Difficulty: Intermediate | Prerequisites: Acid-base definitions (Part 1), logarithms, equilibrium constants
Water autoionises into H₃O⁺ and OH⁻ with a constant product Kw = 1.0 × 10⁻¹⁴ at 25 °C, which anchors the pH and pOH scales. Strong acids and bases dissociate completely; weak ones reach equilibrium described by Ka (acids) or Kb (bases), with smaller values meaning weaker dissociation. A strong acid's conjugate base is inert, a weak acid's conjugate base is itself weakly basic, and when acids and bases react, the resulting salt's pH depends on which partner was stronger.
Autoionization of water
The reversible reaction in which two water molecules react to form one hydronium ion (H₃O⁺) and one hydroxide ion (OH⁻). Think of it as water acting as both a Brønsted acid and a Brønsted base with itself.
Kw (water dissociation constant / ion-product constant)
The equilibrium constant for water's autoionization: Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C. It changes only with temperature, not with concentration, pressure, or volume.
pH
The negative base-10 logarithm of the hydrogen ion concentration: pH = −log₁₀[H⁺]. In simple terms, it is the scale that tells you how acidic a solution is, from 0 (very acidic) through 7 (neutral at 25 °C) to 14 (very basic).
pOH
The negative base-10 logarithm of the hydroxide ion concentration. At 25 °C, pH + pOH = 14.
Strong acid
An acid that dissociates completely in aqueous solution. The six to memorise: HCl, HBr, HI, H₂SO₄, HNO₃, HClO₄.
Strong base
A base that dissociates completely in aqueous solution. The main examples: NaOH, KOH, and other Group I metal hydroxides.
Weak acid
An acid that only partially dissociates in water, reaching an equilibrium state. Its strength is quantified by Ka.
Weak base
A base that only partially dissociates in water, reaching an equilibrium state. Its strength is quantified by Kb.
Ka (acid dissociation constant)
The equilibrium constant for the dissociation of a weak acid in water: Ka = [H⁺][A⁻] / [HA]. A smaller Ka means a weaker acid.
Kb (base dissociation constant)
The equilibrium constant for the reaction of a weak base with water: Kb = [BH⁺][OH⁻] / [B]. A smaller Kb means a weaker base.
pKa / pKb
The negative logarithm of Ka or Kb. Larger pKa means weaker acid; larger pKb means weaker base.
Conjugate acid
The species formed when a base gains a proton.
Conjugate base
The species formed when an acid loses a proton.
Neutralisation
A reaction between an acid and a base that produces a salt and (often) water. These reactions typically go to completion.
Hydrolysis
The reverse of neutralisation: the salt reacts with water to regenerate the acid or base, shifting the pH away from 7.
Water autoionises: 2 H₂O ⇌ H₃O⁺ + OH⁻
Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C
In pure water, [H₃O⁺] = [OH⁻] = 1.0 × 10⁻⁷ M
The two concentrations are equal only in a neutral solution; in any solution at 25 °C, their product is always 1.0 × 10⁻¹⁴
Kw is an equilibrium constant: it changes only with temperature. It increases when temperature increases (autoionization is endothermic). Changes in concentration, pressure, or volume do not affect Kw.
pH = −log₁₀[H⁺]; therefore [H⁺] = 10⁻ᵖᴴ
pOH = −log₁₀[OH⁻]
pH + pOH = 14 (at 25 °C)
pH 7 is neutral only at 25 °C. At higher temperatures Kw is larger, so neutral pH drops below 7.
As pH increases by one unit, pOH decreases by one unit
If [H⁺] is a clean power of ten, pH is simply the exponent made positive: [H⁺] = 1.0 × 10⁻³ gives pH = 3
Same logic for pKa, pKb: if Kb = 1.0 × 10⁻¹², then pKb = 12
Dissociate completely in aqueous solution
Strong acids to memorise: HCl, HBr, HI, H₂SO₄, HNO₃, HClO₄
Strong bases to memorise: NaOH, KOH, and other Group I hydroxides
Example: NaOH dissociates fully, so a 1 M solution gives pH 14
When the concentration of a strong acid or base is near 1.0 × 10⁻⁷ M, the autoionization of water contributes meaningfully to [H⁺] or [OH⁻], and you must account for it
Only partially dissociate in aqueous solution, reaching equilibrium
Weak acid dissociation: HA ⇌ H⁺ + A⁻, described by Ka = [H⁺][A⁻] / [HA]
Water is left out of the Ka expression because it is a pure liquid
Smaller Ka = weaker acid (more undissociated HA remains)
Weak base reaction: B + H₂O ⇌ BH⁺ + OH⁻, described by Kb = [BH⁺][OH⁻] / [B]
Smaller Kb = weaker base
Weak acids and bases have K values less than 1.0
Conjugate acid: the species formed when a base accepts a proton
Conjugate base: the species formed when an acid donates a proton
Inverse relationship: if Ka is large (strong acid), its conjugate base has a small Kb (weak base), and vice versa
The conjugate base of a strong acid is effectively inert (unreactive in water)
Weak acids and bases tend to have weak conjugates
Key relationship: Ka × Kb = Kw for any conjugate pair
Bicarbonate ion (HCO₃⁻) is thermodynamically favoured to form, which makes it an effective buffer
Neutralisation: acid + base → salt + water (usually goes to completion)
Hydrolysis is the reverse: the salt reacts with water to regenerate acid or base, pulling pH away from 7
Predicting pH of salt solutions:
Strong acid + strong base → pH = 7
Strong acid + weak base → pH < 7 (example: HCl + NH₃ → NH₄⁺ + Cl⁻; the conjugate acid NH₄⁺ is stronger than the conjugate base Cl⁻)
Weak acid + strong base → pH > 7 (example: CH₃COOH + NaOH → CH₃COO⁻ + Na⁺ + H₂O; the conjugate base CH₃COO⁻ is stronger than the conjugate acid Na⁺)
Weak acid + weak base → pH near 7, direction depends on relative Ka and Kb; if Kb > Ka, the solution is basic
Stomach acid (HCl, a strong acid) maintains a pH around 1.5 to 2.0 for protein digestion. Antacids use weak bases to partially neutralise excess acid. Understanding Ka and Kb lets pharmacologists predict how much of a drug will be ionised at physiological pH, which determines whether it can cross cell membranes.
