The solvent system definition of acids and bases allows solutes to be defined as acids and bases by considering the autoioniza- tion products of the solvent. Most solvents are also either elec- tron pair acceptors or donors and hence are either Lewis acids or bases. The chemical consequences of solvent acidity and basicity are considerable, as they help to account for the differ- ences between reactions in aqueous and nonaqueous media. It follows that a displacement reaction often occurs when a solute dissolves in a solvent, and that the subsequent reactions of the solution are also usually either displacements or metatheses.
For example, when antimony pentafluoride dissolves in bro- mine trifluoride, the following displacement reaction occurs:
SbF5(s) + BrF3(l) → BrF2+(sol) + SbF6−(sol)
In the reaction, the strong Lewis acid SbF5 abstracts F− from BrF3. A more familiar example of the solvent as participant in a reaction is in Brứnsted theory, where the Lewis acid (H+) is always regarded as complexed with the solvent (as in H3O+ if the solvent is water) and reactions are treated as the trans- fer of the acid, the proton, from a basic solvent molecule to another base. Only the saturated hydrocarbons among com- mon solvents lack significant Lewis acid or base character.
Solvents with Lewis base character are common. Most of the well-known polar solvents, including water, alcohols, ethers, amines, dimethyl sulfoxide (DMSO, (CH3)2SO), dimethylformamide (DMF, (CH3)2NCHO), and acetonitrile (CH3CN), are hard Lewis bases. Dimethyl sulfoxide is an interesting example of a solvent that is hard on account of its O donor atom and soft on account of its S donor atom.
Reactions of acids and bases in these solvents are generally displacements:
N H
HH S N
O O H
HH S
O O
O S Me Me
+ O S
Me Me +
TABLE 5.7 Hammett acidity functions H0 for some strong acids, either as pure substances or as mixtures
Acid H0
HSbF6
(50 per cent HF, 50 per cent SbF5)
−31.3
‘Magic acid’
(25 per cent SbF5, 75 per cent HSO3F)
−21.5
HSO3F −15.1
CF3SO3H −14.1
H2S2O7 −14.1
HClO4 −13.0
H2SO4 −12.0
HF −11
Among Lewis acids, liquid sulfur dioxide is a good soft solvent for dissolving the soft base benzene. Unsaturated hydrocarbons may act as acids or bases by using their π or π* orbitals as frontier orbitals. Alkanes with electro- negative substituents, such as haloalkanes (e.g. CHCl3), are significantly acidic at the hydrogen atom, although saturated fluorocarbon solvents lack Lewis acid and base properties.
EXAMPLE 5.11 Accounting for properties in terms of the Lewis basicity of solvents
Silver perchlorate, AgClO4, is significantly more soluble in benzene than in alkane solvents. Account for this observation in terms of Lewis acid–base properties.
Answer We need to consider how the solvent interacts with the solute. The π electrons of benzene, a soft base, are available for complex formation with the empty orbitals of the cation Ag+, a soft acid. The Ag+ ion is thus solvated favourably by benzene.
The species [Ag–C6H6]+ is the complex of the acid Ag+ with π electrons of the weak base benzene.
Self-test 5.11 Boron trifluoride, BF3, a hard acid, is often used in the laboratory as a solution in diethyl ether, (C2H5)2O which is a hard base. Draw the structure of the complex that results from the dissolution of BF3(g) in (C2H5)2O(l).
(a) Liquid ammonia
KEY POINTS Liquid ammonia is a useful nonaqueous solvent. Many reactions in liquid ammonia are analogous to those in water.
Liquid ammonia is widely used as a nonaqueous solvent.
It boils at −33°C at 1 atm and, despite a somewhat lower relative permittivity (εr = 24) than that of water, it is a good solvent for inorganic compounds such as ammonium salts, nitrates, cyanides, and thiocyanides, and organic compounds such as amines, alcohols, and esters. It closely resembles the aqueous system as can be seen from the autoionization
2 NH3(l) NH4+(sol) + NH2−(sol)
Solutes that increase the concentration of NH4+, the solvated proton, are acids. Solutes that decrease the concentration of NH4+ or increase the concentration of NH2− are defined as bases. Thus, ammonium salts are acids in liquid ammonia and amines are bases.
Liquid ammonia is a more basic solvent than water and enhances the acidity of many compounds that are weak acids in water. For example, acetic acid is almost completely ionized in liquid ammonia:
CH3COOH(sol) + NH3(l) → NH4+(sol) + CH3COO−(sol)
Many reactions in liquid ammonia are analogous to those in water. The following acid base neutralization can be car- ried out:
NH4Cl(sol) + NaNH2(sol) → NaCl(sol) + 2 NH3(l) Liquid ammonia is a very good solvent for alkali and alkali earth metals, with the exception of beryllium. The alkali metals are particularly soluble and 336 g of caesium can be dissolved in 100 g of liquid ammonia at −50°C. The metals can be recovered by evaporating the ammonia. These solutions are very conducting and are blue when dilute and bronze when concentrated. Electron paramagnetic resonance spectra (see Chapter 8) show that the solutions contain unpaired electrons. The blue colour typical of the solutions is the outcome of a very broad optical absorp- tion band in the near IR with a maximum near 1500 nm.
