6.1.2.2. Effect of Pressure Changes (for reactions involving gases)

Pressure changes are only significant for reactions involving gases where there is a change in the total number of moles of gas. If the number of moles of gaseous reactants equals the number of moles of gaseous products (e.g., H₂(g) + I₂(g) ⇌ 2HI(g), where Δn_gas = 0), then pressure changes do not affect the position of equilibrium.

Pressure changes have different effects depending on the number of gas moles during a reaction. If the total number of moles of gas does not change in the reaction (Δn_gas = 0), adjusting the pressure will not have an impact on the equilibrium position. For example, in the reaction of hydrogen and iodine forming hydrogen iodide, there are 1 mole of H₂ and 1 mole of I₂ producing 2 moles of HI. In this case, changing pressure does not shift the equilibrium.

Think of a room filled with balloons (representing gas molecules). If you keep adding more balloons but the total count at the start and end are the same, squeezing the room tighter (increasing pressure) won’t change how many balloons are in there. But if you change the amounts of balloons, then changing the space around them can have a big impact.

Increasing the total pressure (by decreasing the volume of the container): The system attempts to reduce the pressure. It achieves this by favouring the side of the reaction with fewer moles of gas.

When the pressure in a closed system is increased by reducing the volume of the container, the reaction will shift towards the side where there are fewer moles of gas. This is the system's natural response to counter the pressure change. For example, if we have a reaction producing more gas moles on one side than the other, the equilibrium will shift towards the side with fewer moles to help balance the pressure.

Imagine trying to expand a crowded subway train. If you decrease the space (volume) inside the train, the people (gas molecules) will bunch up more. The train can only hold a set number of people comfortably, so they will move towards less crowded sections, just as gas molecules shift towards fewer moles to reduce pressure.

Decreasing the total pressure (by increasing the volume of the container): The system attempts to increase the pressure. It achieves this by favouring the side of the reaction with more moles of gas.

When the pressure is decreased in a closed system by increasing the volume, the equilibrium shifts toward the side with more gas moles. In trying to counteract the change in pressure, the system favors producing more gas to fill up the space and increase pressure.

Think of opening a door to let in fresh air into a room. When you increase the size of the room (volume), more air (gas molecules) can come in, filling it up and increasing the pressure. Similarly, in chemical reactions, reducing pressure will encourage the formation of more moles of gas to increase it back.

For N₂(g) + 3H₂(g) ⇌ 2NH₃(g), there are 4 moles of gaseous reactants and 2 moles of gaseous products. Therefore, increasing the pressure shifts the equilibrium to the right, increasing the yield of ammonia. Industrially, high pressures (e.g., 200 atm) are used.

In the reaction to synthesize ammonia (the Haber Process), we start with four moles of reactants (one mole of nitrogen and three moles of hydrogen) compared to two moles of products (ammonia). Increasing the pressure favors the production of ammonia, thus optimizing the reaction's yield. Industries utilize this by maintaining high pressure during the reaction.

Consider making a smoothie. If you have too many fruits (reactants) in a blender (reaction vessel), you won't reach the smooth consistency (products) you want. However, if you add more pressure by sealing the blender tight and blending longer, you make a smoother drink (more products), similar to how increased pressure pushes a gas reaction to favor products.

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