An acid is a chemical species that can participate in reactions by donating a proton, accepting an electron pair in the Lewis definition, or increasing the concentration of hydrogen ions in aqueous solution depending on the theoretical framework being used. Three major definitions are commonly taught. The Arrhenius definition describes an acid as a substance that increases the concentration of hydrogen ions in water. The Bronsted Lowry definition describes an acid as a proton donor.
The Lewis definition describes an acid as an electron pair acceptor. These definitions overlap but are not identical in scope. Hydrogen chloride acts as an Arrhenius acid in water and as a Bronsted Lowry proton donor, while species such as boron trifluoride can act as Lewis acids without containing a hydrogen atom that they donate as a proton. Acids can therefore be understood through several complementary models rather than through a single characteristic. Common acidic substances include hydrochloric acid, sulfuric acid, nitric acid, acetic acid and citric acid, although their strengths and chemical behaviours differ substantially.
In the Bronsted Lowry framework, an acid donates a proton to a base. A proton in ordinary acid base chemistry means a hydrogen ion, H plus, although in aqueous solution free H plus does not exist as an isolated bare particle for meaningful periods and is strongly associated with water molecules, commonly represented through hydronium ions. When an acid donates a proton, it becomes its conjugate base.
For example, hydrochloric acid donates a proton to water, producing hydronium and chloride ions. The acid base relationship is therefore described through conjugate pairs. A stronger proton donor generally has a weaker conjugate base under the same conditions because the equilibrium favours the side containing the relatively more stable acid base pair. Water can act as either an acid or a base depending on its reaction partner, demonstrating that acidity is relational rather than an isolated property. Amphiprotic species such as bicarbonate can donate or accept protons and therefore participate in multiple acid base equilibria.
In the Arrhenius framework, acids increase the concentration of H plus in aqueous solution. Strong mineral acids such as hydrochloric acid, nitric acid and sulfuric acid are commonly encountered in laboratory and industrial chemistry. Their aqueous behaviour depends on dissociation, solvent interactions and concentration. Sulfuric acid is diprotic, meaning that it can donate two protons through successive ionisation steps, although the first ionisation is much stronger than the second.
Carbonic acid is a weak acid associated with dissolved carbon dioxide and participates in important environmental and biological equilibria. Organic acids often contain carboxyl groups and include acetic acid, formic acid, lactic acid and citric acid. The presence of a hydrogen atom in a molecule does not automatically make the molecule an acid. What matters is whether the relevant hydrogen can participate in proton transfer under the specified conditions and how stable the resulting conjugate base is.
Acid strength is a thermodynamic concept describing the extent to which an acid donates protons in a particular solvent, commonly water. For a weak acid HA in water, the equilibrium can be represented as HA plus H2O reversible reaction H3O plus plus A minus. The acid dissociation constant Ka quantifies the equilibrium position. A larger Ka corresponds to greater dissociation and therefore a stronger acid in the same solvent and under comparable conditions.
The pKa is defined as negative logarithm base ten of Ka, so a lower pKa generally indicates a stronger acid. Strength must be distinguished from concentration. A concentrated weak acid can contain more acid molecules per litre than a dilute strong acid, yet the strong acid has a greater tendency to dissociate. Similarly, acid strength should not be confused with corrosiveness. Corrosive effects depend on concentration, chemical reactivity, temperature, contact conditions and the specific material involved.
The molecular structure of an acid strongly influences its acidity. Factors such as bond strength, electronegativity, resonance stabilisation, inductive effects, solvent interactions and molecular geometry can alter the stability of the conjugate base. An acid becomes more favourable as a proton donor when the resulting conjugate base is relatively stable. Resonance can distribute negative charge across several atoms and therefore stabilise certain conjugate bases. Electronegative atoms can withdraw electron density through inductive effects, sometimes increasing acidity of nearby functional groups.
In oxoacids, the number and electronegativity of oxygen atoms can strongly influence acidity. In aromatic and organic chemistry, substituents can alter acidity through electronic effects. These structural principles explain why molecules containing superficially similar hydrogen atoms can have very different acid strengths. Acidity is therefore not determined simply by the presence of hydrogen but by the energetics of proton transfer and the stability of the resulting species in a particular environment.
Acids have extensive practical and biological functions. Gastric acid helps create the strongly acidic environment of the stomach and participates in digestion and defence against microorganisms. Organic acids participate in metabolism, fermentation and biochemical regulation. In industry, acids are used for metal processing, mineral treatment, fertiliser production, polymer manufacture, battery chemistry, cleaning and synthesis. Acid base reactions are also fundamental to analytical chemistry, including titration methods used to determine unknown concentrations.
Acids react with bases in neutralisation reactions and can react with certain metals, carbonates and other substances. Strong acids require careful handling because concentrated solutions can cause severe chemical burns and can react vigorously with water or other materials. Acidity also plays major roles in environmental chemistry, including atmospheric processes, freshwater chemistry, soil chemistry and ocean carbonate equilibria.
Acids can be classified in several different ways. Monoprotic acids donate one proton per molecule under the relevant reaction scheme, while diprotic and polyprotic acids can donate multiple protons in successive stages. Strong acids are substantially dissociated in water, whereas weak acids establish appreciable equilibria between dissociated and undissociated forms. Some acids are oxidising agents, while others are not, so acidity and oxidising power are separate properties.
Lewis acids include positively charged ions and electron deficient molecules that accept electron pairs from Lewis bases. Some Lewis acids are also involved in catalysis, coordination chemistry and organic synthesis. Acidity is also solvent dependent. A substance that behaves as a strong acid in one solvent may behave differently in another because solvent molecules influence proton transfer and solvation. This is known as the solvent effect and is one reason why acid strength must always be interpreted in context.
The scientific understanding of acids developed from practical observations of sour taste, corrosive behaviour and reactions with metals and carbonates, followed by increasingly formal definitions. The word acid comes from the Latin acidus, meaning sour, reflecting the sensory property historically associated with many acidic substances. Robert Boyle studied characteristic properties of acids and bases in the seventeenth century, while Svante Arrhenius later developed the ionisation theory that provided the Arrhenius definition.
Johannes Bronsted and Thomas Lowry independently developed the proton transfer definition in 1923, and Gilbert N Lewis developed the electron pair definition of acids and bases. Modern chemistry uses all three frameworks where appropriate. An acid is therefore best understood not as simply a substance that tastes sour or burns materials, but as a chemical species whose interactions can be described through proton donation, hydrogen ion activity or electron pair acceptance depending on the theoretical context.
