International Baccalaureate IB Chemistry

Brønsted–Lowry acid is a proton donor and a Brønsted–Lowry base is a proton acceptor.

Deduce the Brønsted–Lowry acid and base in a reaction.

A pair of species differing by a single proton is called a conjugate acid–base pair.

Deduce the formula of the conjugate acid or base of any Brønsted–Lowry base or acid.

Some species can act as both Brønsted–Lowry acids and bases.

Interpret and formulate equations to show acid–base reactions of these species.

The pH scale can be used to describe the $[H^+]$ of a solution: $pH = -\log_{10}[H^+]$; $[H^+] = 10^{-pH}$.

Perform calculations involving the logarithmic relationship between pH and $[H^+]$.

The ion product constant of water, $K_w$, shows an inverse relationship between $[\mathrm{H}^+]$ and $[\mathrm{OH}^-]$: $K_w = [\mathrm{H}^+][\mathrm{OH}^-]$.

Recognise solutions as acidic, neutral and basic from the relative values of $[\mathrm{H}^+]$ and $[\mathrm{OH}^-]$.

Strong and weak acids and bases differ in the extent of ionisation.

Recognise that acid–base equilibria lie in the direction of the weaker conjugate.

Acids react with bases in neutralisation reactions.

Formulate equations for the reactions between acids and metal oxides, metal hydroxides, hydrogen‑carbonate and carbonates. 

Identify the parent acid and base of different salts.

pH curves for neutralization reactions involving strong acids and bases have characteristic shapes and features. 

Sketch and interpret the general shape of the pH curve. 

Interpretation should include the intercept with the pH axis and equivalence point. 

Only monoprotic neutralization reactions will be assessed.

The $pOH$ scale describes the $[\mathrm{OH}^-]$ of a solution. $pOH = -\log_{10}[\mathrm{OH}^-]$; $[\mathrm{OH}^-] = 10^{-\text{pOH}}$.

Interconvert $[\mathrm{H}^+]$, $[\mathrm{OH}^-]$, $pH$ and $pOH$ values.

The equations for $pOH$ are given in the data booklet.

The strengths of weak acids and bases are described by their Ka, Kb, pKa or pKb values.

Interpret the relative strengths of acids and bases from these data.

For a conjugate acid–base pair, the relationship $K_a \times K_b = K_w$ can be derived from the expressions for $K_a$ and $K_b$.

Solve problems involving these values.

The pH of a salt solution depends on the relative strengths of the parent acid and base.

Construct equations for the hydrolysis of ions in a salt, and predict the effect of each ion on the pH of the salt solution. 

Examples should include the ammonium ion $ \mathrm{NH_4^+}$, the carboxylate ion $ \mathrm{RCOO^-}$, the carbonate ion $ \mathrm{CO_3^{2-}}$, and the hydrogencarbonate ion $ \mathrm{HCO_3^-}$. The acidity of hydrated transition element ions and (aq) is not required. 

$$\mathrm{NH_4^+ + H_2O \rightleftharpoons NH_3 + H_3O^+}$$  

$$\mathrm{RCOO^- + H_2O \rightleftharpoons RCOOH + OH^-}$$  

$$\mathrm{CO_3^{2-} + H_2O \rightleftharpoons HCO_3^- + OH^-}$$  

$$\mathrm{HCO_3^- + H_2O \rightleftharpoons H_2CO_3 + OH^-}$$  

$$\mathrm{HCO_3^- + H_2O \rightleftharpoons CO_3^{2-} + H_3O^+}$$

pH curves of different combinations of strong and weak monoprotic acids and bases have characteristic shapes and features.

Interpret the general shapes of pH curves for all four combinations of strong and weak acids and bases. 

Interpretation should include: intercept with the pH 

Tool 1—When collecting data to generate a pH axis, equivalence point, buffer region, points where curve, when should smaller volumes of titrant be added between each measurement? $pH = pKa$ or $pOH = pKb$.

Acid–base indicators are weak acids, where the components of the conjugate acid–base pair have different colours. The $p\mathrm{H}$ of the end point of an indicator, where it changes colour, approximately corresponds to its $pK_a$ value.

Construct equilibria expressions to show why the colour of an indicator changes with $p\mathrm{H}$. 

The generalized formula $\mathrm{HInd_{(aq)}}$ can be used to represent the undissociated form of an indicator. 

The equilibrium can be written as 
$$\mathrm{HInd_{(aq)} \rightleftharpoons H^+ + Ind^-}$$ 
and the acid dissociation constant as 
$$K_a = \frac{[H^+][\mathrm{Ind^-}]}{[\mathrm{HInd}]}.$$ 

Examples of indicators with their $p\mathrm{H}$ range are given in the data booklet. 

Include universal indicator as a mixture of many indicators with a wide $p\mathrm{H}$ range of colour change.

An appropriate indicator for a titration has an end point range that coincides with the $pH$ at the equivalence point.

Identify an appropriate indicator for a titration from the identity of the salt and the $pH$ range of the indicator. 

Distinguish between the terms “end point” and “equivalence point”.

A buffer solution is one that resists change in pH on the addition of small amounts of acid or alkali.

Describe the composition of acidic and basic buffers and explain their actions. 

Why must buffer solutions be composed of weak acid or base conjugate systems, not of strong acids or bases?

The $pH$ of a buffer solution depends on both: the $pKa$ or $pKb$ of its acid or base; the ratio of the concentration of acid or base to the concentration of the conjugate base or acid.

Solve problems involving the composition and $pH$ of a buffer solution, using the equilibrium constant. Include explanation of the effect of dilution of a buffer.