Monitoring and controlling chemical reactions Study Guide

In reversible reactions, a state of equilibrium is reached when the rate of the forward reaction equals the rate of the backward reaction, and the concentrations of reactants and products remain constant. Le Chatelier's Principle states that if the conditions of a system at equilibrium are changed, the system will shift in a direction that counteracts the change. This principle allows us to predict how changes in temperature, pressure, and concentration will affect the position of equilibrium and, consequently, the yield of products.

For an exothermic forward reaction, increasing the temperature will shift the equilibrium to the left (towards reactants) to absorb the added heat, decreasing product yield. Conversely, decreasing the temperature will shift it to the right (towards products) to release heat. For an endothermic forward reaction, the effects are opposite.

Changes in pressure specifically affect reactions involving gases. Increasing the pressure (by decreasing volume) will shift the equilibrium towards the side with fewer moles of gas to reduce the pressure. Decreasing the pressure will shift it towards the side with more moles of gas. If there are equal moles of gas on both sides, pressure changes have no effect on equilibrium position.

Changing the concentration of a reactant or product also shifts the equilibrium. Increasing the concentration of a reactant will shift the equilibrium to the right (towards products) to consume the added reactant. Increasing the concentration of a product will shift it to the left (towards reactants) to reduce the added product. Catalysts increase the rate of both forward and backward reactions equally, so they do not affect the position of equilibrium, only the speed at which equilibrium is reached. To maximise product yield, industrial processes carefully select conditions based on these principles, considering both the desired outcome and economic factors (e.g., energy costs for high temperatures or pressures).

Chemical reactions occur when reactant particles collide with sufficient energy. This is known as collision theory. The rate of a reaction is determined by how often these successful collisions happen. To increase the rate of reaction, we need to increase the frequency of collisions, the energy of collisions, or both. Factors such as temperature, concentration, and pressure (for gases) significantly influence these collision parameters.

Temperature directly affects the kinetic energy of particles. Higher temperatures mean particles move faster, leading to more frequent collisions. Crucially, a greater proportion of these collisions will also have energy equal to or greater than the activation energy, making them successful.

Concentration refers to the number of reactant particles in a given volume. Increasing the concentration of reactants means there are more particles packed into the same space, leading to a higher frequency of collisions. While the average energy of collisions doesn't change, the likelihood of a successful collision occurring increases simply because there are more opportunities for collision.

Pressure specifically applies to reactions involving gases. Increasing the pressure of a gas means compressing the same number of gas particles into a smaller volume. This effectively increases the concentration of the gas, leading to more frequent collisions between reactant particles and therefore a faster reaction rate. The kinetic energy of the particles themselves is not directly affected by the pressure change, but the proximity of particles increases the collision frequency.