Significance Alan G. Sharpe University Chemical Laboratory. Cambridge. England When an alkali metal halide MX dissolves in a large amount of water, the enthalpy changes involved may be represented as:
e,,,, a*,,
isvery small Whatever MX, the enthalpy of solution, and the sum of compared with the lattice enthalpy, so ionthe hydration energies of M+(g) and X-(g), EMhyd, water interactions in solution must be approximately equal to ion-ion interactions in the solid salt. Modern studies of ion solvation began with the classic work on the energies of solvation by Born (1)and on the structures of liquid water and hydrated ions by Bernal and Fowler (2). Until 1960 interest was concentrated on aqueous solutions, but since then i t has spread to nonaqueous media (especially those of high relative permittivity, in which ion association is least) and to experimental and theoretical studies of gaseous aggregates of ions and solvent molecules. Metal ion solvation has been thoroughly reviewed by Burgess (3), but there is no comparable treatment of anions, though much scattered material is available in a comprehensive account of ion solvation by Marcus (4).Detailed knowledee of solvated anions is confined to a few species, of which hakde ions are by far themost important. M& of this article is concerned with structural and thermodynamic aspects of halide ion solvation, first by water and then by other solvents; but, to show the significance of these matters to experimental chemistry, we end by briefly discussing some subjects in which anion solvation plays a key part--the thermodynamics of complex formation and of halogen exchange reactions and the kinetics of substitution reactions of alkyl halides. Characterlzatlon and Structures of Hydrated Halide Ions The now-accepted model of a hydrated ion in solution in a dvnamicone. At anv instant. the ion is surrounded by a first &ell of suitably oiiented water molecules, their number being the primary solvation number of the ion. Far away from the ion there is the unperturbed structure of bulk water. In between there is a secondary solvation shell, a few molecules thick, in which the structure of the solvent is modified to an intermediate extent. (As we shall see later, this secondary solvation shell is responsible for as much as half the total hydration energy of a singly charged ion.) Water molecules and ions alike are in ceaseless motion, and the identities of the individual molecules in the different solvation shells are time-variable. Where exchange of water between the primary and secondary solvation shells is slow (e.e.. for manv transition metal ions) isoto~icdilution analvnicoi NMR peak area ratioscan be used todetermine primarv solvation numbers. For halide ions, however, solvent exchange rates are among the fastest known: the mean time a
given water molecule is in the primary solvation shell of C1is onlv about 5 X 10-12 s (5). . . far too short for anv NMR structural investigation. Nor is electronic spectroscopy useful for halide ions insolution: onlv charge transfer to sohent hands are found in their spectra: So piacticslly all modem exoerimental information on these ions comes from diffraction methods or studies of their effect on the infrared vibrational sDectrum of the solvent. ~ e f o i ediscussing this information we must mention the two limitine structures (both Dlanar) Drooosed for a X-Hz0 entity.l'hese are
or, if we show also the dipole of the water molecule,
Structure 1 (symmetry C,) was first put forward by Bernal and Fowler (2); we would now describe the interaction as ion-dipolar reinforced by hydrogen bonding. In structure 2 (symmetry CzJ, first suggested explicitly by Buckingham (6), the interaction is essentially ion-dipolar, the geometry being unfavorable to hydrogen bonding. (In a more sophisticated model (6), the water molecule is considered as a quadrupole consisting of two positive charges near the hydrogen atoms and two neeative charees corresoondine to the lone pairs oftheoxygenatom; weshall return to this model later.) The definitive structural studies of an aaueous halide ion, which show that in the case of chloride s&ucture 1 is the correct one, are those of Enderbv and his colleames (7,8), made by neutron diffraction usini a difference mkthod. The coherent neutron scattering.lengths . of ' T I and 3-CI differ by a factor of 4 5 , and measurement of the difference between scattering by 35C1and 37C1- in D 2 0 eliminates the waterwater scattering (which would otherwise dominate the total scattering) and permits accurate determination of the timeaveraged C1-0 and C1-D terms. Table 1 (8)lists values for the C1-0 and C1-D distances and for the departure of the C1D-0 angle from linearity, J., derived by assuming the O-D distance is 1.00 A. I t will he seen that interatomic distances are nearly independent of environment, though the primary solvation number is smaller in concentrated solutions. A 4.35 rn solution of NiClz has nearly the same molar ratio as the known solid hydrate MgC12.12Hz0 (9.10) and X-ray diffraction shows solution and solid to be isostructural(l1). The positions of the protons in the solid magnesium salt can only be inferred from the interatomic distances, but the presence of distorted octahedra of oxygen atoms round the
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Volume 67 Number 4
April 1990
309
Table 1. Chloride Ion Hydration Cl:Dil)
Solute Molality
ClLO
A
A
CI-92)
Ji
A
deg.
