IB DP Chemistry: Reactivity 3.1 Proton Transfer Reactions Revision Notes

Comprehensive practice flashcards for IB DP Chemistry Reactivity 3.1: Proton Transfer Reactions, covering Brønsted–Lowry definitions, conjugate pairs, neutralisation reactions, strong vs. weak acids, the pH scale, weak-acid equilibria, buffer systems, pH titration curves, and salt hydrolysis.

How does the Brønsted–Lowry theory define an acid and a base?

An acid is defined as a proton ($H^+$) donor, and a base is defined as a proton ($H^+$) acceptor.

What is a conjugate acid–base pair, and by how much do the formulas of the two species differ?

A conjugate acid–base pair consists of two species that differ by exactly one proton ($H^+$). The charge changes by $+1$ going from the base to its conjugate acid.

What is the distinction between an amphiprotic species and an amphoteric species?

An amphoteric substance can react with both acids and bases (e.g., $Al_2O_3$, $ZnO$). An amphiprotic substance is the narrower, proton-specific case that specifically donates or accepts a proton ($H^+$). Every amphiprotic species is amphoteric, but not the reverse.

What products are formed when an acid reacts with a metal carbonate or hydrogencarbonate?

The reaction produces a salt, water ($H_2O$), and carbon dioxide ($CO_2$).

What is the net ionic equation and approximate enthalpy change ($\Delta H$) for the neutralisation of any strong acid with any strong base?

The net ionic equation is $H^+(aq) + OH^-(aq) \rightarrow H_2O(l)$ with $\Delta H \approx -57\,kJ\,mol^{-1}$.

How is acid strength distinguished from acid concentration?

Strength measures the degree of dissociation (ionisation) of the acid in water (complete, approximately $100\%$ for strong; partial, often $< 5\%$ for weak), whereas concentration measures the amount of solute dissolved per unit volume.

What are three experimental observations that distinguish equal-concentration solutions of a strong acid and a weak acid?

Compared to a weak acid at equal concentration, the strong acid has a lower $pH$, a higher electrical conductivity, and a faster reaction rate (faster fizzing) with reactive metals (such as $Mg$) or carbonates.

Why does a fixed volume of an equimolar weak monoprotic acid require the same volume of alkali to neutralise as a strong monoprotic acid?

Neutralisation depends on the total amount (moles) of acidic hydrogen present, not acid strength. As the weak acid is neutralised, its equilibrium shifts to the right until all of its acidic hydrogen has reacted.

What is the expression for the ionic product of water ($K_w$) and its numerical value at $298\,K$?

$K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$ at $298\,K$.

Why is pure water at $50\,^\circ\text{C}$ still considered neutral even though its $pH$ falls to approximately $6.6$?

The self-ionisation of water is endothermic, so $K_w$ increases as temperature rises. Water remains neutral because $[H^+] = [OH^-]$, even though the neutral $pH$ is below $7$.

How are $K_a$ and $K_b$ (as well as $pK_a$ and $pK_b$) related for any conjugate acid–base pair at $298\,K$?

$K_a \times K_b = K_w$ and $pK_a + pKb = 14.00$.

Which two approximations allow the calculation of $[H^+]$ for a weak acid using the formula $[H^+] \approx \sqrt{K_a \times c}$?

(1) $[H^+] = [A^-]$ (both originate almost entirely from acid dissociation), and (2) equilibrium $[HA] \approx c$ (the extent of dissociation is negligible, typically $< 5\%$).

What components are required to prepare an acidic buffer solution?

An acidic buffer requires a weak acid combined with a comparable amount of its salt / conjugate base (for example, $CH_3COOH$ and $CH_3COONa$).

How does an ethanoic acid/ethanoate buffer neutralise added small amounts of acid ($H^+$) and base ($OH^-$)?

Added $H^+$ is mopped up by the conjugate base ($CH_3COO^- + H^+ \rightarrow CH_3COOH$), and added $OH^-$ is mopped up by the weak acid ($CH_3COOH + OH^- \rightarrow CH_3COO^- + H_2O$).

According to the Henderson–Hasselbalch equation, what is the $pH$ of a buffer when $[\text{salt}] = [\text{acid}]$?

$pH = pK_a$, because $\log(1) = 0$ in the relationship $pH = pK_a + \log\left(\frac{[\text{salt}]}{[\text{acid}]}\right)$.

What main conjugate pair functions as the primary buffer system in human blood to maintain a steady $pH$ near $7.4$?

The carbonic acid / hydrogencarbonate pair ($H_2CO_3 / HCO_3^-$).

What key relationship occurs at the half-equivalence point in a weak acid–strong base titration curve?

At the half-equivalence point, exactly half the weak acid has been neutralised so $[\text{salt}] = [\text{acid}]$, meaning $pH = pK_a$.

What criterion determines whether an acid–base indicator is suitable for a titration?

The indicator's colour-change range (approximately $pK_a(In) \pm 1$) must fall completely within the steep, near-vertical jump section of the titration pH curve.

Why is phenolphthalein chosen over methyl orange for the titration of ethanoic acid with sodium hydroxide?

The equivalence point of a weak acid–strong base titration is basic ($pH \approx 8.7$) with a vertical jump around $pH\,7\text{–}11$. Phenolphthalein (range $8.3\text{–}10.0$) changes colour within this jump, while methyl orange (range $3.1\text{–}4.4$) would change too early in the buffer region.

Why does an aqueous solution of ammonium chloride ($NH_4Cl$) produce an acidic solution ($pH < 7$)?

$NH_4Cl$ is the salt of a strong acid ($HCl$) and a weak base ($NH_3$). The ammonium cation ($NH_4^+$) hydrolyses by donating a proton to water ($NH_4^+ + H_2O \rightleftharpoons NH_3 + H_3O^+$), producing excess $H_3O^+$.