chapter 5 : chemical energetics

• Energy changes in chemical reactions

  • Energy changes occur, generally in the form of thermal energy
  • The energy change is either endothermic or exothermic
• Endothermic reaction

  • Heat energy is absorbed during the reaction
  • Positive enthalpy change
  • More energy is used to break bonds than is released in making bonds
• Exothermic reaction

  • Heat energy is released during the reaction
  • Negative enthalpy change
  • More energy is released making bonds than is used to break bonds
• Activation energy

  The minimum amount of energy required to start a reaction
• Standard conditions

  • Pressure: 101 kPa (or 100 kPa) or 1 atm
  • Temperature: 298 K or 25℃
  • Solution concentrations: 1 moldm-3
• Standard enthalpy change of reaction, ΔrHθ

  The enthalpy change that accompanies a reaction in the molar quantities expressed in an equation under standard conditions, with all reactants and products in their standard states
• Standard enthalpy change of formation, ΔfHθ

  The enthalpy change that occurs when one mole of product is formed from its constituent elements under standard conditions, with all reactants and products in their standard states
• Standard enthalpy change of combustion, ΔcHθ

  The enthalpy change that occurs when one mole of compound completely reacts with oxygen under standard conditions with all reactants and products in their standard states
• Standard enthalpy change of neutralisation, ΔneutHθ

  The enthalpy change that occurs when one mole of water is formed in a neutralisation reaction under standard conditions
• Standard enthalpy change of atomisation, ΔaHθ

  The enthalpy change that occurs when 1 mole of gaseous atoms are formed from an element in its standard state, under standard conditions
• Standard enthalpy change of hydration, ΔHhyd

  θ

  The enthalpy change that occurs when one mole of gaseous ions are completely hydrated by water to form one mole of aqueous ions, under standard conditions
• Standard enthalpy change of solution, ΔHsol

  θ

  The enthalpy change that occurs when one mole of ionic compound is completely dissolved in water to form aqueous ions, under standard conditions
• Bond energy

  • The energy in a covalent bond
  • A measure of the strength of a chemical bond
• Lattice enthalpy (A level only)

  The enthalpy change that occurs when 1 mole of a solid ionic compound is formed from its gaseous ions under standard conditions
• Directly calculating enthalpy change of a reaction

  1. q = mc∆T

  2. m - mass of the solution that changes temperature (1g = 1cm3)

  3. c - specific heat capacity (usually of water)

  4. ∆T - temperature change

  5. q - heat energy taken in/ released (joules)

  6. Divide q (in kJ) by the number of moles of the limiting reactant

  7. Add a sign to show whether ∆Hr

  θ is exothermic (-) or endothermic (+)
• Factors affecting lattice enthalpy (A level only)

  • Ionic charge
  • Ionic radius
• Ionic radius and lattice enthalpy (A level only)

  • A smaller ionic radius means the ions are closer together
  • As a result, the positive and negative ions are more strongly attracted together
  • Lattice enthalpy is more exothermic
• Ionic charge and lattice enthalpy (A level only)

  • A greater ionic charge means greater electrostatic attraction between positive and negative ions
  • Lattice enthalpy is more exothermic
• Hess' law


  The enthalpy change of a reaction is independent of the route it takes
• Using a Hess cycle to indirectly determine an enthalpy change

  1. There are two routes shown for this reaction

  2. The enthalpy change for ∆fHθ is equal to ∆Hθ1 + ∆Hθ2 - ∆Hθ3

  3. Make sure you multiply the enthalpy change by the stoichiometric ratio, i.e. 6 x ∆Hθ1

  4. Going with the arrow: add the value

  5. Going against the arrow: subtract the value
• First ionisation energy

  The energy required to remove an electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions under standard conditions
• First electron affinity

  The enthalpy change that takes place when each atom in one mole of gaseous atoms gains an electron to form one mole of gaseous 1- ions under standard conditions
• Average bond enthalpy

  The average enthalpy change that occurs when one mole of gaseous covalent bonds are broken
• Entropy

  The measure of the disorder of a system
• A system is more stable when its energy is spread out/ in a more disordered state. The higher the entropy, the more stable the system.
• Entropy increases from Solid → Liquid/Aqueous → Gas. This is because the system becomes more disordered.
• As temperature increases, the particles in a substance gain kinetic energy and become more disordered. As a result, entropy increases.
• Entropy increases if the number of gas molecules increases. This is because the gaseous state has the highest entropy as there is more disorder in the system.
• Calculate the entropy change of the system (given the standard entropies of reactants and products)


  ∆S = ΣSθ

  products - ΣSθ

  reactants
• ΔS will be positive when disorder increases, e.g. state change from solid to liquid/liquid to gas, increase in temperature, increase in the number of gaseous molecules.
• Gibbs free energy

  The overall change in energy during a reaction, consisting of entropy change (ΔS) and enthalpy change (ΔH) at a given temperature (T). ΔG = ΔH - TΔS