Alevel Chemistry Revision PC 05
Revision Material
Duration: 2nd – 8th March
Topic 5 Chemical Energetics
Enthalpy changes and entropy changes accompany chemical reactions. This topic demonstrates why some reactions and processes are spontaneous and
others are not.
Learning outcomes
Candidates should be able to:
5.1 Enthalpy change, ∆𝐻 (a) explain that chemical reactions are accompanied by energy changes, principally in the form
of heat energy; the energy changes can be exothermic (ΔH is negative) or endothermic (ΔH
is positive)
(b) explain and use the terms:
(i) enthalpy change of reaction and standard conditions, with particular reference to:
formation, combustion, hydration, solution, neutralisation, atomisation
(ii) bond energy (ΔH positive, i.e. bond breaking)
(c) calculate enthalpy changes from appropriate experimental results,
including the use of the relationship ΔH = –mcΔT
5.2 Hess’ Law (a) apply Hess’ Law to construct simple energy cycles, and carry out calculations involving such
cycles and relevant energy terms, with particular reference to:
(i) determining enthalpy changes that cannot be found by direct experiment, e.g. an
enthalpy change of formation from enthalpy changes of combustion
(ii) average bond energies
(b) construct and interpret a reaction pathway diagram, in terms of the enthalpy change of the
reaction and of the activation energy
, 5.1 Enthalpy change, ∆𝑯
(i) chemical reactions
Exothermic reactions Endothermic reactions
- energy given out - energy given in
- surrounding warmer - surrounding cooler
- bond making - bond breaking
- ∆𝐻 negative - ∆𝐻 positive
Ereactants > Eproducts Ereactants < Eproducts
e.g. combustion, e.g. photosynthesis,
respiration, neutralization, electrolysis, Thermal
crystallization, fermentation decomposition
(ii) Enthalpy change
Standard conditions: 298K, 1 atm
The standard enthalpy change of formation: The enthalpy change when one mole of a compound is formed from its
element. Both of compound and the element are at their standard. +/-
The standard enthalpy change of combustion: The enthalpy change when one mole of flammable substance
combusts completely in O2 under standard conditions. Both the reactants and products are at their standard states. -
∆H exothermic
The standard enthalpy change of atomization: enthalpy / energy / heat change when one mole of gaseous atoms is
produced [1] from the element in its standard state [1] under standard conditions
The standard enthalpy change of reaction: enthalpy change when the amount/moles of reactants as shown in a
(reaction) equation react together to give products [1] measured at standard conditions [1]
The standard enthalpy change of neutralization: the enthalpy change when 1 mol of H+ is completely neutralized by
OH- to form 1 mol of H2O under standard conditions
Bond energy: energy needed to break a mole of covalent bonds in the gaseous state
The standard enthalpy change of solution: 1 mole of a solute is dissolved in a solvent to form an infinitely dilute
solution under standard conditions in their standard states
The standard enthalpy change of hydration 1 mole of ions in the gas phase are dissolved in water
(fomration) (combustion)
Example
One of these is DME or dimethyl ether, CH3OCH3.
Define, with the aid of an equation which includes state symbols, the standard enthalpy change of combustion, , for
DME at 298 K.
Solution
CH3OCH3(l) + 3O2(g) → 2CO2(g) + 3H2O(l)
the enthalpy change/heat change/heat evolved when one mole of CH3OCH3/a compound is completely burned or
burned in an excess of air/oxygen
(iii) Calculate enthalpy changes from appropriate experimental results
Including the use of the relationship ΔH = –mcΔT
Revision Material
Duration: 2nd – 8th March
Topic 5 Chemical Energetics
Enthalpy changes and entropy changes accompany chemical reactions. This topic demonstrates why some reactions and processes are spontaneous and
others are not.
Learning outcomes
Candidates should be able to:
5.1 Enthalpy change, ∆𝐻 (a) explain that chemical reactions are accompanied by energy changes, principally in the form
of heat energy; the energy changes can be exothermic (ΔH is negative) or endothermic (ΔH
is positive)
(b) explain and use the terms:
(i) enthalpy change of reaction and standard conditions, with particular reference to:
formation, combustion, hydration, solution, neutralisation, atomisation
(ii) bond energy (ΔH positive, i.e. bond breaking)
(c) calculate enthalpy changes from appropriate experimental results,
including the use of the relationship ΔH = –mcΔT
5.2 Hess’ Law (a) apply Hess’ Law to construct simple energy cycles, and carry out calculations involving such
cycles and relevant energy terms, with particular reference to:
(i) determining enthalpy changes that cannot be found by direct experiment, e.g. an
enthalpy change of formation from enthalpy changes of combustion
(ii) average bond energies
(b) construct and interpret a reaction pathway diagram, in terms of the enthalpy change of the
reaction and of the activation energy
, 5.1 Enthalpy change, ∆𝑯
(i) chemical reactions
Exothermic reactions Endothermic reactions
- energy given out - energy given in
- surrounding warmer - surrounding cooler
- bond making - bond breaking
- ∆𝐻 negative - ∆𝐻 positive
Ereactants > Eproducts Ereactants < Eproducts
e.g. combustion, e.g. photosynthesis,
respiration, neutralization, electrolysis, Thermal
crystallization, fermentation decomposition
(ii) Enthalpy change
Standard conditions: 298K, 1 atm
The standard enthalpy change of formation: The enthalpy change when one mole of a compound is formed from its
element. Both of compound and the element are at their standard. +/-
The standard enthalpy change of combustion: The enthalpy change when one mole of flammable substance
combusts completely in O2 under standard conditions. Both the reactants and products are at their standard states. -
∆H exothermic
The standard enthalpy change of atomization: enthalpy / energy / heat change when one mole of gaseous atoms is
produced [1] from the element in its standard state [1] under standard conditions
The standard enthalpy change of reaction: enthalpy change when the amount/moles of reactants as shown in a
(reaction) equation react together to give products [1] measured at standard conditions [1]
The standard enthalpy change of neutralization: the enthalpy change when 1 mol of H+ is completely neutralized by
OH- to form 1 mol of H2O under standard conditions
Bond energy: energy needed to break a mole of covalent bonds in the gaseous state
The standard enthalpy change of solution: 1 mole of a solute is dissolved in a solvent to form an infinitely dilute
solution under standard conditions in their standard states
The standard enthalpy change of hydration 1 mole of ions in the gas phase are dissolved in water
(fomration) (combustion)
Example
One of these is DME or dimethyl ether, CH3OCH3.
Define, with the aid of an equation which includes state symbols, the standard enthalpy change of combustion, , for
DME at 298 K.
Solution
CH3OCH3(l) + 3O2(g) → 2CO2(g) + 3H2O(l)
the enthalpy change/heat change/heat evolved when one mole of CH3OCH3/a compound is completely burned or
burned in an excess of air/oxygen
(iii) Calculate enthalpy changes from appropriate experimental results
Including the use of the relationship ΔH = –mcΔT
