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Chapter 12: Problem 52

Predict the direction in which each of the following equilibria will shift if the pressure on the system is decreased by expansion. (a) \(\mathrm{Ni}(s)+4 \mathrm{CO}(g) \rightleftharpoons \mathrm{Ni}(\mathrm{CO})_{4}(g)\) (b) \(\mathrm{CI} \mathrm{F}_{5}(g) \rightleftharpoons \mathrm{Cl} \mathrm{F}_{3}(g)+\mathrm{F}_{2}(g)\) (c) \(\mathrm{HBr}(g) \rightleftharpoons \frac{1}{2} \mathrm{H}_{2}(g)+\frac{1}{2} \mathrm{Br}_{2}(g)\)

Short Answer

Expert verified
Question: Predict the direction of the following chemical equilibria when pressure on the system is decreased by expansion: a) Ni(s) + 4 CO(g) ⇌ Ni(CO)4(g) b) ClF5(g) ⇌ ClF3(g) + F2(g) c) HBr(g) ⇌ 1/2 H2(g) + 1/2 Br2(g) Answer: When pressure is decreased by expansion, the equilibrium for the given reactions will shift as follows: (a) to the left (b) to the right (c) no shift

Step by step solution

01

a) Ni(s) + 4 CO(g) ⇌ Ni(CO)4(g)

First, count the number of moles of gas on each side of the equation. On the left side of the equilibrium, there are 4 moles of CO gas. On the right side, there is 1 mole of Ni(CO)4 gas. Since there are more moles of gas on the left side, the equilibrium will shift to the left when pressure is decreased by expansion.
02

b) ClF5(g) ⇌ ClF3(g) + F2(g)

Now, count the number of moles of gas in this reaction. On the left side of the equilibrium, there is 1 mole of ClF5 gas. On the right side, there are 1 mole of ClF3 gas and 1 mole of F2 gas, for a total of 2 moles of gas. Since there are more moles of gas on the right side, the equilibrium will shift to the right when pressure is decreased by expansion.
03

c) HBr(g) ⇌ 1/2 H2(g) + 1/2 Br2(g)

Finally, count the number of moles of gas in this last reaction. On the left side of the equilibrium, there is 1 mole of HBr gas. On the right side, there are 1/2 mole of H2 gas and 1/2 mole of Br2 gas, for a total of 1 mole of gas. Since there is an equal number of moles of gas on both sides, the equilibrium is not influenced by the decrease in pressure, and there will be no shift in the equilibrium in this case. In summary, when pressure is decreased by expansion, the equilibrium for the given reactions will shift as follows: (a) to the left (b) to the right (c) no shift

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Most popular questions from this chapter

Consider the decomposition of ammonium hydrogen sulfide: $$\mathrm{NH}_{4} \mathrm{HS}(s) \rightleftharpoons \mathrm{NH}_{3}(g)+\mathrm{H}_{2} \mathrm{~S}(g)$$ In a sealed flask at \(25^{\circ} \mathrm{C}\) are \(10.0 \mathrm{~g}\) of \(\mathrm{NH}_{4} \mathrm{HS}\), ammonia with a partial pressure of \(0.692 \mathrm{~atm}\), and \(\mathrm{H}_{2} \mathrm{~S}\) with a partial pressure of \(0.0532 \mathrm{~atm}\). When equilibrium is established, it is found that the partial pressure of ammonia has increased by 12.4\%. Calculate \(K\) for the decomposition of \(\mathrm{NH}_{4} \mathrm{HS}\) at \(25^{\circ} \mathrm{C}\).

. Consider the equilibrium $$\mathrm{N}_{2}(g)+\mathrm{O}_{2}(g) \rightleftharpoons 2 \mathrm{NO}(g)$$ At a certain temperature, the equilibrium constant for the reaction is \(0.0255\). What are the partial pressures of all gases at equilibrium if the initial partial pressure of the gases (both reactants and products) is \(0.300 \mathrm{~atm} ?\)

At a certain temperature, nitrogen and oxygen gases combine to form nitrogen oxide gas. $$\mathrm{N}_{2}(\mathrm{~g})+\mathrm{O}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NO}(g)$$ When equilibrium is established, the partial pressures of the gases are: \(P_{\mathrm{N}_{2}}=\) \(1.2 \mathrm{~atm}, P_{\mathrm{O}_{2}}=0.80 \mathrm{~atm}, P_{\mathrm{NO}}=0.022 \mathrm{~atm} .\) (a) Calculate \(K\) at the temperature of the reaction. (b) After equilibrium is reached, more oxygen is added to make its partial pressure \(1.2\) atm. Calculate the partial pressure of all gases when equilibrium is reestablished.

The following data are for the system $$\mathrm{A}(g) \rightleftharpoons 2 \mathrm{~B}(g)$$ $$\begin{array}{lcccccc}\hline \text { Time (s) } & 0 & 20 & 40 & 60 & 80 & 100 \\ P_{\mathrm{A}} \text { (atm) } & 1.00 & 0.83 & 0.72 & 0.65 & 0.62 & 0.62 \\ P_{B} \text { (atm) } & 0.00 & 0.34 & 0.56 & 0.70 & 0.76 & 0.76 \\\\\hline\end{array}$$ (a) How long does it take the system to reach equilibrium? (b) How does the rate of the forward reaction compare with the rate of the reverse reaction after 30 s? After 90 s?

Iodine chloride decomposes at high temperatures to iodine and chlorine gases. $$2 \mathrm{ICl}(g) \rightleftharpoons \mathrm{I}_{2}(g)+\mathrm{Cl}_{2}(g)$$ Equilibrium is established at a certain temperature when the partial pressures of ICI, \(\mathrm{I}_{2}\), and \(\mathrm{Cl}_{2}\) are (in atm) \(0.43,0.16\), and \(0.27\), respectively. (a) Calculate \(K\). (b) If enough iodine condenses to decrease its partial pressure to \(0.10 \mathrm{~atm}\), in which direction will the reaction proceed? What is the partial pressure of iodine when equilibrium is re-established?
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