definitions

• Enthalpy change:

  Heat energy change in a reaction at constant pressure (kJmol​ ).
• Standard conditions:​


  100 kPa,

  298 K,

  (1 moldm​ concentrations of ions)
• Exothermic​:

  Heat given off;

  −​v​e value of enthalpy change.
• Endothermic​:


  Heat absorbed;

  +ve value of enthalpy change.
• Standard enthalpy change of reaction:

  Enthalpy change when reaction occurs in the molar quantities shown in the chemical equation under standard conditions.
• Standard enthalpy change of formation:

  Enthalpy change when 1 mole of a compound is formed from its constituent elements in their standard states under standard conditions of 100kPa & 298K.
• Standard enthalpy of combustion: 


  Enthalpy change when 1 mole of a substance is completely burned in oxygen under their standard states under standard conditions of 100kPa & 298K.
• Standard enthalpy of neutralisation:

  Enthalpy change when an acid and alkali react together under standard conditions (100kPa & 298K) to form 1 mole of water.
• Equation for calorimetry calculations:

  heat change ​Q = mCΔT​,

  m = mass, C = specific heat capacity, ΔT = change in temperature.
• Specific heat capacity:​

  The amount of energy needed to raise a temperature of 1 g of a substance by 1 degree.
• Hess’s law:​


  The total enthalpy change is independent of the reaction pathway taken.
• Bond Enthalpy:


  Amount of energy required to break 1 mole of the stated bond in the gas phase.
• In terms of bond enthalpies:


  Enthalpy Change of reaction = Sum of bond enthalpies of reactants – Sum of bond enthalpies of products.
• Mean bond enthalpy: 


  a​verage amount of energy needed to break a specific type of bond, measured over a variety of different molecules.