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Chapter 14: Problem 19

Predict the effect of each of the following on the indicated equilibrium system in terms of which reaction (forward, reverse, or neither) will be favored. $$\mathrm{H}_{2}(g)+\mathrm{Cl}_{2}(g) \rightleftarrows 2 \mathrm{HCl}(g)+184 \mathrm{kJ}$$ a. addition of \(\mathrm{Cl}_{2}\) b. removal of \(\mathrm{HCl}\) c. increased pressure d. decreased temperature e. removal of \(\mathrm{H}_{2}\) g. decreased pressure g. increased temperature h. decreased system volume

Short Answer

Expert verified
a. Forward reaction is favored. b. Forward reaction is favored. c. Neither reaction is favored. d. Forward reaction is favored. e. Reverse reaction is favored. g. Neither reaction is favored. g. Reverse reaction is favored. h. Neither reaction is favored.

Step by step solution

01

Effect of Addition of Cl2

Adding more \(Cl_2\) increases the concentration of one of the reactants. According to Le Chatelier's principle, this will shift the equilibrium to the right i.e. favor the forward reaction to consume the added Cl2.
02

Effect of Removal of HCl

The removal of \(HCl\) decreases the concentration of the product. This will cause the equilibrium to shift to the right, favoring the forward reaction to replace the removed \(HCl\).
03

Effect of Increased Pressure

An increase in pressure will favor the side with fewer moles of gas. In this case, there are 2 moles of gas on the reactant side and 2 moles of gas on the product side. Therefore, increasing the pressure will not favor either the forward or reverse reaction.
04

Effect of Decreased Temperature

Decreasing the temperature will favor the exothermic reaction i.e. the reaction that releases heat. In this case, the forward reaction is exothermic (since it releases 184 kJ), so lowering the temperature will favor the forward reaction.
05

Effect of Removal of H2

The removal of \(H_2\) will decrease the concentration of one of the reactants, causing the equilibrium to shift to the left to replace the removed \(H_2\). This will favor the reverse reaction.
06

Effect of Decreased Pressure

Decreasing the pressure will favor the side with more moles of gas. As the number of moles of gas is equal on both sides, decreasing the pressure will not favor either the forward or the reverse reaction.
07

Effect of Increased Temperature

Increasing the temperature will favor the endothermic reaction i.e. the reaction that absorbs heat. In this case, the reverse reaction is endothermic (since the forward reaction is exothermic), so raising the temperature will favor the reverse reaction.
08

Effect of Decreased System Volume

Decreasing the system volume is equivalent to increasing the pressure. Similar to step 3, the equilibrium will not shift either to the right or to the left because the number of moles of gas is equal on both sides.

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Most popular questions from this chapter

The solubility of cobalt \((\mathrm{II})\) sulfide, coS, is \(1.7 \times 10^{-13} \mathrm{M}\) . Calculate the solubility product constant for \(\mathrm{CoS} .\)

Imagine the following hypothetical reaction, taking place in a sealed, rigid container, to be neither exothermic nor endothermic. $$\mathrm{A}(g)+\mathrm{B}(g) \rightleftarrows \mathrm{C}(g) \quad \Delta H=0$$ Would an increase in temperature favor the forward reaction or the reverse reaction? (Hint: Recall the gas laws.)

Vinegar-a solution of acetic acid, \(\mathrm{CH}_{3} \mathrm{COOH},\) in water \(-\) is used in varying concentrations for different household tasks. The following equilibrum exists in vinegar. $$\begin{array}{c}{\mathrm{CH}_{3} \mathrm{COOH}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightleftarrows} \\ {\mathrm{H}_{3} \mathrm{O}^{+}(a q)+\mathrm{CH}_{3} \mathrm{COO}^{-}(a q)}\end{array}$$ If the concentration of the acetic acid solution at equilibrium is 3.00 \(\mathrm{M}\) and the \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=\left[\mathrm{CH}_{3} \mathrm{COO}^{-}\right]=7.22 \times 10^{-3},\) what is the \(K_{e q}\) value for acetic acid?

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Develop a model that shows the concept of equilibrium. Be sure that your model includes the impact of Le Chatelier's principle on equilibrium.
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