Chemistry 242 - Inorganic Chemistry II
Chapter 7 - Solvents, Solutions, Acids and Bases

This chapter contains a very useful summary of definitions related to the title topics, notably extentions of acid bases definitions to non-aqueous systems.

Solvent Properties

Their usefulness is a function of:

  1. Liquid temperature range
  2. Dielectric constant
  3. Its Lewis acid/base properties
  4. Its Brønsted/Lowry acid base properties
  5. Autodissociations properties
Table 7.1 gives the properties of a range of common solvents

Liquid Range

The most useful are liquid at room temperature but we often need solvents to go upto high or low temperatures. The range is large for DMF and propane-1,2-diol carbonate. Ammonia is a good solvent for many reactions but it boils at about -33oC. Hydrogen fluoride, which boils at about 20 oC is also useful, but very corrosive.

Dielectric Constant

This factor is important if the sovent is to dissolve ionic compounds. Water happens to have a very high dielectric constant (82) and is probably the best all-round solvent for ionic compounds that we have. The key formula here is:

F = q+q-/4per2

Lewis Acid/Base Properties

The extent to which the sovent can act as a ligand towards the ions of an ionic solute, with the dielectric constant which is related but not directly, determines how good a solvent it will be. In the crystal lattice with simple ions, the anions effectively complex the cations and vice versa. The solvent will have to do a better job.

Generally, because the cations are smaller than the anions, the greatest gains are to be made by effective complexation of the cations, so it is the Lewis basicity of the solvent that is more important. For common solvents:

(CH3)2SO > H(CO)N(CH3)2 » H2O > (CH3)2CO >

(CH3CHCH2)O2CO > (CH3)2SO2 > CH3NO2 > C6H5NO2 » CH2Cl2

Protic Solvents

These solvents contain protons which can be inozed, in other words the solvent is a Brønsted acid. Examples are H2O, HF, H2SO4, HCN and even MH3.


Also known as autoionization, so examples in protic solvents are:

2H2O      H3O+   +   OH-
3HF      H2F+   +   HF2-
2H2SO4      H3SO4+   +   HSO4-
2H2O      H3O+   +   OH-
2H3N      H4N+   +   NH2-
The cations and anions produced can interact with the solutes helping to explain certain reaction products. Also, the autodissociation reactions are not really this simple - consider water: The more general equation is:

(n+m+1)H2O      [H(H2O)n]+   +   [OH(H2O)m]-
The ions [H3O]+, [H(H2O)2]+ and [H(H2O)3]+ have been characterized, the first two in crystal structures.

K'25oC = [H+][OH-]/[H2O] = (1.0x10-14)/55.56 and K25oC = [H+][OH-] = 1.0x10-14

Aprotic Solvents

Such solvents fall into three broad classes:
  1. Non-polar, or weakly polar solvents which do not dissociate and are not strongly coordinating. Examples are hydrocarbons and CCl4. They are poor solvents for everything except like substances, i.e. other nonpolar molecules.

  2. Strongly solvating (usually polar) solvents which do not dissociate. Examples are acetonitrile ( CH3CN), N,N-dimethyformamide (DMF, H(CO)N(CH3)2), dimethyl sulphoxide (DMSO (CH3)2SO), tetrahydrofuran (THF C4H8O) and liquid sulphur dioxide (SO2).

    They are similar in being very strongly coordinating towards cations although SO2 forms an adduct with acetonitrile. The dielectric constants can very quite a bit (DMSO = 45, THF = 7.6) which will govern their ability to dissolve ionic materialas.

  3. Highly polar and autoionizing solvents. Examples are:

    2BrF3      BRF2+   +   BrF4-

    2NO2      NO+   +   NO3-

    2Cl3PO      Cl2PO+   +   Cl4PO-

    Molten Salts

    These are the extreme case of autoionizing solvents. Generally they will be very high temperature systems but lower temperatures can be attained by the right eutectic mixtures, e.g.

    LiNO3/NaNO2/KNO3 mixtures can reach as low as 130 oC

    (C2H5)2NH2Cl melts at 215 oC

    [N,N-RR'N2C3H3]+Cl- with AlCl3 can be liquid at room temperature.


    Aluminum is made by electrolyis of Al2O3 in a cryolite {Na3AlF6]

    Re3Cl9   +   Et2NH2Cl      [Et2NH2][Re2Cl8] with its quadruple bond.

