March, 1962
ACIDIONIZATION CONSTANTS IN H20 AND D 2 0
explained as a manifestation of the “synergistic effect”I2 existing between c and T bonds. (12) L.E, Orgel, “An Introduction t o Transition Metal Chemistry,” John Wiley and Sons, New York, N. Y., 1960.
429
Acknowledgment.-We wish to thank the National Science Foundation and the Wisconsin Research Foundation for grants which supported this and related work.
R,ELATIVE HYDROGEN BONDING OF DEUTERIUM. 11. ACID IONIZATION CONSTANTS IN H20 AND DzO’ BY A. 0. MCDOUGALL AND F. A. LONG Department of Chemistry, Cornell University, Ithaca, N . Y . Received August $0, 1961
Studies have been made of the ionization constants of a variety of weak acids in the solvents HzO and DzO by glass electrode or b y conductometric procedures. Emphasis has been on acids which in either the acid or conjugate base form would be expected to form intramolecular hydrogen bonds. These data, when combined with data of other workers, offer support to the proposal that for weak acids in general ( ~ K D A ~ K H Aincreases ) with KEA. However, the data also support the suggeEtion that the reference line of such a correlation is different for acids o f different types. Several of the present results are consistent with earlier studies on maleic acid in that intramolecular hydrogen bonding, relative to bonding to the solvent, appears t o be weaker for deuterium than for hydrogen. However, there are enough exceptions to this rule to suggest that a true picture will not be obtained without more explicit consideration of the properties of particular solutes as well as of the solvent differences in the two cases.
Reasons for interest in the relative abilities of hydrogen and deuterium to form hydrogen bonds have been summarized in Part I of this series.2 Part I also reported measurements of the ionization constanta of acids, one of which involved an intramolecular hydrogen bond; it was found that the ratio of these constants in water and deuterium oxide, for any particular acid, varied according to whether an intramolecular hydrogen bond was concerned in the ionization. The comparison in such a case is of the competition between internal hydrogen bonding and hydrogen bonding with the solvent, and it is not a direct measure of the strength of the intramolecular hydrogen bond invol~ed. The present work extends this comparison to other ,acids, some of which are hydrogen-bonded, and considers the results on the basis of a general correlation for the dissociation of acids in water and deuterium oxide. Such a relationship first was postulated by Rule and LaMer,3who found for the small. number of acids at that time investigated that log (KHAIKDA) was proportional to log KHA, KHAand KDAbeing the ionization constants of H- and D- acids in water and deuterium oxide, respectively. This relation has been support’ed by later moinkers4 but HEgfeldt and Bigeleisen5 have suggested recently that the type of acid (e.g., phenol, carboxylic acid, etc.) may change the constant of proportionality, i e . , the slope of the line, and this proposal is supported by the present work. Most of the conventional methods for determination of ionization constants have been applied to studies in deuterium oxide; Dahlgren and Long,2 for example, used e.m.f. measurements with a quinhydrone electrode. A much simpler method is suggesLed by the work of Glasoe and Long6 and (1) Work supported by a grant from the Atomic Energy Commission. (2) G. Dahlgren, Jr., and F. A. Long, J. Am. Chem. Soc., 82, 1303 (1960). (3) C. K.Rule and V. K. LaMer, ibid., 60, 1974 (1938). (4) P. Ballinger a n d F. A. Long, ibid., 82, 795 (1960). (5) E. Hiigfeldt and J. Bigeleisen, i b i d . , 84, 15 (1960). (6:) P. K.IGlasoe and F. A. Long, J . Phys. Chem., 64,188 (1960).
