Table of Contents
What Is a Neutralization Reaction?
In Brønsted acid–base theory, a neutralization reaction is a reaction in which a proton ($\mathrm{H^+}$) transferred from an acid is taken up by a base, forming a conjugate acid–base pair and, in aqueous solution, typically leading to the formation of water and a salt.
Idealized for strong acid and strong base in water:
$$
\mathrm{H^+(aq) + OH^-(aq) \rightarrow H_2O(l)}
$$
A general molecular form:
$$
\mathrm{HA + B \rightarrow BH^+ + A^-}
$$
In aqueous solution with a hydroxide base:
$$
\mathrm{HA + BOH \rightarrow BA + H_2O}
$$
Here $\mathrm{HA}$ is an acid, $\mathrm{BOH}$ a base containing hydroxide, and $\mathrm{BA}$ the salt produced.
“Neutralization” does not always mean the final solution has exactly $\mathrm{pH}=7$, but it does mean that the acid’s $\mathrm{H^+}$ and the base’s $\mathrm{OH^-}$ (or other basic sites) have reacted according to stoichiometry.
Typical Examples of Neutralization
Strong Acid + Strong Base
These reactions are essentially complete in aqueous solution.
Example:
$$
\mathrm{HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)}
$$
Net ionic equation:
$$
\mathrm{H^+(aq) + OH^-(aq) \rightarrow H_2O(l)}
$$
Key features:
- Reaction goes (for practical purposes) to completion.
- At equivalence, the solution is (approximately) neutral: $\mathrm{pH \approx 7}$ at $25^\circ\mathrm{C}$.
- The heat released per mole of water formed is nearly constant for all strong acid–strong base pairs (in dilute aqueous solution).
Strong Acid + Weak Base
Example with ammonia:
$$
\mathrm{HCl(aq) + NH_3(aq) \rightarrow NH_4Cl(aq)}
$$
Ionic form:
$$
\mathrm{H^+(aq) + NH_3(aq) \rightarrow NH_4^+(aq)}
$$
Characteristics:
- The base is weak, so the conjugate acid (e.g. $\mathrm{NH_4^+}$) is relatively stronger and can donate $\mathrm{H^+}$ to water.
- The solution at equivalence is acidic ($\mathrm{pH<7}$).
Weak Acid + Strong Base
Example:
$$
\mathrm{CH_3COOH(aq) + NaOH(aq) \rightarrow CH_3COONa(aq) + H_2O(l)}
$$
Ionic net form:
$$
\mathrm{CH_3COOH(aq) + OH^-(aq) \rightarrow CH_3COO^-(aq) + H_2O(l)}
$$
Features:
- The conjugate base $\mathrm{CH_3COO^-}$ is basic.
- At equivalence, the solution is basic ($\mathrm{pH>7}$).
Weak Acid + Weak Base
Example:
$$
\mathrm{CH_3COOH(aq) + NH_3(aq) \rightarrow CH_3COO^-(aq) + NH_4^+(aq)}
$$
Depending on the relative strengths of the weak acid and weak base, the final $\mathrm{pH}$ at equivalence may be acidic, neutral, or basic.
Ionic Perspective: Spectator Ions and Net Ionic Equations
In aqueous neutralization, many ions do not participate chemically; they are spectator ions.
Example:
$$
\mathrm{HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)}
$$
Dissociated form:
$$
\mathrm{H^+ + Cl^- + Na^+ + OH^- \rightarrow Na^+ + Cl^- + H_2O}
$$
Removing spectator ions ($\mathrm{Na^+, Cl^-}$) yields the net ionic equation:
$$
\mathrm{H^+(aq) + OH^-(aq) \rightarrow H_2O(l)}
$$
Net ionic equations highlight the essential proton-transfer step in neutralization.
Stoichiometry of Neutralization
Equivalence of Acid and Base
In a neutralization, the amount of substance of acidic protons equals the amount of basic sites at the equivalence point.
For monoprotic strong acid $\mathrm{HA}$ and monobasic strong base $\mathrm{BOH}$:
$$
n(\mathrm{H^+}) = n(\mathrm{OH^-})
$$
or
$$
c_\text{acid} \cdot V_\text{acid} = c_\text{base} \cdot V_\text{base}
$$
Where:
- $c$ = concentration (e.g. in $\mathrm{mol\,L^{-1}}$),
- $V$ = volume (e.g. in L),
- $n$ = amount of substance (mol).
Polyprotic Acids and Polybasic Bases
If acids or bases can donate/accept more than one proton, the stoichiometry changes.
Example: Sulfuric acid $\mathrm{H_2SO_4}$ (diprotic) with sodium hydroxide:
$$
\mathrm{H_2SO_4 + 2\,NaOH \rightarrow Na_2SO_4 + 2\,H_2O}
$$
Stoichiometric relationship:
- 1 mol $\mathrm{H_2SO_4}$ neutralizes 2 mol $\mathrm{NaOH}$.
In general:
$$
n_\text{acid} \cdot z_\text{acid} = n_\text{base} \cdot z_\text{base}
$$
Where:
- $z_\text{acid}$ = number of ionizable protons per acid molecule (basicity of the acid),
- $z_\text{base}$ = number of $\mathrm{OH^-}$ (or equivalent) per base formula unit (acidity of the base).