    Solvents for Electrochemistry

    Solvent which are useful for electrochemistry must have two characteristics: a highish dielectric constant so that they are good solvents for ionic compounds, and they must be redox resistant. Water is actually not ideal because, although its dielectric constant is very high, it is susceptible to both oxidation and reduction at relatively small potentials:

    H2O    O2   +   4H+(10-7 M)   +   4e-      E = -0.82 V

    i.e. ½X2   +   e-      X-      E > +0.82 V might generate oxygen from water

    H+(10-7 M)   +   e-      ½H2      E = -0.41 V

    i.e. M      M+   +   e-      E > +0.41 V might generate H2 from water

    Acetonitrile or DMF are often used as an electrochemistry solvent for example in cyclic voltammetry for organometallic substances usually with a redox inert supporting electrolyte such as t-butylammonium perchlorate.

    Purity of Solvents

    Water and oxygen are th most common contaminants and they can be very difficult to remove. Many research labs have several continuously operating stills from which the solvents can be removed under cover of nitrogen or argon. Hydrocarbon solvents and ethers can be distilled over sodium (or potassium) in the presence of benzoquinone as an intense blue indicator (of dryness). Various hydrides can also be used.

    More reactive solvents can be dried effectively over molcular sieves and vacuum distilled to remove dissolved oxygen.

    It is important to remember that even traces of contaminents will mess up sensitive reactions because the solvent is present in such excess.

    Definitions of Acids and Bases

    The Brønsted/Lowry Defintions specifies an acid as a proton donor and a base as a proton acceptor which applies to aqueous systems.

    The General Solvent System Definition is an extension to any autoionizing solvent. An acid is defined as a substance increasing the concentration of the characteristic cation of the solvent. One that increases the concentration of the characteristic anion (or decreases the concentrqtion of the cation) is a base:

    2H2O      H3O+   +   OH-

    HCl is an acid
    NaOH and Na2O are bases

    2H3N      NH4+   +   NH2-

    NH4Cl or urea, H2N(CO)NH2 are acids
    NaNH2 and Na2N are bases

    2NO2      NO+   +   NO3-

    NOCl is an acid
    NaNO3is a base

    2H2SO4      H3SO4+   +   HSO4-

    CH3COOH is an actually a base which is protonated to CH3COOH2+
    HAsF6 might be an acid if it does not react.
    NaHSO3 is a base

    The Lux/Flood Definition covers things which would become acids or bases if dissolved in water.
    CO2   +   CaO      CaCO3
    Here CO2 is considered the acid - carbonic acid anhydride and CaO is considered the base since it woud give Ca(OH)2 in water.

    The Lux/Flood definition defines an acid as an oxide ion acceptor and a base as an oxide ion donor and is mainly used for high temperature anhydrous systems for example in steel-making (in acidic or basic "slags"):

    CaO   +   SiO2      CaSiO3

    2Na2O   +   P2O5      2Na3PO4

    The Lewis Definition considers an acid to be an electron pair acceptor and a base as an electron pair donor. When they react, the resulting compound is called an adduct and the bond a dative bond. It encompasses the Brønsted/Lowry definition and many other reaction types:

    H+   +   :OH-      H2O
    F3B   +   :NH3      F3BNH3
    All the usual ligands can be viewed as Lewis bases and the metal ions as lewis acids.

    Acid/Base Strength and Electonic Effects

    (CH3)3N:   >   H3N:   >   F3N:
    (CH3)3B   <   "H3B"   <   F3B
    In the above series, F is the most electron withdrawing and CH3 the most electron releasing. However, there can be unexpected effects. Consider the observed order:

    BF3   <   BCl3   <   BBr3
    The reduced strength of the fluoride compound is due to internal B-F p-bonding which goes some way towards satisfying the electron deficiency of the boron. A similar effect is observed in the B(OR)3 compounds (relative to the S or Se analogues).

    Acid/Base Strength and Steric Effects

    Towards protons, pyridine is a weaker base than 2-methyl or 4 methyl pyridine which are about equally stronger. However, towards B(CH3)3 the 4-methyl pyridine is the significantly stronger base. The methyl group of the 2-methyl pyridine gets in the way. A similar effect is seen with triethylamine and quinuclidine (CH(CH2CH2)3N.

    Also, B(CH3)3 is a stronger acid than the highly hindered B(C(CH3)3)3.

    Hard and Soft Acids and Bases

    The basic principle is that like prefers like:

    Complexes of
    Type (a) Metals
    Ligands Complexes of
    Type (b) Metals
    Strongest bonding R3N R2O F- Weakest bonding
    R3P R2S Cl-
    R3As R2Se Br-
    Weakest bonding R3Sb R2Te I- Strongest bonding
    Type (a) Hard Acids

    • Alkali and alkaline earth metals
    • Other lighter and highly charged cations e.g. Ti4+, Fe3+, Co3+, Al3+

    Type (b) Soft Acids

    • Heavier transition metal ions e.g. Hg22+, Hg2+, Pt2+, Pt4+
    • Metals in lower oxidation states e.g. Ag+, Cu+ and zero metals

    Hardness is associated with low polarizability and a tendency towards ionic bonding.

