acids and bases

acids and bases, two related classes of chemicals; the members of each class have a number of common properties when dissolved in a solvent, usually water.

Properties

Acids in water solutions exhibit the following common properties: they taste sour; turn litmus paper red; and react with certain metals, such as zinc, to yield hydrogen gas. Bases in water solutions exhibit these common properties: they taste bitter; turn litmus paper blue; and feel slippery. When a water solution of acid is mixed with a water solution of base, water and a salt are formed; this process, called neutralization, is complete only if the resulting solution has neither acidic nor basic properties.

Classification

Acids and bases can be classified as organic or inorganic. Some of the more common organic acids are: citric acid, carbonic acid, hydrogen cyanide, salicylic acid, lactic acid, and tartaric acid. Some examples of organic bases are: pyridine and ethylamine. Some of the common inorganic acids are: hydrogen sulfide, phosphoric acid, hydrogen chloride, and sulfuric acid. Some common inorganic bases are: sodium hydroxide, sodium carbonate, sodium bicarbonate, calcium hydroxide, and calcium carbonate.

Acids, such as hydrochloric acid, and bases, such as potassium hydroxide, that have a great tendency to dissociate in water are completely ionized in solution; they are called strong acids or strong bases. Acids, such as acetic acid, and bases, such as ammonia, that are reluctant to dissociate in water are only partially ionized in solution; they are called weak acids or weak bases. Strong acids in solution produce a high concentration of hydrogen ions, and strong bases in solution produce a high concentration of hydroxide ions and a correspondingly low concentration of hydrogen ions. The hydrogen ion concentration is often expressed in terms of its negative logarithm, or pH (see separate article). Strong acids and strong bases make very good electrolytes (see electrolysis), i.e., their solutions readily conduct electricity. Weak acids and weak bases make poor electrolytes.

See buffer; catalyst; indicators, acid-base; titration.

Acid-Base Theories

There are three theories that identify a singular characteristic which defines an acid and a base: the Arrhenius theory, for which the Swedish chemist Svante Arrhenius was awarded the 1903 Nobel Prize in chemistry; the Brönsted-Lowry, or proton donor, theory, advanced in 1923; and the Lewis, or electron-pair, theory, which was also presented in 1923. Each of the three theories has its own advantages and disadvantages; each is useful under certain conditions.

The Arrhenius Theory

When an acid or base dissolves in water, a certain percentage of the acid or base particles will break up, or dissociate (see dissociation), into oppositely charged ions. The Arrhenius theory defines an acid as a compound that can dissociate in water to yield hydrogen ions, H⁺, and a base as a compound that can dissociate in water to yield hydroxide ions, OH^- . For example, hydrochloric acid, HCl, dissociates in water to yield the required hydrogen ions, H⁺, and also chloride ions, Cl^- . The base sodium hydroxide, NaOH, dissociates in water to yield the required hydroxide ions, OH^-, and also sodium ions, Na⁺.

The Brönsted-Lowry Theory

Some substances act as acids or bases when they are dissolved in solvents other than water, such as liquid ammonia. The Brönsted-Lowry theory, named for the Danish chemist Johannes Brönsted and the British chemist Thomas Lowry, provides a more general definition of acids and bases that can be used to deal both with solutions that contain no water and solutions that contain water. It defines an acid as a proton donor and a base as a proton acceptor. In the Brönsted-Lowry theory, water, H₂O, can be considered an acid or a base since it can lose a proton to form a hydroxide ion, OH^-, or accept a proton to form a hydronium ion, H₃O⁺ (see amphoterism). When an acid loses a proton, the remaining species can be a proton acceptor and is called the conjugate base of the acid. Similarly when a base accepts a proton, the resulting species can be a proton donor and is called the conjugate acid of that base. For example, when a water molecule loses a proton to form a hydroxide ion, the hydroxide ion can be considered the conjugate base of the acid, water. When a water molecule accepts a proton to form a hydronium ion, the hydronium ion can be considered the conjugate acid of the base, water.

The Lewis Theory

Another theory that provides a very broad definition of acids and bases has been put forth by the American chemist Gilbert Lewis. The Lewis theory defines an acid as a compound that can accept a pair of electrons and a base as a compound that can donate a pair of electrons. Boron trifluoride, BF₃, can be considered a Lewis acid and ethyl alcohol can be considered a Lewis base.

A Lewis acid is a chemical compound, A, that can accept a pair of electrons from a Lewis base, B, that acts as an electron-pair donor, forming an adduct, AB.

A + :B → A—B

Gilbert N. Lewis proposed this definition, which is based on chemical bonding theory, in 1923. Brønsted-Lowry acid-base theory was published in the same year. The two theories are distinct but complementary to each other as a Lewis base is also a Brønsted-Lowry base, but a Lewis acid need not be a Brønsted-Lowry acid.

The classification into hard and soft acids and bases (HSAB theory) followed in 1963. The strength of Lewis acid-base interactions, as measured by the standard enthalpy of formation of an adduct can be predicted by the Drago-Wayland two-parameter equation.

