Enthalpy changes

Enthalpy change

The heat energy transferred during a reaction at constant pressure

Exothermic energy profile

How can you tell an energy profile is exothermic?

The products are at a lower energy than the reactants

Examples of Exothermic Reactions

• Combustion

• Respiration

• Neutralisation always -57kJ

Endothermic Energy Profile

How can you tell an energy profile is endothermic ?

The products are at a higher energy than the reactants

Examples of Endothermic reactions

• Thermal decomposition

• Photosynthesis

Activation energy

The minimum energy required for a reaction to take place

Standard enthalpy change of a reaction

• The enthalpy change for a given reaction in the molar quantities shown in the equation

• Can be either exothermic or endothermic

Standard enthalpy change of combustion

• The enthalpy change when one mole of a substance completely combusts at 298K and 100KPa

• Exothermic reaction

Standard enthalpy change of formation

• The enthalpy change when 1 mol of a compound is formed from its elements at 298K and 100KPa

• Can be either exothermic or endothermic

Standard enthalpy change of neutralisation

• The enthalpy change when 1 mol of H₂0 is formed from a reaction of H⁺ and OH^- at 298K and 100KPa

• H⁺ + OH^- → H₂0

• Always exothermic

Calorimetry

• q=mcΔt

  • q= Energy Change

  • m= Mass of solution

  • ΔT= Change in temperature

• ΔH = ± q/n

  • ΔH= Enthalpy Change

  • q= Energy change in kJ

  • n= Moles of reactant

Combustion Calorimetry: What acts as the surroundings?

The water

Combustion Calorimetry: What acts as the chemical system?

The fuel burning

Why is the value for ΔH values from experiments different from the data book values

• Heat loss to surroundings

• Incomplete combustion can occur

How to reduce sources of error in combustion calorimetry

• Ensure a ventilated room

• Move beaker closer to the flame

• Copper beaker instead of glass beaker

• Add a lid to the cover

• Cover the wick

Calorimetry with a reaction in solution: What acts as the surroundings?

The solution

Calorimetry with a reaction in solution: What acts as the chemical system

The chemicals reacting or dissolving

Sources of error in calorimetry reaction in solution

• Heat loss/gain to/from surroundings

• Water evaporates from the beaker

• Incomplete reaction

• Concentration of solution is not the same as water

How to minimise the errors in calorimetry reaction in solution

• Use a polystyrene cup

• Use a lid to prevent loss/gain

• Ensure standard conditions

Predicting temperature changes with different quantities

• ΔH stays the same

  • If n is doubled, q is doubled

  • Volume is doubled/ Volume is the same

  • So ΔT stays the same/ ΔT is doubled

• ΔH stays the same

  • If n is halved, q is halved

  • Volume is halved

  • ΔT stays the same

Average bond enthalpy

The enthalpy change for the breaking of one mole of bonds in gaseous molecules

How does the bond enthalpy indicate the strength if the bond?

The more positive the bond enthalpy the larger the amount of energy needed to break the bond and so the stronger the bond

Why are bond enthalpies endothermic

• Bond breaking requires energy, it is endothermic

• Bond making releases energy, it is exothermic

Why is a reaction exothermic?

The energy required to break the bonds is less than the energy released when making bonds

Why is a reaction endothermic?

The energy required to break the bonds is more than the energy released when making bonds

Bond enthalpy calculation

Δ_rH= ΣΔ_BeH (bonds broken) - ΣΔ_BeH (bonds formed)

Why is it not always possible to measure enthalpy changes of reactions directly from experiments?

• A very high activation energy

• A very slow rate of reaction

• Occurrence of side reactions

Enthalpy of combustion cycle

Enthalpy of formation cycle