Monitoring and controlling chemical reactions Revision Notes

Temperature and Reaction Rate

When the temperature of a reaction mixture increases, the reactant particles gain more kinetic energy. This increased energy means the particles move faster and collide more frequently. Crucially, a larger proportion of these collisions will have energy equal to or greater than the activation energy, making them successful collisions. Therefore, an increase in temperature generally leads to an increase in the rate of reaction. Conversely, decreasing the temperature reduces the kinetic energy of particles, leading to fewer and less energetic collisions, and thus a slower reaction rate.

Concentration and Pressure and Reaction Rate

For reactions occurring in solutions, increasing the concentration of reactants means there are more reactant particles per unit volume. This leads to a higher frequency of collisions between reactant particles, resulting in a faster rate of reaction. For reactions involving gases, increasing the pressure has a similar effect. When pressure increases, gas particles are forced into a smaller volume, increasing their concentration. This close proximity means they collide more frequently, leading to a faster reaction rate. Decreasing concentration or pressure will have the opposite effect, slowing down the reaction.

Surface Area and Reaction Rate

When a solid reactant is involved in a reaction, only the particles on its surface are available to collide with other reactant particles. If the solid is broken into smaller pieces or ground into a powder, its total surface area increases significantly. This larger exposed surface allows more reactant particles to come into contact and collide with the other reactants simultaneously. As a result, the frequency of successful collisions increases, and the reaction proceeds at a faster rate. A classic example is a sugar cube dissolving more slowly than granulated sugar in water.

Collision Theory Fundamentals

For a chemical reaction to occur, reactant particles must collide with each other. Not all collisions lead to a reaction; only collisions with energy equal to or greater than the activation energy, and with the correct orientation, are successful. The faster the rate of reaction, the more frequent these successful collisions are. Anything that increases the number of successful collisions per unit time will increase the reaction rate.

Effect of Temperature

Increasing the temperature of a reaction mixture increases the kinetic energy of reactant particles. This has two key effects:

1. Particles move faster, leading to a higher frequency of collisions.
2. A greater proportion of these collisions will have sufficient energy (activation energy) to be successful. Both factors contribute significantly to a faster reaction rate. Conversely, decreasing temperature slows down the reaction.

Effect of Concentration

For reactions in solutions, increasing the concentration of a reactant means there are more reactant particles dissolved in the same volume of solvent. This results in the particles being closer together and therefore colliding more frequently. While the energy of individual collisions does not change, the increased frequency of collisions leads to a greater number of successful collisions per unit time, thus increasing the reaction rate. Diluting reactants slows the reaction.

Effect of Pressure (for Gases)

For reactions involving gaseous reactants, increasing the pressure means compressing the gas into a smaller volume. This effectively increases the concentration of gas particles. The particles are therefore closer together and collide more frequently. This increased frequency of collisions leads to a higher rate of successful collisions and thus a faster reaction rate. Decreasing pressure has the opposite effect.

What is Activation Energy?

Activation energy is the minimum amount of energy that reactant particles must possess for a chemical reaction to occur upon collision. It represents the energy barrier that must be overcome to break existing bonds and initiate the rearrangement of atoms into products. Without sufficient activation energy, collisions between reactant particles, even if frequent, will not lead to a chemical transformation.

How Catalysts Work

Catalysts speed up chemical reactions by providing an alternative reaction pathway or mechanism that has a lower activation energy. This means that a smaller amount of energy is required for reactant particles to successfully collide and form products. The catalyst itself is not consumed in the reaction and can be recovered unchanged afterwards.

Impact of Lowered Activation Energy

When a catalyst lowers the activation energy, a greater proportion of reactant particles will have the necessary energy to react at a given temperature. This increases the frequency of effective collisions, leading to a faster rate of reaction. It's important to note that catalysts do not affect the overall energy change of the reaction, only the speed at which it reaches completion.

Properties of Catalysts

Catalysts are highly specific, meaning that one catalyst typically works for a particular reaction or a small group of reactions. They are not used up during the reaction, suggesting they can be reused. They are vital in many industrial processes to lower energy costs and increase efficiency by allowing reactions to proceed faster or at lower temperatures and pressures.

Le Chatelier's Principle: Predicting Shifts (Higher)

Le Chatelier's Principle states that if a change of condition is applied to a system in equilibrium, the system will shift in a direction that counteracts the change. This principle is crucial for understanding and predicting how to maximise yields in industrial chemical processes. For instance, in the Haber process (N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = -92 kJ/mol), conditions are chosen to favour ammonia production.

Effect of Temperature on Equilibrium (Higher)

If the forward reaction is exothermic (releases heat, ΔH is negative), increasing temperature shifts the equilibrium to the left (towards reactants) to absorb the extra heat, reducing the yield of products. Decreasing temperature shifts it to the right (towards products) to release heat, increasing product yield. If the forward reaction is endothermic (absorbs heat, ΔH is positive), the opposite effects occur: increasing temperature shifts to the right, decreasing temperature shifts to the left. Optimal temperature selection balances yield with reaction rate.

Effect of Pressure on Equilibrium (Higher)

Pressure changes primarily affect reactions involving gases. Increasing the pressure (by reducing volume) shifts the equilibrium towards the side with fewer moles of gas to reduce the pressure. Decreasing the pressure shifts it towards the side with more moles of gas. If the number of moles of gas is equal on both sides of the equation, pressure changes have no effect on the equilibrium position. For example, in the Haber process, there are 4 moles of gas on the left and 2 moles on the right, so high pressure favours ammonia production.

Effect of Concentration and Catalysts (Higher)

Increasing the concentration of a reactant shifts the equilibrium to the right (towards products) to use up the added reactant. Conversely, increasing the concentration of a product shifts the equilibrium to the left (towards reactants) to reduce the added product. Removing a product as it forms will therefore shift the equilibrium to the right, continuously producing more. Catalysts speed up both the forward and backward reactions equally. They do not change the position of equilibrium or the product yield, but they help the system reach equilibrium faster, which is often beneficial for industrial efficiency.