others’vs which has established that satisfactory measurements with deuterium oxide solutions can be made with a glass electrode. The determination of hydrogen ion concentration with a glass electrode is by no means as accurate as that with a hydrogen or quinhydrone electrode but it was hoped that satisfactory pK differences could be obtained readily for a series of acids. Studies now have been made for a number of carboxylic acids and phenols, several of which involve intramolecular hydrogen-bonding. In addition some data have been obtained for certain other acids (not hydrogen-bonded species) by conductance measurements, which method of course is considerably more accurate than the glass electrode procedure. Experimental Materials.-Reagent grade inorganic chemicals were used throughout. Deuterium oxide was supplied by the Liquid Carbonic Company; it contained at least 99.5% DzO. Salicylic acid (m.p. 159-160”), o-nitrophenol (m. 46”), p-nitrophenol (m.p. 112-1 14”), 2,4-dinitrophenof (m.p. lll’), 2,6-dinitrophenol (m.p. 62-63’), and y-resorcylic acid (2,6-dihydroxybenzoic acid) (m.p. 167”), were purified by recrystallization from water. Glycolic, oxalic, iodic, and chloroacetic acids were used without further purification except that the oxalic acid first was dehydrated over sulfuric acid. Reagent grade phosphoric acid (containing not less than 85% H,PO,) was employed. Sodium hydroxide solutions were British Drug Houses “Concentrated Volumetric Solutions” diluted appropriately. A stock solution of sodium deuteroxide in deuterium oxide was made up by allowing metallic sodium in toluene to react with boiled-out deuterium oxide in a separatory funnel. This stock solution was diluted according to requirements. All solutions were made up with boiled-out distilled water or deuterium oxide, and were estimated either by titration against standard base or in the case of certain of the phenols by a bromination procedure or by spectrophotometry. pH Measurements.-The potential was measured with either a Beckman Model G pH meter or (in some of the later experiments) a Cambridge Instrument Company Research Model Electron-ray pH meter. The electrode assembly, consisting of a Beckman No. 30167 glass electrode and a
.
(7) R. Lumry, E.L. Smith, and R. R. Glante, J. Am. Chem. Soc., 18, 4330 (1951). (8) K. Mikkelson and 8. 0. Nielsen, J. Phys. Chem., 64,632 (1960).
A. C. MCDOUGALL A N D F.8.LONG
430
Vol. 66
TABLIF. I P K H AAXD ApK VALUESFROM GLASSELECTRODE MEASUREMENTS, 25” ~KEIA
-pResorcylic acid Salicylic acid m-Nitrobenzoic acid Glycolic acid 2,6-Dinitrophenol 2,4-Dinitrophenol Oxalic acid ( K z ) o-Nitrophenol p-Nitrophenol
1.87 f 0.04 2.94 f .01 3.62 f .04 3.90 f .02 3.92 2c .03 4.12 f .02 4.30 7.19 f .02 7.26 f .03
~ K D A
2.33 f 0.04 3.69 f .03 4.12 f ,132 4.37 f .01 4 42 f .03 4.82 =I=.02 4.79 7.94 f .02 7.74 z t .01
KO. 39168 calomel reference electrode, was suspended in a test-tube containing about 5 ml. of the sample solution, which was immersed in an oil-bath maintained at 25.0 f 0.05”. The meter was standardized with either potassium acid phthalate solution ( p H 4.005) or a phosphate buffer solution (pH 6.860), according to the range required. The actual measurements were made on appropriate buffer solutions of the acid under investigation; these solutions normally were made up by partial neutralization of the acid solution with sodium hydroxide or deuteroxide and sodium chloride solution was added to make a series of varying ionic strengths-usually between 0.01 and 0.1 M. However, in a few cases the solutions were made up directly from the acid and its conjugate base. The thermodynamic dissociation constant was evaluated from the equation where pH is defined as -log [ CH+yt], CHAand CA- are the added molar concentrations of free acid and anion, respectively, and yt is the mean activity coefficient of the dissociated acid. I n the correction term, C H + was taken as antilog ( - p H ) . Log y* mas established from the approximate Debye-Huckel formula -logy*
0.509di 1 dl
+
= ___
I being the ionic strength of the solution.