Coordination number
Table 2. Some Propetiles ot Solvents Solvent Water Methanol Acetone Formamide Dimethyiformsmide Propylene carbonate Nilromethane Acelonitrile Dimethylsulfoxide
Formula H20
CHSOH
(CHd&O HCONHI
HCON(CHah CH3CHCH29
L CH3N02
CH&N (CH&SO
~
bpl°C
Density1 g cm+
p,lD
cs
100 65 56 210 153 242
1.00 0.79 0.78 1.13 0.94 1.20
1.8 2.9 2.7 3.3 3.9 5.0
78 33 21 111 37 66
3.6 3.5 3.9
36 37 47
-
~
101 82 189
~ 1.13 0.76 1.10
anion a t 3.11-3.26 A now leaves no room for doubt that the environments of chloride ions in MgCly12HzO and in 4.35 m NiClz solution are substantially the same. The difference neutron diffraction method can he used only for elements that have two relatively abundant isotopes having different scattering powers. These factors, together with financial considerations, severelv restrict its application to other ions, for which we still rely on scanty daia-from X-ray studies. The structures of solid KF.2H20 (lo), (CH3)4NF.4H20 (12), and NaBrdHzO (13) all strongly imply X--HOH hydrogen honding, and X-ray diffraction by aaueous LiBr (14). CaBr7iI5). - . .. and NaI (16) indicates octahLdral coordinki& of the anion with oxygen atoms of six water molecules a t 3.4.3.4. and 3.7 A. resoectivelv. distances again compatible w i t h x - ~ H O Hhydrogen bonding. Evidence of a different kind is available from infrared spectroscopy. I t is well established that for the O-H stretching vihration, an increase in the strength of hydrogen honding is accompanied by a shift to lower frequency. Because of the breadth of the fundamental O-H stretching band for water, it in difficult to deduce much from the smill changes that occur when metal halidesare dissolved in it. However, a considerable imnrovement can he made bv studvine the 0H stretching vidration for HOD in DzO o;the 0 I ~ - s t r e t c h ine vihration for HOD in H9O. - . since the three vibrational bands are then well separated and are not coupled mgether. The soectra of solutions of alkali metal halides are closelv similar and, as expected, frequency changes increase slight& with increasing concentration. Typical results (17-19) are those for v(0H) in 3 M KF, 3295; 9 M NaC1,3425; 6 M KI, 3460 em-'. Values for gaseous and liquid HOD are 3703 and 3405 cm-I, respectively, and it is inferred that the strength of the hydrogen bonding decreases in the sequence The evidence considered so far indicates that structure 1 is probably essentially correct for all hydrated halide ions in solution, but the structures of gaseous hydrates are also of considerable interest, and i t is appropriate to mention here 310
Journal of Chemical Education
the results of two sets of ah initio molecular orbital calculations. For an isolatedx-.Hz0 entity, the first set show structure I, with LXHO = 172O for F-and 154O for C1-, to he more stable than structure 2, hut only by 25 and 5 kJ mol-1, respectively (20). In the second set the most stable LXHO is obtained as 169,147,124, and 115O for F-, C1-, Br-, and I-; thus structure 2 is a better representation of isolated Br--Hz0 and I-.Hz0 species (21). Does structure 2 ever occur in solution? In this connection. it mav he noted that a partial structural chanee occurs for some hidrated cations on dilution: the angle m i e by the olane of the water molecule with the M-0 axis of MOHg is ;bout 50' in concentrated solution (implying i n t e ~ a c t i o ~ o f only one lone pair with the cation) and much less in dilute solution (implying ion-dipole interaction) (7). This change occurs with no appreciable enthalpy of dilution. However, it is not accompanied by any obviouidecrease in the ahilitiof molecules in the primary solvation sphere to hydrogen-bond to those bevond it. For anions. on the other hand. s
