Neutralization (chemistry)

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Neutralization reaction between sodium hydroxide and hydrochloric acid. Indicator agent bromothymol blue.

In chemistry, neutralization (or neutralisation, see spelling differences) is a chemical reaction in which an acid and a base react to form a salt. Water is frequently, but not necessarily, produced as well. Neutralizations with Arrhenius acids and bases always produce water where acid–alkali reactions produce water and a metal salt.

Often, neutralization reactions are exothermic (the enthalpy of neutralization). For example, the reaction of sodium hydroxide and hydrochloric acid. However, forms of endothermic neutralization do exist, such as the reaction between sodium bicarbonate (baking soda) and acetic acid (vinegar).

Neutralization reactions do not necessarily imply a resultant pH of 7.[1] The resultant pH will vary based on the respective strengths of the acid and base reactants.

Arrhenius acids and bases[edit]


Svante Arrhenius defined an acid as a substances that produces hydrogen ions, or protons (H+
) in aqueous solutions. Hydrochloric acid (HCl) and sulfuric acid (H
) are common examples of Arrhenius acids because, when they dissociate into their ions, they increase the amount of H+ in the solution:

(aq) + Cl
(aq) + HSO

Arrhenius defined a base as a substance that produces hydroxide (OH
) in aqueous solutions. Arrhenius's definition of a base was limited because along with produced hydroxide ions in solution, the base must contain the hydroxide group in its formula. Potassium hydroxide (KOH) and caesium hydroxide (CsOH) are common examples of Arrhenius bases because, when they dissociate into their ions, they increase the amount of OH- in the solution:

KOH(aq) → K+
(aq) + OH
CsOH(aq) → Cs+
(aq) + OH

Net ionic equation[edit]

When an acid reacts with an equal amount of a base, the word "neutralization" is used to describe the result because the acid and base properties of H+ and OH- are destroyed or neutralized. In the reaction, H+ and OH- combine to form HOH - more commonly written as H2O, the water molecule. Thus, acid–base neutralization reactions can be simplified to the net ionic equation:

+ OH

It is important to realize that this representation is somewhat inaccurate, as it is known that the hydrogen ion (H+) does not actually occur by itself, but instead as the hydronium ion (H3O+). As seen in the following reaction, the H+, or protons, cause molecules of water undergo protonation to form the hydronium ion:

H+ + H2O → H3O+

Considering the hydronium ion, the actual net ionic reaction occurring is:

H3O+ + OH- → 2H2O

General neutralization equation[edit]

A neutralization reaction is a type of double replacement reaction. Typically, the resulting solution produced by the reaction consists of a salt and water. The general formula for acid–base neutralization reactions can be written as

acid + base → salt + water
HA + BOH → BA + H2O

where HA represents the Arrhenius acid, BOH represents the Arrhenius base, and BA is the salt produced. Notice how, typical of a double replacement reaction, the cations and anions of the substances merely switch places.

An example reaction of this form is the reaction between hydrochloric acid (HCl) and sodium hydroxide (NaOH):

NaOH + HCl → NaCl + H2O

Water and sodium chloride, or common table salt are produced.

The following are other examples of acid-base neutralization reactions

H2SO4 + 2NH4OH → (NH4)2SO4 + 2H2O
H2CO3 + 2NaOH → Na2CO3 + 2H2O
3HCl + Al(OH)3 → AlCl3 + 3H2O


An acid–alkali reaction is a neutralization reaction which is considered a special case of an acid–base reaction, where the base used is also an alkali. An alkali is a base that contains a metal from Group 1 (the alkali metals: lithium, sodium, potassium, rubidium, or caesium) or from Group 2 (the alkaline earth metals: beryllium, magnesium, calcium, strontium, or barium). When an acid reacts with an alkali it forms a metal salt and water.

Alkalis may be defined as soluble bases, which means they must be able to dissolve in water. Therefore, one may also speak of hydroxide bases that dissolve in water, and thus these would also be alkalis. Some examples, then, of alkalis would be sodium hydroxide (NaOH), potassium hydroxide (KOH), magnesium hydroxide (Mg(OH)2), and calcium hydroxide (Ca(OH)2). Note that only hydroxides with an alkali metal — column 1 — are very soluble in water; hydroxides with an alkaline earth metal — column 2 — are not as soluble. Some sources will even say the alkaline earth metal hydroxides are insoluble.[2]

However, alkalis may also have a broader definition that includes carbonates (CO32−) bonded to a Group 1 metal, an ammonium ion (NH4+), or an amine (NHx radical) as the positive ion. Examples of alkalis would then also include Li2CO3, Na2CO3, and (NH4)2CO3.

Non-aqueous reactions[edit]

In non-aqueous reactions, water is less likely to be formed; however, there is always a donation of protons (see Brønsted-Lowry acid-base theory). Since a variety of definitions of acids and bases exist, a variety of reactions may be considered neutralization reactions. All of the following may be considered neutralization reactions under different definitions:

HCl + NaOH → NaCl + H2O
2HCl + Mg → MgCl2 + H2
2HCO2H + MgO → Mg(HCO2)2 + H2O
HF + NH3 → NH4F

Resultant pH[edit]

Neutralization reactions do not necessarily imply a resultant pH of 7.[1] In the case that a strong acid and strong base participate in a neutralization reaction, the resultant pH will be 7. For example, the strong acid, HCl, and the strong base, NaOH, react to give water and a salt, NaCl:

HCl + NaOH → H2O + NaCl

Since there is no net change in the concentrations of either H3O+ or OH-, the end pH is 7.

