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The rate of the following reaction in aqueous solution is monitored by measuring the number of moles of Hg_{2}Cl_{2} that precipitate per liter per minute. The data obtained are listed in the table.
2HgCl_{2}(aq) + C_{2}O_{4}^{2–}(aq)
2Cl^{–}(aq)
+ 2CO_{2}(g) + Hg_{2}Cl_{2}(s)



mol L^{–1} min^{–1} 
















(a) Determine the order of reaction with respect to HgCl_{2},
with respect to C_{2}O_{4}^{2–} and overall.
(b) What is the value of the rate constant k?
(c) What would be the initial rate of reaction if [HgCl_{2}]
= 0.094 M and [C_{2}O_{4}^{2–}] = 0.19 M?
(d) Are all four experiments necessary to answer parts (a)  (c)? Explain.
Use the method of initial rates to find the orders of reaction in each component. This will allow evaluation of the rate constant and the initial rate of reaction at any other condition.
(a) Rate = k[HgCl_{2}]^{m}[C_{2}O_{4}^{2–}]^{n}
Compare the rates in experiments 1 and 2 (or 3 and 4) to find the order
in oxalate ion:
Compare the rates in experiments 2 and 3 (or 1 and 4) to find the order
in mercury(II) chloride:
Therefore, the reaction is first order with respect to mercury(II)
chloride and second order with respect to oxalate. The overall order is
the sum of these, 2 + 1 = 3, third order.
(b) To find the rate constant, use the rate equation and solve
for k:



mol L^{–1} min^{–1} 





















(c) Rate = k[HgCl_{2}][C_{2}O_{4}^{2–}]^{2}
Rate = 7.6 ×10^{–3} M^{–2}min^{–1}[0.094
M][0.19 M]^{2} = 2.6 ×10^{–5} M^{ }min^{–1}
(d) Since there are only two reactants, three experiments are the minimum required to find the rate equation and rate constant. Experiments 1  3 would have sufficed to answer the questions posed.