In chemistry, neutralization or neutralisation (see spelling differences), is a chemical reaction in which an acid and a base react quantitatively with each other. In a reaction in water, neutralization results in there being no excess of hydrogen or hydroxide ions present in solution. The pH of the neutralized solution depends on the acid strength of the reactants. Neutralization is used in many applications.
Meaning of “Neutralization”
- acid + base → salt + water
- HCl + NaOH → NaCl + H2O
The statement is still valid as long as it is understood that in an aqueous solution the substances involved are subject to dissociation, which changes the substances ionization state. The arrow sign, →, is used because the reaction is complete, that is, neutralization is a quantitative reaction. A more general definition is based on Brønsted–Lowry acid–base theory.
- AH + B → A + BH
Electrical charges are omitted from generic expressions such as this, as each species A, AH, B, or BH may or may not carry an electrical charge. Neutralization of sulphuric acid provides a specific example. Two partial neutralization reactions are possible in this instance.
- H2SO4 + OH− → HSO4−+ H2O
- HSO4− + OH− → SO42−+ H2O
- Overall: H2SO4 + 2OH− → SO42−+ 2H2O
After an acid AH has been neutralized there are no molecules of the acid (or hydrogen ions produced by dissociation of the molecule) left in solution.
When an acid is neutralized the amount of base added to it must be equal the amount of acid present initially. This amount of base is said to be the equivalent amount. In a titration of an acid with a base, the point of neutralization can also be called the equivalence point. The quantitative nature of the neutralization reaction is most conveniently expressed in terms of the concentrations of acid and alkali. At the equivalence point:
- volume (acid) × concentration (H+ ions from dissociation) = volume (base) × concentration (OH− ions)
In general, for an acid AHn at concentration c1 reacting with a base B(OH)m at concentration c2 the volumes are related by:
- n v1 c1 = m v2 c2
An example of a base being neutralized by an acid is as follows.
- Ba(OH)2 + 2H+ → Ba2+ + 2H2O
The same equation relating the concentrations of acid and base applies. The concept of neutralization is not limited to reactions in solution. For example, the reaction of limestone with acid such as sulfuric acid is also a neutralization reaction.
- [Ca,Mg]CO3(s) + H2SO4(aq) → (Ca2+, Mg2+)(aq) + SO42−(aq) + CO2(g) + H2O
Such reactions are important in soil chemistry.
Strong acids and strong bases
- HCl(aq) → H+(aq) + Cl−(aq)
- NaOH(aq) → Na+(aq) + OH−(aq)
Therefore, when a strong acid reacts with a strong base the neutralization reaction can be written as
- H+ + OH− → H2O
For example, in the reaction between hydrochloric acid and sodium hydroxide the sodium and chloride ions, Na+ and Cl−take no part in the reaction. The reaction is consistent with the Brønsted–Lowry definition because in reality the hydrogen ion exists as the hydronium ion, so that the neutralization reaction may be written as
- H3O+ + OH− → H2O + H2O
When a strong acid is neutralized by a strong base there are no excess hydrogen ions left in the solution. The solution is said to be neutral as it is neither acidic nor alkaline. The pH of such a solution is close to a value of 7; the exact pH value is dependent on the temperature of the solution.
Weak acids and strong bases
- AH + H2O ⇌ H3O+ + A−
Acetic acid is an example of a weak acid. The pH of the neutralized solution is not close to 7, as with a strong acid, but depends on the acid dissociation constant (pKa) of the acid. The pH at the end-point or equivalence point in a titration may be easily calculated. At the end-point the acid is completely neutralized so the analytical hydrogen ion concentration, TH, is zero and the concentration of the conjugate base, A−, is effectively equal to the analytical concentration of the acid; writing AH for the acid, [A−] = TA. Defining the acid dissociation constant, pKa, as
- [HA] = Ka[A−][H+]; pKa = −log10Ka
and the self-dissociation constant for water, Kw, as
- Kw = [H+][OH−]; pKw = −log10Kw
the equation for mass-balance in hydrogen ions is easy to write down.
- TH = [H+] + Ka[A−][H+] − Kw/[H+]
The term Kw/[H+] is equal to the concentration of hydroxide ions. At neutralization, TH is zero.
- [H+] + Ka[A][H+] − Kw/[H+] = 0
- [H+]2 + KaTA[H+]2 − Kw = 0
- [H+]2 = Kw/1 + KaTA
- log [H+] = 1/2 log Kw − 1/2 log (1 + KaTA)
- pH = 1/2 pKw − 1/2 log (1 + TA/Ka)
In most circumstances the term 1 + TA/Ka is much larger than 1, and is equal to TA/Ka to a good approximation.
- pH ≈ 1/2 (pKw + pKa − log TA )
This equation explains the following facts:
- The pH at the end-point depends mainly on the strength of the acid, pKa.
- The pH at the end-point also depends on the concentration of the acid, TA.
- The pH rises more steeply at the end-point as the acid concentration increases.
When a weak acid is titrated with a strong base the end-point occurs at pH greater than 7. Therefore, the most suitable indicator to use is one, like phenolphthalein, that changes color at high pH.
Weak bases and strong acids
The situation is analogous to that of weak acids and strong bases.
- H3O+ + B ⇌ H2O + BH+
The pH of the neutralized solution depends on the acid dissociation constant of the base, pKa, or, equivalently, on the base association constant, pKb.
The most suitable indicator to use for this type of titration is one, like Methyl orange, that changes color at low pH.
Weak acids and weak bases
When a weak acid reacts with an equivalent amount of a weak base complete neutralization does not occur.
- AH + B ⇌ A− + BH+
The concentrations of the species in equilibrium with each other will depend on the equilibrium constant, K, for the reaction, which can be defined as follows.
- [A−][BH+] = K[AH][B]
Given the association constants for the acid (Ka) and the base (Kb).
- A− + H+ ⇌ AH; [AH] = Ka[A−][H+]
- B + H+ ⇌ BH+; [BH+] = Kb[B][H+]
it follows that K = Ka/Kb.
A weak acid cannot be neutralized by a weak base, and vice versa.