Acids, bases and salts

Acids and bases have been defined in three ways, each wider than the last:

• Arrhenius: an acid gives H⁺ ions (really H₃O⁺, hydronium) in water; a base gives OH⁻ ions in water. This works only for aqueous solutions.
• Brönsted-Lowry: an acid is a proton donor, and a base is a proton acceptor. In HCl + H₂O → H₃O⁺ + Cl⁻, HCl gives a proton (acid), and water takes it (base). This also covers ions, not just molecules.
• Lewis: an acid accepts an electron pair; a base donates one. Lewis acids include BF₃, AlCl₃, H⁺ and metal ions such as Ag⁺ and Fe³⁺; Lewis bases include NH₃, OH⁻ and F⁻. This covers reactions with no protons at all.

In the Brönsted-Lowry picture, acids and bases come in conjugate pairs that differ by one proton. When HCl gives up a proton, what is left (Cl⁻) is its conjugate base; when NH₃ takes a proton, it becomes its conjugate acid NH₄⁺. Take away H⁺ to get the conjugate base; add H⁺ to get the conjugate acid.

A strong acid gives up its proton very readily, so its conjugate base has little wish to take it back: a strong acid has a weak conjugate base, and vice versa.

Some species can act either way, as acid or base, depending on what they meet. They are amphoteric (or amphiprotic): water (H₂O can lose H⁺ to give OH⁻, or gain H⁺ to give H₃O⁺), HCO₃⁻, HSO₄⁻ and NH₃.