The second ionization constant of oxalic acid was established from solutions containing measured amounts of oxalic acid and sodium oxalate (i.e., the disodium salt and the pK calculated from the expression
where CaZoxand Cox= are the added molar concentrations of oxalic acid and the disodium salt, respectively, and y* is the mean 1 :1 activity coefficient. No correction proved to be necessary to take account of the first dissociation since a t the pH involved there vas essentially no un-ionized oxalic acid present. The procedure was substantially the same for the deuterium oxide solutions except that the p H meter reading was converted t o “pD” by means of the equation pD = p H 0.40
+
(9) C . T. Abichandari and S. X. K. Jatkar, J . Indian Inst. Sci., A23 77 (1941). (10) S . Korman and V. K. Lahler, J . Am. Chem. Soc., 58, 1396 (1936). (11) J. F. J. Dippy and R. H. Lewis, J . Chern. Soc., 1008 (1937). (12) L. F. Nims, J . Am. Chem. Soc., 58, 987 (1936). (13) L. J . Minnick and M. Kilpatriok, J . Phys. Chem., 43, 259 (1939). (14) D. C. Martin and J. A. V. Butler, J . Chem. Soc., 1366 (1939). (15) R. G. Bates and G. Schwarzenbach, H e h . Chirn. Acta, 37, 1069 (1954). (16) G. D. Pinching and R. G. Bates, J . Research h’atl. Bur. Standards, 40, 405 (1948). (17) H. S. Harned and L. D. rallon, J . A m . Chem. Soc., 61, 341 (1939). , 542 (18) H. N. Parton and R. C. Gibbons, Trans. Faraday S O C .35, (1939). (19.) A. I. Biggs, zbid., 50, 800 (1954).
ApK
0.46
.75 .50 .47 .50 .70 .49 .75 .48
Previous results
~ K H =A 1.22 a t 30’0 ~ K H =A 2.97; ApK = 0.61’0 ~ K H =A 3.49311 ~ K H =A 3.82112;3.8015 ~ K H =A 3.58; ApK = 0.45 a t 18014 ~ K H =A 4.1116; ~ K H=A4.02; ApK = O.52at 18”14 ~ K H =A 4.266l6; 4.29‘7; 4.301* ~ K H =A 7.22lQ;~ K H=A7.25; ApK = 0.57 at 18”14 ~ K H =A 7.15l5; ~ K H =A 7.24; ApK = 0.56 at 1S0l4
established by Glasoe and Long.6 All studies in deuterium oxide were with solutions containing 99-99.574 DzO and the results were not corrected to 100% D 2 0 . Some idea of the accuracy of the procedures can be gained from the following typical results for glycolic acid. Separate measurements were made on six different buffer solutions with acid and anion concentrations each about 0.02 M , and with values of the ionic strength varying from 0.03 to 0.12. The calculated pK values ranged from 3.75 to 3.82. After activity coefficient corrections, the calculated ~ K H values A varied from 3.82 to 3.95, with an average value of 3.90 f 0.02 (standard deviation). A very similar study in deuterium oxide led to p K n ~= 4.37 & 0.01. Hence ApK = ~ K D A ~ K H =A 0.47 f 0.02. Table I summarizes the ~ K Hand A ApK values which were measured by the glass electrode technique and compares the ~ K H A values with those of other workers. Conductance Measurements.-The conductivity cell was of a type similar t o that described by Baker and LaMerzn; the electrodes were lightly platinized and the cell had a capacity of 10 ml. The cell constant was 1.846 cm.3, determined with carefully prepared solutions of potassium chloride. The electrical apparatus comprised a 1-kc. oscillator, a Campbell-Shackleton ratio bridge, a Leeds and Sorthrup precision a.c resistance box, an amplifier and earphones. Measurements of the conductance of the acid solutions were made at a t least two concentrations chosen t o provide a compromise between a lorn ionic concentration (to minimize the activity coefficient correction) and not too high a degree of dissociation (since the method is considerably less accurate for higher CY values). In the case of phosphoric acid and, more particularly, oxalic acid i t was necessary to use a concentration sufficiently high to minimize the effect of the second dissociation. The calculation of KHAwas carried out essentially by the method of MacInnes and Shedlovsky.21 Values of An, the equivalent conductivity OB the acid at infinite dilution, were obtained from the literature and Ai, the equivalent conductivity at the ionic concentration concerned, calculated from the Qnsager equation by a successive approximation procedure ,-
where B , BI and Bz are known constants, c is the stoichiometric acid concentration (in moles per liter), CY is the degree of dissociation and d is the ionic size parameter. The values of K, calculated from the relation C Y “ / ( 1 - a),are not unduly sensitive to choice of AO or of d (which was taken as 3.0 A. unless otherwise mentionedh. The activity coefficient term necessary to convert the “concentration” ionization constant to the “thermodynamic” value was evaluated from the modified Debye-Hiickel equation -Adcuc logy, = 1 B d d / c A being a constant and the other symbols having their usual meanings. For the concentrations involved in these studies the difference between j , the rational activity coefficient (for which this expression is derived) and y, the
+-
(20) W. N. Baker and V. K. LaMer, J . Chem. Phys., a, 406 (1935). (21) D. A . IlacInnes and T. Shedlovsky, J . A m . Chem. Soc., 54, 1429 (1932).
