iinirradiated, :ire in sufficient agreement to suggest that the same chemical reaction occurs after the slow reaction. Acknowledgment.-The authors wish to thank
the Council for Scientific and 1riclristri:tl lirsearch (South Africa) for a grant to cover the cost of‘ irradiation and for a scholarship held by P. J. H. during the investigation.
THE ZERO POIST OF CHAIiGE OF OXIDES’ BY
G. A. P A R K S AXD 1’. L.
DE B B U S S 2
1)cpnrlnirnt of” 3Trtallurgy, Massachusetts Institute of Technology, Cambridge 39, l~fassarhiksctls Receiied June PG, 196’1
The zero point of charge (z.P.c.) of the electrical double layer of frre charges a t reversible interfaces is definrd, and the espcviniental methods for dcterniining this parameter arc discussed. Adsorption isotherms of I3 + and OH- ions on crystalline fcrric oxide precipitates as obtained by acid and base potentiometric titrations are reported. From these studies at various ionic strrngths (fix08indifferent electrolyte), the Z.P.C.of a-hematite is established at p H 8.6. A . model of the doublr layer ai the Fe&-aqucoiis solution interface is developrd. This model is based on the dissociation of surface hydroxyl groups forrncd by hydration of the virgin surfacc. In addition, it is shown that the Z.P.C.may be drtcrmined by considering thr adsorption of metal hydroso-coinplcx ions from solution. According to this analpis it is found that for Fr,03the 2.p.c. qgc’es with the PIT of niiriimurn solubility of the oxide and with the isoelectric p H as determined by the positively and negatively charged Fc(II1) hydroso-complcsrs in solution. i\lthough limited exprrimental data are availablr, this analysis may br applird to other ovidc systenis.
Introduction arbitrariness is introduced in assigning the potrnDepending or1 the mechanism by which free tial-determining ions only to the solid side of the charges are distributed across a solid-solution interface.8 By allowing the concentration of the interface, two basic types of electrochemical double indifferent electrolyte in the solution to esceed layers may be distinguished. A reversible double greatly that of the potential-determining eleclayer is the result of transfer of potential-determin- trolyte, the uncertainty in the physical significance ing ions across the two-phase boundary. I n of gais avoided. The potential difference, $o, across the double contrast, the completely polarizable double layer is established by imposing by an external poten- layer of free charges may be related to the c.m.f. tiometer circuit a potential difference (to the inter- ( E ) of reversible cell which uses as one of the face) so that all charge species have meaningfid reversible electrodes the solid to be studied. By concentrations only in one of the two bulk phases. suitable choice of :Lrcfercncc electrode, one may Equilibriiim studies of double layers on AgI,304 write Ag2S,586arid liquid mercury’ in contact with $0 = E - Eo (2) aqueous electrolyte solutions have illustrated the similarity in structiire and in the magnitude where Eo is the e.m.f. of the cell a t the z.p.c. of the electrical properties of thew t x o types of From (2) it follows thatg double layer systems. The most important parameter which may be tiscd to describe a double layer of free charges is its zero point of surface charge (z.P.c.). For Both relations 2 and 3 are only valid in the absence of any specific interaction between the solid siirfacc :I reversible double layer system the surface charge, and the ions derived from the supporting electro(u8)may be expressed by the relation lyte or between the surface and non-ionized = F ( Z +r+ + Z - r - ) (1) molecules in solution. In the presence of finite where F is the Yarnday constant; x + ( x - ) the val- amoiints of such solute species, the z.p.c. also will ( m e including sign; and I’+ (r-) the adsorption be a function of their activities. density of the potential-determinirig ionic species. To locate the z.p.c. for a rcversiblc double layer ‘The Z.P.C. then is expressed by the condition that system, the activity (a,(”,) of the potential us be zero and will be determined by a particular determining species (i) must be determined by \ d u e of the activity (UL(0) also a- ( 0 , ) of the potential- experiment. From potentiometric titrations of a rniinirig