For the reaction, 2SO2(g) + O2(g) => 2SO3(g), suggest the optimal conditions to maximise yield of SO3 when the forward reaction is exothermic.

So the most important thing to answer this question is to understand how an equilibrium works: it will always try to counteract the change you are making. For example, if you add more SO₃ into the mixture, it's going to try and reduce the amount of SO₃ by making more SO₂ and O₂. With this particular reaction we can see that there are more moles of gas on the left hand side of the equation (3 moles) than on the right (2 moles). We can therefore see that the equilibrium is going to try and shift to the right to counteract the higher pressure by making more SO₃ - which is the product that we want more of. One thing you could therefore do to make more SO₃ is to increase the pressure because the equilibrium is going to try and shift to the right to make more products, since they have a lower pressure.The forward reaction is also exothermic, meaning heat is lost to the environment (think exo- sounds like exit, and endo- sounds like enter) when the reaction happens. This means that if you increase the temperature, the reaction is going to try and lose more heat to the surroundings by carrying out more of the forward reaction, which will result in a higher yield of SO₃. Remember with equilibriums that a catalyst will not increase the yield of either the reactants or products! It will meant that equilibrium is reached faster but it speeds up the forward and reverse reactions equally, so there is no change to the overall amount of products or reactants.