Edexcel Chemistry - Topic 8: Energetics I

• Enthalpy change is the amount of heat energy taken in or given out during any change in a system provided the pressure is constant
• If enthalpy change occurs, energy is transferred between the system (chemicals) and the surroundings (everything outside the chemicals)
• In an exothermic change, energy is transferred from the system to the surroundings; the products have less energy than the reactants (∆H is negative)
  e.g. combustion of fuels
• In an endothermic change, energy is transferred from the surroundings to the system (an input of heat energy is required); the products have more energy than the reactants (∆H is positive)
  e.g. Thermal decomposition of Calcium Carbonate
  

• Standard conditions are:
  Pressure - 100 kPa
  Temperature - 25C or 298K
  Solutions at 1mol dm-3
  All substances at their normal state at 298K
• Standard Enthalpy Change of Formation (∆Hf) - The enthalpy change when 1 mole of the compound is formed from its elements under standard conditions.
• Standard Enthalpy Change of Combustion (∆Hc) - The enthalpy change when 1 mole of a substance is combusted completely in oxygen under standard conditions.
• Standard Enthalpy Change of Neutralisation (∆Hn) - The enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce 1 mole of water
• Standard Enthalpy Change of Atomisation (∆Ha) - The enthalpy change when 1 mole of gaseous atoms is formed from its element in its defined physical state under standard conditions
• Standard Enthalpy Change of Reaction (∆Hr) - The enthalpy change when equation quantities of materials react under standard conditions

• Hess's law states that the total enthalpy change for a reaction is independent of the route by which the chemical change takes place.
• Hess's law is a version of the first law of thermodynamics, which is that energy is always conserved
• On an energy level diagram the direction of the arrows can show the different routes a reaction can proceed by
• Hess's Cycles can be used to measure the enthalpy change for a reaction that cannot be measured directly by experiments
  

• ΔH can be measured by following the paths of two different reactions
• If the reaction for the hydrogenation of an alkene needs to be measured then the reaction of the reactants with oxygen and the reaction of the products with oxygen can be used to calculate the enthalpy change of reaction by the following calculation below.

• Calorimetry is the process of measuring the amount of heat released or absorbed during a chemical reaction
• To calculate energy change for a reaction in a solution we can use the equation:
  Energy Change = Mass of Solution x Heat Capacity x Temperature Change
  Q (J) = m (g) x C (J g-1K-1) x ∆T (K)
• The enthalpy change can be calculated by dividing the Energy Change by the number of moles reacting.
• Remember to always include a sign and unit (Jmol-1/kJmol-1)
• Calorimetry can measure the amount of energy released by dissolving a solid to form a solution or by reacting two solutions together

The errors with Calorimetry include:

• Heat transfer from surroundings (usually loss)
• Approximations in Specific Heat Capacities of solutions (methods assume all solutions have the specific heat capacity of water - 4.18 Jg-1K-1)
• Neglecting the Specific Heat Capacity of the calorimeter/energy absorbed by the apparatus
• Incomplete or slow reactions

• Wash the equipment with the solutions to be used
• Dry the polystyrene cup after washing
• Put the polystyrene cup in a beaker for insulation and support
• Measure out the desired volumes of the solutions with volumetric pipettes and transfer to the insulated cup
• Clamp thermometer into place making sure the thermometer bulb is immersed in the solution
• Measure the initial temperature of the solution
• Add the second reagent to the cup and if using a solid reagent weigh the solid in a weighing pot before and after it has been added to the solution
• Stir the mixture to ensure that all of the solution is at the same temperature
• Record the temperature at regular intervals for five to ten minutes
• 

• Calorimetry can be used to calculate enthalpies of combustion by burning a fuel and using the energy released to heat up water in a metal cup
• Enthalpy Change for Combustion is calculated by:
  Calculating energy change used to heat up the water: Q = mc∆T
  Calculating the number of moles of fuel used: moles = mass/Mr
  Dividing the energy change by the number of moles used: ∆H = Q/mol
• Remember to add the correct sign (-ve for exothermic reaction) and units

Errors in this method include

• Heat losses from calorimeter
• Incomplete combustion of fuel
• Incomplete transfer of heat
• Evapouration of fuel after weighing
• Heat capacity of the calorimeter not included
• Measurements not carried out at standard conditions as H2O is a gas not a liquid in the experiment

• The mean bond enthalpy is the enthalpy needed to break the covalent bond into gaseous atoms, averaged over different molecules
• Mean bond enthalpies are used because every single bond in a compound has a slightly different bond energy e.g. In CH4, there are 4 C-H bonds. Breaking each one will require a different amount of energy so an average for the C-H bond is given
• The value for the bond enthalpy for the C-H bond in methane is: 
  1/4 CH4(g) --> C(g) + H(g)
• These values are positive because energy is required to break a bond (endothermic)
• The substances must be in a gaseous state at the start and end of the reaction
• In an exothermic reaction, the sum of the bonds in the reactant molecules will be less than the sum of the bonds in the product molecules
  ∆H = Σ bond enthalpies broken - Σ bond enthalpies made
• ∆H values calculated from the equation above are less accurate than using formation or combustion data because the mean bond energies are not exact

• When comparing the heats of combustion for successive members of a homologus series such as the alcohols, there is a constant rise in the size of energies of combustion as the number of carbon atoms increases
  Ethanol:            1 C-C, 5 C-H, 1 C-O, 1 O-H and 3 O=O bonds are broken
                           4 C=O and 6 O-H bonds are made
                           ∆Hc = -1365 kJmol-1
  Propan-1-ol:     2 C-C, 7 C-H, 1 C-O, 1 O-H and 4.5 O=O bonds are broken
                           6 C=O and 8 O-H bonds are made
                          ∆Hc = -2016 kJmol-1
• There is a constant amount and type of extra bonds being broken and made so the enthalpy of combustion increases by a constant amount.
• If the results are worked out experimentally with a calorimeter the experimental values will be lower than the calculated ones because there will be significant heat loss and incomplete combustion leading to less energy being released
• Calculated values of enthalpy of combustions will be more accurate if calculated from enthalpy of formation data than from mean bond enthalpies because the mean bond enthalpies are averaged over various compounds