A base is a chemical species that can accept a proton, increase the concentration of hydroxide ions in aqueous solution under the Arrhenius definition, or donate an electron pair under the Lewis definition. As with acids, several definitions provide complementary descriptions. The Arrhenius definition describes a base as a substance that increases hydroxide ion concentration in water. The Bronsted Lowry definition describes a base as a proton acceptor. The Lewis definition describes a base as an electron pair donor.
Hydroxide ions, ammonia, carbonate ions, bicarbonate ions and many amines can act as bases under appropriate conditions. Bases are not necessarily substances that contain hydroxide groups. Ammonia, for example, contains no hydroxide ion but reacts with water by accepting a proton and producing ammonium and hydroxide ions. Basicity is therefore fundamentally connected to chemical interactions rather than to the mere presence of a particular structural group.
In the Bronsted Lowry model, a base accepts a proton from an acid. When a base accepts a proton, it becomes its conjugate acid. For example, ammonia accepts a proton from water to form ammonium. The water molecule in that reaction acts as the acid because it donates a proton. This demonstrates that water is amphiprotic and can act either as an acid or as a base depending on the reaction partner.
The strength of a Bronsted Lowry base is related to the stability of the species after protonation and to the corresponding conjugate acid. A stronger base generally has a weaker conjugate acid under the same solvent conditions. Proton transfer reactions are therefore described as coupled acid base processes rather than as isolated properties of individual molecules. The direction of proton transfer depends on relative acid strengths, solvent effects, concentration and thermodynamic stability.
In the Arrhenius framework, bases increase hydroxide ion concentration in aqueous solution. Soluble metal hydroxides such as sodium hydroxide and potassium hydroxide dissociate substantially in water and provide high concentrations of hydroxide ions. Calcium hydroxide is also a strong base in the portions that dissolve, although its limited solubility means that a saturated solution contains less total dissolved material than a similarly concentrated soluble hydroxide. The terms strong base and concentrated base are therefore distinct.
A strong base has a high degree of ionisation or proton accepting capability in the relevant solvent, whereas concentration refers to the amount of solute present per unit volume. A dilute strong base can have a lower hydroxide concentration than a concentrated weak base under some circumstances. This distinction parallels the difference between acid strength and acid concentration.
The Lewis definition broadens the concept of bases beyond proton transfer. A Lewis base is an electron pair donor, meaning that it can provide a pair of electrons to a Lewis acid to form a coordinate covalent bond. Ammonia is a classic Lewis base because its nitrogen atom has a lone pair that can be donated to an electron deficient species. Water, hydroxide, chloride and many organic molecules can also behave as Lewis bases.
Lewis basicity is central to coordination chemistry, catalysis and organic reaction mechanisms. Nucleophiles are often Lewis bases because they donate electron density to electrophilic centres, although the terms are not identical in every context. A Lewis base can participate in reactions even when no proton is transferred, which illustrates why the Lewis definition encompasses a broader range of chemical behaviour than the Bronsted Lowry definition.
Basicity is influenced strongly by molecular structure, solvent and the availability of electron density. A species with a readily available lone pair can be a strong Lewis base, but its Bronsted Lowry basicity also depends on how favourable protonation is in the relevant solvent. Resonance can reduce basicity by delocalising a lone pair, while electron donating substituents can sometimes increase electron density and enhance proton accepting behaviour.
Solvation can stabilise either the base or its conjugate acid and thereby alter the observed basicity. In aqueous solution, the behaviour of small ions is strongly affected by hydration. Steric effects can also influence how easily a base interacts with a proton or other electrophile. Consequently, basicity cannot always be predicted from a simple rule such as more lone pairs meaning a stronger base. The identity of the solvent and the specific reaction must be considered when making quantitative comparisons.
Bases have numerous practical applications. Sodium hydroxide is widely used in chemical manufacturing, soap production, paper processing, cleaning and water treatment. Ammonia is important in fertiliser production and also acts as a chemical base in many industrial and laboratory reactions. Carbonate and bicarbonate compounds are used in buffering, baking, water chemistry and numerous manufacturing processes. In biology, bases participate in maintaining acid base balance, nucleic acid chemistry, enzyme mechanisms and proton transfer reactions.
A buffer contains a weak acid and its conjugate base or an analogous pair and resists large changes in pH when limited quantities of acid or base are added. Basic substances can neutralise acids, but neutralisation does not necessarily produce a solution with exactly pH seven because the resulting solution may contain conjugate ions that themselves react with water. The final pH depends on the identities and concentrations of the reacting species.
Strong bases can be highly corrosive because hydroxide ions and other strongly basic species can react with biological tissues and chemical materials. Contact with concentrated sodium hydroxide, potassium hydroxide or other caustic substances can cause severe chemical injury. The word caustic refers particularly to the destructive chemical action of strong bases, while basic simply describes acid base behaviour. These concepts should not be treated as synonyms. Weak bases can also be hazardous depending on concentration, volatility and chemical properties.
Base solutions can be measured using pH, although pH is not a direct universal measure of base strength. A strong base may produce a very high pH when sufficiently concentrated, but pH also depends on dilution and solvent conditions. Base strength is more fundamentally related to equilibrium constants such as Kb or the pKa of the conjugate acid.
The historical understanding of bases developed alongside observations of substances that neutralised acids and produced characteristic reactions. The term alkali historically referred to substances that produce alkaline solutions, particularly soluble bases, and is derived through Arabic and later European terminology associated with ashes and alkaline materials. Arrhenius formalised the aqueous definition of bases in the late nineteenth century, Bronsted and Lowry expanded the concept to proton transfer in 1923, and Lewis broadened it further to electron pair donation.
Modern chemistry therefore treats base as a context dependent chemical concept rather than merely a synonym for hydroxide. Bases can be proton acceptors, hydroxide producing substances in water or electron pair donors. Understanding these distinctions is essential because some species that are weak or negligible bases in one environment can behave as strong bases or Lewis bases under different conditions, while a substance's concentration, corrosiveness and alkalinity are separate properties that should not be conflated with fundamental base strength.
