# Effect of Temperature on Chemical Reaction Equilibrium

here is the question (part I):

For part (i), I need some assistance, I cannot figure out how to do the question. I know eventually what to do, it's just working out the equilibrium constants I'm having trouble with. So to start, I want to work out the K values at $\pu{600K}$ and $\pu{800K}$; I do this by considering the total pressure:

$$\text{Total Pressure} = 1 \text{ bar}$$

Initially:

$$P_\ce{O2} + P_\ce{HCl} = \pu{1bar}$$ Since there are $4$ moles of $\ce{HCl}$ per $\ce{O2}$ we can write:

$$5P_\ce{O2} = \pu{1bar} \to P_{\ce{O2}} = \pu{0.2bar}$$

$85\%$ conversion means that in the end, $85\%$ of the oxygen has reacted; therefore, there is $\pu{0.2bar}*0.15 = \pu{0.03bar}$ oxygen remaining, therefore $\pu{0.12bar}~\ce{HCl}$ remaining. In addition to this, each mole of oxygen reacted results in $4$ new moles being created, and therefore there are $(0.85*0.2*2)\,\mathrm{bar}$ of water (and also chlorine) produced. With this information, I can calculate my equilibrium constant, right?

• What is the equation for determining the temperature dependence of the equilibrium constant for a reaction? Apr 21 '17 at 11:21
• No. The total pressure does not have to remain constant at 1 bar. What you need to do is review the derivation of the equilibrium constant from $\Delta G^0$ Apr 21 '17 at 11:57
• Are you familiar with $\Delta G=\Delta G^0+RT\ln{Q}$? Apr 21 '17 at 13:01
• So, when you use the equation $K=\exp({-\Delta G^0/RT})$, does the total pressure have to be 1 bar? Apr 21 '17 at 13:07
• Then I can't understand why you feel that you cannot use the equation involving $\Delta H^0$ to determine the value of K at 600K. Apr 21 '17 at 13:21