2228
STANLEY BRUCEENSTEIX AYD L. M. MUNHERJEE
Vol. 66
that it is associated with a surface group. The absence of a deuterium isotope ahift indicates that it is not a hydrogen containing group. The frequency of the band is in a range which could be assigned to an A10 vibration, possibly an overtone or combination band, The disappearance upon the addition of an adsorbate suggests a shift t o lower frequency where it cannot be observed due to overlapping with the strong Si0 fundamental, The absence of the band in dehydrated KSA implies that it is associated with acidic sites on SA which can be poisoned by K. On the basis of these observations, this band is tentatively assigned t o a vibration of a surface A10 group, probably an overtone or combination band. If this interpretation were correct, it would be consistent with the suggestion that, the acidic surface AI atom changes co6rdination number from four to five upon interaction with an H20 adsorbate molecule, since
the A10 force constant would be smaller for the higher coordination number.23 The assignments of the bands in the region 40001250 cm.-l in the spectrum of dehydrated SA are reviewed in Table 11. Conclusions The foregoing discussion of the spectroscopic data has led to the following tentative conclusions: 1. In highly dehydrated samples, the surface hydroxyl groups are predominantly attached to silicon atoms, 2. The fixedly adsorbed water added at 150' is held on acidic surface sites which can be poisoned by K. 3, The fixedly adsorbed mater added at 150' retains its molecularity and is located on sites far enough removed from the surface hydroxyl groups that essentially no hydrogen-bonding to these groups occurs. TABLE I1 4. A weak band located at 1394 em.-' may be ASSIGNMENT OF WE BANDSIN THE DEHYDRATED SILICA- due to a vibration of the A10 linkage, possibly ALIXIXASPECTRUM an overtone or combination, in acidic surface Y groups. (om. -1) Assignment Acknowledgments.-The author is grateful t o 3745 OH stretch in surface SiOH groups Mr. T. R. Kantner for assistance with the experi1975 Si0 combination mental work, and to Dr. D. S. MacJver for several 1866 Si0 combination informative discussions. 1633 Si0 overtone 1394
Surface A10 overtone or combination (1)
(23) Reference 14, p, 45,
EQUILIBRIA I N ET:TIHYLENEDIA&IINE. 11. HYDROGEN ELECTRODE STUDIES OF SOME ACIDS -4ND SODIUM SALTS BY STANLEY BRUCKENBTEIN AND L.M, MUKHERJEE~ School of Chemistrv, Urbiverszly of Minne8ota, Minneapolie, Minn, RUEWE& May $1, 19UB
Hydrogen electrode studies of a series of pure acid solutions have yielded the dieeociation constants of four phenols relative t o hydrochloric acid. The behavior of these phenols indicates that the reaction HX X- = HXg-(KaxP-)occurs. Values of &m- found were 15 (phenol), 7 (thymol), 40 (0- henylphenol), and 46 (p-phenylphenol). Studies of the acid-Rodium salt mixtures permitted the determination of the reztive dissociation constants of these salts The pH of varioud 8OdlUm
+
salts has been found t o be independent of concentration and has Fielded another means of determining the relative dissociation constants for various acids. The results obtained for pKvx - pKacl are 5 40 (thymol), 4.35 (a-phenylphenol), 4,30 (phenol), 4.20 (p-phenylphenol), 2.05 (sodium thymolate), 2.10 (sodium o-phenylphenolate), 0.10 (phenylacetic acid), -0.35 (3-methyl-4phenylazophenol), -0.60 (hydrobromic acid), and - 1.4 (hvdriodic acid). The value of K x a x / K ~ xfor thymol is 2.1 X loaand 2.1 X 101for 0-phenylphenol. The difference between the negative logarithm of the nutoprotolysis constant of EDA and that of the sodium salt of EDA is 7,O.
