Theories of acids and bases

Example: HF + NH3 ↔ NH4

+ + F-

The HF transfers a H+ to the NH3, so it acts as an acid; the NH3 accepts the H+ so it acts as a base.

In Bronsted-Lowry theory, an acid can only behave as a proton donor if there is also a base present to act as a proton acceptor.

Consider the acid-base reaction between a generic acid HA and base B:HA + B ↔ A- + BH+

We can see that HA acts as an acid, while B acts as a base. But, if we look at the reverse reaction, BH+ acts as an acid (donates a proton) and A- acts as a base (accepts a proton). So, acids react to form bases and bases react to form acids. Acid-base pairs that are related to each other in this way are called conjugate acid-base pairs. Conjugate acid-base pairs differ by one proton.

One example of a conjugate pair is H2O and H3O+, which is found in all acid-base reactions in aqueous solution.

H2O + H+ ↔ H3O+ -- this reaction takes place when an acid dissolves in water. H3O+ is called the hydroxonium ion, the oxonium ion, or the hydronium ion.

For most reactions it is convenient to simply write H+ (aq).

So, H3O+ is the conjugate acid and H2O is the conjugate base

Water can act as a base or as an acid, depending what it is reacting with.

Substances that can act as acids or bases are said to be amphoteric or amphiprotic. To act as a Bronsted-Lowry acid, substances

must be able to dissociate and release H+. To act as a Bronsted-Lowry base, substances

must be able to accept H+, which means, they must have a lone pair of electrons.

Example: HCO3- acts as acid and base

HCO3-(aq) + H2O(l) ↔ CO3

2-(aq) + H3O+(aq) HCO3

-(aq) + H2O(l) ↔ H2CO3(aq) + OH-(aq)

Lewis acids and Bases

The Lewis definition of acids is broader than the Bronsted-Lowry theory.

By the Lewis definition, an acid is any species capable of accepting a lone pair of electrons (no longer restricted to just H+).

Lewis acid-base reactions result in the formation of a covalent bond which will always be a dative bond, because both electrons come from the base.

For example:

BF3 has an incomplete octet, so it is able to act as a Lewis acid by accepting a pair of electrons; NH3 acts as a Lewis base by donating its lone pair of electrons. The arrow shows the bond is a coordinate covalent bond (dative bond)

Other good examples of Lewis acid-base reactions are found in the chemistry of transition elements.

Transition metals often form ions with vacant orbitals. Thus they are able to act as Lewis acids and accept lone pairs of electrons.

The Lewis base, which donates the lone pair, is called a ligand, and usually surrounds the ion in a fixed number ratio.

Dative bonds form between the metal ion and the ligands, resulting in a complex ion that has a characteristic color.

Typical ligands found in complex ions include H2O, CN- and NH3. Note that these all possess lone pairs of electrons.

Theory Definition of acid Definition of baseBronsted-Lowry Proton donor Proton acceptorLewis Electron pair acceptor Electron pair donor

Although all Bronsted-Lowry acids are Lewis acids, not all Lewis acids are Bronsted-Lowry acids, so the term Lewis acids are usually reserved for those species which can only be described by Lewis theory (that is those that do not release H+).