HYDRATION OF IONS IN
THE
GASPHASE
1483
Hydration of OH- and 0; in the Gas Phase. Comparative Solvation of OH- by Water and the Hydrogen Halides.
Effects of Acidity by M. Arshadi and P. Kebarle Chemistry Department, Uninersity of Alberta, Edmonton, Canada
(Received September 8, 1969)
+
The thermodynamic quantities AHn-l,n! and ASon-l.nfor the reactions OH-(H20)n-1 HzO = OH-(HaO), and the analogous reaction but involving 02- have been determined for n up to 5,(0H-) and n up to 3,(02-) with a mass spectrometer utilizing a high-pressure source made for the study of clustering equilibria. Comparison of the AH,-1,, values with those of 3'- which were determined previously shows that both ions have very similar hydration interactions. The - A H c , , ~of F- is slightly higher than that of OH-; however the -AH,.+ of OH- becomes slightly higher for larger n. The results show a specifically high AH0,l for both E'- and OH- which must be due to the partial covalent bonding in the monohydrates. The results indicate that the total enthalpies of hydration of F- and OH- should be very similar with -AH,,(OH-) probably higher. The changes of AH for the gas-phase solvation reactions OHHX = (0HHX)where X = OH, F, C1, Br, I, can be calculated from available experimental data. The exothermicity of the reactions increases in the order given above. The energy values can be correlated with the acidity of HX.
-
+
A. Hydration of OH- and O p - in the Gas Phase A recent mass spectrometric study1 of the gas-phase hydration reactions of the halide ions has been extended to the OH- and Oa- ions. The mass spectrometer which has a high-pressure ion source and utilizes 2000-eV electrons for the ionization has been described in ref 1. All techniques used were similar to those given in ref 1. The OH- and its hydrates OH-(H20)n were observed in water vapor containing traces of hydrogen peroxide and in pure water vapor. The hydrates of 02were observed in oxygen containing a known partia.1 pressure of water. Total pressures were in the 1-6 Torr range. The water vapor partial pressures in the determinations of O2-(HzO), varied from 0 , l to 0.8 Torr. Equilibrium constants K,-,,, for the reactions OH-(HzO).-i
+ H2O = OH-(H20).
which were constant with water pressure in the experimental range were obtained for both ions. Figure 1 shows a van't Hoff plot of the equilibrium constants for the OH- hydration reactions. The A H n - ~ , nand corresponding AGO and AS" data for both ions are shown in Table I. The A H O J value for OH- can be compared with the results of an electron detachment study by Golub and Steiner.p These authors obtained 2.95 f 0.15 eV as the required energy for the reaction
H30z---+ O H
+ H20 + e
Using 1.S eV for the electron affinity of OH- they calculated a value of 1.2 eV or 27.6 kcal/mol for the sum of the kinetic energy of the separating particles Table I : Experimental Thermodynamic Values for the Reactions Reaction n - 1,n
- AH*n-l,n, kca1/ mo1
OD-(DzO),,i 011
1,2 273 314 495 091 112
2,3
22.5 16.4 15.1 14.2 14.1
Oa-(HzO),-i 18.4 17.2 15.4
-
AU0n-1,n~m),*
kcal/mol
-
AS"n-l,n(loI),b
eu
+ DzO * OD-(DzOIna 16.9 10.7 7.73 5.45 4.22
19.1 19.3 24.8 29.5 33,2
+ HzO * Oz-(HzO), 12.6 9.71 7.02
20.1 25.1 28.2
a In order to avoid overlap of the mass of OH-(H20) (35) with the mass of C1- (35) which was present as a small impurity, OD- and DzO were used in the experiments. The thermodynamic values for the protonated species should be practically the same. * Standard state 1 atm.
and the binding energy between H302- relative to OH- and HzO. The present experimental value AHo,l = -22.5 kcal/mol would indicate that the kinetic energy contribution was 5 kcal. (1) M. Arshadi, R. Yamdagni, and P.Kebarle, J . Phz(s. Chew., 74, 1475, (1970). (2) S.Golub and B. Steiner, J . Chem. Phys., 49, 5191 (1968).
Volume 74, Number 7
April 8 , 1070
M. ARBHADIAND P. K E B A R L ~
1484 25
1
20
-0 0,
c
E
\
o
0
U
C
Y
Y
5
(3
15
I '
3
?
-1
I
-7
1.0
I
1.8
I
I
I
2.6
I
3.4
1031~~~
Figure 1. van't Hoff type plots of the equilibrium constants &-I,* for the gas-phase hydration of OH-. Numbers in figure correspond to n.
The enthalpies AH%-I,% of OH- and 02- are shown in Figure 2 in function of n. Included in the figure for comparison are the corresponding values for Fwhich is a small ion of the size of OH- and I- which is a large negative ion, The AH values of the halide ions come from ref 1. Comparing first OH- and Fwe notice that -AHo,l for F- is somewhat higher but that for higher n the -AH%-I,~of OH- become slightly bigger. Both ions show a very large dropoff between - AHo,1and AH1,2,which indicates specially strong bonding in the monohydrate complex. The strong bonding in the monohydrate complex in Fwas explained as partly due to covalent bonding. Similar reasons should apply also for the OH- ion. The resonance structures probably contribut'ing to the bonding in both cases are shown below
On the basis of the AHn-I,nvalues for F- and OHshown in Figure 2, one would have expected that the total single ion heats of hydration AHh(F-) and A H h (OH -) , corresponding to the enthalpy changes for the process: ion(gas phase) -+ ion(aqueous soluT h e Journal of Physical Chemistry
I
OJ
1,2
I
233
I
3,4
I
4!5
(n-l,nl
Figure 2. Plots of for hydration reactions in the gas phase. E'- and I- are added to the figure for comparison purpoBes. 0, F-; 0, OH-; 0 , 02-; A,I-.
