(a)(i) Outline a suitable procedure for the preparation of ZnCl\(_2\) crystals stating from zinc granules. (ii) Write a balanced equation for the reaction i...
Question 1 Report
(a)(i) Outline a suitable procedure for the preparation of ZnCl\(_2\) crystals stating from zinc granules.
(ii) Write a balanced equation for the reaction in (a)(i).
(b) Considei' the following solutions: Na\(_2\)SO\(_{4(aq)}\), CH3COOK\(_{(aq)}\), Pb(NO\(_3\))\(_{2(aq)}\) and MgCl\(_2\). Which of them has/have a pH
(i) greater than 7;
(ii) equal to 7;
(iii) less than 7?
(c)(i) Determine the oxidation number of Mn in I. MnO\(_2\); II. KMnO\(_4\).
(ii) State one laboratory use of each of the .compounds in (c)(i).
(d) Explain why oxidation and reduction processes are complementary.
(e) Consider the reaction represented by the equation: Fe\(_{(s)}\) + 2Ag\(^+_{(aq)}\) \(\to\) Fe\(^{2+}_{(aq)}\) + 2Ag\(_{(s)}\)
(i) Write a balanced ionic half equations for the reaction.
(ii) Which of the species is I. oxidized; II. reduced?
(iii) State the change in the oxidation number of silver during the reaction.
Answer Details
(a)(i) Preparation of ZnCl2 crystals from zinc granules
Add excess zinc granules to warm dilute hydrochloric acid in a beaker. The zinc reacts, giving off hydrogen gas (effervescence). When effervescence stops (all the acid is used up), filter off the excess unreacted zinc. Evaporate the filtrate over a water bath to the point of crystallization, allow it to cool so that crystals of zinc chloride form, then filter off and dry the crystals between filter papers.
(a)(ii) Equation
\[Zn_{(s)} + 2HCl_{(aq)} \to ZnCl_{2(aq)} + H_{2(g)}\](b) pH of the solutions
\(Na_2SO_4\) is the salt of a strong acid and strong base; \(CH_3COOK\) is the salt of a weak acid and strong base; \(Pb(NO_3)_2\) and \(MgCl_2\) are salts of strong acids and weak bases.
- (i) pH greater than 7: \(CH_3COOK\).
- (ii) pH equal to 7: \(Na_2SO_4\).
- (iii) pH less than 7: \(Pb(NO_3)_2\) and \(MgCl_2\).
(c)(i) Oxidation number of Mn
- I. In \(MnO_2\): \(Mn + 2(-2) = 0 \Rightarrow Mn = +4\).
- II. In \(KMnO_4\): \(+1 + Mn + 4(-2) = 0 \Rightarrow Mn = +7\).
(c)(ii) One laboratory use of each
- \(MnO_2\): used as a catalyst (e.g. in the decomposition of hydrogen peroxide) and as an oxidizing agent in the preparation of chlorine.
- \(KMnO_4\): used as an oxidizing agent in volumetric analysis (redox titrations), where it acts as its own indicator.
(d) Why oxidation and reduction are complementary
Oxidation is the loss of electrons and reduction is the gain of electrons. The electrons lost by the species being oxidized are exactly those gained by the species being reduced. Therefore the two processes must occur together and simultaneously (in equal number of electrons); one cannot take place without the other.
(e) The reaction \(Fe + 2Ag^+ \to Fe^{2+} + 2Ag\)
- (i) Half equations: \(Fe \to Fe^{2+} + 2e^-\) (oxidation); \(Ag^+ + e^- \to Ag\) (reduction).
- (ii) I. Oxidized: iron (\(Fe\)); II. Reduced: silver ion (\(Ag^+\)).
- (iii) The oxidation number of silver changes from \(+1\) to \(0\) (a decrease of 1).
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