-
Introduction A large number of compounds have been determined by titration as acids in ethylenediamine (EDA) as solvent using potentiometric methods to detect the equivalence point.3 However, only one quantitative potentiometric equilibrium study (1) This work was supported by the Office of Ordnance Research,
U. 5. Army. (2) From a thesis submitted by L. M. Mukherjee to the Graduate School of the University of Minnesota in partial fulfillment of the requirements for the degree of Doctor of Philosophy, August, 1961. (3) (a) M. L. Moss, J. H. Elliot, and R. T. Hall, Anal. Chem.. 20, 784 (1948); (h) M. K a t s and R. A. Glenn, i h i d . , 24, 1167 (1952); (0) V. Z. Deal and G . E. A. Wyld, z h d , 27, 47 (1955); (d) A. J. Martin, ibid., 29, 79 (1957); ( e ) H.Brockman and E. ,Meyer, ~'alurzuzseenachaflen, 4 0 , 2 4 2 (1953);(f) H. Brookman and 5,Meyer, Chem. Ber., 87,
81 (1964).
of an acid and its conjugate base in EDA has been reported in the l i t e r a t ~ r e despite ,~ the fact that
such &dies are necessary to assess the limitations and advantages of EDA as a solvent for acidbase titrations. In this study, Schaap and coworkers titrated hydrogen bromide with sodium ethanolamine in the presence and absence of excess sodium bromide and were able to explain their results satisfactorily in terms of the ion-pair dissociation constants of hydrobromic acid, sodium bromide, and sodium ethanolamine using 5 X as the autoprotolysis constant ( K s )of EDA. Earlier potentiometric5 and conductometric5 stud(4) W. B. Sohaap, R. E. Bayer, J. R. Siefker, J. L. Kim P. W. Brewster, and F. C. Schmidt, Rec. Chem. P r o p . . 2 2 , 197 (1961).
HYDROGEN ELECTRQDE STUDIESOF ACIDSAND SODIUM SALTS
Nov,, 1962
EHX =
Escm
2229
+ 0.0296 log KHX+
0.0246 log [CHx(l 4- KHX,-CHX)I (5a) C H X is the equilibrium concentration of undissociated HX. Three limiting cases of eq. 5a are Reference electrode
(S,C.fl.E, 2~5)
considered below. H X ( C H ~ ) Hz (1atm.), 1. K H X Very Small.-If KHXis very small, the and/or Pt (I) analytical concentration of HX, (CKX)~,differs negligibly from the equilibrium concentration, NaX (C,,x> CHX, and the former may be substituted for the in EDA (1) latter in eq. 5a. At low acid concentrations I >> KHX,-CHX, and 5%becomes
EHX=
+ 0.0296 log KHX(CHX)~(5b)
Escs~
and a plot of ERXus. log (C,,), has a slope of 0.0296. Eq. 5b also is applicable at all concentrations for those acids which have ho tendency to form H X 2 ions, such as hydrochloric acid, At high concentrations of acid, if KHXCHX>>1, eq. 5a becomes
EHx
Escs,
-t
0.059 16 log ( K ~ x K a x-),'"(CHX)~ ( 5 ~ ) A plot of EEXvs. log (CH& has a slope of 0.05916. This result corresponds to the dissociation 2HX H+ HXz- and is similar to that found previously6 with silver cyanide in EDA. I n the concentration range where KHX,-* (CHX)~, the slope is intermediate of the pldt of EHx vs. log between 0.0296 and 0.05916. Evaluation of KHXz-.--Frorn eq. Sa, 5b, and 5c, it follows that the intersection of the two limiting slopes occurs a t (CHX)~= ~ / K H x ~ -Behavior . of this sort has been observed with four phenols. A more rigorous procedure for evaluating &x,uses the method of least squares by transforming eq. 5a to eq. 5d 10Eax/0.0296
+
+
I{HX*=
a m l -/aHxnx -
(3b)
where H + is used to represent the solvated proton in EDA. At 2 5 O , the e.m.f. of cell I is given by
E
CHX
+
ICHX 1 0 ~ a c a ~ l 0 . 0 ~ 9 ~ 1 0 E ~ ~ ~ E ~ o ~KHXKHX~-CHX 02ee (5d)
Thus a plot of 10EAX/0002Qe/CHX us. CHXis a straight line. K H X *is- the ratio of the slope to the intercept for eq. 5d. where' 2. KHX Very Large.-If K H X becomes very large, B H + = (CHX)~ and a plot of E us. log (CH,), ESCSE 1 j 1 . j . &Ref. Electrode (4b) would approach a limiting slope of 0.05916 v.; and the standa,rd potential of the hydrogen elec- thus the observed e.m.f.'s would exceed those trode is izssumed t o be zero. Assuming that g 1 . j found with all other acids. This behavior has is constaiit and that the activity coefficients of all not been observed in EDA and would not be exunivalent, ions are the same, while that of HX is one, pected in a uolvent with a dielectric constant substitution of eq. 3a and 3b into the rule of electro- of 12.9, except in very dilute solution. True neutrality for pure acid solutions (4c) yields an ex- strong acid behavior is to be distinguished from pseudo strong acid behavior arising from HX2[H+] = [X-] IITX,-] formation on the basis of the magnitude of the ob(44 e.m.f. values. pression for a H t which may be substituted into eq. served 3. KHX Intermediate in Value, KHXt-Very 4a to yield eq. 5a a t 25". Small.