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Chapter 15: Problem 34

Starting with \(0.280 \mathrm{mol} \mathrm{SbCl}_{3}\) and \(0.160 \mathrm{mol} \mathrm{Cl}_{2},\) how many moles of \(\mathrm{SbCl}_{5}, \mathrm{SbCl}_{3},\) and \(\mathrm{Cl}_{2}\) are present when equilibrium is established at \(248^{\circ} \mathrm{C}\) in a 2.50 L flask? $$\begin{aligned} \mathrm{SbCl}_{5}(\mathrm{g}) \rightleftharpoons \mathrm{SbCl}_{3}(\mathrm{g})+\mathrm{Cl}_{2}(\mathrm{g}) & \\ K_{\mathrm{c}}=& 2.5 \times 10^{-2} \mathrm{at} \ 248^{\circ} \mathrm{C} \end{aligned}$$

Short Answer

Expert verified
To find the final moles of all compounds at equilibrium, create an ICE table, express and solve equilibrium constant expression in terms of X, and substitute the value of X back into the ICE table.

Step by step solution

01

Identify Initial Moles of Compounds

From the problem, we know there are 0.280 mol of SbCl3 and 0.160 mol of Cl2 initially. The problem doesn’t provide the initial amount of SbCl5, so we assume there is a zero mol of SbCl5. This is the 'I' in our ICE table.
02

Create an ICE Table

We then create an ICE table that represents the Initial moles (I), Change in moles (C) and Equilibrium moles (E) for each compound in the reaction. When the system reaches equilibrium, some amount X (in mol) of SbCl5 forms by using up same amount X of SbCl3 and Cl2. Thus, the 'C' part of ICE table will be +x for SbCl5 and -x for SbCl3 and Cl2.
03

Calculate Final Moles at Equilibrium

We then fill out the 'E' part of the ICE table. At equilibrium, there will be x mol of SbCl5, (0.280-x) mol of SbCl3, and (0.160-x) mol of Cl2. These values are got by adding the initial moles and change in moles.
04

Use KC to Find the Value of X

Next, plug the equilibrium concentrations into the expression for the equilibrium constant and solve for x. The equilibrium molar concentration of a compound is its moles at equilibrium divided by the volume of the flask. The equilibrium constant KC is given as 2.5 x 10^-2.
05

Calculate the Moles of Compounds at Equilibrium

After finding the value of X, substitute it back into the 'E' part of the ICE table to find the final moles of all three compounds at equilibrium.

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Most popular questions from this chapter

In your own words, define or explain the following terms or symbols: (a) \(K_{\mathrm{p}} ;\) (b) \(Q_{\mathrm{c}} ;\) (c) \(\Delta n_{\text {gas }}\)

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Starting with \(\mathrm{SO}_{3}(\mathrm{g})\) at \(1.00 \mathrm{atm},\) what will be the total pressure when equilibrium is reached in the following reaction at \(700 \mathrm{K} ?\) \(2 \mathrm{SO}_{3}(\mathrm{g}) \rightleftharpoons 2 \mathrm{SO}_{2}(\mathrm{g})+\mathrm{O}_{2}(\mathrm{g}) \quad K_{\mathrm{p}}=1.6 \times 10^{-5}\)

We can represent the freezing of \(\mathrm{H}_{2} \mathrm{O}(1)\) at \(0^{\circ} \mathrm{C}\ \mathrm{as} \mathrm{H}_{2} \mathrm{O}\) \(\left(1, d=1.00 \mathrm{g} / \mathrm{cm}^{3}\right) \rightleftharpoons \mathrm{H}_{2} \mathrm{O}\left(\mathrm{s}, d=0.92 \mathrm{g} / \mathrm{cm}^{3}\right) . \quad \mathrm{Ex}\) plain why increasing the pressure on ice causes it to melt. Is this the behavior you expect for solids in general? Explain.
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