Acid base equilibrium describes how acids and bases transfer protons and settle into a balance between reactants and products. It matters because pH controls many chemical systems, including blood chemistry, ocean chemistry, food preservation, and laboratory titrations. At equilibrium, reactions have not stopped, but the forward and reverse reaction rates are equal.
This makes acid base chemistry a dynamic process rather than a one-way change.
A weak acid such as acetic acid only partially ionizes in water, so both acid molecules and ions remain in solution. The strength of an acid or base is measured by equilibrium constants such as Ka and Kb, which show how far the reaction favors products. Buffers use a weak acid and its conjugate base to resist pH changes when small amounts of acid or base are added.
Understanding these relationships helps students predict pH, compare acid strength, and explain titration curves.
Understanding Acid-Base Equilibrium
Conjugate pairs explain why acid and base strength are linked. When an acid loses a proton, the particle left behind can accept that proton again. This particle is the conjugate base.
A strong acid has a very weak conjugate base because it releases its proton readily and the reverse change is unlikely. A weak acid has a stronger conjugate base. The same pattern applies to bases and their conjugate acids.
This idea helps when comparing chemicals without memorising every reaction. A lower pKa means a larger acid ionisation constant and a stronger acid. The pKa scale is logarithmic, so a difference of one pKa unit represents a tenfold change in the acid ionisation constant.
Most weak acid and weak base calculations begin with an initial change equilibrium table. First, write the amounts present before the reaction. Next, use one variable for the amount that reacts.
Then express the equilibrium amounts using that variable. Substituting those values into the equilibrium constant gives an equation that can be solved. In many school problems, the amount reacting is very small compared with the starting concentration.
The starting concentration can then be treated as nearly unchanged. This shortcut must be checked after solving. If the change is more than about five percent of the initial amount, the shortcut is not reliable.
Units matter too. Equilibrium expressions use concentration in moles per litre, not simply the number of moles added to a flask.
Water connects acid and base calculations. At twenty five degrees Celsius, the acid ionisation constant of an acid multiplied by the base ionisation constant of its conjugate base equals the ion product of water. This relationship lets students find the strength of one member of a conjugate pair when the other is known.
It also explains why a solution cannot have independently chosen hydronium and hydroxide concentrations. If one concentration rises, the other falls. The pH scale compresses a huge range of concentrations into manageable numbers.
Because it is logarithmic, a change of one pH unit means a tenfold change in hydronium concentration. Small-looking pH differences can therefore be chemically important.
Buffers work best when appreciable amounts of both members of a conjugate pair are present. Added acid is mainly removed by the conjugate base. Added base is mainly removed by the weak acid.
The Henderson Hasselbalch equation states that pH equals pKa plus the logarithm of the concentration of conjugate base divided by the concentration of weak acid. It is most accurate when the two concentrations are not extremely small and their ratio is reasonably close to one. At equal concentrations, pH equals pKa.
During a titration, this condition occurs at the half equivalence point for a weak acid titrated with a strong base. Students should distinguish this point from the equivalence point.
At equivalence, the original acid has been used up, so the pH depends on the reaction of its conjugate base with water. Blood, cells, medicines, soil, and aquarium water all depend on controlled pH, which is why these calculation rules have practical value.
Key Facts
- Acid dissociation: HA + H2O ⇌ H3O+ + A-
- Base reaction: B + H2O ⇌ BH+ + OH-
- Acid ionization constant: Ka = [H3O+][A-] / [HA]
- Base ionization constant: Kb = [BH+][OH-] / [B]
- Water ion product at 25°C: Kw = [H3O+][OH-] = 1.0 × 10^-14
- pH = -log[H3O+] and pOH = -log[OH-]
Vocabulary
- Equilibrium
- A state in which the forward and reverse reactions continue at equal rates, so concentrations stay constant.
- Weak acid
- An acid that only partially donates protons in water and forms an equilibrium mixture.
- Conjugate base
- The particle left after an acid donates a proton.
- Ka
- The acid ionization constant that measures how strongly an acid produces H3O+ in water.
- Buffer
- A solution that resists changes in pH because it contains a weak acid and its conjugate base or a weak base and its conjugate acid.
Common Mistakes to Avoid
- Treating equilibrium as a stopped reaction is wrong because particles continue reacting in both directions even when concentrations are constant.
- Assuming all acids fully ionize is wrong because weak acids only partially dissociate and must be described using Ka.
- Forgetting to include water correctly is wrong because water can act as an acid or base, but it is usually omitted from Ka and Kb expressions because it is a pure liquid.
- Using initial concentrations as equilibrium concentrations is wrong because reaction changes must be accounted for, often with an ICE table.
Practice Questions
- 1 A solution has [H3O+] = 2.5 × 10^-4 M. Calculate its pH.
- 2 For the equilibrium HA ⇌ H+ + A-, the equilibrium concentrations are [H+] = 1.0 × 10^-3 M, [A-] = 1.0 × 10^-3 M, and [HA] = 0.20 M. Calculate Ka.
- 3 A weak acid solution is diluted with water. Explain how the equilibrium shifts and why the percent ionization may increase even though the solution becomes less acidic.