ACIDIONIZATIOX CONSTAKTS IN HzO AND DzO
March, 1962
431
TABLE I1 IONIZATION CONSTANTS FROM CONDUCTANCE MEASUREMEXTS, 25"
a
d
Ao(H20)
L(Dz0)
PKHA
~KDA
Iodic acid Oxalic acid (K1)
3.0 4.0
390.7" 390.0b
284.2 283.1
0.848 1.270
1.151 1.666
0.303 .396
Phosphalric acid (K1) Chloroacetic acid
3.0 3.0
382.8' 389.6d
277.6 282.8
2.128 2.851
2.362 3.339
.234 .488
Ref. 32.
b
Ref. 27.
Ref. 33.
ApK
Previous results
~ ~ ~ ~
K K K K
H =A 0.772*6 H= A 1.62; ApK = 0.0226 H= A 1.271'1 H =A 2.1328; 2.14P9 KEA = 2.85431; ~ K H =A 2.76; ApK = 0.4480
Ref. 31.
molar activity coefficient, is negligible. Moreover any errors arising from the use of this approximate form of the Debye-Huckel equation probably will not be significant for a situation where the ratio of ionization constants is of greater interest than an accurate value of K H A . For the deuterium oxide solutions,.values of A0 were in most, cases not available but the relation22
TABLE I11 IONIZATION CONSTANTS IN WATERAND DEUTERIUM OXIDE, ApK = ~ K D A~KHA Symbols for methods used: C, conductance; Hz, e.m.f.-hydrogen/deuterium electrode: PH, glass electrode; QH, qumhydrone electrode; S, solubility of salts and transfer data; SP, spectrophotometry. Acid
was assumed to be applicable. Since as mentioned above the final value of K is not very sensitive to choice of Ao, errors from the use of this equation probably are not serious. Values of AO(Cl-) and Ao(D+) in deuterium oxide were obtained from the work of Chittum and LaMerZ3and of Longsworth and M a c I n n e ~ . ' ~The constants A , B , B1 and BZ in the Debye-Hiickel and Onsager equations were adjusted to compensate for the changed viscosity and dielectric constant of deuterium oxide. A summary of t,he results of the conductivity experiments is given in Table 11.
