. The measurements were limited to this narrow pH range because of the relatively high solubility (10 -s moles per liter and higher) of the oxide. In order to account for changes in solubility in the potentiometric titration of the suspension, it was necessary to make direct measurements of the solubility differences with pH. Although it is not expected that the point of minimum solubility will differ much from that calculated from available thermodynamic data, the actual solubility of the precipitates deviates substantially from the thermodynamic calculations (Fig. 3). The pzc of the zinc oxide precipitates used in this study was established by the intersection of fast adsorption isotherms at low ionic strengths. In agreement with observations of previous investigators (11-13), the pzc was observed to vary over a wide pH range (about 2 p i t units). It is now possible to correlate this variation with the particular conditions used in the preparation of the oxide from aqueous solutions. A pze value approaching that of the point of minimum solubility (pH 9.7) is indicated for precipi-
525
tares made by reacting ZnCI2 solutions with a large excess of base. If the precipitation is carried out with less than the stoichiometric amount of base, the pzc drops to a value as low as pH 8. A low pzc (below pH 9) is also found when precipitation occurs from a concentrated zinc chloride or nitrate solution. These observations suggest that the incorporation of foreign anionic species in the solid during precipitation may exercise an important control on the surface behavior of this oxide. Further consideration will be given this conclusion in a second paper. In contrast to results from adsorption studies with TiO2 (~), the pzc of ZnO precipitates is also influenced by the nature of the supporting electrolyte. The precipitates consistently showed a lower pze in NaNOs solutions than in NaG1 solutions. InsufiL cient experiments were made with other electrolytes to establish any further trends. The magnitude of the observed shift is always less than that observed with precipitates prepared by different methods. A small shift in pzc with increasing concentration (especially above 0.05 M) of supporting electrolyte is also observed. This effect cannot be ascribed simply to specific adsorption of the anion because the observed shift is in a direction (increasing pH) opposite to that expected. ACKNOWLEDGMENTS The authors gratefully acknowledgethe counsel and advice offeredby ProfessorJ. Th. G. Overbeek and also the helpful discussions with Dr. Y. G. B~rub6. This investigation was made possible by the liberal financial support provided by the U. S. Army Research Office, Durham, North Carolina. REFERENCES 1. I:}ARKS, G. A., AND DE BRUYN, P. L., or. Phys.
Chem. 66, 967 (1962). 2. ONODA, G. Y., JR., AND DE BRUYN, P. L.,
Surface Sci. 4, 48 (1966). 3. ATKINSON,R. J., POSNEI~, A. M., ANDQVlRI<, J. P., J. Phys. Chem. 71,550 (1967). 4. BI~RUB]~, Y. G., AND DE BRUYN, P . L., or.
Colloid Interfac. Sci. 27, 305 (1968). Journal of Colloid and Interface Science, Vol. 32, No. 3, March 1970
526
BLOK AND DE BlZUYN
5. B]~RUB]~, Y. G., AND DE Bt~UYN, P. L., Or. Colloid Interfac. Sei. 28, 92 (1968). 6. TADROS, TH. F., AND LYKLEMA, J., J. Electroanal. Chem. 17, 267 (1968). 7. MATIJE¥I¢, E., COUCH, J. P., AND KERKER, M., J. Phys. Chem. 66, 111 (1962). 8. LATIMER,W. M., "Oxidation PotentiMs." 2nd ed. Prentice Hall, Englewood Cliffs, New Jersey (1956). 9. FULTON, J. W., AND SWINEHART, D. F., Or. Am. Chem. Soc. 76, 863 (1954). 10. SILL~N, G. L., AND MARTELL, A. E., "Stability
Journal of Colloid and Interface Science,
Vol. 32, No. 3, March 1970
11. 12. 13.
14.
Constants of Metal-Ion Complexes," Chem. Soc. (London) Spec. Publ. 17 (1964). MATTSON, S., AND PUGH, A. J., Soil Sci. 38, 229 (1934). PAm~s, G. A., Chem. Rev. 65, 177 (1965). HERczYNsKA, E., AND PROSZYSKA, K., Polish Academy of Sciences Institute of Nuclear Research, Warsaw, Report No. 372/V, 1 (1962). ROBINSON, R. A., AND STOKES, R. I-I., "Electrolyte Solutions," Butterworths, London, (1965).