Chemical reactions Revision Notes

What is Activation Energy?

Activation energy is the minimum amount of kinetic energy that reactant particles must have for a collision to be successful and result in a chemical reaction. Without this minimum energy, particles will simply bounce off each other without any chemical change occurring. It acts as an energy barrier that must be overcome for the reaction to proceed.

Role of Activation Energy in Reactions

Every chemical reaction, whether it releases energy (exothermic) or absorbs energy (endothermic), requires an input of energy to start. This initial energy input is the activation energy. For example, a match needs to be struck (providing activation energy) to ignite, even though burning is an exothermic process. Similarly, endothermic reactions need continuous energy input to overcome this barrier and proceed.

Relating Activation Energy to Reaction Rate

If molecules collide with less than the activation energy, they will not react. Only collisions with energy equal to or greater than the activation energy will lead to the formation of products. Therefore, a higher activation energy means that a smaller proportion of reactant particles will have enough energy to react at a given temperature, generally leading to a slower reaction rate. Conversely, a lower activation energy means a larger proportion of particles can react, increasing the reaction rate.

What is Electrolysis?

Electrolysis is a process that uses electrical energy to break down or decompose a compound. This compound, called an electrolyte, must contain free-moving ions, either by being molten (liquid) or dissolved in a solvent (aqueous solution). Without these mobile ions, the substance cannot conduct electricity and thus cannot undergo electrolysis. The electrical energy causes chemical reactions that would not occur spontaneously.

Ions and Electrodes

In an electrolytic cell, there are two electrodes connected to a power supply. The positive electrode is called the anode, and the negative electrode is called the cathode.

• Anions are negatively charged ions that are attracted to the positive anode. At the anode, anions lose electrons (oxidation) to become neutral atoms or molecules.
• Cations are positively charged ions that are attracted to the negative cathode. At the cathode, cations gain electrons (reduction) to become neutral atoms or molecules.

Reactions at the Electrodes

The reactions at the electrodes are fundamental to understanding electrolysis.

• At the Anode (positive electrode): Oxidation occurs. Anions (negative ions) lose electrons. For example, if halide ions (Cl⁻, Br⁻, I⁻) are present, they are oxidised to their respective halogens (e.g., 2Cl⁻ → Cl₂ + 2e⁻). If no halide is present and the solution is aqueous, hydroxide ions (OH⁻) from water are oxidised, producing oxygen gas and water (4OH⁻ → O₂ + 2H₂O + 4e⁻).
• At the Cathode (negative electrode): Reduction occurs. Cations (positive ions) gain electrons. Reactive metals (e.g., K⁺, Na⁺, Ca²⁺, Mg²⁺, Al³⁺) are not discharged in aqueous solutions; instead, hydrogen ions (from water) are reduced (2H⁺ + 2e⁻ → H₂). Less reactive metal ions (e.g., Cu²⁺, Ag⁺) are reduced to form the metal (e.g., Cu²⁺ + 2e⁻ → Cu).

Defining Oxidation and Reduction by Oxygen Transfer

Oxidation is a chemical process where a substance gains oxygen. For instance, when magnesium reacts with oxygen to form magnesium oxide (2Mg + O₂ → 2MgO), the magnesium atoms gain oxygen and are thus oxidised. The magnesium metal changes into an ionic compound, losing electrons to the oxygen to form positive magnesium ions and negative oxide ions. Reduction is the opposite process, where a substance loses oxygen. An example is the reduction of iron(III) oxide by carbon monoxide in a blast furnace (Fe₂O₃ + 3CO → 2Fe + 3CO₂). Here, the iron(III) oxide loses oxygen atoms to become iron metal and is therefore reduced.

Identifying Oxidised and Reduced Species

In any redox reaction involving oxygen transfer, it is essential to identify which specific substance is oxidised and which one is reduced. The substance that gains oxygen is the one that has been oxidised. The substance that loses oxygen is the one that has been reduced. For example, in the reaction CuO + H₂ → Cu + H₂O, copper(II) oxide (CuO) loses oxygen to become copper (Cu), so CuO is reduced. Hydrogen (H₂) gains oxygen to become water (H₂O), so H₂ is oxidised. These processes always occur simultaneously; you cannot have oxidation without reduction, and vice versa.

Simultaneous Nature of Oxidation and Reduction

Oxidation and reduction reactions, when defined by oxygen transfer, always occur together. This is why they are often referred to as 'redox' reactions. If one reactant gains oxygen (is oxidised), there must be another reactant that provides that oxygen and therefore loses it (is reduced). Consider the combustion of methane: CH₄ + 2O₂ → CO₂ + 2H₂O. While this is more complex due to multiple products, the core idea holds: the methane loses hydrogen and gains oxygen (it is oxidised to carbon dioxide), and the oxygen reactant itself is incorporated into products. Focus on clear cases like metal oxides. In the reaction of lead(II) oxide with carbon, 2PbO + C → 2Pb + CO₂, the lead(II) oxide loses oxygen and is reduced, while the carbon gains oxygen and is oxidised.

Bond Breaking and Bond Making (Higher)

All chemical reactions involve two fundamental energy changes: bond breaking and bond making. Energy must be put into a system to overcome the forces holding atoms together in a chemical bond; this is an endothermic process. Conversely, energy is released from a system when new chemical bonds are formed, as atoms achieve a more stable state; this is an exothermic process. Understanding these opposing energy changes is crucial for calculating the overall enthalpy change of a reaction.

Calculating Overall Energy Change (Higher)

The overall energy change (enthalpy change, ΔH) for a reaction is calculated by subtracting the total energy released during bond formation from the total energy absorbed during bond breaking. The formula is: ΔH = Σ(bond energies of bonds broken) - Σ(bond energies of bonds formed). It's important to list all bonds in the reactants and products, including multiple bonds (e.g., C=O) and ensuring the correct number of each bond type is accounted for in balanced chemical equations. Average bond energies are used for these calculations, meaning the calculated ΔH is an approximation.

Interpreting Energy Change Results (Higher)

The sign of the calculated overall energy change (ΔH) indicates whether a reaction is exothermic or endothermic. If ΔH is negative, it means that more energy was released during bond formation than was absorbed for bond breaking, resulting in a net release of energy to the surroundings. This signifies an exothermic reaction, where the temperature of the surroundings increases. If ΔH is positive, it means more energy was absorbed for bond breaking than was released during bond formation, resulting in a net absorption of energy from the surroundings. This signifies an endothermic reaction, where the temperature of the surroundings decreases.