Given is a following reaction at equilibrium: N2(g) + 3H2(g) ⇄ 2NH3(g), ΔH < 0. What will be the effect of changing the following conditions on the system? 1. Increasing pressure. 2. Decreasing temperature. 3. Adding a catalyst. 4. Adding HCl(g).

The correct approach to this question is to use the Le Chatelier's principle, which states that when you change conditions of a system at equilibrium, the system counteracts the change - a new equilibrium is established. The conditions that can be changed are: pressure, temperature, concentration, volume or adding/removing certain reagents from the reaction mixture. Let's have a look at all the conditions in the question.

1. We increase the pressure of the system. The system wants to do something to decrease the pressure (counteract). What can it do? Well, if a molecule of nitrogen reacts with three molecules of hydrogen, two molecules of ammonia are formed - this reduces the no. molecules in the system and hence the pressure (remember Avogadro's law: equal volumes of all gases, at the same temperature and pressure, have the same number of molecules - so no matter what gas, less molecules mean lower pressure at constant temperature and volume). So the answer here is: equilibrium is shifted towards ammonia. 2. The enthalpy of the reaction is negative - it's an exothermic reaction. If heat is produced in the reaction and we remove the heat (by decreasing temperature), the system wants to produce more heat, so ammonia will be formed. Equilibrium shifts towards ammonia as well. 3. A catalyst doesn't shift equilibrium at all - it just accelerates reaching the equilibrium by a system. No shift here! 4. Gaseous HCl reacts with gaseous ammonia to produce ammonium chloride (solid). The formation of solid drives the reaction strongly towards the salt. So basically, ammonia is removed from the reaction mixture. That means that the system wants to produce more ammonia! Equilibrium shifts towards ammonia.