History

Lewis had suggested in 1916 that two atoms are held together in a chemical bond by sharing a pair of electrons. When each atom contributed one electron to the bond is was called a covalent bond. When both electrons come from one of the atoms is was called a dative covalent bond or coordinate bond. The distinction is not clear-cut as the diagram at the right shows; although the ammonia molecule donates a pair of electrons to the hydrogen ion, the identity of the electrons is lost in the ammonium ion that is formed. Nevertheless, Lewis suggested that an electron-pair donor be classified as a base and an electron-pair acceptor be classified as acid.

The modern definition of a Lewis acid is an atomic or molecular species that has an empty atomic or molecular orbital of low energy (LUMO) that can accommodate a pair of electrons, as illustrated in the molecular orital diagram at the right.

Comparison with Brønsted-Lowry theory

A Lewis base is usually a Brønsted-Lowry base as it can donate a pair of electrons to a proton; the proton is a Lewis acid as it can accept a pair of electrons. The conjugate base of a Brønsted-Lowry acid is is also a Lewis base as loss of a proton from the acid leaves those electrons which were used for the A—H bond as a lone pair on the conjugate base. However, a Lewis base can be very difficult to protonate, yet still react with a Lewis acid. For example, carbon monoxide is a very weak Brønsted-Lowry base but it forms a strong adduct with BF₃.

In another comparison of Lewis and Brønsted-Lowry acidity by Brown and Kanner, 2,6-di-t-butylpyridine reacts to form the hydrochloride salt with HCl but does not react with BF₃. This example demonstrates that for pyridine bases, HCl (typically thought of as a is a Brønsted-Lowry acid) is a "stronger" acid than BF₃ (a Lewis acid).

A Brønsted-Lowry acid is a proton donator, not an electron-pair acceptor.

Lewis acids

Acceptor orbitals of a Lewis acid are as in the following acid + base reactions.

1s orbital

• H⁺ + NH₃: → NH₄⁺

p orbitals: elements in groups 1—3

• B₂H₆ + 2H^- → 2BH₄^-
• BF₃ + F^- → BF₄^-
• Al₂Cl₆ + 2Cl^- → 2AlCl₄^-

d orbitals: elements in the second and lower rows of the periodic table

• AlF₃ + 3F^- → AlF₆^3-
• SiF₄ + 2F^- → SiF₆^2-
• PCl₅ + Cl^- → PCl₆^-
• SF₄ + F^- → SF₅^-
• Metal ions forming solvates, such as [Mg(H₂O)₆]²⁺, [Al(H₂O)₆]³⁺, etc. where the solvent is a Lewis base.

A typical example of a Lewis acid in action is in the Friedel-Crafts alkylation reaction. The key step is the acceptance by AlCl₃ of a chloride ion lone-pair, forming AlCl₄^- and creating the strongly acidic, that is, electrophilic, carbonium ion.

RCl +AlCl₃ → R⁺ + AlCl₄^-

Lewis Bases

A Lewis base is an atomic or molecular species that has an lone pair of electrons in the HOMO. Typical examples are

• compounds of N, P, As, Sb and Bi in oxidation state 3
• compounds of O, S, Se and Te in oxidation state 2, including water, ethers, ketones, sulphoxides
• molecules like carbon monoxide

An easy way to remember this is that nearly all of the compounds formed by the transition elements are coordination compounds wherein the metal or metal ion is a Lewis acid.

Hard and soft classification

Considerations concerning the strength of acid base adducts lead R.G. Pearson to propose, in 1963, the classification of both acids and bases into hard and soft. Within each category he established an order of binding strengths such as

• hard acids: R₃P << R₃N, R₂S << R₂O
• soft acids: R₂O << R₃N, R₂S << R₃P

For example, an amine will displace a phosphine from the adduct with the acid BF₃. In the same way, bases could be classified. For example, bases donating a lone pair from an oxygen atom are harder than bases donating through a nitrogen atom. Although the classification was never quantified it proved to be very useful in predicting the strength of adduct formation, using the key concepts

• hard acid — hard base interactions are stronger than hard acid — soft base or soft acid — hard base interactions.
• soft acid — soft base interactions are stronger than soft acid — hard base or hard acid — soft base interactions.

Later investigation of the thermodynamics of the interaction suggested that hard—hard interactions are enthalpy favoured, whereas soft—soft are entropy favoured. If the interaction between acid and base in solution results in an equilibrium mixture the strength of the interaction can be quantified in terms of an equilibrium constant. An alternative quantitative measure is the standard heat (enthalpy) of formation of the adduct in a non-coordinating solvent. Drago and Wayland proposed a two-parameter equation which predicts the formation of a very large number of adducts quite accurately.

–ΔH^O (A—B) = E_AE_B + C_AC_B

Value of the E and C parameters can be found in Drago et. al.

Another quantitative system as been proposed, in which Lewis acid strength is based on gas-phase affinity for fluoride.

References

Further reading

Jensen, W.B. (1980). The Lewis acid-base concepts : an overview. New York: Wiley. Yamamoto, Hisashi (1999). Lewis acid reagents : a practical approach. New York: Oxford University Press.

See also

• acid
• base
• Acid-base reaction
• Brønsted-Lowry acid-base theory