Brønstedhappoja

Brønstedhappoja, in Brønsted–Lowry acid–base theory, are substances that donate protons (H+) to bases during chemical reactions. The concept, named after the Danish-Norwegian chemist Johannes Brønsted, provides a general definition of acidity that applies beyond water and across different solvents. In a typical Brønsted–Lowry reaction, an acid transfers a proton to a base, which accepts the proton.

Common examples of Brønstedhappoja include strong acids such as hydrochloric acid (HCl), sulfuric acid (H2SO4), and nitric acid (HNO3), as well as weaker acids like acetic acid (CH3CO2H) and carbonic acid (H2CO3). The substance that remains after donating a proton is called the conjugate base; the base that accepts the proton forms its conjugate acid. Acid–base reactions are thus described by conjugate acid–base pairs.

Acid strength is quantified by the dissociation constant (Ka) or its logarithmic form, pKa. Strong Brønstedhappoja have large Ka values and low pKa, dissociating readily to donate protons, especially in water. Weak Brønstedhappoja dissociate only to a limited extent. The measured strength of an acid depends on the solvent; the same substance can behave as a stronger or weaker acid in different solvents.

Relation to other concepts: Brønstedhappoja provide a broader framework than the Arrhenius definition, which is specific to hydrogen ion production in water. The Brønsted–Lowry view is compatible with Lewis acid–base theory, in which acids are electron-pair acceptors, though the focus here is on proton transfer.

Historically, the theory was developed in the 1920s by Brønsted and Lowry and remains foundational for understanding acid–base chemistry in chemistry and biochemistry.