What Is an Acid in Chemistry? Definition and Examples

What Is an Acid in Chemistry
There are different types of acids. By definition, an acid donates hydrogen ions or protons or accepts an electron pair.

In chemistry, an acid is a chemical species that donates hydrogen ions or protons or accepts an electron pair. Acids react with bases and some metals via a neutralization reaction that forms a salt. They have a pH less than 7 and taste sour. The word acid comes from the Latin word acidus, which means “sour.” Take a closer look at the definition of acids, examples, and their properties.

  • An acid is a hydrogen ion or proton donor or an electron pair acceptor.
  • Not all compounds containing hydrogen are acids.
  • Acids have a pH less than 7, turn litmus paper red, taste sour, and react with bases.
  • Examples of acids include hydrochloric acid (HCl), sulfuric acid (H2SO4), and acetic acid (CH3COOH).

Acid Definition and Examples

There are three ways of defining an acid, based on the three main acid-base theories. Some chemicals are acids under one definition, but not another.

  • Arrhenius acid: An Arrhenius acid increases the hydrogen ion (H+) concentration of an aqueous solution. Since hydrogen ions attach to water molecules, what this really means is an Arrhenius acid increases the hydronium ion (H3O+) concentration. An Arrhenius acid has the element hydrogen (H) as part of its chemical formula. Examples include hydrochloric acid (HCl), nitric acid (HNO3), and acetic acid (CH3COOH).
  • Brønsted-Lowry acid: A Bronsted-Lowry acid is a proton donor. Since a hydrogen ion and a proton are essentially the same, all Brønsted acids contain hydrogen. The difference between these acids and Arrhenius acids is that they can react in solvents besides water.
  • Lewis acid: A Lewis acid accepts an electron pair to form a covalent bond. All Arrhenius and Bronsted-Lowry acids are Lewis acids. But, there are Lewis acids which are not Arrhenius or Bronsted-Lowry acids. For example, BF3, AlCl3, and Mg2+ are Lewis acids, but are not acids by the other definitions. Boric acid (H3BO3) has hydrogen in its formula, but it is only a Lewis acid because it does not dissociate in water, but does accept an electron pair.

Most of the time, when chemists refer to an acid, they mean a Brønsted-Lowry acid. This definition includes all of the Arrhenius acids, plus it extends to solvents besides water.

Amphoteric Species

An amphoteric compound acts as either an acid or a base, depending on the situation. Examples include water, amino acids, and metal oxides. For example, water donates a proton when it reacts with a base, but accepts a proton when it reacts with water.

Strong and Weak Acids

The two broad categories of acids are strong acids and weak acids.

  • Strong acids completely dissociate into their ions in water (or other solvent, for Brønsted-Lowry acids). Examples include hydrochloric acid (HCl) and nitric acid (HNO3). There are only seven common strong acids.
  • Weak acids incompletely dissociate into their ions in a solvent, so the solution contains both the weak acid and the ions. There are numerous weak acids. Examples include acetic acid (CH3COOH), nitrous acid (HNO2), and formic acid (HCOOH).
Common Strong AcidFormula
hydrochloric acidHCl
nitric acidHNO3
sulfuric acidH2SO4
hydrobromic acidHBr
hydroiodic acidHI
perchloric acidHClO4
chloric acidHClO3

Monoprotic vs Polyprotic

A monoprotic or monobasic acid only donates one proton per molecule. An example is hydrochloric acid (HCl).

HA (aq) + H2O (l) ⇌ H3O+ (aq) + A (aq)

A polyprotic or polybasic acid can donate more than one proton per acid molecule. There are diprotic (dibasic) acid and triprotic (tribasic acids). For example, sulfuric acid (H2SO4) is a diprotic acid that has two protons it can donate.

H2A (aq) + H2O (l) ⇌ H3O+ (aq) + HA (aq)     Ka1

HA (aq) + H2O (l) ⇌ H3O+ (aq) + A2− (aq)       Ka2

The equilibrium constant of the first dissociation (Ka1) usually is greater than the second dissociation constant (Ka2).


A superacid is any acid that is stronger than sulfuric acid. The strongest acid is fluoroantimonic acid (HSbF6). It donates protons about a billion times better than sulfuric acid.

Properties of Acids

Acids display several characteristic properties:

  • Most taste sour. (Don’t test this.)
  • Most are corrosive.
  • They have pH values less than 7.
  • Acids turn litmus paper red.
  • In water, Arrhenius acids are electrolytes. In other words, they conduct electricity in aqueous solution.
  • Arrhenius acids react with bases to form salt and water.
  • Arrhenius acids react with most metals to release hydrogen gas.


  • Finston, H.L.; Rychtman, A.C. (1983). A New View of Current Acid-Base Theories. New York: John Wiley & Sons. doi:10.1002/ciuz.19830170211
  • Hall, Norris F. (March 1940). “Systems of Acids and Bases”. Journal of Chemical Education. 17 (3): 124–128. doi:10.1021/ed017p124
  • IUPAC (1997). “Acid.” Compendium of Chemical Terminology (2nd ed.). Oxford: Blackwell Scientific Publications. doi:10.1351/goldbook
  • Jensen, W.B. (1980). The Lewis Acid-Base Concepts: An Overview. New York: Wiley. ISBN 0-471-03902-0.
  • Masterton, William; Hurley, Cecile; Neth, Edward (2011). Chemistry: Principles and Reactions. Cengage Learning. ISBN 978-1-133-38694-0.