Discussion Table 111 lists the data for ionization of acids in water and deuterium oxide obtained up to the present time, excluding only some very early figureP; indication is given of the experimental methods used. The temperature is 2.5' unless otherwise indicated. (22) J. P. Chittum and V. K. LaMer, J . Am. Chem. Soc., 69, 2455 (1937). (23) V. K. LaMer and J. P. Chittum, ibid., 63, 1642 (1936). (24) L. C. Longsworth and D. A. MacInnes, ibid., 59, 1666 (1937). (25) R. M. Fuoss and C. A. Kraus, ibid., 56, 476 (1933). (26) J. C. IiIornel and J. A. V. Butler, J . Chem. Soe., 1361 (1936). (27) L. S. Darken, J. A m . Chsm. SOC.,63, 1007 (1941). (28) L. F. :Nims, ibid.,66, 1110 (1934). (29:)R. G. Bates, G. L. Siegel, and S. F. Acree, J . Research Natl. Bur., Standards, 30, 129 (1943). (30) G. N. Lewis and P. W. Schutz, J . Am. Chem. Soc., 56, 1913 (1934). (31:l B. Saxton and T. W. Langer, ibid., 55, 3638 (1933). (32) E. Bock, Can. J . Chem., 37,3888 (1959). (33) C. M. Mason and J. E. Culvern, J . Am. Chem. SOC.,71, 2387 (1949). (34) The recent results of Hogfeldt and Bigeleisen,s and Hyman, Kaganove, and Katzas from indioator experiments have not been included partly because estimation of their probable accuracy is dilficult and partly because the acids examined are largely of different charge types (e.&, amino acids) from those we have considered. As a different point, we should note that some additional studies of relative ionimtion constants by glass electrode procedures have been reported recen.tly.sa (35) H. H. Hyman, -4. Kaganove, and J. J. Katz, J . Phys. Chem., 64, 1653 (1960). (36) N. C. Li, P. Tang, and R. Mathur, Presented a t the 139th National Meeting of the American Chemical Society, St. Louis, Spril, 1961.
(37) T. Riley and F. il. Long, unpublished work. (38) C. Drucker, Trans. Faraday Soe., 33, 660 (1937). (39) G. Schwarzenbach, A. Epprecht, and H. Erlenmeyer, Hela. Chim. Acta, 19, 1292 (1936). (3%) D. Bunn, F. 8. Dainton, and S. Duckworth, Trans. Faraday Soc., 57, 1131 (1961). (40) J. Curry and Z. Z. Hugus, Jr., J . Am. Chem. SOC.,66, 653 (1944). (41) T. Riley and F. A. Long, t o be published.
Picric Iodic Oxalic ( K l ) Sulfuric ( K z ) 7-Resorcylic (K,) Maleic (K1) Phosphoric (K1) m-Nitroaniline Chloroacetic Salicylic ( K 1 ) Ethyl hydrogen maleate Fumaric Ethyl hydrogen fumarate m-Nitrobenzoic Formic
Method
~ K H A ApK
Ref.
SP C C C PH QH C PH C
0.38 0.44 0.848 .303 1.270 .396 1.90 .30 1.87 .46 1.91 .62 2.128 ,234 .48 2.48 2.851 .488 2.94 .75
37 This work This work 38 This work 2 This work 6 This work This work
PH QH
3.08 3.10
.46 .46
3.40 3.62 H.r 3.75 PH 3.75 Glycolic PH 3.90 2,6-Dinitrophenol PH 3.92 2,4-Dinitrophenol PH 4.12 Benzoic QH 4.21 Oxalic (Kz) PH 4.30 Aniline PH 4.55 Fumaric (Kp) QH 4.60 PH 4.68 Hydrazoic (20') Acetic PH 4.73 4.74 QH C 4.74 Maleic ( K z ) QH 6.33 Carbonic ( K 1 ) Hz 6.35 o-Nitrophenol PH 7.19 Phosphoric ( K z ) QH 7.19 p-Nitrophenol PH 7.26 Ammonia Hz 9.26 Methyl acetylacetone PH 9.35 2-Acetylcyclohexanone P H 9.47 Trimethylamine Hz 9.90 Glycine (K", +) Hz 9.90 Carbonic ( K z ) Hz 10.25 PH 10.33 Hydroquinone QH 10.58 2,2,2-TriRuoroethano1 C 12.37 2-Chloroethanol C 14.31 Water S 15.72 a Recalculated by Glasoe and Long.
.45 .50 .40 .45 .47 .50 .70 .50 .49 .58 .42 .33 .52 .52 .52 .38 .43 .75 .56 .48 .49 .40 .5l .59
QH QH PH
.53 .64"
.63 .62 .65 .70 .84
2 2 2 This work 39 6 This work This work This work 3 This work 6 2 39a 6 10 23 2 40 This work 3 This work 39 41 41 39 39 40 6 3 4
4 42
(42) R. W. Kingerly and V. K. LaMer, J . Am. Chem. Soc., 63,3256 (1941).
