Inorganic nonaqueous solvent
An inorganic nonaqueous solvent is a solvent other than water, that is not an organic compound. Common examples are liquid ammonia, liquid sulfur dioxide, sulfuryl chloride and sulfuryl chloride fluoride, phosphoryl chloride, dinitrogen tetroxide, antimony trichloride, bromine pentafluoride, hydrogen fluoride, pure sulfuric acid and other inorganic acids. These solvents are used in chemical research and industry for reactions that cannot occur in aqueous solutions or require a special environment.
Noble gas chemistry
The reactions of the compounds containing xenon are mostly conducted in hydrogen fluoride or bromine pentafluoride, which dissolve readily both xenon difluorides and its multiple derivatives, although sulfuric solvents are also used sometimes, in particular sulfuryl chloride fluoride for strong oxidants.
Nonaqueous acid-base chemistry
The acid-base reactions in non-aqueous solvents are typically described by means of the solvent-system definition, although the regular Brønsted-Lowry theory may be applied for the protic solvents, which possess a hydrogen atom that can dissociate. According to the solvent-system definition, acids are the compounds that increase the concentration of the solvonium (positive) ions, and bases are the compounds that result in the increase of the solvate (negative) ions, where solvonium and solvate are the ions found in the pure solvent in equilibrium with its neutral molecules:
protic solvents autodissociation
- 2NH3 ⇌ NH4+ (ammonium) + NH2− (amide)
- 3HF ⇌ H2F+ + HF2− (hydrogen difluoride)
- 2H2SO4 ⇌ H3SO4+ + HSO4−
aprotic solvents autodissociation
- N2O4 ⇌ NO+ (nitrosonium) + NO3− (nitrate)
- 2SbCl3 ⇌ SbCl2+ (dichloroantimonium) + SbCl4− (tetrachloroantimonate)
- POCl3 ⇌ POCl2+ + POCl4−
Thus NaNH2 is a base and NH4Cl is an acid in liquid ammonia, and they react, producing the salt and the solvent:
- NaNH2 + NH4Cl → 2NH3 + NaCl
or, for an aprotic example,
- NaNO3 + NOCl → N2O4 + NaCl
Limiting acids and limiting bases
The limiting acid in a given solvent is the solvonium ion, such as H3O+ (hydronium) ion in water. An acid which has more of a tendency to donate a hydrogen ion than the limiting acid will be a strong acid in the solvent considered, and will exist mostly or entirely in its dissociated form. Likewise, the limiting base in a given solvent is the solvate ion, such as OH− (hydroxide) ion, in water. A base which has more affinity for protons than the limiting base cannot exist in solution, as it will react with the solvent.
For example, the limiting acid in liquid ammonia is the ammonium ion, which has a pKa value in water of 9.25. The limiting base is the amide ion, NH2−. This is a stronger base than the hydroxide ion and so cannot exist in aqueous solution. The pKa value of ammonia is estimated to be approximately 34 (c.f. water, 14).
Any acid which is a stronger acid than the ammonium ion will be a strong acid in liquid ammonia. This is the case for acetic acid, which is completely dissociated in liquid ammonia solution. The addition of pure acetic acid and the addition of ammonium acetate have exactly the same effect on a liquid ammonia solution: the increase in its acidity: in practice, the latter is preferred for safety reasons.
Bases can exist in solution in liquid ammonia which cannot exist in aqueous solution: this is the case for any base which is stronger than the hydroxide ion but weaker than the amide ion. Many carbon anions can be formed in liquid ammonia solution by the action of the amide ion on organic molecules (see sodium amide for examples).
- SbF5 + HF ⇌ H+ + SbF6−
The limiting base, the hexfluoroantimonate anion SbF6−, is so weakly attracted to the hydrogen ion that virtually any other base will bind more strongly: hence, this mixture can be used to protonate organic molecules which would not be considered bases in other solvents.
Comparisons of acidity and basicity between solvents
There exists a large corpus of data concerning acid strengths in aqueous solution (pKa values), and it is tempting to transfer this to other solvents. Such comparisons are, however, fraught with danger, as they only consider the effect of solvation on the stability of the hydrogen ion, while neglecting its effects on the stability of the other species involved in the equilibrium. Gas phase acidities (normally known as proton affinities) can be measured, and their relative order is often quite different from that of the aqueous acidities of the corresponding acids. Few quantitative studies on acidities in nonaqueous solvents have been carried out, although some qualitative data are available. It appears that most acids which have a pKa value of less than 9 in water are indeed strong acids in liquid ammonia. However, the hydroxide ion is often a much stronger base in nonaqueous solvents (e.g. liquid ammonia, DMSO) than in water.
pH values are at present undefined in aprotic solvents, as the definition of pH assumes presence of hydronium ions. In other solvents, the concentration of the respective solvonium/solvate ions should be used, such as pCl in POCl3.
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