ions in the solution phase. For re- finely divided suspcmion of t: heteroionic solid, \-crsible intcrfaccs the concept of a surface charge the surface charge and thus also the: z.p.c. may is loss iiri:imhigiious than that of a completely be determined. In addition l o this adsorption polarized system. This follows because a certain method, the e.p.c. also may lie located by first determining the zero point of any property which (1) Parts of this paprr are baaed on a dissertation hy G. A. Parks in partial fulmlinent of t h e ioquiremonts for the degree of Doctor of depends on the presence of an electrical double Philosophy, M.I.T. layer. Thus the absence of the suspension effect (2) Associate Piofessor of Metallurgy, M.I.T. between a settled siispension and the supernatant (3) E 1,. >lackor. Rec. Lrau. chrm., 70, 763 (1951). (4) J. Liiklema, Kollozd-2.. 176, 129 (1961). ( 5 ) W. I,. lhyberger and P. L. de Bruyn, J . Phys. Chem., 61, 586 (1%7). (6) I. Inasirki and P L de Rruyn, zbzd., 63, 5 q l (1958). ( 7 ) I) C Ciaharnr, .I A m . Chem Sor., 76, 4810 (1934).
( 8 ) J. Th. G. Overbeek, Nntl. Bur. Standards (v. S,), Cirr 524, 1953, p. 213. (9) Equation 3 should not be intrrpreted to indicate a zero potrntial diffrrenrc aciosa the reversible interfncr: it only drfinw the m i o point of potential across a double layer of free charges.
solution, the zero point of any electrokinetic property, and the maximum flocculation rate also may be used to evaluate the z.P.c., if specific adsorption effects are eliminated.i0 This paper deals with the evaluation of the 2.p.c. of crystalline oxides and in particular with Fc203. I n contrast to studies with AgI and AgsS solids, most inorganic solids do not function satisfactorily as reversible electrodes. A direct identification of thc potential-dctennining species is, therefore, not possible. Howcvcr, by definition the potential-determining species for Fe201 are the ionic constituents of the lattice (Fe3+and 02-) and ions, in particular, II+ and OH-, which arc in equilibrium with them. Accordingly, a potentiometric titration of a ferric oxide precipitate in a reversible cell which employs a glass electrode as one of the electrodes should yield information on the 2.p.c. The Zero Point of Charge of Ferric Oxide.-The z.p.c. of ferric oxide was determined by the adsorption method which was used so successfully for AgI and AgSS. A potentiometric titration of an aqueous suspension of FelO, was perto 1M ) and at 21’ in formed a t several ionic strengths the reversible cell g l ~ s selectrode 1 aqueous suspension satd. KNOa saturated of FezOa;supporting salt bridge calomel electrolyte KN03 electrode
I
I
1
As titrants, 0.1 N solutions of KOH and IIN03 were used. These solutions were prepared from Acculute standard volumetric solutions (,Inachemica Chemicals Jltd.) by dilution with double-distillcd conductivity water. The alkaline solution was stored in polycthylcne bottles to avoid silicate contamination. Potassium nitrate wa9 chosen as supporting electrolyte because the anions NOa- and C10,- arc least lihely to compleu Fe(ITI).ll A Ikckman saturated calomel electrode connected t o the titration cell through a KNO, saturated salt bridge and a Beckman eneral purpose glms electrode served as the two electrodes. !’he salt bridge, introduced to prevent chloride contamination, was identical in design to thc asbestos-wickrd bridge used in Beckman calomel clectrotles except that the lower end of the bridge was filled with solid KNO, crystallixcd in situ. Leakage of KNO, into the ccll from the bridge was rapid enough so that below 1 0 - 2 ,If control of ionic strength was possible only by rcmoval of the bridge after each mcasi~remcnt. .I Rlodel G I3eckman 1111 meter was used to detcrmine the plI of the sy?tem. l’he titration ccll w w :t flat-bottomed 500-mi. I’yrcx jar with a rubber stopper through which passed the electrodes, a thermometer, a microburet, and an inlct and outlet for purified nitrogen gas. Xitrogen n‘as used to purge the ccll and buret of C02. hlising of the suspension was provided by a Teflon-covercd magnetic stirrer. The apparatus \vas kept a t constant temperature by pumping thermostuted water through a j:tckrt around the cell and through a rubber tube wrapped around the calomel elcctrode. Owing to the