Methanol is broken down in the body during digestion - AQA - GCSE Chemistry - Question 9 - 2018 - Paper 2

According to the stoichiometry of the reaction, 1 mole of carbon monoxide reacts with 2 moles of hydrogen. Therefore, to find the moles of CO required for 4.0 x 10^3 moles of H2, we can set up the ratio:

$n_{CO} : n_{H2} = 1 : 2$

From this, we calculate:

n_{CO} = rac{4.0 imes 10^3 ext{ moles H}_2}{2} = 2.0 imes 10^3 ext{ moles CO}

Thus, 2.0 x 10^3 moles of carbon monoxide are needed.

If the temperature is raised above 250 °C, the yield of methanol will decrease. This occurs because the forward reaction is exothermic; increasing the temperature shifts the equilibrium position to favor the endothermic reverse reaction, thereby producing less methanol.

Increasing the pressure results in a greater yield of methanol due to Le Chatelier's principle. In this reaction, there are fewer gas molecules on the product side (1 mole of CH3OH) compared to the reactant side (1 mole of CO and 2 moles of H2, totaling 3 moles). A higher pressure shifts the equilibrium towards the side with fewer molecules, increasing the yield of methanol. Additionally, higher pressure increases the rate of reaction by raising the concentration of gaseous reactants, leading to more frequent collisions.

A catalyst increases the rate of a reaction by providing an alternative reaction pathway with a lower activation energy. This means that more reactant molecules have sufficient energy to overcome the barrier, leading to an increase in the frequency of successful collisions and thus accelerating the reaction.

A catalyst is used in this industrial process to enable the reaction to occur at lower temperatures and pressures, which reduces the overall energy costs and makes the process more economically viable.

Using a catalyst does not affect the position of equilibrium, meaning that while it speeds up the rate at which equilibrium is reached, it does not change the equilibrium yield of methanol.