Sim D. Lessley and Ronald 0. Ragsdale University of Utah
Sail Lake Cny 8 4 1 12
I Trends in the Acidities of Some Binary Hydrides in Aqueous Solution I
For the nonmetallic binary hydrides ( I ) in aqueous solution, there is no general relation between acidity and electronegativity. Nevertheless, electronegativity has been a popular explanation for the increase in acidity of the acids in the sequence, H-CH3 < H-NH2 < H-OH < H-F since the H-X bond becomes more polar as the electronegativity of the central element increases; thus dissociation of the hydrogen ion is facilitated. Using this argument, i t could be argued (incorrectly) that hydrofluoric acid is the strongest of the hydrogen halides in aqueous solution. Consequently, a thermodynamic analysis is usually shown which illustrates the factors involved in the dissociation of the halogen acids in aqueous solution and to account for their relative acidities. A thermochemical treatment of a series of binary hydrides in aqueous solution in a horizontal row has not heen made previously due to limited data. Sufficient data are
now available so in this paper we will use a thermodynamic cycle to analyze the separate enthalpy t e r n s which contrihute to the dissociation of H-CH3, H-NH2, H-OH, and H-F. Another purpose for this analysis was to ascertain whether the electronegativity argument was a valid explanation for the second period binary hydrides or whether its use was dictated by its agreement with the experimental observations. Of course, one disadvantage with electronegativity is that i t is a term which has involved much controversy. Electronegativity was originally defined as an invariant property of atoms ( 2 ) .but it is now generally agreed that the value for an atom will depend upon its environment in a molecule (3-5).Finally, by using thermochemical cycles to analyze the ionization of H-X acids in both a period and a group we are presenting to the chemistry students a consistent picture. The processes which determine the ability of a binary hydride to dissociate can be given by the cycle
I
H-XaQ, where
m,xn = AH,.,,
AH.."
H+,,,, + X-,.,,,,
+ AH,,, + I + E.A. + AH,,,,,, + AHhydi-l (1)
Irr."d
Prri.d
minrrr
wrmr3
A plot of tha emhalpy changes which determine the acidities of the second period binary hydrides In aqueous solution.
The above symbols have the following meanings: m d e h y d is the enthalpy of dehydration of the neutral molecule H-X; AHdi.. is the bond-dissociation energy for breaking the first H-X bond (it is not the mean thermochemical bond energy); I is the ionization energy of the hydrogen atom; E. A. is the electron affinity of the X group; M h y d ( + ) is the enthalpy of hydration of the H+ ion; and M h y d ( - ) is the enthalpy of hydration of the X- ion. The sign conven-
Volume 53,Number 1, January 1976 1 19
Enthslpy D a t a for the D i s s o c i a t i o n of S a m e Binary Hydride. in Aqueous Solution I k c a l l m o l e l
H-X a c i d H-F H-CI H-Br H-1 H--DH H-NH2 H-CH3 H-SH
AHdehyd o f H-X
AHdiss
of H-X
E.A. (181, X
1 (181
A H h y d (-)
A H h y d (+)
(20)
AHrxn
11.7 (15) 4.20 5.00 5.9 10.5 (15) 8.2 (15) 3.18 (16) 4.6 (15)
-4.6,
a Reference (6) T h e r e e s t i m a t e s are S u P n o r t e d b y t n e f o l l o w i n g o b s e r v a t i o n s : he h e a t s o f s o l u t i o n for PH, H,S, a n d H,se are -3.46, a n d -2.5 k c a l / m o l e , r e r ~ e c t i v e l y(Ref. (15)). S i n c e H C I i s not m u c h l a r g e r t h a n PH, or H,S, i t s n e a t of s o l u t i o n i s e x p e c t e d t o b e s i m i l a r to t h a t of PH, a n d H,S, l i k e w i s e H B r w o u l d be s i m i l a r t o H,Se. b l n t e * P o l a t e d f r o m a g r a p h of t h e n u m b e r o f l o n e P a i r s o f e l e c t r o n s a r o u n d t h e c e p t r a l a t o m versus t h e h y d r a t i o n e n e r g i e s for t h e i r o e l e c t l O n i C i o n s . F-. O H - , a n d BH,-. T h e v a l u e of -80 k c a l w a r used for t h e h y d r a t i o n e n e r g y of BH,- a n d t h i s v a l u e w a r c a l c u l a t e d f r o m t h e K s p u r t i n r k i i a n d Born e q u a t i o n s (8).