high rcsistance of the cell a t low ionic strength, all parts of the equipment including the p11 mcter caSe were carefully giounded. Thr ferric osiclc prccipitatc \vas prepared by hydrolyzing Baker and Adamson reagent grade Fe(N03), a t the boiling Doint. .Ibout 100 L‘. of the salt added to one liter of conhuct.ivi\y wntcr w& boiled under reflux conditions for 18 days. l o cnsure that the hydrolysis product was a-Fe20a and not &’e0011 (goethite), the IINOa formed by hydrolysis was prevcnted from escaping by vaporization.’Z The prccipitatc W‘IH washed by centrifuging and decantation with conductivity wat,cr. Vsually the washing was done under flocculated conditions by adding KOII t o give a p H of about 9. Examination of the washed ferric oxide by X-ray diffraction verified it, to be rhombohedral a-hematitc.
__-
(10) IT. R. Kruyt, “Colloi~l Scienre,” Vol. I , Elsevier I’ul~l. Co..
New York, N. Y., 3952, p. 231. (11) J. C. Bailar, “The Cliemistry of t.he Coordination Compounds,” Reinhold I’ubl. Corp., N e w York, N. Y.,1996. (12) F. C. Smith and D. J. Kidd, A m . Mineralogist, 34,403 (1949).
Spectrographic analysis of the precipitate revealed no major impurity. The titration procedure consisted of adding betwecn 250 and 300 g. of a stock suspension of FaOa (about 3 t o 5 g. of Fe203)to the cell. Most commonly a series of titrations was started a t ionic strength IO-, and p H 9.3, and small increments of acid were made. The p1-I range 5 to 10 was covered. The approach to eqnilibrium was fast (lcss than 20 min.) i n the high alkaline and in the acid p€I range but in the vicinity of the Z.P.C. and near the equivalence point (pH 7) equilibration times as high as 4 hr. were experienced. The adsorption density (I’) of OH- and H+or rather the excess of one species over the other ( r H + - r o e ) ~ a determined a by the difference between total added base or acid and the equilibrium OH- and €I+ concentration in solution. The reproducibility and revcrsibility of the adsorption curves waa well established. I n all, a large number of testa were made a t different ionic strengths in the narrow p H range 9.5 to 7 and a limited number of testa over the wider range p€I 5 to 10. A typical adsorption curve is shown in Fig. 2, where the adsorption density of OH-and H+ ions in moles/g. is plotted against pH. From the results presented in Fig. 1 one may conclude that OH- and H+ ions are potential-determining for Feo03 and that I of the two surface sites is about unity, then it follows from (9) that the magnitude of KI is determined by the pH of the Z.P.C. (pH 8.5) as obtained from Fig. 1. Thus, K 1is equal to 10-17, a result which shows that the proton is strongly bonded t,o the surface. It is interesting to note that this approach to the evaluation of the Z . P . C . of ferric oxide is similar to the procedure used for determining the isoelectric pH of proteins in solution. Furthermore, the electrochemical behavior of the FezOa-solution interface a t different ionic strengths, as illustrated by Fig. 1, is very similar to that of the titration curves of proteins in aqueous solutions.25 The Relationship between Z.P.C. and the Point of Minimum Solubility.-As written, the surface reaction 8 also suggests that as an alternative to the surface dissociation mechanism, the establishment of the surface charge may be explained by the adsorption from solution of the hydroxo-complex species, FeOz- (or Fe(OH)4-) and Fe(OH)2+. Indeed for negligible differences in the adsorbability of these two species the activity ratio of the surface sites may be replaced by the activity of the species in the aqueous solution which is in equilibrium with solid Fen03. Subject to these conditions the Z.P.C. will be determined by the pH of the aqueous solution for which the activities of the aqueous species, Fe(OH)4- and Fe(OH)z+,are equal. At this point it is worthwhile to interrupt the discussion of the Z.P.C.of FezOJ to recall an important characteristic of complex solutions. If a group of positive and negative complexes of the same central ion is considered independently of any other ionic species in solution, this group will be electroneutral when &+e+ = zz-e-
n e u t r a I hydra l e d i n t e r f a c e
Fig. 3.-Schematic
illustration of an uncharged hydrated ferric oxide surface.
