Energetics

• Reactions can be endothermic or exothermic.
• Enthalpy change ($\Delta H$) is the heat energy change measured under conditions of constant pressure.
• Standard enthalpy changes refer to standard conditions, ie 100 kPa and a stated temperature.
• The enthalpy of combustion ($\Delta_cH^{\theta}$) is the enthalpy change when 1 mole of a substance is burned completely in excess oxygen with all reactants and products in their standard states under standard conditions.
• The enthalpy of formation ($\Delta_fH^{\theta}$) is the enthalpy change when 1 mole of a compound is formed from its constituent elements with all reactants and products in standard states under standard conditions.
• The heat change, q, in a reaction is given by the equation q = mc∆T
• q = mc∆T

  q = Change in heat energy in joules (J).

  m = mass in grams of the substance to which the temperature change occurs (usually water).

  c = Specific heat capacity.

  ∆T = Temperature change in K
• q = mc∆T
• $\Delta H=$$\frac{q}{n}$
• Enthalpy change can be calculated by dividing change in heat energy (q) by the amount of moles (n).
• Hess's Law states that the enthalpy change for a chemical reaction is independent of the route taken and depends only on the initial and final states.
• Hess's law can be used to calculate enthalpy changes for chemical reactions from the enthalpy change of other reactions.
• Hess's law diagram:

  A) Elements

  B) Products

  C) Reactants

  D) ∆H

  E) ∆fH

  F) ∆fH

  
• Hess's Law equation:

  $\Delta_cH=$$\left(\Sigma\Delta_fH\left(products\right)-\Sigma\Delta_fH\left(reactaNts\right)\right)$
• The means bond enthalpy is a measure of the energy required to break one mole of covalent bond measured in the gaseous state in $kJ\ mol^{-1}$ averaged across many compounds containing the bond.
• General points for bond enthalpies:

  Mean bond enthalpies relate to the strength of a covalent bond. A higher bond enthalpy value means a stronger covalent bond.

  • The shorter the covalent bond, the stronger the bond.
  • Triple covalent bonds are generally shorter than double covalent bonds which are shorter than single covalent bonds.
  • So triple covalent bonds are stronger than double covalent bonds that are stronger than single covalent bonds.
• Use mean bond enthalpies to calculate an approximate value of ∆H for reactions in the gaseous phase:

  ∆H = sum of mean bond enthalpies of bond broken - sum of mean bond enthalpies of bonds made.
• Values from mean bond enthalpy calculations differ from those determined using Hess's law because mean bond enthalpies are average values across different molecules, while Hess's law uses specific enthalpy data for a particular reaction, leading to more precise calculations.