432
-
A. C. NCDOUGALL AND F. A. LONG
Vol. 66
than the “hydrogen €1-bond” in similar circumstances. Maleic acid is of course a favorable case 2.6-DI~TROPHENOL 2-CH LOROETHANOL for this kind of study since the work of Westheimer 070 and B e i i f e ~ *already ~ has established that in A/A 2.2-TRIFLUOROaqueous solutions the bimaleate ion exists primarily in the internally hydrogen bonded form, Evidence for the work of other investigators for hydrogen bonding of some of the acids and anions of the present study is considered below. A serious difficulty is that much of the evidence for intramolecular bonding in these compounds comes “2, from studies with crystals or with solutions in *CARBOXYLIC ACIDS non-aqueous solvents of low polarity. Since water 030 I O D ~ s\LFUR~C(K?) OlNORGANlC OXYACIDS as solvent will compete strongly for both donor and AALCOHOLS 8 PHENOLS PHOOSPHORIC(K~) V A M M O N I U M IONS acceptor sites for hydrogen bonding, one may expect intramolecular hydrogen bonding to be con0200 WHA. 8 10 12 14 16 siderably less in water than in such solvents. UnFig. 1.-Plot of ApK against pKHA:1, m-nitroaniline; 2, fortunately for present purposes the central point chloroacetic; 3, ethyl hydrogen maleate; 4, fumaric ( K ! ) ; 5, ethyl hydrogen fumarate; 6 , m-nitrobenzoic; 7, glycolic; is whether internal hydrogen bonding exists in aqueous solution. 8, 2,6-dinitrophenol; 9, benzoic; 10, oxalic ( K z ) . With the two ,&diketone enols, we really do not When ApK is plotted against the ~ K H for A all know the situation for solutions in water. In these results (Fig. l), a graph in the form of an non-aqueous media, evidence from both infrared approximate stright line is obtained. The line of and ultraviolet studies points definitely to internal Fig. 1 is for the equation hydrogen So also does the high enol A p K = 0.41 + 0 . 0 2 p K ~ ~ content found a t equilibrium for these ketones in solvents.47 The enol content is usually as suggested by Bell.43 If the results are grouped inert sharply lower in aqueous solutions but still large according to the structural type of the acid con- enough to suggest some internal hydrogen bonding cerned, then straight lines, similar but differing o i the enols. size of the pK values for these slightly in slope, are obtained for carboxylic acids, enols in water The is also consistent with considerable inorganic acids, and phenols, although the paucity intramolecular hydrogen but this arguof results in some of these categories makes the ment is rather uncertain. bonding The relatively ApK lines uncertain. If the figures of Li, Tang, and values of Fig. 1 are, of course, as expected low from the Mathur36 for the second and third dissociation proposal that the internal bonding persists into constants of citric and tricarballylic acids are in- water since from the maleic acid studies the excluded in the carboxylic acid plot, and if these are pectation is that the ApK for an H-bonded acid assumed to be “normal” acids, then the slope of the mill be abnormally best line is somewhat greater (0.05) than if these Salicylic acid and y-resorcylic acid are both results are ignored (0.03) or than if the phenols are plotted on the same diagram (0.02). However, carboxylic acids with o-hydroxy groups and might whichever plot is used, some points lie a consider- be expected to behave similarly. Branch and able distance on either side of the line drawn. Yabroff49 showed by comparison of the dissociaSince the uncertainty in many of the ApK measure- tion constants of hydroxy- and methoxybenzoic ments is comparatively large, the significance of a acids that the anion of salicylic acid undoubtedly particular point falling away from the drawn line is is hydrogen bonded in aqueous solution. More recently Scheraga and Hermansso have studied small unless the discrepancy is striking. I n a number of cases large discrepancies may be salicylic acid in both HzO and DzO and interpret