negatively charged surface sites are exposed. The surface reactions involved in the establishment of a surface charge and the electrical double layer may be represented formally as Fe(OH)z (surface) +
+
2H20 Fe(OH)3 (surface)
Fe(OH), (surface) FeOz- (surface)
+ HaO+(aq.) (6)
+ Ha0 +(aq.)
(7)
I n these reactions Fe(OH)3(surface) represents the uncharged surface or rather a surface adsorption site which by adsorbing a proton becomes positive (Fe(OH)2+) or by desorbing a proton becomes negative (FeOz-). From ( G ) and (7) it follows that the surface site (Fe(OH)3) has amphoteric properties, thus Fe(OH)3may be regarded as the conjugate base of the acid site, Fe(OH)*+,or as the conjugate acid of the base, FeOz-. By adding reactions G and 7, the important result is obtained that Fe( OH)z+ (surface)
FeOz+ (surface)
+ 2H
(8)
[FeOz- (surface)] x [Fe(OH)z+(surface)]
(E)x
where z+(x-) is the net charge of a positive (negative) complex and c+(c-) the concentration of the positive (negative) complex. The ligand concentration (or activity, to be exact) at which this condition is satisfied determines the isoelectric point of the complex solution.2E J’ Furthermore, when the complex system includes a solid phase in equilibrium with the solution, the isoelectric point (i.e.p,) of the solution corresponds to the point of minimum solubility of the solid. A clear and lucid proof of this observation has been given by Beck.*’ Figure 4 illustrates the relation between the i.e.p. of an aqueous solution in contact with solid AgzO and the point of minimum solubility of the solid. The existence of a number of positive Fe(II1) hydroxo-complexes has been well established and their stability constants are known. Although the presence of a negative complex, Fe(0H)e(or FeOz-), in highly alkaline solutions has been known, it is only recently that Lengweiler, Buser, and Feitknechtzs were able to determine from total (25) R . K. Cannan, -4.H. Palmer, end
for which the equilibrium constant
K1
Vol. 66
$. C. Kibrick, J . Biol. Chem.,
148, 803 (1942). UHte
(9)
measures the strength with which the ferric oxide surface binds protons. If the ratio of the activity
(26) L. P. Hammett, “Solution of Electrolytes,” MaGraw-Hill Book Go., New York, N. Y., 1936. (27) 11. T. Beck, Acta Chzm. Acad. Sci. Huno., 4 , 227 (1954). (28) H. Lengweiler, W. Buser, and W, Feitknecht, Helv. Chtm Acta 44, 796 (1961).
