Enthalpy changes

ΔH

The heat energy transferred during a reaction at constant pressure

Exothermic reaction

• ΔH is negative

• Energy is transferred from the system to the surroundings

Endothermic reaction

• ΔH is positive

• Energy is transferred from the surroundings to the system

How can you tell an energy profile represents an exothermic reaction?

The products are at a lower energy than the reactants

Examples of exothermic reactions

• Combustion

• Respiration

• Neutralisation - ΔH= -57 kJmol^-

How can you tell an energy profile represents an endothermic reaction?

The products are at a higher energy than the reactants

Examples of endothermic reactions

• Thermal decomposition

• Photosynthesis

• Melting of ice

Activation energy

The minimum energy required for a reaction to take place

Standard enthalpy changes

• Temperature - 298K

• Pressure - 100kPa

• All solutions at 1 moldm^-

Standard enthalpy change of a reaction

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

• Either endothermic or exothermic reaction

Standard enthalpy change of combustion

The enthalpy change when 1 mole of a substance completely combusts at 298k and 100kPa

• Exothermic reaction

Standard enthalpy change of formation

The enthalpy change when 1 mole of a compound is formed from its elements at 298k and 100kPa

• Either endothermic or exothermic reaction

Standard enthalpy of neutralisation

The enthalpy change when 1 mol of H2O is formed from reaction of H+ and OH- at 298k and 100kPa

• H⁺ + OH^- → H₂0

• Always exothermic

Calorimetry

q=mcΔt

• q = energy change - J

• m = mass of solution - g

• c = specific heat capacity - Jg^-1K^-1

• Δt = Change in temperature of solution ^.C or K

Calorimetry 2

ΔH = +-q/n

• ΔH = enthalpy change - kJmol^-1

• q = energy change - kJ

• n = moles of reactant

  If there are 2 reactants we would use the limiting reactant

Combustion Calorimetry

• The water acts as the surroundings

• The fuel burning acts as the chemical system

How come the calculated value from an experiment is different to the data book in combustion calorimetry?

• Heat loss to surroundings

• Incomplete combustion will 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 beaker

• Cover the wick

Calorimetry with a reaction in solution

• The solution acts as the surroundings

• The chemical reacting or dissolving acts as the chemical system

Sources of error in calorimetry reaction in solution

• Heat loss/gain to/from surroundings

• Water evaporates from the beaker

• Incomplete combustion reaction

• C of solution is not the same as water

How to minimise errors in calorimetry reaction in solution

• Use a polystyrene cup

• Use a lid to prevent heat loss/gain

• Ensure standard conditions

Predicting temperature changes 1

ΔH is the same, If n is doubled q is doubled, Volume is doubled, So Δt is doubled

Predicting temperature changes 2

ΔH is 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 1 mole of bonds in gaseous molecules

How does the bond enthalpy indicate the strength of a covalent bond?

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

Why is a reaction exothermic?

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

Why is a reaction endothermic?

The energy required to make the bond is more the energy released when making bonds

Mexican Ben

ΔrH

ΣΔBeH (bonds broken) - ΣΔBeH (bonds formed)

Hess’ Law: Enthalpy of formation cycle

Hess’ Law: Enthalpy of combustion cycle