Brønsted–Lowry acid–base theory
|Independently, Johannes Nicolaus Brønsted and Thomas Martin Lowry formulated the idea that acids are proton (H+) donators while bases are proton acceptors.|
|Acids and bases|
In chemistry, the Brønsted–Lowry theory is an acid–base reaction theory, proposed independently by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923. The fundamental concept of this theory is that an acid (or Brønsted acid) is defined as being able to lose, or "donate" a proton (the hydrogen cation, or H+) while a base (or Brønsted base) is defined as a species with the ability to gain, or "accept," a proton.
Properties of acids and bases 
The Brønsted-Lowry model of proton donors and proton acceptors in acid-base reactions is an improvement over the Arrhenius theory, which was limited for it stated that bases had to contain a hydroxide ion. The main effect of the Brønsted-Lowry definition is to identify the proton (H+) transfer occurring in the acid-base reaction.
In the Brønsted-Lowry theory, an acid donates a proton and the base accepts it. The ion or molecule remaining after the acid has lost a proton is known as that acid's conjugate base, and the species created when the base accepts the proton is known as the conjugate acid. This is expressed in the following reaction:
- acid + base conjugate base + conjugate acid.
Notice how this reaction can proceed in either forward or backward direction; in each case, the acid donates a proton to the base.
With letters, the above equation can be written as:
- HA + Z A- + HZ+
The acid, HA, donates a H+ ion to become A-, its conjugate base. The base, Z, accepts the proton to become HZ+, its conjugate acid. In the reverse reaction, A- it accepts a H+ from HZ+ to recreate HA in order to remain in equilibrium. In the reverse reaction, as HZ+ has donated a H+ to A-, it therefore recreates Z and remains in equilibrium.
Consider the following acid-base reaction, seen in the image to the right:
- CH3COOH + H2O CH3COO− + H3O+
Acetic acid, CH3COOH, is an acid because it donates a proton to water (H2O) and becomes its conjugate base: the acetate ion (CH3COO−). In the same sense, H2O is the base because it accepts a proton from CH3COOH and becomes its conjugate acid: the hydronium (H3O+).
Hydronium, H3O+, is the conjugate acid of water because, in the reverse reaction, it donates a proton to the acetate ion, CH3COO−, and becomes water. The acetate ion, CH3COO−, is the conjugate base of acetic acid because, in the reverse reaction, it accepts an proton from H3O+ to become the acid.
Both of these processes demonstrate the equilibrium nature of the acid-base reaction.
Amphoteric substances 
- CH3COOH + H2O CH3COO− + H3O+
Water can also act as an acid, for instance when it reacts with ammonia. The equation given for this reaction is:
- H2O + NH3 OH− + NH4+
Acid strength 
A strong acid, such as hydrochloric acid, dissociates completely. A weak acid, such as acetic acid, may be partially dissociated; the acid dissociation constant, pKa, is a quantitative measure of the strength of the acid.
A wide range of compounds can be classified in the Brønsted–Lowry framework: mineral acids and derivatives such as sulfonates, phosphonates, etc., carboxylic acids, amines, carbon acids, 1,3-diketones such as acetylacetone, ethyl acetoacetate, and Meldrum's acid, and many more.
Brønsted concept and Lewis acids/bases 
A Lewis base, defined as an electron-pair donor, can act as a Brønsted–Lowry base as the pair of electrons can be donated to a proton. This means that the Brønsted–Lowry concept is not limited to aqueous solutions. Any donor solvent S can act as a proton acceptor.
- AH + S: A− + SH+
Typical donor solvents used in acid-base chemistry, such as dimethyl sulfoxide or liquid ammonia have an oxygen or nitrogen atom with a lone pair of electrons that can be used to form a bond with a proton.
Some Lewis acids, defined as electron-pair acceptors, also act as Brønsted–Lowry acids. For example, the aluminium ion, Al3+ can accept electron pairs from water molecules, as in the reaction
- Al3+ + 6H2O → Al(H2O)63+
The aqua ion formed is a weak Brønsted–Lowry acid.
The overall reaction is described as acid hydrolysis of the aluminium ion.
However not all Lewis acids generate Brønsted–Lowry acidity. The magnesium ion similarly reacts as a Lewis acid with six water molecules
- Mg2+ + 6H2O → Mg(H2O)62+
but here very few protons are exchanged since the Brønsted–Lowry acidity of the aqua ion is negligible (Ka = 3.0 × 10−12).
Boric acid also exemplifies the usefulness of the Brønsted–Lowry concept for an acid that does not dissociate but does effectively donate a proton to the base, water. The reaction is
- B(OH)3 + 2H2O B(OH)4− + H3O+
Here boric acid acts as a Lewis acid and accepts an electron pair from the oxygen of one water molecule. The water molecule in turn donates a proton to a second water molecule and, therefore, acts as a Brønsted acid.
See also 
- R.H. Petrucci, W.S. Harwood, and F.G. Herring, General Chemistry (8th edn, Prentice-Hall 2002), p.666
- G.L. Miessler and D.A. Tarr, Inorganic Chemistry (2nd edn, Prentice-Hall 1998), p.154
- K.W. Whitten, K.D. Gailey and R.E. Davis, "General Chemistry" (4th edn., Saunders College Publishing 1992) p.750