rationalized by assumption of internal hydrogen their results in terms of extensive intramolecular bonding in the acidic or anionic group. The work bonding for the anion. Thus the high position of of Dahlgren and Long2 suggests that intramolecular this acid in Fig. 1 (ApK = 0.75) is entirely consisthydrogen-bonding will affect the relation between ent with the data for maleic acid and suggests a PKHAand ApK in the sense that if the anion of an similar explanation. For resorcylic acid on the acid (e.g., maleic acid) contains an internal hydro(44) I?. H. Westheimer and 0. T. Benfey, J . Ana. Chem. Soc., 78, gen bond then the ApK is predicted to fall above 5309 (1956). R. 9. Rasmussen, D. D. Tunnialiff, and R. R. Brattain, ibdd.? the “normal,” Le., above the value which would 71,(45) 1068 (1949). be expected from consideration of other non(46) S. Bratoz, D.Hadzi, and G. Rossmy, Trans. Faraday Soc., Sa, hydrogen-bonded acids. Conversely, if the acid 464 (1956). (47) For a discussion, see G. 8. Hamrnond in “Steric Effects in form contains an intramolecular hydrogen bond Organic Chemistry,” edited by M. 8. Newman, 1856, John Wiley and (e.g., bimaleate ion) the ApK should be lower than Sons, New York, N. Y., Chap. 9. expected. This implies that the H-anion in its (48) One further acid which lies well below the line on the A p K hydrogen bonded form is relatively more stable in ~ R I diagram ~ A is phosphorio (KI).No reason can be advanced for water than is the D-anion in DtO. Consequently thin, although i t is true that comparatively few inorganio acids have the strength of the “deuterium H-bond” relative to been studied and those that have may not be sufficientlyreprenentato provide a good base line. intermolecular bonding with the solvent is weaker tive(49) G. E. K. Branch and D. L. Yabroff, J. Am. Chem. SOC.,66, SA~ICYCLIC
P-NITROPHENOL
FThANnl
on* LLI \r,
(43) R. P. Bell, “The Proton in Chemistry,” Cornell University Press, lthaca, N. Y., 1958,p. 188.
2568 (1934). (50) H. A. Scheraga and J. Hermans, unpublished works.
March, 11962
LATTICEENERGY AND STABILITY OF CHROMIUM MONOHALIDES
433
other hand the ApK value is very much lower and stabilized by hydrogen bonding, the only effect of such bonding can be to lower the acid strength, is jn fact close to the line of Fig. 1. ’The infrared spectrum of the crystalline potas- and thus it seems quite possible that no intrasium salt, indicates the existence of two symmetrical molecular hydrogen bonding is present in aqueous internal hydrogen bridges between the carboxylate solutions of nitrophenols. Astle and M ~ C o n n e l l ~ ~ anion and the hydroxyl groups,5l and moreover a found that the polarographic reduction of o-nitrocomparison of the ionization constants of dihy- phenol in aqueous solution was much slower than droxybenzoic acids5’ shows the 2,6-dihydroxy iso- that of the m- and p-isomers and was dependent on mer to he by far the most acidic, which strongly the p H of the solution and they suggested that suggests hydrogen bonding of the anion in aqueous hydrogen-bonding of the acidic form was responsisolution. I n view of this evidence the relatively ble for this, but steric hindrance is a fairly obvious “normal” ApK value for resorcylic acid implies alternative explanation. It will be seen from Fig. 1 tbst the strengths of the hydrogen and deuterium that while o-nitrophenol and 2,4-dinitrophenol have bonds are about equal. It is not clear, however, relatively high ApK values, those for 2,Bdinitrowhy there should be any difference between these phenol and picric acid are approximately as would two acids which both have a hydroxyl substituent have been predicted for normal acids of comparable strength. Thus while there is little support for ortho to the carboxylic group. The situation for the nitrophenols is also uncer- hydrogen bonding in the acid forms of these phenols tain. Although the o-substituted nitrophenols in aqueous