Heat of Neutralization
For strong acid–strong base reactions in dilute aqueous solution, the molar enthalpy of neutralization (per mole of water formed) is approximately constant:
$$
\Delta_\mathrm{n} H \approx -57\ \mathrm{kJ\,mol^{-1}}
$$
Example:
$$
\mathrm{HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)} \quad \Delta_\mathrm{r}H \approx -57\ \mathrm{kJ\,mol^{-1}}
$$
Reason:
- Both reactants are almost completely dissociated.
- The essential process is simply:
$$
\mathrm{H^+ + OH^- \rightarrow H_2O}
$$
For reactions involving weak acids or bases, part of the enthalpy is used (or released) in dissociation or association steps, so the observed heat of neutralization differs from this value.
Consequences:
- Neutralization is typically exothermic.
- Temperature of the solution rises during neutralization, which can be observed experimentally.
Neutralization and pH at Equivalence
Whether a neutralized solution has $\mathrm{pH=7}$ at the equivalence point depends on the strengths of the acid and base.
- Strong acid + strong base:
- Equivalence solution is approximately neutral ($\mathrm{pH \approx 7}$).
- Strong acid + weak base:
- Conjugate acid of weak base hydrolyzes, giving $\mathrm{pH<7}$.
- Weak acid + strong base:
- Conjugate base of weak acid hydrolyzes, giving $\mathrm{pH>7}$.
- Weak acid + weak base:
- $\mathrm{pH}$ at equivalence depends on the relative strengths; no general value.
Thus, “neutralization” in the acid–base sense concerns consumption of reactants according to stoichiometry, not necessarily a final $\mathrm{pH}$ of exactly 7.
Neutralization in Titrations (Conceptual Overview)
Neutralization is the underlying reaction in many acid–base titrations, which are used to determine unknown concentrations.
Basic idea:
- A solution of known concentration (titrant) is added to a solution of unknown concentration (analyte).
- The volume of titrant required to reach the equivalence point (stoichiometric neutralization) is measured.
- From the stoichiometric relationship, the concentration of the analyte is calculated.
Indicator choice and $\mathrm{pH}$ curve shape depend on whether the neutralization is:
- strong acid–strong base,
- strong acid–weak base,
- weak acid–strong base,
- or weak acid–weak base.
(Details of titration curves and indicators are covered elsewhere; here, they are only mentioned as an application of neutralization.)
Neutralization in Everyday Life and Technology
Neutralization principles are widely used:
- Stomach acid relief
- Antacids (e.g. $\mathrm{Mg(OH)_2}$, $\mathrm{CaCO_3}$) neutralize excess $\mathrm{HCl}$ in gastric juice, forming salts, water, and sometimes $\mathrm{CO_2}$:
$$
\mathrm{2\,HCl + CaCO_3 \rightarrow CaCl_2 + H_2O + CO_2 \uparrow}
$$ - Soil treatment
- Acidic soils are neutralized with lime ($\mathrm{CaCO_3}$ or $\mathrm{CaO}$) to create better conditions for plants.
- Wastewater treatment
- Industrial effluents often have to be adjusted to near-neutral $\mathrm{pH}$ by adding acids or bases before discharge.
- Corrosion and spills
- Acid spills may be treated with basic materials (e.g. sodium bicarbonate).
- Basic spills (e.g. lye) can be cautiously neutralized with weak acids (e.g. acetic acid).
In all these applications, the same fundamental process occurs: an acid and a base react so that protons are transferred and acidic/basic properties are reduced or balanced.
Precipitation During Neutralization
Sometimes the salt formed in a neutralization reaction is sparingly soluble and precipitates.
Example:
$$
\mathrm{2\,HCl(aq) + Ca(OH)_2(aq) \rightarrow CaCl_2(aq) + 2\,H_2O(l)}
$$
(no precipitate; $\mathrm{CaCl_2}$ is soluble)
In contrast:
$$
\mathrm{2\,HNO_3(aq) + Ca(OH)_2(aq) \rightarrow Ca(NO_3)_2(aq) + 2\,H_2O(l)}
$$
also fully soluble.
However, if a neutralization involved anions/cations forming an insoluble salt (e.g. $\mathrm{BaSO_4}$), a solid would form. Though usually classified separately as precipitation reactions, they can still involve neutralization when acids and bases are present.
The key point: neutralization is defined by proton transfer, regardless of whether a precipitate forms.
Summary
- Neutralization is the acid–base reaction in which $\mathrm{H^+}$ from an acid and a basic site (often $\mathrm{OH^-}$) from a base react, commonly forming water and a salt.
- The net ionic reaction for strong acid–strong base neutralization in water is:
$$
\mathrm{H^+(aq) + OH^-(aq) \rightarrow H_2O(l)}
$$ - Stoichiometric equality of acidic and basic equivalents defines the equivalence point.
- Strong acid–strong base neutralizations yield solutions with $\mathrm{pH \approx 7}$ at equivalence; other combinations do not necessarily.
- Neutralization reactions are usually exothermic, with a nearly constant enthalpy for strong acid–strong base pairs.
- They play an essential role in titrations and in numerous everyday and industrial processes such as antacid action, soil and wastewater treatment, and spill neutralization.