solubility mcasurcments of 17'c6°((OH)a the eqiiilihrium oonstlantfor tha roaction Fe(OH)j(s) OH-
+
Fe(OH),-(itrl.); ph' = 5 rt 0.15 (10)
in the presence of 3 A4 NaC104. From a, knowledge of the standard frce enargy of form:it,ion (AG?) of solid Ice(OII):3 and I"(~rOa,2" ihc cqiiilibriiun constant, for the rcact.ion
+ s/11-120 + OH-
I/2Fe203(s)
Fe(OI-I)4-(fbq.); pK
=
1o.ti
(11)
may bc calculated. The corrcsponding rc:iction for the monovalcnt positive complex is l/~F('zOa(~)t H.' '/zH*O I-
+
Fe(OH)Z+(:iq.); p K = 7.6 (12)
-
PH I n Fig. 5 the composition of t,he aqucous solut'ion Fig. 4.-Vttri:ition of the concentration of silvci-bearing eqiiilihriiim n.it,h solid lj'e&):i is demonst'rat>cdby a plot of log co1icentratiori3" against pI-1. A11 sprrirs with pH in an aqueous solution in eqiiilihriuin with Ag,O. (Thc data used in constructing this figure wcrc pcrtinant thermodynamic dat,s iiscd in con- solid taken from ref. 27 and 46.) struct,ing Fig. 5 are summarized in Table 11. The 2 isoclect,ric point> of t hc complcx solution is stfen I " " I to lic at, a pH of 8.5. E'urthcrmore, in the Yicinity of this pIi, the concentratioii of thc two monovalent complex specics far exceeds that of thc polyvale~it~species. The minimum solubility of Ik203also falls at pH 8.5. Hecaiisc of tha high conccntration of thc undissocintcd ~ p c c i e s( Fe(OH)3), ~~ the minimum is not as sharply defined as t'hat of hg2O* iii
TABLE I1 ~ ' I I E R Y O D Y N A M I C DATAUSEI) IS COXSTRCCTISG F I G .
5 FeOHz H+; pK = 2 . 1 7 2 I?C~(OH)~~ 2H + + ; pK = 2.8532 21'c3i2H20 F'e(OI-I)2+ II+; pK = 4.7033 I+'rOFI2" HQO A G p ( l W + ) = -2.53 kcal./niolc29 A G p ( H 2 0 ) = -.56.W kcsal./niolcZQ AGP( Fe&) = - 177.1 kcrtl./inole29 AClp(F'e(OH)3)= - l(i(i.0 kcnl./niole29 concn. of I+(OH)3(aq.) = 2.9 X 1 0 - 7 rnolc/l.3a Fc3+
+ I-IzO
+
+
+ + +
v m o u s Fc( 111)-bearing species 8
0
IO
2
DH
Fig. 5.--Concentration in :In aqueous solution (a-E'e 07).
of ill
equilibrium
11
ith solid hernatitc>
(combiniLtion of w a r t ions 11 and 12) in the blilk solution ph:Lsc arc approximately equal. As The above thcrrnodyriainic information wis usotl t,o drtcr- pointed out previously, this concliision in turn suggwts that the adsorbabilities of the specics mine the constants for the reactions l/2E'ejOa(H) + 211' Jr FeOW+ + 1/21-120; PI\: = 2 . ~ Ik(OH)2+ and I?e02- arc approximatcly eqiial and '/ZFe2O3(8) + H + + I/?H20 Fr(OH)?+;p K = i.(i thus thc x.p.(*.,i.e.p., and ihc point of minimum coliibility fall at the same pI1 ~iiliie. Fe20,(s) + 4 H + I+2(OH)24+ + HzO; pK = 4.31 Since the solubility of a solid in a liquid phase Fe3+ + 3/?H20 '/?Fe2Os + 3 H + ; pK = -0.7% l/?F&(S) + 3 / 2 1 T d ) + OHFe(OH),--; pK = 10.6 and thc cstablishmcrii of un clertrical double layer of free charges at the solid-liquid interface The most striking fenture of Fig. 5 is thc fact. both arc d r p c u h i t on ihe same physico-chcmical that the i.e.p. of t,he complex Fe(111) solution agrccs process-traiisfer of vharged specks across the inso well with the mcasured Z . P . C . The exact terface, the roincidencc of the Z.P.C. and the point agreement suggested by t,he experiment,al data of minimurn solubility perhaps is not unexpected. must, however, be viewed as a fortuitous circum- The mnximum in the intcrfaci:il tension (or the stance. Scvert'hclcss, it may be conclitded that electroc'apil1:u.y maximiim for :i po1:irizable interthe available experimcntsl data suggest that tha face) lies at thc x.p.ca. t)cc:iiisc, diic to clectrostntic equilibrium constmts for rcaction 8 and its m : h g rc.pulsion, lew work