solution (which would tend to lower (Le., o-nitrophenol, 2,4-dinitrophenol, 2,6-dinitro- the ApK) the problem of explaining the observed phenol, and picric acid) have infrared spectra in ApK values remains. There are clearly enough uncertainties in the pK inert solvents which suggest that the undissociated the existence data for these hydrogen bonded systems as to form contains a hydrogen of such bonding in water is much less definite. preclude general conclusions about the relative When io.nization constants of various nitrophenols H-bonding of hydrogen and deuterium. Three of are plotted against their Ilammett a-values (ob- tLhesystems do agree with the previous conclusion tained from figures quoted by Taft5’ and Jaff6558) from the maleic acid work that under the particular and a line of slope p = -2.11357 drawn, the 0- competitive situation of H20 and DzO as solvents substituted compounds fall. below this line, which an internal D-bond is weaker than an H-bond. indicates that they are stronger acids than would However, the results for resorcylic acid and the have been predicted on the basis of the Hammett nitrophenols point t o the existence of complicating equation Since the phenolate anion cannot be factors and suggest that detailed consideration must be given both to the particular species con(51) H. Musso, Chem. Be?., 88, 1917 (1955). (52) W. Baker, Nature, 137,236 (1936). cerned and to changes in solvent effects. Much (53) 0. R. Wulf and U. Liddel, J . Am. Chem. Soc., 67, 1464 (1935). the same sort of qualified statement applies to the (54) G. E. Hilbert, 0. R. Wulf, S. B. Hendrioks, and TJ. Liddel, other trends in the values of ApK. There is some Nature, 136, 147 (1938). support for the proposal that ApK is an increasing (55) R. J. Franoel, J . Am. Chem. SOC.,74, 1265 (1952). function of pK but here also there is so much (56) A. hi. Buswell, V. Deitz, and W. N. Rodebush, J . Chem. Phys., 6, 501 (1937). evidence for specific effects as to cast doubt on al(57) R. Vir. Taft, Jr., in “Steric Effects in Organic Chemistry,” most any generalization. edited by M. S. Newman, 1956, John Wiley and Sons, New York, N. Y., Chap. 13. (58) H. H. Jaffh, Chem. Revs.,53, 191 (1953).
(59) M. J. Astle and W. V. MoConnell, J . Am. Chem. Soc., 65, 35 (1943).
LATTICE ENERGY AND STABILITY OF CHROMIUM MONOHALIDES BY LEONIDAS PETRAKIS~ Department of Chemistry, University of California, Berkeley 4, California Received Auuust $1, 1861
T:he lattice energies of the crystalline chromium monohalides are calculated assuming ionic bonding. These energies are combined in the Born-Haber cycle with empirical heats of formation of the ions and with measured and estimated absolute entropies to yield enthalpies and free energies of formation of the crystalline monohalides. The stability of these monohalides is further considered with respect to dissociation and disproportionation.
Introduction Consideration of the electronic configuration of the elements indicates the possible stability of the plus one (+1) oxidation state for certain transition elements in addition to the groups which ordinarily have monoVale’lt compou’lds. Due to the extra stability associated with half-filled d-shells Cr. MOand have a single n s electron in addition to ‘the inoomplete (% - Y)d shell outside the ap-
w
(1) National Research Council, Ottawa, Canada.
propriate inert gas configuration [(inert gas) (n- 1) d6% 4 4 6 61. Indeed, claims have been made to have prepared certain such monovalent compounds.’ m‘oreover, TToodbury and c o - ~ o r k e r s ~ recently reported electron spin measurements for (2) (a) M. J. Udy, “Chromium,” Vol. 2, Am. Chem. Soo. Monograph Series, Reinhold Publ. Corp.. New York, N. Y., 1956, p. 113; (b) J. W. Mellor, “A Comprehensive Treatise of Inorgania and Theoretical Chemistry," Vol. XI, Longmans, Green and Co., New York, N. Y., 1931, p. 366. (3) H. H. Woodbury and G. W. Ludwig, Phys. Rev., 117, 102 (1960).
