Sergey Bylikin, Gary Horner and, Brian Murphy Solutions for Chapter: Equilibrium, Exercise 7: Questions

The equation for a reversible reaction used in industry to convert methane to hydrogen is shown below.

${\mathrm{CH}}_4\left(\mathrm{g}\right)+{\mathrm{H}}_2\mathrm{O}\left(\mathrm{g}\right)\rightleftharpoons \mathrm{CO}\left(\mathrm{g}\right)+ 3{\mathrm{H}}_2\left(\mathrm{g}\right), ∆{\mathrm{H}}^{⊝}=+210 \mathrm{kJ}$

Which statement is always correct about this reaction when equilibrium has been reached?

Sulphur dioxide and oxygen react to form sulphur trioxide according to the equation:

$2{\mathrm{SO}}_2\left(\mathrm{g}\right)+{\mathrm{O}}_2\left(\mathrm{g}\right)\rightleftharpoons 2{\mathrm{SO}}_3\left(\mathrm{g}\right)$

How are the amount of ${\mathrm{SO}}_2$ and the value of the equilibrium constant for the reaction affected by an increase in pressure?

What will happen to the position of equilibrium and the value of the equilibrium constant when the temperature is increased in the following reaction?

${\mathrm{Br}}_2\left(\mathrm{g}\right)+{\mathrm{Cl}}_2\left(\mathrm{g}\right)\rightleftharpoons 2\mathrm{BrCl}\left(\mathrm{g}\right), ∆\mathrm{H}=+14\mathrm{kJ}$.