The metal is ionized in ammonia solution to give ‘solvated electrons’:
Na(s) + NH3(l) → Na+(sol) + e−(sol)
The blue solutions survive for long times at low temperature but decompose slowly to give hydrogen and sodium amide, NaNH2. The exploitation of the blue solutions to produce compounds called ‘electrides’ is discussed in Section 11.14.
(b) Hydrogen fluoride
KEY POINT Hydrogen fluoride is a reactive toxic solvent that is highly acidic.
Liquid hydrogen fluoride (b.p. 19.5°C) is an acidic solvent with a relative permittivity (εr = 84 at 0°C) comparable to that of water (εr = 78 at 25°C). It is a good solvent for ionic substances. However, as it is both highly reactive and toxic, it presents handling problems, including its ability to etch glass. In practice, liquid hydrogen fluoride is usually con- tained in polytetrafluoroethylene and polychlorotrifluoro- ethylene vessels. Hydrogen fluoride is particularly hazardous because it penetrates tissue rapidly and interferes with nerve function. Consequently, burns may go undetected and treat- ment may be delayed. It can also etch bone and reacts with calcium in the blood.
Liquid hydrogen fluoride is a highly acidic solvent as it has a high autoprotolysis constant and produces solvated protons very readily:
3 HF(l) → H2F+(sol) + HF2−(sol)
Although the conjugate base of HF is formally F−, the abil- ity of HF to form a strong hydrogen bond to F− means that the conjugate base is better regarded as the bifluoride ion, HF2−. Only very strong acids are able to donate protons
and function as acids in HF, for example, fluorosulfonic acid:
HSO3F(sol) + HF(l) H2F+(sol) + SO3F−(sol)
Organic compounds such as acids, alcohols, ethers, and ketones can accept a proton and act as bases in HF(l). Other bases increase the concentration of H2F− to produce basic solutions:
CH3COOH(l) + 2 HF(l) CH3COOH2+(sol) + H2F−(sol) In this reaction acetic acid, an acid in water, is acting as a base.
Many fluorides are soluble in liquid HF as a result of the formation of the HF2– ion; for example
LiF(s) + HF(l) → Li+(sol) + HF2−(sol) (c) Anhydrous sulfuric acid
KEY POINT The autoionization of anhydrous sulfuric acid is complex, with several competing side reactions.
Anhydrous sulfuric acid is an acidic solvent. It has a high relative permittivity and is viscous because of extensive hydrogen bonding. Despite this association the solvent is appreciably autoionized at room temperature. The major autoionization is
2 H2SO4(l) H3SO4+(sol) + HSO4−(sol)
However, there are secondary autoionizations and other equilibria, such as
H2SO4(l) H2O(sol) + SO3(sol)
H2O(sol) + H2SO4(l) H3O+(sol) + HSO4−(sol) SO3(sol) + H2SO4(l) H2S2O7(sol)
H2S2O7(sol) + H2SO4(l) H3SO4+(sol) + HS2O7−(sol) The high viscosity and high level of association through hydrogen bonding would usually lead to low ion mobilities.
However, the mobilities of H3SO4+ and HSO4− are compara- ble to those of H3O+ and OH− in water, indicating that simi- lar proton transfer mechanisms are taking place. The main species taking part are H3SO4+ and HSO4−:
O S
O HO O H
O S
O
–O OH
O S
O HO O–
O S
O O OH H
O S
O HO OH
O S
O HO OH
O S
O HO OH
O S
O HO OH H
+ H +
Most strong oxo acids accept a proton in anhydrous sulfu- ric acid and are thus bases:
H3PO4(sol) + H2SO4(l) H4PO4+(sol) + HSO4−(sol) An important reaction is that of nitric acid with sulfuric acid to generate the nitronium ion, NO2+, which is the active species in aromatic nitration reactions:
HNO3(sol) + 2 H2SO4(l) NO2+(sol) + H3O+(sol) + 2 HSO4−(sol) Some acids that are very strong in water act as weak acids in anhydrous sulfuric acids, for example, perchloric acid, HClO4, and fluorosulfuric acid, HFSO3.
(d) Dinitrogen tetroxide
KEY POINTS Dinitrogen tetroxide autoionizes by two reactions. The preferred route can be enhanced by addition of electron-pair donors or acceptors.
Dinitrogen tetroxide, N2O4, has a narrow liquid range with a freezing point at −11.2°C and boiling point of 21.2°C.