-EDA is a very basic solvent and levels ( 5 ) 8. Brutkenstein and L. M. Mukherjee, J . P h y s . Chem., 64, 1601 acids with a (PKHX)H%O Q 5.00. One criterion (1960). used for leveling is the inability to distinguish a (6) (a) W.H. Bromley and W. B. Ludw, J . A m . Chem. Soc., 66, difference of acid strength on the basis of potentio107 (1944); (b) B. B. Hibbard with F. C. Schmidt, ibid., 11,225 (1955). (7) The natation used is analogous to thai, used previously.6 metric measurements. I n a solvent such as EDA,
E
=
+
RSCSE 0.05918 log U H +
I -
+
+
(4a)
STANLEY BRVCKEX'BTEIN K N R L. Lt. MCKHERJEE
2230
leveling corresponds to the virtually complete
conversion of the non-ionized acid to the ion-pair
[H+]
+ [wail = [S-1 +-
according to
[X-I
HS
+ IIX
HSHI-X-
(6a)
while the hydrogen ion concentration is determined by the dissociation reaction
HBH-kXwhere HS
Vol. 66
4=. HSH+ + X-
(6b)
+ [HX2-1
(9)
Using eq. 9 and assuming the activity coefficients of all univalent ions to be the same and those of the undissociated species to be unity, it falloffs from eq. 2a, Zb, Sa, and Sb that aHt =
(loa)
KNaXCNaX
EDA. Defining the equilibrium
+
KHXCHX constants for reactions 6a and 6b as KIHXand KdHX, respectively, it hae been whom8 that For any EDA solution of an acid Rrhich can be titrated successfully with baae, in the absence of X H =~ KaHXKiHX/(l R I H X ) (GO) excem KaX, Xa is negligibly small as compared to the second term in the numerator of eq. loa. The value of KdHXdepends upon steric factors but In the experiments reported below with thymol would not be expected t o vary widely for different and o-phenylphenol, K N ~ X>> KHX and 1 < acids, R , H X is estimated to be between i o - 6 arid therefore using the Bjerrum ion-pair relation#, Thus, KN~XCN~X/KHXCHX; for leveled acids, K L H X >> 1, KHX F ~ T K d H X l asld it is not a good approximation to assume that (CHX)~= CHX when KdHX = because appreciable dissociation will occur. If 1 >> (Cax. and the e,m,f, of cell I containing scid-salt mixtures KHX~-)! the slope of a, plot of E vs. log CHx will be is between the limiting slope of 0.0298 for a very weak electrolyte and 0.05916 for a Strong electroK %x ~ XE 0 o s ~ 0.0296 log1lyte. I n addition, EHx for leveled acids will be E H X , N = K,,x more positive than for non-leveled acids. Three acids, hydrochloric, acetic, and phenylacetic, i R,x,-Cnx 0.0298 log C'HX (114 were studied anticipating behavior of the type CT\'aX outlined above. The results obtained with these acids indicate the model used probably is an over- Viider the experimental conditions used, analytical simplification in EDA and every acid must be con- concentrations, Gt, may be used in place of equilibrium concentrations, and eq. l l a describes the sidered individually. For those acids whose behavior is described by behavior of mixtures of thymol and o-phenylphenol eq, 5d, ~ K H X P K H ~is) obtained from the dif- with their sodium rsalts, For the purpose of treatihg the experimental ference of the logarithm of the intercepts obtained data by the method of least squaiqes, it is convenient with the two acids HX and HX'. Equilibria in BX-NaX Solutions.--E,m.f, meas- to transform eq, I l a to l l b urements obtained with cell I using solutions containing NaX and HX (thymol and o-pheiiylphenol) have been interpreted in t e r m of reactions 2a and 2b, and the equilibria
+
+
[
NaX 2 S a 7
4- X-
+
]
(7a)
HS 1-r H + + 6Sa9 Sa+ 8-
Sb)
-+
(7c)
The corresponding equilibrium constants arz
Thus, a plot of (CNax/CZEx) l O E a x . ~ n x / Q * Q11s. ~Q~ CHX in a straight line. As in the case of eq. bd, KHXb-is the ratio of the slope to the intercept, while the ratio of the intercept (or slope) of eq. 5d -- t o the intercept (or slope) of eq. l l b yields K N~XIKHX. Equilibria in NaX Solutions.--In EDA solutions of pure sodium salts, the principal reaction between NaX and solvent is I - -
Sax
arid
+ HS 7-f NaS + HX
(12a)
The equilibrium constant for reaction 12a is the reciprocal of the formation constant, KfNaX,of the salt where The rule of electroneutrality in a mixture of an acid and its sodium salt is (8) I. M. Kolthoff and (1956).