    Softness is associated with greater polarizability and more covalent bonding.

    The Drago-Wayland Equation

    This empirical equation can be used to obtain a fairly quantitative value for DHab, the enthalpy of formation of and Lewis adduct:

    -DHAB   =   EA.EB   +   CA.CB   +   W
    The factors E are a measure of the acid or base's tendency to form a bond by electrostatic interaction which is roughly correlated with hardness. Similarly the factors C go with the tendency to form a covalent bond, and is correlated with softness. Both factors are considered to be constant regardless of the partner. The factor W is almost always zero (always in the data given!) and is a fudge factor for use if it is suspected that and acid or a base makes a constant contribution to DHAB which is not dependent on the partner and so should not be included in the product terms.

    Table 7.2 lists the experimental E and C parameters for a variety of acids and bases. It must be understood that every adduct of a "new" acid or base whose DEAB is measured allows a new entry to be added to the table. In the beginning four parameters, had to be given arbitarily assigned values (remember the Pauling electronegativity scale). They are marked b

    As an example, consider:

    (CH3)3B   +   :N(CH3)3      (CH3)3B:N(CH3)3
    -DE = 5.79x1.19 + 1.57x11.20 = 24.47 kcal mol-1
    (CH3)3B   +   :P(CH3)3      (CH3)3B:P(CH3)3
    -DE = 5.79x1.11 + 1.57x6.51 = 14.80 kcal mol-1
    "(CH3)3Al"   +   :N(CH3)3      (CH3)3Al:N(CH3)3
    -DE = 17.32x1.19 + 0.94x11.20 = 31.14 kcal mol-1
    "(CH3)3Al"   +   :P(CH3)3      (CH3)3Al:P(CH3)3
    -DE = 17.32x1.11 + 0.94x6.51 = 25.35 kcal mol-1
    The amine adducts are both stronger than the phosphine adducts, but the difference is much more marked for the harder boron. Unfortunately the table of acids does not contain any really soft ones where the order of strength of the dative bond might be reversed.

    Some Common Protic Acids

    Sulphuric acid (H2SO4)

    Nitric Acid (HNO3)

    Aqua Regia

    This is a mixture of 3 volumes of concentrated hydrochloric acid to 1 volume of concentrated nitric acid. It contains free Cl2 and ClNO and has very powerful oxidizing properties as well as being a strong acid. It is the acid of last resort to get metal into solution for analysis. It will dissolve Au and Pt which form [AuCl4]- and [PtCl6]2-. Mercuric sulphide will dissolve because the sulphide is oxidized to SO2 and the mercuric ion is in the form of [HgCl4]2-

    Perchloric Acid

    Hydrohalic Acids (except HF)

    Hydrofluoric Acid

    The Strengths of the Oxyacids

    Consider compounds of type: XOn(OH)m e.g. PO(OH)3 Te(OH)6P

    Super Acids

    The strongest acid in aqueous solution is H3O+ since all the strong acids protonate water to 100% of their capacity. In non-aqueous solution it is possible to generate more powerful proton donors.

    The strength of such strong acids is measured by the Hammett acidity function:

    Ho = pKBH+ - log([BH+]/[B])
    Here B and BH+ usually refer to a colorimetric indicator, that is either B or (more likely) BH+ is strongly coloured, and pKAB+ is small or negative. Typical indicators are:

    Indicator pKAB+
    (in H2SO4)
    m-nitroaniline +2.5*
    p-nitroaniline +0.99*
    o-nitroaniline -0.29
    2,4-dinitroaniline -4.53
    3-methyl-2,4,6-trinitroaniline -8.22
    2,4,6-trinitroaniline -10.10
    *The first two provide overlap with the normal range of the pH scale (~0 - 14).

    Notice that the Hammett scale is actually equivalent to the pH scale in the case of aqueous solutions:

    A table of selected Ho values is given below. In each case, the equilibrium leading to the proton donor species is shown.

    Ho = -log([B][H+]/[BH+] - log([BH+]/[B]) = -log[H+]

    Acid Ho
    2H2SO4 H3SO4+ + HSO4- -12
    H2SO4 + H2S2O7 H3SO4+ + HS2O7- (in oleum) -15
    3HF H2F+ + HF2- -11
    2HF + SbF5 H2F+ + HSBF6 -12
    2HSO3F H2SO3F+ + SO3F- -15
    2HSO3F + SbF5 H2SO3F+ + SbF5SO3F- -19

    The super acids can be used in a variety of reactions leading to unusual cations:

    (CH3)3COH   +   super acids      (CH3)3C+   +   H2O
    I2   +   super acids      I2+ or I3+