\text{pH} = -\log_{10}[\text{H}^+][\text{H}^+] = 10^{-\text{pH}}\text{pH} + \text{pOH} = 14 \quad (\text{at } 25\,°\text{C})K_w = [\text{H}_3\text{O}^+][\text{OH}^-] = 1.0 \times 10^{-14} \quad (\text{at } 25\,°\text{C})K_a = \frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}K_b = \frac{[\text{BH}^+][\text{OH}^-]}{[\text{B}]}K_a \times K_b = K_wStudents often think pH 7 is always neutral. It is neutral only at 25 °C. At higher temperatures Kw is larger, pure water has [H⁺] > 10⁻⁷, and neutral pH falls below 7.
A common error is assuming that a smaller Ka means the acid produces fewer H⁺ ions in total. In fact, smaller Ka means the acid dissociates to a lesser extent at equilibrium. The total H⁺ produced also depends on the initial concentration.
Students frequently forget to exclude water from the Ka or Kb expression. Water is a pure liquid and does not appear in the equilibrium expression.
Many students assume that the conjugate base of a weak acid is a strong base. The conjugate of a weak acid is a weak base (small but nonzero Kb). Only the conjugate of a strong acid is effectively inert.
⚠️ You will be asked to calculate pH from [H⁺] and vice versa. Be comfortable moving between logarithmic and exponential forms.
⚠️ Know the six strong acids and Group I hydroxides by heart. Everything else is weak unless told otherwise.
⚠️ Ka × Kb = Kw is a high-yield relationship. Expect problems that give you Ka for an acid and ask for the Kb of its conjugate base.
⚠️ Predicting whether a salt solution is acidic, basic, or neutral is a standard exam question. The decision tree: identify the acid and base that formed the salt, determine whether each is strong or weak, and apply the rules.
⚠️ When [H⁺] or [OH⁻] is near 10⁻⁷ M, you must account for water's autoionization. This is a common trick question.
True or false: Kw changes when you add acid to pure water. (Answer: False. Kw is an equilibrium constant that changes only with temperature.)
Fill in the blank: A solution with [H⁺] = 1.0 × 10⁻⁵ M has a pH of ______. (Answer: 5)
True or false: The conjugate base of HCl is a strong base. (Answer: False. Cl⁻ is inert because HCl is a strong acid.)
Fill in the blank: For any conjugate acid-base pair, Ka × Kb = ______. (Answer: Kw = 1.0 × 10⁻¹⁴ at 25 °C)
True or false: A solution of sodium acetate (CH₃COONa) in water will have pH < 7. (Answer: False. Sodium acetate is the salt of a weak acid and a strong base, so pH > 7.)
Q: A weak acid HA has Ka = 1.8 × 10⁻⁵. What is the Kb of its conjugate base A⁻?
A: Kb = Kw / Ka = (1.0 × 10⁻¹⁴) / (1.8 × 10⁻⁵) = 5.6 × 10⁻¹⁰.
Q: You dissolve NH₄Cl in water. Will the resulting solution be acidic, basic, or neutral? Explain.
A: Acidic (pH < 7). NH₄Cl is the salt of a weak base (NH₃) and a strong acid (HCl). The conjugate acid NH₄⁺ partially donates protons to water, producing excess H⁺. Cl⁻ is inert.
Q: Calculate the pH of a 0.010 M HCl solution.
A: HCl is a strong acid and dissociates completely, so [H⁺] = 0.010 M = 1.0 × 10⁻². pH = −log(1.0 × 10⁻²) = 2.0.
Q: At a temperature above 25 °C, Kw = 5.5 × 10⁻¹⁴. What is the pH of pure water at this temperature?
A: In pure water, [H⁺] = [OH⁻] = √Kw = √(5.5 × 10⁻¹⁴) = 2.35 × 10⁻⁷. pH = −log(2.35 × 10⁻⁷) ≈ 6.63. Note that this is below 7 but the solution is still neutral (equal [H⁺] and [OH⁻]).
Q: A solution has a pOH of 3. What is its pH and is it acidic or basic?
A: pH = 14 − 3 = 11. The solution is basic.
This material builds directly on Part 1 (Definitions and Nomenclature): you need to know the difference between strong and weak before you can write the correct equilibrium expression. It feeds into Part 3 (Titration and Buffers), where Ka and Kb determine the shape of the titration curve and the effectiveness of a buffer. The Henderson-Hasselbalch equation (Part 3) is just the logarithmic form of the Ka expression rearranged.
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