Chemistry · MU Pleven

Autoionization of water, ion product of water, pH, methods to measure pH.

Covers the auto-ionisation of water and how folding the constant water concentration into the equilibrium expression gives the ion product Kw, then defines pH as the negative base-ten logarithm of hydronium concentration. Moves on to weak acids, their partial dissociation, Ka and pKa, and derives the square-root approximation for calculating hydronium concentration and pH of a weak acid solution.

Inside: 5 sections · 10 formulas · 16 flashcards · 2 min

Topic 16 of 47 in the MU Pleven syllabus for Chemistry (Medicine, Year 1). The free AceMed note covers it in 5 sections, about 326 words (2 min of reading), with 16 flashcards.

In this note

  1. Autoionization of water, ion product of water, pH, methods to measure pH.
  2. Auto-ionisation of water
  3. The pH scale
  4. Weak acids
  5. Calculating the pH of weak acids

Auto-ionisation of water

Water can behave both as an acid and a base: in aqueous solution a small fraction of the water molecules undergo acid-base reactions with each other to produce the equilibrium:

This is known as the auto-ionisation of water. The equilibrium expression for this reaction can be written:

The concentration of water under these conditions is not changed significantly, since the proportion of water which dissociates into its ions is small. The water concentration can thus be assumed to be constant and it can be incorporated into Kc:

The pH scale

The acidity and alkalinity of a solution is often measured by its pH.

Strong Acids The pH of a solution is the negative logarithm to the base ten of the H₃O⁺ concentration. pH

Weak acids

A weak acid is one which only dissociates partially in water. It reaches equilibrium as follows:

C-x x x

A weak acid is one which is a worse proton donor than H₃O⁺ Examples are CH₃COOH (ethanoic acid), NaHSO₄ (sodium hydrogensulphate) and CO₂ (carbon dioxide).

The equilibrium expression for the dissociation of a weak acid is as follows:

The water is in large excess, so its concentration is not significantly affected by this equilibrium and so can be assumed to remain constant. The equilibrium expression can thus be re-written:

Ka is known as the acid dissociation constant for the acid and has units of moldm-3.

It is often quoted as pK_a = -\log_{10}K_a.

The larger the Ka, the greater the degree to which the acid dissociates into its ions and the stronger the acid.

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