Revise Proton transfer reactions for IB Chemistry (R3.1)

In 1923 Johannes Brønsted and Thomas Lowry independently proposed that acid-base behaviour is really about proton transfer. A Brønsted-Lowry acid is a species that donates a proton ($H^+$); a Brønsted-Lowry base is a species that accepts a proton. A 'proton' here is a hydrogen ion — a hydrogen atom that has lost its electron. The model is not limited to water and does not require a base to contain hydroxide, which is why it explains far more than the older Arrhenius definitions.

$\text{Acid} + \text{Base} \rightleftharpoons \text{Conjugate base} + \text{Conjugate acid}$

When hydrogen chloride dissolves in water, the HCl molecule donates a proton to a water molecule, so HCl is the acid and $H_2O$ is the base. In water the free proton is carried by a water molecule as the hydronium (oxonium) ion, $H_3O^+$.

$HCl(g) + H_2O(l) \rightleftharpoons Cl^-(aq) + H_3O^+(aq) \\ \text{(acid)} \quad\; \text{(base)} \quad\; \text{(conj. base)} \quad \text{(conj. acid)}$

Because a proton is transferred, every Brønsted-Lowry reaction produces two conjugate pairs. When an acid donates its proton, what remains is its conjugate base; when a base accepts a proton, the species formed is its conjugate acid. The two members of a pair always differ by exactly one $H^+$, so their formulae differ by one H and their charges differ by one unit.

Some species can act as either a Brønsted-Lowry acid or a base, depending on their reaction partner. These are amphiprotic. To behave this way a species needs both a hydrogen atom it can donate and a lone pair with which to accept a proton. Water is the classic case: it donates a proton to become $OH^-$, or accepts one to become $H_3O^+$. The hydrogencarbonate ion, $HCO_3^-$, is another key example, along with $HSO_4^-$ and $H_2PO_4^-$.

Hydrogen ion concentrations in solution span an enormous range, so we compress them onto a logarithmic scale: $\text{pH} = -\log_{10}[H^+]$, where $[H^+]$ is in mol dm⁻³. Rearranged, $[H^+] = 10^{-\text{pH}}$. The key feature is that the scale is logarithmic, so each whole pH unit corresponds to a TEN-FOLD change in $[H^+]$: a solution of pH 2 has ten times the hydrogen ion concentration of one at pH 3, and a hundred times that of one at pH 4. A lower pH means a higher $[H^+]$ and a more acidic solution.

$\text{pH} = -\log_{10}[H^+] \qquad\Longleftrightarrow\qquad [H^+] = 10^{-\text{pH}}$

'Strong' and 'weak' describe the EXTENT to which an acid or base reacts with water — nothing to do with how concentrated it is. A strong acid dissociates essentially completely: in HCl(aq) virtually every molecule has donated its proton, so we write a single arrow. A weak acid only partially dissociates and reaches an equilibrium, so we write a reversible arrow and most of it remains as intact molecules. The same distinction applies to bases: a strong base such as NaOH is fully dissociated into ions, whereas a weak base such as NH₃ is only partially protonated at equilibrium.

$\text{Strong acid: } HCl(aq) \rightarrow H^+(aq) + Cl^-(aq) \\ \text{Weak acid: } CH_3COOH(aq) \rightleftharpoons H^+(aq) + CH_3COO^-(aq)$

This is why, at the SAME concentration, a strong acid has a higher $[H^+]$ and therefore a LOWER pH than a weak acid: the strong acid has released far more of its available protons. It is also why 'strong' and 'concentrated' must never be swapped — you can have a dilute strong acid (low concentration, fully dissociated) or a concentrated weak acid (high concentration, barely dissociated).

Acids show three signature reactions you must be able to write and recognise. Each is a proton transfer, and each has an observation you can quote in the exam.

Neutralisation is the exact cancellation of an acid by a base: hydrogen ions from the acid combine with hydroxide ions (or another base) to form water, and the remaining ions form the salt. Because the underlying reaction is always $H^+ + OH^- \rightarrow H_2O$, neutralisation of any strong acid by any strong alkali releases the same enthalpy change per mole of water formed.

At HL the strong/weak distinction is made quantitative. For a weak acid $HA \rightleftharpoons H^+ + A^-$, the acid dissociation constant is $K_a = \dfrac{[H^+][A^-]}{[HA]}$; a larger $K_a$ means more dissociation and a stronger acid. For a weak base the equivalent is $K_b$. Because these constants span many orders of magnitude, we usually quote $pK_a = -\log_{10}K_a$ and $pK_b = -\log_{10}K_b$. Watch the direction: the stronger the acid, the LARGER its $K_a$ but the SMALLER its $pK_a$. For a conjugate acid-base pair at 298 K, $pK_a + pK_b = 14$.

Acid-base questions on Paper 2 mix a short calculation with a worded explanation, and the marks are analytic: each is tied to a specific line of working (a method mark, M, or an answer mark, A), while error-carried-forward (ECF) means a wrong number early on need not cost every later mark. The explanation marks are awarded for distinct, correct points — you earn them one idea at a time. Study how each mark below is won.

Proton transfer is the foundation for everything that follows in acid-base chemistry. Titration curves are just pH plotted as one solution neutralises another; the sharp jump at the equivalence point is a large change in $[H^+]$ shown on the logarithmic pH scale. At HL, $K_a$ and $pK_a$ turn the strong/weak idea into calculations of pH for weak acids, and buffers apply Le Châtelier's principle to a conjugate pair. Master the two habits here — track the single proton, and reason from extent of dissociation rather than concentration — and the rest of the topic becomes variations on ideas you already own.