Brønsted–Lowry acids and bases

• Brønsted–Lowry acid = proton donor.

• Brønsted–Lowry base = proton acceptor.

• In aqueous chemistry, a proton can be written as $H^+(aq)$ or $H_3O^+(aq)$.

• To identify acid/base roles in an equation, track which species loses $H^+$ and which species gains $H^+$.

• Base does not always mean alkali: an alkali is a soluble base that produces $OH^-(aq)$ in water.

• Typical example: $NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$

  • $NH_3$ acts as the base.

  • $H_2O$ acts as the acid.

This diagram shows a proton transfer equilibrium between ammonium/ammonia and water/hydroxide, making the acid, base, conjugate acid, and conjugate base easy to identify. It is ideal for practicing how to label Brønsted–Lowry species in both forward and reverse directions. Source

Conjugate acid–base pairs

• A conjugate acid–base pair consists of two species differing by one proton.

• When an acid loses $H^+$, it forms its conjugate base.

• When a base gains $H^+$, it forms its conjugate acid.

• Quick rule:

  • acid $\to$ conjugate base + $H^+$

  • base + $H^+$ $\to$ conjugate acid

• Example: in $HCl + H_2O \to H_3O^+ + Cl^-$

  • $HCl/Cl^-$ is one conjugate pair.

  • $H_2O/H_3O^+$ is the other conjugate pair.

• Exam tip: species on opposite sides of the equation that differ by one proton are the conjugate pair.

Amphiprotic species

• Some species can act as both acids and bases: these are amphiprotic.

• An amphiprotic species can donate $H^+$ in one reaction and accept $H^+$ in another.

• Important example: water.

  • As a base: $HCl + H_2O \to H_3O^+ + Cl^-$

  • As an acid: $NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$

• Other common amphiprotic ions at HL include $HCO_3^-$ and sometimes $HSO_4^-$ in broader acid–base chemistry.

The figure shows one water molecule donating a proton to another, producing hydronium and hydroxide. It is useful for understanding why water is amphiprotic and how $K_w$ arises. Source

pH and hydrogen ion concentration

• pH describes the hydrogen ion concentration of a solution.

• Equations you must use confidently:

  • $pH = -\log_{10}[H^+]$

  • $[H^+] = 10^{-pH}$

• Lower pH = higher $[H^+]$ = more acidic.

• Higher pH = lower $[H^+]$ = more basic/alkaline.

• A change of 1 pH unit means a tenfold change in $[H^+]$.

• Be careful with calculator use: use 10 to the power of minus pH, not just the negative sign.

• Practical methods:

  • Universal indicator / pH paper = estimate pH.

  • pH probe / meter = precise pH measurement.

This visual represents how hydronium ion concentration changes in acidic solution, helping students connect the abstract idea of $[H^+]$ to particle-level meaning. It supports the logarithmic definition of pH. Source

Water ionization and $K_w$

• Water undergoes self-ionization:

  • $H_2O(l) \rightleftharpoons H^+(aq) + OH^-(aq)$

  • or $2H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)$

• The ion product constant of water is:

  • $K_w = [H^+][OH^-]$

• At 298 K, $K_w = 1.0 \times 10^{-14}$.

• Classifying solutions:

  • acidic: $[H^+] > [OH^-]$

  • neutral: $[H^+] = [OH^-]$

  • basic: $[H^+] < [OH^-]$

• In neutral water at 298 K:

  • $[H^+] = [OH^-] = 1.0 \times 10^{-7}$ mol dm$^{-3}$

  • therefore pH = 7.

• Common exam move: use $K_w$ to calculate the missing ion concentration first, then convert to pH or pOH if needed.

Strong vs weak acids and bases

• Strong acids/bases ionize almost completely in water.

• Weak acids/bases ionize partially in water.

• Strength is about extent of ionization, not concentration.

• Concentrated vs dilute describes how much solute is dissolved, not how fully it ionizes.

• Acid–base equilibria lie in the direction of the weaker conjugate acid–base pair.

• Strong acids commonly treated in IB include $HCl$, $HNO_3$, and $H_2SO_4$.

• Group 1 hydroxides are strong bases.

• For equal concentration:

  • strong acid has lower pH than weak acid.

  • strong base has higher pH than weak base.

Neutralization and salts

• Neutralization = acid + base $\to$ salt + water.

• You must be able to write equations for acids reacting with:

  • metal oxides

  • metal hydroxides

  • carbonates

  • hydrogencarbonates

• General patterns:

  • acid + metal oxide $\to$ salt + water

  • acid + metal hydroxide $\to$ salt + water

  • acid + carbonate $\to$ salt + water + carbon dioxide

  • acid + hydrogencarbonate $\to$ salt + water + carbon dioxide

• Bases in this topic include ammonia, amines, soluble carbonates, and hydrogencarbonates.

• You should be able to identify the parent acid and parent base from a salt.

  • Example: $NaCl$ comes from $HCl$ and $NaOH$.

  • Example: $NH_4NO_3$ comes from $HNO_3$ and $NH_3$.

Strong acid–strong base titration curves

• IB expects the general shape of the pH curve for a strong acid–strong base neutralization.

• Key features to interpret:

  • initial pH

  • gradual rise before equivalence

  • steep vertical section near equivalence

  • equivalence point at pH 7 (for strong acid–strong base)

• The equivalence point is the stoichiometric point where the amount of acid equals the amount of base according to the balanced equation.

• For monoprotic titrations, use mole ratio 1:1 when appropriate.

• Exam trap: end point is where the indicator changes colour; equivalence point is the true stoichiometric point.

This illustration shows how a basic buffer contains a weak base and its conjugate acid, and how added acid or base is removed by equilibrium reactions. It is a strong visual aid for explaining why buffers resist pH change. Source

The image shows the colour changes of universal indicator across the pH scale, linking indicator colour to acidity and alkalinity. It is useful when revising estimated pH values, indicator ranges, and practical identification of acidic, neutral, and basic solutions. Source

Ultra-short exam reminders

• Conjugate pairs differ by exactly one proton.

• Strong means fully/almost fully ionized; concentrated means high amount dissolved.

• Acidic solutions have higher $[H^+]$ than $[OH^-]$.

• For strong acid–strong base titrations, equivalence point = pH 7.

• At HL, always check whether a question is really about equilibrium, hydrolysis, indicator choice, or buffer action.