tion that is used in this paper is the accepted thermodynamic one of endothermic changes being positive and exothermic changes negative. For comparison purposes the enthalpy data for the hydrogen halides are given in the table along with that of the HX molecules analyzed in this paper. I t can be seen that the small enthalpy change for the ionization of H F is the result of several factors. For a complete analysis of acid strengths the entropy term must also be considered (6, 7). In fact, H F is also a weak acid because of entropy effects. A simple explanation for relative acidities is not appropriate ( 8 ) as one needs to consider a complete thermochemical cycle. The AH,. values for the second period binary hydrides have been calculated using eqn. (1). From the table it can values decrease in the sequence be seen that the AHr, H-CH3 > H-NH2 > H-OH > H-F which also agrees with the relative acidities. Now i t is of interest to analyze the variation in the enthalpy changes for AHd;, AHhyd(-), E.A., and A H d e h y d . These are shown in the figure. The values for I and A H h y d ( + ) are not shown because they are constant. Here i t can be seen that the high enthalpy of dissociation of the H-F bond is counterhalanced by a high enthalpy of hydration of the fluoride ion. In fact the slopes for the lines representing AHdi.. and - A H h y d are very similar so that these two effects essentially cancel each other. The important factor which makes H F the strongest acid is the high electron affinity of the fluorine atom versus the low electron affinities of OH, NH2, and CH3. The trend in the electron affinities can be accounted for by the fluorine's higher effective nuclear charge.. Each of the s ~ e c i e sis isoelectronic and apparently an additional protbn(s) in the nucleus is more effective in increasing the attraction of the species for an electron than the existence of one, two, and three extranuclear protons in OH, NHp, and CH3, respectively. Even though electronegativities have been calculated from the average of the ionization potential and the electron affinity of the individual atoms, (9, 10) it suffers from the fact that it can not he determined by a direct experimental measurement. Also as already alluded to above, much controversy surrounds the use of this term. Consequently, it is best to use a thermochemical cycle to examine the factors which influence the acidities of the binary hydrides. From this analysis the main differences in acidities of the second period binary hydrides must be attributed to the electron affinities. Electron affinities are discussed in many elementary
20 / Journal of Chemical Education
texts as values which have questionable accuracy and something measured for only a few elements. Much progress has been made in the exoerimental measurement of electron affinities as noted by some recent re\,iewn (11-121 hlost o i the accuratelv known electnm affinities have heen determined optically by photodetachment or radiative captures (11). Appearance potential measurements are also used in evaluating electron affinities (13). The use of a continuously tunable laser and an ion cyclotron resonance spectrometer offers a new technique which may he very valuable in obtaining accurate values ( 1 4 ) . It is also of interest to see if the explanation for the acidity of the hydrogen halides is applicable to Group VI. Because of the lack of data, the comparison needs to be restricted to Hz0 and H2S. These results are given in the table. There is a reasonable parallel between the pair, Hz0 and H2S and the pair, H F and HCI. In both cases large A H d i s s values (strong H-X bonds) make H F and Hz0 weaker acids than HCI and H2S. The larger electron affinity of SH versus OH is also a significant factor in the increased acidity of HpS as compared to H20. A final comparison to consider is that for the third row binary hydrides HC1 and H2S. Again there is a parallel with the second row hydrides. The high electron affinity is an important factor in contributing to the acidity of HCI as compared to H2S. Lierature Clted (11 Bell, R. P.. "The Proton in Chemistry," 2nd Ed.. Cornell University Press. Ithaea. New York. 1973. 121 Pauline. L.. "The Nature of the Chemical Rond." 3rd Ed.. Carnell Universify Pmrr. New Ynrk. ,960. lthara. ~~~, (91 Huheey. J.. J. Phyr. Chrm.. 69,3284 119651. 141 Wa1sh.A. D..Discusainnr FmraduvSoc.. 2.18119471. (51 Sandcnon.d.T., J. CHEM. ED~C.,31; ~'(19451. (61 McCoubrey, J. C., Trans Forodoy Sot.. 51,743 119551. (71 Myw8.R. T.. J. CHEM. EDUC.53.17 (19761. (81 Ilasenl, W. E., "lnorganie Enerpties? Penguin Books. Baltiml,re. Md.. l 9 7 L (9) Mu1liken.R.S.. J. Chrm. Phus.. 2.782 119341. (LO1 Mulliken. R.S. J . Chem Phys.. 3.573 (19351. i l l 1 Rerry,R. S.,Chem. Re". 69.533 (19691. 1121 Blaunsfcin. R. P..endChrirtophorou. L. G..Rudief. Re*. R s u . 3.89 119711. 1131 Dillsrd, J. G.. Chrm. Re"., 73.589 (18791. (141 Smith. K. C.. Mclvel. 1 . . R. T..Rrsuman. J. L a n d Wallare. R. W.. J. Chem. Ph."*..
=. *,"" " 7 c o , ,,.-, " " , , .,. "",
(15) Wsgman. D. D., o t 81.. "Sei~ctedValves of Chemical Thermodynamic Prnpertiel." Tech. Note 270. U S . Gout Printine Offke, 196%1971. 1161 Rutlec. J. A. V.. Trans. F o m d o u S a i . 33.229119371. 1171 Mortimer. C. T.. '"Reaction Heair Bond StrenFhr? Pergsmon Press. New
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