Brønsted–Lowry acid–base theory

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Johannes Brønsted.jpgThomas Martin Lowry2.jpg
Independently, Johannes Nicolaus Brønsted and Thomas Martin Lowry formulated the idea that acids are proton (H+) donors while bases are proton acceptors.

In chemistry, the Brønsted–Lowry theory is an acid–base reaction theory, proposed independently by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923.[1][2] 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[edit]

The Brønsted–Lowry model of proton donors and proton acceptors in acid–base reactions, proposed in 1923[3][4] is an improvement over the Arrhenius theory, which was limited for it stated that bases had to contain the hydroxyl group. The main effect of the Brønsted–Lowry definition is to identify the proton (H+) transfer occurring in the acid–base reaction. According to Arrhenius theory substances having OH group were called bases but this theory showed that substances which have one or more lone pair of electrons and are proton acceptors are bases.[4]

In the Brønsted–Lowry theory, an acid donates a proton and the base accepts it.[5][6] 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.[7] This is expressed in the following reaction:

acid + base is in equilibrium with 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 + B is in equilibrium with A + HB+

The acid, HA, donates a H+ ion to become A, its conjugate base. The base, B, accepts the proton to become HB+, its conjugate acid. In the reverse reaction, A it accepts a H+ from HB+ to recreate HA in order to remain in equilibrium. In the reverse reaction, as HB+ has donated a H+ to A, it therefore recreates B and remains in equilibrium.


Acetic acid, CH3COOH, is composed of a methyl group, CH3, bound chemically to a carboxylate group, COOH. The carboxylate group can lose a proton and donate it to a water molecule, H20, leaving behind an acetate anion CH3COO– and creating a hydronium cation H3O . This is an equilibrium reaction, so the reverse process can also take place.
Acetic acid, a weak acid, donates a proton (hydrogen ion, highlighted in green) to water in an equilibrium reaction to give the acetate ion and the hydronium ion. Red: oxygen, black: carbon, white: hydrogen.

Consider the following acid–base reaction, seen in the image to the right:

+ H
is in equilibrium with CH
+ H

Acetic acid, CH
, is an acid because it donates a proton to water (H
) and becomes its conjugate base: the acetate ion (CH
). In the same sense, H
is the base because it accepts a proton from CH
and becomes its conjugate acid: the hydronium (H

Hydronium, H
, is the conjugate acid of water because, in the reverse reaction, it donates a proton to the acetate ion, CH
, and becomes water. The acetate ion, CH
, is the conjugate base of acetic acid because, in the reverse reaction, it accepts an proton from H
to become the acid.

Both of these processes demonstrate the equilibrium nature of the acid–base reaction.

Amphoteric substances[edit]

Water as both base and acid. One H2O acts as a base and gains H+ to become H3O+; the other H2O acts as an acid and loses H+ to become OH.

Water is amphoteric and can act as an acid or as a base. In the reaction between acetic acid, CH3CO2H, and water, H2O, discussed in the above section, water acts as a base.

CH3COOH + H2O is in equilibrium with CH3COO + H3O+

The acetate ion, CH3CO2, is the conjugate base of acetic acid and the hydronium ion, H3O+, is the conjugate acid of the base, water.

Water can also act as an acid, for instance when it reacts with ammonia. The equation given for this reaction is:

H2O + NH3 is in equilibrium with OH + NH4+

in which H2O donates a proton to NH3. The hydroxide ion is the conjugate base of water acting as an acid and the ammonium ion is the conjugate acid of the base, ammonia.

Acid strength[edit]

Main article: 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, Ka, is a quantitative measure of the strength of the acid.[8]

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.


This theory requires the transfer of a proton and requires a solvent which can dissociate the substance into respective ions[9] and fails to explain why the reactions between acidic oxides and basic oxides occur.[10] An example of reaction between acidic and basic oxides is this:

CaO + SO3 = CaSO4

CaO is a basic oxide while SO3 is an acidic oxide. This reaction doesn't involves the transfer of a proton.

There are several acids like BF3 and AlCl3 which don't have a proton.

Non-protonic acid base reactions cannot be explained on the basis of this theory.[11] Examples of non-protonic acid-base reactions are:

SO2 + SO2 is in equilibrium with SO2+ + SO32-

SO3 + H2O is in equilibrium with H2SO4

Brønsted concept and Lewis acids/bases[edit]

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: is in equilibrium with 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.

Al(H2O)63+ + H2O is in equilibrium with Al(H2O)5OH2+ + H3O+...........Ka = 1.2 × 10−5 [12]

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).[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 is in equilibrium with 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[edit]


  1. ^ R.H. Petrucci, W.S. Harwood, and F.G. Herring, General Chemistry (8th edn, Prentice-Hall 2002), p.666
  2. ^ G.L. Miessler and D.A. Tarr, Inorganic Chemistry (2nd edn, Prentice-Hall 1998), p.154
  3. ^ Darrell Ebbing; Steven D. Gammon (2010). General Chemistry, Enhanced Edition. Cengage Learning. pp. 644–. ISBN 0-538-49752-1. 
  4. ^ a b Kenneth Whitten; Raymond Davis; Larry Peck; George Stanley (2013). Chemistry. Cengage Learning. pp. 350–. ISBN 1-133-61066-8. 
  5. ^ "Brønsted-Lowry Concept". Clackamas Community College. Retrieved 28 June 2014. 
  6. ^ "Brønsted Acids and Bases". Retrieved 28 June 2014. 
  7. ^ A. Ramakrishna (2014). Goyal’s IIT FOUNDATION COURSE CHEMISTRY: For Class-10. Goyal Brothers Prakashan. pp. 85–. GGKEY:DKWFNS6PECF. 
  8. ^ "Acid Dissociation Constant (Ka)". Boundless. Retrieved 28 June 2014. 
  9. ^ "The Lewis theory of acids and bases". Retrieved 28 June 2014. 
  10. ^ "LIMITATIONS: THE THEORIES (of Arrhenius & Bronsted-Lowry)". Retrieved 28 June 2014. 
  11. ^ Srivastava, H. C. (2003). Nootan - ISC Chemistry (7 ed.). India: Nageen Prakashan. p. 323. ISBN 978-93-80088-89-1. 
  12. ^ a b K.W. Whitten, K.D. Gailey and R.E. Davis, "General Chemistry" (4th edn., Saunders College Publishing 1992) p.750