Two autoionization reactions occur:
N2O4(l) NO+(sol) + NO3−(sol) N2O4(l) NO2+(sol) + NO2−(sol)
The first autoionization is enhanced by addition of Lewis bases, such as diethyl ether:
N2O4(l) + :X XNO+(sol) + NO3−(sol)
Lewis acids such as BF3 enhance the second autoionization reaction:
N2O4(l) + BF3(sol) NO2+(sol) + F3BNO2−(sol)
Dinitrogen tetroxide has a low relative permittivity and is not a very useful solvent for inorganic compounds. It is however a good solvent for many esters, carboxylic acids, halides, and organic nitro compounds.
(e) Ionic liquids
KEY POINT Ionic liquids are polar, nonvolatile solvents able to pro- vide very high concentrations of Lewis acids or bases as catalysts for many reactions.
Ionic liquids are salts with low melting points, usually below 100°C, which typically comprise asymmetric quater- nary (alkyl) ammonium cations and complex anions such as AlCl4− and carboxylates of various chain lengths. Ionic liquids are characterized by low volatility, high thermal
stability, inertness over a wide range of electrode potentials, and high conductivity, making possible numerous appli- cations such as providing alternative solvents for organic syntheses and electrochemistry. The ability of these ionic compounds to exist as liquids under ambient conditions is attributed to the large size and conformational flexibility of the ions—properties that give rise to small lattice ener- gies and large entropy increases accompanying melting. The cation and/or anion may be selected to impart chirality and acid–base properties to the solvent. Ionic liquids may them- selves serve as catalysts: for example, the chloroaluminate ionic liquid formed in reactions such as that shown here, provides the strong Lewis acid Al2Cl7− in very high concen- tration at ambient temperatures.
N N+ R Cl–
+ 2 AlCl3 N N+ R
Al2Cl7–
The anion of ionic liquids is often a Lewis base, such as the dicyanamide ion ((NC)2N−) which is a catalyst for acetylation reactions. In some cases the cation may pos- sess a basic group, an example being 1-alkyl-1,4-diazabi- cyclo[2.2.2]octane, known as [Cndabco]+ (26), which contains a tertiary-N atom capable of forming hydrogen bonds and conferring water solubility. Salts of [Cndabco]+ with the bis(trifluoromethane)sulfonimide anion, TFSI− (26), are miscible with water for n = 2 but immiscible with water for n = 8, while the melting point also drops with increasing n. Salts such as Cu(NO3)2 are usually insoluble in ionic liquids but dissolve in [Cndabco]+ TFSI because Cu2+ is complexed by the tertiary-N donors.
N+ N Cn
F3C S N–
S CF3 O O O O
26 [Cndabco]+ TFSI− (f) Supercritical fluids
KEY POINT Supercritical fluids have special properties as solvents and are finding increasing use in environmentally benign industrial processes.
A supercritical (sc) fluid is a state of matter where the liquid and vapour phases are indistinguishable: it has low viscosity combined with high dissolving capability for many solutes, and many gases are completely miscible. Supercritical fluids are produced by applying a combination of temperature and pressure that exceed the critical point (Fig. 5.16).
The most important example is supercritical carbon diox- ide (scCO2) which has a critical point at Pc = 72.8 atm and Tc = 30.95°C. The CO2 molecule is bipolar and can act as a Lewis base (27) or a Lewis acid; indeed, both types of inter- action can occur at a single CO2 molecule (28). As a solvent,
scCO2 has some important applications, such as decaffeina- tion of coffee and an increasing number of green industrial processes where scCO2 is replacing organic solvents that cause so much environmental concern. Unlike organic sol- vents, scCO2 can be removed at the end of a process by depressurization and then recycled (Section 25.5).
O C O
Al Cl ClCl
27 CO2:AlCl3 C O
O O
H R H H
28 CO2:OC(R)CH3
Compared to normal water, supercritical water (scH2O) is an excellent solvent for organic compounds and a poor sol- vent for ions. Close to the critical conditions (Pc = 218 atm, Tc = 374°C) water undergoes a remarkable change in prop- erties, from having a very high extent of autoprotolysis as it approaches the critical point (pKw is approximately 11, compared to 14) to having a much lower value above the critical point (pKw is approximately 20 at 600°C and 250 atm). Pressure and temperature can therefore be used to optimize the solvent for specific chemical reactions. A particularly important application of scH2O is the oxida- tion of organic waste materials, a process that exploits the complete miscibility of both organic compounds and O2 in this solvent (Section 25.5).
FIGURE 5.16 Pressure–temperature phase diagram for carbon dioxide showing the conditions under which it behaves as a supercritical fluid (1 atm = 1.01 × 105 Pa).
Solid
Liquid
Gas
200 250 300 350
10 102 103 104
1
Temperature / K pressure /10−5Pa
Critical point Supercritical fluid
Applications of acid–base chemistry
The Brứnsted and Lewis definitions of acids and bases do not have to be considered separately from each other. In fact, many applications of acid–base chemistry utilize both Lewis and Brứnsted acids or bases simultaneously.