Monitoring and controlling chemical reactions Study Guide

Many chemical reactions proceed in one direction, creating products from reactants in an irreversible way. However, some reactions are reversible, meaning that the products formed can react together under certain conditions to reform the original reactants. This is represented by a double arrow (⇌) in a chemical equation, indicating that the reaction can proceed in both the forward and reverse directions. A classic example is the hydration of anhydrous copper(II) sulfate. When water is added to white anhydrous copper(II) sulfate, it turns blue, forming hydrated copper(II) sulfate. This is the forward reaction. If the blue hydrated copper(II) sulfate is heated, the water is driven off, and the white anhydrous copper(II) sulfate is reformed. This is the reverse reaction. The key to controlling reversible reactions often lies in altering the reaction conditions. Factors like temperature, pressure (for gaseous reactions), and concentration of reactants or products can all influence the direction in which a reversible reaction is favoured. For example, heating often favours the endothermic direction (the reaction that absorbs heat), while cooling favours the exothermic direction (the reaction that releases heat). By understanding these principles, we can manipulate reaction conditions to maximise the yield of desired products or to reverse a reaction entirely.

Chemical reactions require a certain amount of energy input to start, known as the activation energy. This energy is needed to break existing bonds in the reactant molecules before new bonds can form to create products. Imagine a hill that reactants must climb over to become products; the height of this hill represents the activation energy. Without enough energy, reactant particles might collide but lack the force to break bonds and react.

A catalyst is a substance that speeds up the rate of a chemical reaction without being used up itself. Catalysts achieve this by providing an alternative reaction pathway, or mechanism, which has a lower activation energy than the uncatalysed reaction. Think of our hill analogy again; a catalyst in effect digs a tunnel through the hill or creates a much shallower slope, making it easier for reactants to turn into products.

By lowering the activation energy, more reactant particles will possess the minimum required energy to react when they collide, even at the same temperature. This leads to a higher frequency of successful collisions, resulting in a faster reaction rate. It is crucial to understand that a catalyst does not change the overall energy difference between reactants and products (the enthalpy change); it only affects the pathway and therefore the rate at which the reaction occurs. Catalysts are highly specific, meaning a particular catalyst often works for only one type of reaction or a specific set of reactions.