Chemistry · Ch 6 — Equilibrium
Ionic Equilibrium in Solution
Ionic Equilibrium in Solution
6.9 Ionic Equilibrium in Solution
You have already seen how changing the concentration of a reactant or product can shift the position of an equilibrium. One example you encountered involved ions:
But there are many equilibria that involve only ions in solution. Before we study those, we need to understand what makes a substance conduct electricity in water.
A simple experiment shows the difference: sugar dissolved in water does not conduct electricity, but common salt (sodium chloride) dissolved in water does. Moreover, the conductance increases as you add more salt. This behaviour led Michael Faraday to divide substances into two broad classes based on their ability to conduct electricity in aqueous solution.
Electrolytes are substances whose aqueous solutions conduct electricity. Non-electrolytes are those whose aqueous solutions do not.
Faraday further split electrolytes into two types: strong electrolytes and weak electrolytes.
The key difference is the extent of dissociation (splitting into ions) when the substance dissolves.
A strong electrolyte dissociates almost completely into ions when it dissolves in water. For example, when sodium chloride (NaCl) dissolves, nearly every NaCl unit separates into Na⁺ and Cl⁻ ions. The solution contains essentially no unionized NaCl molecules.
A weak electrolyte, on the other hand, dissociates only partially. Most of the substance remains as neutral molecules, with only a small fraction forming ions. Acetic acid (CH₃COOH) is a classic example: in water, less than 5% of the acetic acid molecules ionize to form acetate ions (CH₃COO⁻) and hydronium ions (H₃O⁺). The vast majority stay as unionized CH₃COOH molecules.
Do not confuse "strong" with "concentrated." A strong electrolyte can be dilute; it is still fully dissociated. A weak electrolyte can be concentrated; it is still only partially dissociated.
Why does this happen? The polar water molecules surround and stabilize the ions, pulling them apart from the crystal lattice (for salts) or from each other (for molecular acids and bases). Figure 6.10 in the textbook illustrates this: Na⁺ and Cl⁻ ions become hydrated — each ion is surrounded by water molecules oriented with the appropriate charged end toward the ion — which lowers the energy of the ions and makes the dissociated state more stable.
For a weak electrolyte, an equilibrium is established between the unionized molecules and the ions they produce. This type of equilibrium, involving ions in aqueous solution, is called ionic equilibrium. …