tion), should be about the same, the -AHh(OH-) being perhaps larger. However, the accepted tabulations of the total heats of hydration3give -AHh(OH-) = 101 and -AHh(F-) = -113.3 kcal/mol; i.e., the -AHh of F- is given as being larger. An examination of the origin of the tabulated data shows that they are based on a calculation by Halliwell and Nyburg4 based on crystal lattice energies of the alkali hydroxides calculated by Waddingtoms However, Waddington's lattice energies lead t o the electron affinity of OH, EA(OH) = 65 kcal/mol which is equal to a value determined by Page.6 However, the value of Page is incorrect. The presently accepted EA(OH) = 42 kcal/moL7 Therefore the lat,tice energies of Waddington are probably too low by some 23 kcal, which would lead to a -AHh(OH-) = 124 kcal/mol, a value that appears too high. On basis of the present experimental results we think that the best estimate for the -AHh(OH-) is that it should be a few kilocalories higher than that of F-. The higher hydration interactions of OH- for high n need not be surprising if one assumes that this ion fits better in the total structure of the water cluster. (3) D. 1%. Rossinsky, Chem. Re%.,65, 467 (1955); J. E. Desnoyers in "Modern Aspects of Electrochemistry," Val. 5 , J. O'M. Bookris, Ed., Plenum Press, New York, N. Y., 1969, Chapter 1. (4) H. F.Halliwell and S. C. Nyburg, Trans. Faraday Soc., 59, 1126 (1963). (5) T. C. Waddington, Adv. Inorg. Chem. Radiochem., 1, 158 (1959). (6) F. M. Page, Discussions Faraday Soc., 19, 87 (1955). (7) L. M. Branscomb, Phys. Rev., 148, 11 (1966).
1485
HYDRATION OF IONS IN THE GASPHASE 100
I
I
I
I
I
I
evaluation of these enthalpies allows one to make
I
HI
some interesting comparisons between the acidity and the solvating ability of the hydrogen halides and water. In order to evaluate AHo,l(OH-, H X ) one needs the heat of formation of the reactants participating in (1). AHr(OH-) can be calculated from EA(0H) = 42 kcal/moP and &%(OH) = 2.3 kcal/mol.* The heats of formation of the hydrogen halides are well known.* The heat of formation of (0HHX)- can be evaluated from the enthalpies for reaction 2 OH2
I
3 20
1
I
I
340 360 3 80 D [ X - H ) - E A ( X ) + I p ( H ) Kcal/mole
I
I
400
Figure 3. Plot of -AHo,l (OH-,HX) corresponding to enthalpy change for reactioii OHHX -.t ( 0 H H X ) - us. D(XH) - EA(X) I,(H) which can be defined as a measure of the “gas-phase acidity” of HX.
+
+
The AHn-l,mof 0 2 - show a characteristically different change with n than the enthalpies of the OHand F- since the AHo,l and AH1,z of this ion are so close. This result is expected in view of the electronic structure of the 0%ion. According to the simple molecular orbital description, 02-should have its three outer electrons in the doubly degenerate orbital Zpn, which spreads the electronic charge to the two ends of the molecule. The 0 2 - hydration equilibria should be of interest in the atmospheric electricity and lower ionosphere research fields since the hydrated 0 2 - ion is certainly an important negative species in the troposphere and the lower ionosphere.
B. Correlation of Gas-Phase Acidity and Strength of Solvating Interactions of Compounds HR with the OH- Ion The AHo,1 enthalpy values for the hydration of the halide ions can be used together with available thermodynamic data for the evaluation of the enthalpies for the reaction OH-
+ H X -+-( 0 H H X ) -
(1) The AH1 will be called AHo,l(OH-,HX) since it corresponds to the AHO,~ of OH- solvated by HX. The
+ X-
= (0HHX)-
(2)
which corresponds to AH,,,(X-,H~O) and was determined previously.’ The heats of formation of Xwere evaluated from bond dissociation energies in ref 8 and the electron affinities of the halide atom^.^ Shown in Figure 3 are the AHo,l(OH-,HX) determined by the procedure outIined above. Plotted on the abscissa are the values for: D(H-X) - EA(X) T,(H). This quantity corresponds to the heterolytic dissociation energy of HX, Le., the energy required to dissociate H X to H + and X- in the gas phase. This energy may be considered as a measure of the “acidity” of H X in the gas phase. As seen from Figure 3 an approximately linear correlation is obtained between the AHo,I(OH-, HX) and the heterolytic dissociation energy or gas phase acidity. Since the electron affinities of the halogens do not change much, the strongest interaction observed between OH- and H I is due to a gradual decrease of the bond dissociation energies D(H-X) from H F to HI. Of course reaction 1 is in a sense a neutralization reaction and the correlation in Figure 3 is therefore not surprising. Experiments on the solvation of C1by compounds HR to be reported in the futurelo show that the solvation interactions also for this ion follow to a certain extent the acidities of HR. These results are in line with the findings for OH- discussed above. I n the case of C1- one must also consider the negative ion to be in a sense a gas-phase Bronsted base and H R the gas-phase acid. The formation of the complex in the gas phase can be considered as involving only partial proton donation by the H R
+
or HX. (8) V. I. Vedeneyev, L. V. Gurvich, V. N. Kondratyev, V. A. Medredev, and Ye. L. Frankevich, “Bond Energies, Ionization Potentials and Electron Affinities,” Translated from Russian, Edward Arnold, London, 1962. (9) R. S. Berry and C. W. Reimann, J . Chem. Phgs., 38, 1540 (1963). (10) R . Yamdagni and P. Kebarle, submitted for publication.
Volume 7 4 , Number 7 April 3, 1970
