Lesson 7 of 12 · 9 min
Ionic equilibrium and acid-base concepts
NCERT § "Ionic Equilibrium in Solution"; § "Acids, Bases and Salts"
Stir salt into water and a conductivity tester's bulb glows brightly. Stir in acetic acid, the acid in vinegar, at the same concentration and it barely glows. Both dissolve. Why the difference?
The lesson in notes
In short
Faraday classified substances as electrolytes (conduct in solution) and non-electrolytes; strong electrolytes ionise almost completely, while weak electrolytes ionise partly and set up an equilibrium with unionised molecules.
Arrhenius: acids give H⁺ (present as H₃O⁺) and bases give OH⁻ in water; the definition is limited to aqueous solutions.
A bare proton is too reactive to exist alone in water; it binds to water as the hydronium ion H₃O⁺, and larger clusters such as H₉O₄⁺ also form.
Brønsted-Lowry: an acid is a proton donor and a base is a proton acceptor.
An acid and the base formed when it loses a proton are a conjugate pair; a strong acid has a very weak conjugate base, and a weak acid a relatively strong one.
Water is amphoteric: it acts as a base towards HCl and as an acid towards NH₃.
Lewis: an acid accepts an electron pair and a base donates one; BF₃, AlCl₃ and cations like Co³⁺ and Mg²⁺ are Lewis acids, while NH₃, OH⁻ and F⁻ are Lewis bases.
Lewis acids need not contain hydrogen, which is how the theory covers electron-deficient molecules like BF₃.