If a weak acid and a strong base participate in a neutralization reaction, the resultant pH will be greater than 7. For example, the weak acid, CH3COOH, and the strong base, NaOH, react to give water, Na+, and acetate, CH3COO-:

CH3COOH + NaOH → Na+ + H2O + CH3COO-

Na+ behaves as a spectator ion. However, acetate is a weak base that hydrolyzes water to give OH- ions.


Thus, the resultant solution is basic.

If a weak base and a strong acid participate in a neutralization reaction, the resultant pH will be less than 7. For example, the weak base, CN-, and the strong acid, HCl, react to give Cl- and hydrocyanic acid, HCN:

CN- + HCl → Cl- + HCN

Cl- behaves as a spectator ion. However, hydrocyanic acid is a weak acid that hydrolyzes water to give H3O+ ions.

HCN + H2O CN- + H3O+

Thus, the resultant solution is acidic.

Finally, if a weak acid and a weak base participate in a neutralization reaction, the resultant pH will depend on the relative strength of the acid and base reactants. For example, the weak base, CN-, and the weak acid, CH3COOH, react to give HCN and CH3COO-. Because CH3COOH (pKa=4.75) is a stronger acid than HCN (pKa=9.2), the equilibrium is driven to the right, assuming equimolar initial concentrations of the weak acid and weak base.


The acetate ions further react with water to give acetic acid and OH-.


In this particular example, the resultant solution is basic. However, this is not a general rule for all neutralization reactions between a weak acid and a weak base.


Equal numbers of moles of acid and base are needed for neutralization reactions. Hence, the formula becomes

a × [A] × Va = b × [B] × Vb

where a is the number of acidic hydrogens and b is the constant that shows how many H3O+ ions the base can accept. [A] denotes the concentration of acid and [B], the concentration of base. Va is the volume of acid and Vb is the volume of base.


Chemical titration methods are used for analyzing acids or bases to determine the unknown concentration. Either a pH meter or a pH indicator which shows the point of neutralization by a distinct color change can be employed. Simple stoichiometric calculations with the known volume of the unknown and the known volume and molarity of the added chemical gives the molarity of the unknown.

In wastewater treatment, chemical neutralization methods are often applied to reduce the damage that an effluent may cause upon release to the environment. For pH control, popular chemicals include calcium carbonate, calcium oxide, magnesium hydroxide, and sodium bicarbonate. The selection of an appropriate neutralization chemical depends on the particular application.

There are many uses of neutralization reactions that are acid-alkali reactions. A very common use is antacid tablets. These are designed to neutralize excess gastric acid in the stomach (HCl) that may be causing discomfort in the stomach or lower esophagus. This can also be remedied by the ingestion of sodium bicarbonate (NaHCO3).

Also in the digestive tract, neutralization reactions are used when food is moved from the stomach to the intestines. In order for the nutrients to be absorbed through the intestinal wall, an alkaline environment is needed, so the pancreas produce an antacid bicarbonate to cause this transformation to occur.

Another common use, though perhaps not as widely known, is in fertilizers and control of soil pH. Slaked lime (calcium hydroxide) or limestone (calcium carbonate) may be worked into soil that is too acidic for plant growth.[3] Fertilizers that improve plant growth are made by neutralizing sulfuric acid (H2SO4) or nitric acid (HNO3) with ammonia gas (NH3), making ammonium sulfate or ammonium nitrate. These are salts utilized in the fertilizer.[4]

Industrially, a by-product of the burning of coal, sulfur dioxide gas, may combine with water vapor in the air to eventually produce sulfuric acid, which falls as acid rain. To prevent the sulfur dioxide from being released, a device known as a scrubber gleans the gas from smoke stacks. This device first blows calcium carbonate into the combustion chamber where it decomposes into calcium oxide (lime) and carbon dioxide. This lime then reacts with the sulfur dioxide produced forming calcium sulfite. A suspension of lime is then injected into the mixture to produce a slurry, which removes the calcium sulfite and any remaining unreacted sulfur dioxide.[5]


  1. ^ a b Lemke, T. L. (2003). Review of Organic Functional Groups: Introduction to Medicinal Organic Chemistry (4th ed.). Lippincott Williams & Wilkins. ISBN 0-7817-4381-8. 
  2. ^ See for example: Group 2: Alkaline Earth Metals,
  3. ^ Neutralisation – 'curing acidity' – Acids, alkalis and salts – Intermediate – Experiments. Practical Chemistry. Retrieved on 2010-12-10.
  4. ^ Reversible Reactions, Chemical Equilibrium, Ammonia & Nitric acid, their Manufacture and Uses e.g. in Artificial Fertilisers
  5. ^ Zumdahl, Steven S., (2000). pp. 226, 228[full citation needed]

Further reading[edit]