Thermochemical Reactions

THERMOCHEMICAL REACTIONS Introducing enthalpy All things contain chemical energy. This energy comes from the attractions between electrons and protons in atoms, repulsions between nuclei, repulsions between electrons, movement of electrons and vibrations and rotations around bonds. A chemical reaction involves the rearrangement of this energy from reactants to products. This change in chemical energy (either up or down) is called enthalpy. It is also referred to as the heat content. It is denoted by the symbol H. There are a number of terms you should be familiar with for this topic. The system refers to the chemical reactions itself, including the chemicals involved. The surroundings refer to everything else external to the reaction. Energy will change within the system and can also be released into or absorbed by the surroundings. As the reactants rearrange bonds to form the products, these different bonds have different chemical energies. Energy will be released or absorbed by the reaction depending on the relative energies of the bonds within the reactants and products. Energy can be released in the surroundings in many forms; light, electricity, movement (kinetic) energy and thermal (heat) energy. A great example of chemical transformations leading to loss of energy is glow-in-the-dark light sticks. When the glow stick is broken, diphenyl oxalate reacts with hydrogen peroxide to produce phenol, carbon dioxide and energy. This energy is transferred to a dye, exciting the electrons which is then released as colourful light as they relax to the ground state. This process is called chemiluminescence - release of light from a chemical reaction. Exothermic & endothermic reactions When the total chemical energy of the products is lower than the total chemical energy of the reactants, it causes a release of energy from the system into the surroundings. This is called an exothermic reaction. This energy can be written into the chemical equation. CH4(g) + 2 O2(g) → CO2(g) + 2H2O(g) + "energy" This energy is released as heat into the surroundings. All combustion reactions release heat energy and are therefore exothermic reactions. When the energy of the products is higher than the energy of the reactants, it requires energy to be absorbed into the system from the surroundings. This is called an endothermic reaction. Endothermic reactions require a constant input of energy. This energy can also be written into the chemical equation. CaCO3(s) + "energy" → CaO(s) + CO2(g) Changes of state (physical changes) also involve energy being absorbed or released e.g. melting of a solid to a liquid. The melting of ice into water requires absorption of energy, and therefore an endothermic process. The boiling of water to produce water vapour is also endothermic. The condensation of a gas to a liquid and freezing a liquid to form a solid release heat into the surroundings and are exothermic processes. Instant cold packs are a great example of an endothermic reaction. When the cold pack is squeezed, ammonium nitrate is dissolved in water which is an endothermic process. This sucks heat from the surroundings and lowers the pack's temperature. Enthalpy change The chemical energy of a substance is called its heat content or enthalpy. The symbol H represents enthalpy. The enthalpy of the reactants is Hr and the enthalpy of the products is Hp. The change in the heat energy between the reactants and products is equal to the energy absorbed by or released into the surroundings. This exchange of heat energy between the system and the surroundings under constant pressure is referred to as the enthalpy change or heat of reaction, and given the symbol ΔΗ, where the A means 'change in'. The units for ΔΗ are most commonly kJ mol-1, but can also be given in kJ g-1. For the general reaction: reactants → products The enthalpy change (ΔΗ) is calculated by: ΔΗ = Hp – Hr Exothermic reactions When Hp is less than Hr, meaning the heat energy of the products is less than the heat energy of the reactants, energy is released from the system into the surroundings. Therefore, the reaction is exothermic. For exothermic reactions (including combustion reactions), ΔΗ < 0. Endothermic reactions When Hp is greater than Hr, energy is absorbed from the surroundings. Therefore, the reaction is endothermic. For endothermic reactions, the system has gained energy so ΔΗ > 0. Writing combustion reactions Combustion reactions are great examples of exothermic reactions. Combustion reactions always release energy. Combustion, or burning, requires three things: 1. Fuel to burn (chemical reagents) 2. Oxygen for the fuel to burn 3. Energy to get the process started Because we are always adding oxygen, a combustion reaction is also considered an oxidation reaction. Complete combustion A complete combustion reaction occurs when there is plentiful oxygen. Carbon dioxide and water are the only products. The combustion of methane is a typical example: CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(I) Incomplete combustion An incomplete combustion reaction occurs when oxygen is in limited supply. Therefore, not all of the carbon can be converted to carbon dioxide. Instead, carbon monoxide and water are produced. Some carbon may also be produced. The incomplete combustion of methane is shown below: 2 CH4(g) + 3 O2(g) → 2CO(g) + 4 H2O(I) We can see the difference between a complete versus incomplete combustion reaction with a Bunsen burner. When the flame is yellow, it is an incomplete combustion reaction. The carbon produced causes the yellow colour. The flame is blue when the collar is open, allowing extra oxygen in and so undergoes a complete combustion reaction.