A solution of ammonia has a concentration of ${\text{0.500 mol}}\,{\text{d}}{{\text{m}}^{ - 3}}$.

Calculate the pH of the ammonia solution using information from Table 15 of the Data Booklet. State one assumption made.

(i)     State the meaning of the term buffer solution.

(ii)     Calculate the concentrations of ammonia and ammonium ion in the buffer solution.

(iii)     Determine the pH of the buffer solution at 25 °C.

(iv)     Explain why the pH of the buffer solution is different from the pH of the ammonia solution calculated in (d).

(v)     Explain the action of the buffer solution when a few drops of nitric acid solution are added to it.

(i)     Identify the property of bromocresol green that makes it suitable to use as an acid–base indicator.

(ii)     State and explain the relationship between the pH range of bromocresol green and its ${\text{p}}{K_{\text{a}}}$ value.

${\text{[O}}{{\text{H}}^ - }{\text{]}} = \left( {\sqrt {0.500 \times 1.78 \times {{10}^{ - 5}}} } \right) = 2.98 \times {10^{ - 3}}{\text{ mol}}\,{\text{d}}{{\text{m}}^{ - 3}}$;

${\text{pOH}} =  - {\log _{10}}{\text{[O}}{{\text{H}}^ - }{\text{]}} = 2.526/{\text{[}}{{\text{H}}^ + }{\text{]}} = \left( {\frac{{1.00 \times {{10}^{ - 14}}}}{{2.98 \times {{10}^{ - 3}}}}} \right) = 3.35 \times {10^{ - 12}}{\text{ mol}}\,{\text{d}}{{\text{m}}^{ - 3}}$;

${\text{pH}} = 11.47$;

Accept correct answer obtained using another method.

Assumption:

${\text{[N}}{{\text{H}}_{\text{3}}}{\text{]}} = 0.500{\text{ mol}}\,{\text{d}}{{\text{m}}^{ - 3}}$ / ${\text{[NH}}_4^ + {\text{]}} = {\text{[O}}{{\text{H}}^ - }{\text{]}}$ / all ${\text{O}}{{\text{H}}^ - }$ ions come from the reaction of ammonia with water and not from the dissociation of water / temperature is 25 °C/298 K / OWTTE;

(i)     resists change in pH when small amounts of a strong base/strong acid/water are added to it;

(ii)     ${\text{[N}}{{\text{H}}_{\text{3}}}{\text{]}} = 0.250{\text{ mol}}\,{\text{d}}{{\text{m}}^{ - 3}}$;

${\text{[NH}}_4^ + {\text{]}} = 0.250{\text{ mol}}\,{\text{d}}{{\text{m}}^{ - 3}}$;

(iii)     ${\text{pOH}} = {\text{p}}{K_b} = 4.75$;

${\text{pH}} = 9.25$;

(iv)     equilibrium shifted left in buffer / OWTTE;

(v)     acid neutralized by hydroxide / most of the added ${{\text{H}}^ + }$ ions react with ${\text{N}}{{\text{H}}_{\text{3}}}$;

more ammonia reacts with water to replace hydroxide ions / more ${\text{NH}}_4^ +$ ions form so there is little change in the pH / OWTTE;

Accept equations.

(i)     colours of HIn and ${\text{I}}{{\text{n}}^ - }$ are different / OWTTE;

(ii)     colour change occurs when ${\text{[HIn]}} = {\text{[I}}{{\text{n}}^ - }{\text{]}}$;

${\text{pH}} = {\text{p}}{K_{\text{a}}}$;

OR

pH range is a range of pH values either side of ${\text{p}}{K_{\text{a}}}$ value;

lower pH when acid colour is seen and upper pH when alkaline colour seen;

Calculation of pH in (d) proved challenging for some and straightforward for others. Those who knew how to perform the calculations generally also correctly stated an assumption.

Most candidates correctly described a buffer solution in (e). Several candidates had difficulty calculating the concentrations of ammonia and ammonium ions in the buffer but managed to calculate the pH correctly (some with ECF). The explanations of why the pH of the buffer differs from the pH of ammonia and the action of the buffer when a few drops of nitric acid are added were poorly done and would have been better with the use of equations and references to equilibrium.

Answers to (f) were quite general. Many candidates simply said that bromocresol green changes colour with no further details, or said that the indicator had different colours in acid and alkaline conditions. Most candidates scored 1 mark for stating that the ${\text{p}}{K_{\text{a}}}$ is in the middle of the pH range.