S. Bruakenstein, J . A m . Chem.
Soc., 1 8 , 1
KHXKN~S - aNax ( 12b) KxaxRs aNasaHx The magnitude of K P X is such that the equilibrium KfNaX =
Nov., 1962
HYDROGEN ELECTRODE ~ T U D I E SOF ACIDSAND SODIUM ~ALTS
concentration of CHXis very small, and 1>> (KHTI-. CHx)in eq. lob, i.e., tho formation of HX2- species is negligikily small. Also, the dissociation of Nag and HX is repressed by the dissociation of XaX and CHX = C N ~ ~Thus . eq. 10b and 12b yield
and
This result, that the e.m.f. (or pH) of pure salt solutions is independent of the salt concentration, also was found in anhydrous acetic acid.g The difference in pK of two acids, HX and HX', whose sodium salts obey eq. 13b is given by
Equation 13c yields results in good agreement with those found using eq. 5d. Using the relationships developed above, it is possible to determine the pK values of acids and their salts relative to one reference acid. I n a subsequent paper, a potentiometric procedure for relating the relative equilibrium constants obtained with the hydrogen electrode t o those obtained with the silver electrode6 will be described, I n addition, absolute values of dissociation constants based upon a spectrophotometric method will be reported, The value of P K N -~ pKs ~ can be obtained by combining data obtained in pure acid solutions with those obtained in solutions of the sodium salt of the acid, L e . , from eq. 13b ESax PKS - P K N ~= B log [intercept cy. 5d] - ____
0.0296
Experimental Reagents.-Commercial 9876 EDA was shaken with Linde Molecular Sieves (70 g./l.) of type 5A and then with a mixture of 50 g. of calcium oxide and 15 g. of potassiun hydroxide per liter of the solvent, followed by subsequent distillation of the supernatant, liquid over a similar batch of molecular tsieve in a current of dry and COS-free nitrogen. Akltjhoughthere was no improvement in the ultraviolet spectra, a,fl compwed to the previously reported procedure,6 the water content of the fraction boiling a t 117.2" was about 0.015 M after this treatment. Ethylenediammonium Chloride .--EDA 2HC1 was prepared bv treating pure EDA with a slight excess of concentrated hyd.rochloric acid; the colorless crystals were repeatedly washed with hot absolute ethanol, air-dried, and, finally, dried in uucuo a t 50'. Titration against a standard silver nitrate indicated the product to be 99.09y> ' ~ ~1,z used in the data in Tables I xnd III using eq. 14, and found t o be 8.97 from thymol data and 7 00 from o-phenylobtaining eq. Ilb. The values of pKtbvmol- ~ K H for X a series of phenol data. Using Gchaap's value of pKs = 15.3,' sodium salts as calculated from the e.m.f. values yields pKNas = 8.35, On the basis of these data, of pure salt aolutions using eq. l a c (Table 111) the monosodium salt of EDA appearfi to be a very haye been mentioned earlier. The value of thikj dif- weak electrolyte, and, if it could be prepared, fcrence for o-phenylphenol is a.99, as compared would be a poor choice as a basic titrant Atto 1.08 (Table I) found from solutions of pure acids. tempts t o prepare this compoulld by reactioii of This agreement is within the experimental error. sodium metal, sodium amide, and sodium hydride with EDA have proved unsuccessful, the ultimate (20) N. van Loov and L. P.Hammett, J . Am. Chem. Soe., 81, 3872 products in all cases being black, tarry residues. (IRAR). 5