is required to increase thc surfaw arca of :i rhargcd interface revrrsibly thim (29) W.M. Latinirr, "Oxidation Potentials," 2nd EWtion, PrcnticeIlall, Inc., New York, N. Y., 1952. that of an uncharged interface. This equality (30) For the positlve cornplexee the values for all confitants have of 2.p.c. nnd the point of maximum surfacc tension been extrapolated to zero NaCIO, concentrution. follows directly from thermodynamics. One might (31) H. Langweiler, W. Buser, and W.Feitknecht ( H e l a . Chim. Acto, reason that the point of minimum solubility of a 44, 805 (1961)) present quite conclusive evidence that the concentmlion of undiesociated Fe(0H)s in solution is lower by a hundredfold solid a h shoiild coincide with the 2.p.c. since the factor than that indicated in Fig. 5. Hoviever, this will h a w rclativcly rcversiblc work required to move a cationic potenlittle effect on the sharpness of the soluhility minimum. tial-determining specks from a positively charged 132) R. M. hlilburn, J . A m . Chem. Soc., 79, 537 (1957). surface (or an anion from :t ncgt1tij.e siirf:icae) into (3.7) A. B. Lamb and .4. G. Jacqiica, i h i d . , 60, 1216 (1838).
__z z__
!I72
G. A .
P ~ ~ R KAsD S
P.
the solution phase should be less than that required for the same transfer process a t a surface which has the same number of oppositely charged sites. Whether the equality of Z.P.C.and i.e.p. has the same fundamentul significance as that of the Z.P.C. and the electrocapillary maximum must, await further experimental vcrifiuition. A comparison between 1he mesmrcd z.p.c. and the i.e.p. of a solution of hydroxo-c~omplcxesmay be extended to all oxides. In Table I11 the avuilable data on a few oxides are collected. 2.p.c. measuremonts made on oxides for which no solubility data are available and cstimatcs of the z.p.c for some oxides based only on solubility measurements and cornplcx stability constants also are included in this table. Table I11 shows the paucity of available information to make this comparison between the Z.P.C. of the oxide surface and the i.e.p. of the complex solution in equilibrium with the solid phase. Some of the natural anhydrous oxides and ignited prccipitates (e.g., A1203, Fe2O3) give Z.P.C. values, as determined by electrokinetic measurements, which do not agree too well with the minimum solubility observations and the z.p.c. of the crystalline precipitates. This may be due to lack of cquilibrium with the solvent due t o incomplete hydration of the surface, a possibility which was avoided in our experiments. The increasing p R of the Z.P.C. with decreasing charge of thc cation in going from the hlOI type oxide to the AI20 t.ype may be explained qualitatively by the decreasing polarizing power of the cations of lower positive charge. Thc more acidic in nature the surface hI0E-I group, the tighter the bonding betmeen hl. and 0-thus t,he acid pH values of z.p.c. shown by the hI02 type osides. Within this oxide group the vari:Aon of the pH of the Z.P.C.may be explained by P a i i l i n g ’ ~concept ~~ of electronegativity. A decreasing differcncc in electronegativity between the oxygen and A 1 atom should correlate with an increasing tendency toward acid dissociation of the hIOH group. According to this criterion the pH of the Z.P.C. for MOs oxides should increase in the order Si02 < SnO2 < Ti02 < ZrOz
r,. I)]:
Yo1. Gti
IZRVYN
TTI Z.P.C., I.E.P.,
TABLE
COMPARISON OF THE pH Solid
To, Si02 SnOt TiOz
A N D Mrxrhr[i>f SOLUBILITY FOX \TARIOUS OXIDES PI1 hlin. 1.e.p. Z.P.C. solubility (calcd.) Comment Ref. OF
35 0.4:1
ao
0.43
37, 38, 1 B
