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Chapter 13: Q11E (page 750)

Benzene is one of the compounds used as octane enhancers in unleaded gasoline. It is manufactured by the catalytic conversion of acetylene to benzene: \(3{{\rm{C}}_2}{{\rm{H}}_2}(g) \to {{\rm{C}}_6}{{\rm{H}}_6}(g)\). Which value of \({K_c}\) would make this reaction most useful commercially?

\({K_c} \approx 0.01,\;{K_c} \approx 1,\;\;{\rm{or}}\;{K_c} \approx 10.\) Explain your answer.

Short Answer

Expert verified

The value that makes the reaction commercially useful is\({K_c} \approx 10\).

Step by step solution

01

Define equilibrium constant

The product of the reactant concentrations to the power of their reaction coefficient (the number before them in the reaction equation) divided by the product of the product concentrations to the power of their reaction coefficients, independent of the reaction, is the equilibrium constant.

02

Find the value of \({K_c}\)

The equilibrium constant for this reaction is expressed as:

\({K_c} = \frac{{c\left( {{C_6}{H_6},g} \right)}}{{c{{\left( {{C_2}{H_2},g} \right)}^3}}}\)

If our goal is to make as much benzene as feasible, the concentration of\({C_6}{H_6}\)should be as high as possible.

Thus, if \({K_c} \approx 10\), this reaction would be most commercially useful.

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Most popular questions from this chapter

Question: At 25 掳C and at 1 atm, the partial pressures in an equilibrium mixture of N2O4 and NO2 are PN2O4= 0.70 atm and PNO2 = 0.30 atm.

(a) Predict how the pressures of NO2 and N2O4 will change if the total pressure increases to 9.0 atm. Will they increase, decrease, or remain the same?

(b) Calculate the partial pressures of NO2 and N2O4 when they are at equilibrium at 9.0 atm and 25 掳C

Question:What is the value of the equilibrium constant at \(50{0^o}C\) for the formation of \(N{H_3}\)according to the following equation? N2(g) + 3H2(g) 鈬� 2NH3(g)

An equilibrium mixture of \(N{H_3}(g)\) \({H_2}(g)\) and \({N_2}(g)\) at \(50{0^o}C\) was found to contain\(1.35M{H_2},1.15M{N_2}\)and \(4.12\)\( \times 1{0^{ - 1}}MN{H_3}\)

Question: The binding of oxygen by hemoglobin (Hb), giving oxyhemoglobin (HbO2), is partially regulated by the concentration of H3O+ and dissolved CO2 in the blood. Although the equilibrium is complicated, it can be summarized as

HbO2(aq) + H3 O+(aq) + CO2(g) 鈬� CO2 鈭扝产鈭扝+ + O2(g) + H2 O(l)

(a) Write the equilibrium constant expression for this reaction.

(b) Explain why the production of lactic acid and CO2 in a muscle during exertion stimulates release of O2 from the oxyhemoglobin in the blood passing through the muscle.

Assume that the change in pressure of \({H_2}S\) is small enough to be neglected in the following problem.(a) Calculate the equilibrium pressures of all species in an equilibrium mixture that results from the decomposition of H2S with an initial pressure of 0.824 atm.

\(2{H_2}S(g) \rightleftharpoons 2{H_2}(g) + {S_2}(g)\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,{K_p} = 2.2 \times {10^{( - 6)}}\)

(b) Show that the change is small enough to be neglected.

What are the concentrations of \(PC{l_5}\), \(PC{l_3}\), and \(C{l_2}\) in an equilibrium mixture produced by the decomposition of a sample of pure \(PC{l_5}\) with \([PC{l_5}] = 2.00M?\) \(PC{l_5}(g) \rightleftharpoons PC{l_3}(g) + {\mathbf{C}}{{\mathbf{l}}_{\mathbf{2}}}(g)\,\,\,\,\,\,\,\;{\mathbf{Kc}} = {\mathbf{0}}.{\mathbf{0}}211\)
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