A chemist adds 0.010 mol of HCl to EACH of the following 1.0 L solutions at 25 °C

CHEMISTRY · Single correct · A Level · Chemistry · Acids and bases · Acids and bases

Problem

A chemist adds 0.010 mol of HCl to EACH of the following 1.0 L solutions at 25 °C. • Solution I: pure water initially at pH 7.00 • Solution II: 0.200 M NH₃(aq) and 0.200 M NH₄Cl(aq) (pK_a of NH₄⁺ = 9.25) Which of the following best explains why the final pH of Solution II is much closer to its initial value than that of Solution I? (a) In Solution II, the strong base NH3\mathrm{NH_3} completely neutralizes the added H+\mathrm{H^{+}}, preventing any appreciable rise in [H+]\mathrm{[H^{+}]}. (b) In Solution II, because [NH3]=[NH4+][\mathrm{NH_3}] = [\mathrm{NH_4^+}], the Henderson–Hasselbalch equation locks the pH at pKapK_a no matter how much acid is introduced. (c) In Solution II, the added H+\mathrm{H^{+}} reacts primarily with NH3\mathrm{NH_{3}}, so the [NH3][NH4+]\dfrac{[\mathrm{NH_3}]}{[\mathrm{NH_4^+]}} ratio and therefore the pH change only slightly. (d) In Solution II, the high ionic strength lowers the activity coefficient of H+\mathrm{H^{+}}, so the pH reading changes very little after the acid is added.

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