3.5 Unit 5 CHEM5 Energetics, Redox and
Inorganic Chemistry
3.5.1 – Thermodynamics (4 pages)
Enthalpy change (ΔH)
Be able to define and apply the terms enthalpy of formation, ionisation enthalpy,
enthalpy of atomisation of an element and of a compound, bond dissociation enthalpy,
electron affinity, lattice enthalpy (defined as either lattice dissociation or lattice
formation), enthalpy of hydration and enthalpy of solution
Enthalpy of formation:
The enthalpy change when one mole of a compound is formed from its
constituent elements under standard conditions, all reactants and
products in their standard states.
Ionisation enthalpy:
First: The standard enthalpy change when one mole of gaseous atoms is
converted into a mole of gaseous ions each with a single positive charge.
Second: The standard enthalpy change when one mole of gaseous
unipositive ions is converted into a mole of gaseous ions each with a
double positive charge.
Enthalpy of atomisation of an element:
Enthalpy change when one mole of gaseous atoms is formed from the
element in its standard state.
Enthalpy of atomisation of a compound:
Enthalpy change when one mole of a compound in its standard state is
converted into its free gaseous atoms.
Bond dissociation enthalpy:
Standard enthalpy change when one mole of gaseous molecules each
breaks a covalent bond to form two free radicals (averaged over a range
of compounds).
Electron affinity:
First: The standard enthalpy change when a mole of gaseous atoms is
converted to a mole of gaseous ions, each with a single negative charge.
Second: The standard enthalpy change when a mole of electrons is added
to a mole of gaseous ions each with a single negative charge to form ions
each with a double negative charge.
Lattice dissociation enthalpy:
The standard enthalpy change when one mole of solid ionic compound is
separated into its free gaseous ions.
Lattice formation enthalpy:
The standard enthalpy change when one mole of solid ionic compound is
formed from its free gaseous ions.
Enthalpy of hydration:
The standard enthalpy change when water molecules surround one mole
of gaseous ions.
Enthalpy of solution:
,The standard enthalpy change when one mole of an ionic solid dissolves
completely in sufficient water to form a solution in which the ions are far
enough apart not to interact with each other.
Be able to construct Born-Haber cycles to calculate lattice enthalpies from experimental
data. Be able to compare lattice enthalpies from Born-Haber cycles with those from
calculations based on a perfect ionic model to provide evidence for covalent character in
ionic compounds
A Born-Haber cycle is a cycle that includes all the enthalpy changes in the
formation of an ionic compound.
The lattice enthalpy of a compound is an indication of the strength of the
ionic bonding. The greater the magnitude of the lattice enthalpy, the
stronger the bonding. Compounds with smaller ions have stronger
attractions and so greater lattice enthalpy.
Lattice enthalpies can be calculated either by:
1. A Born-Haber cycle (which uses experimentally measured values).
2. Theoretical calculation based on the charge and size of positive and
negative ions.
The closer the experimentally measured value, using a Born-Haber cycle,
to the theoretical value, the purer the ionic bonding. If the difference is
large, then the compound has some covalent character. The more
polarising the positive ion, the more the negative ion is distorted in shape
and the more covalent character.
Theoretical values assume Perfect Ionic Bonding: a mathematical
calculation of the lattice formation enthalpy of a compound, which
assumes that the positive and negative ions are perfectly spherical and
that there is no covalent character in the compound (ie. only ionic bonding
present – ions are unpolarised).
Experimental values assume there is partial covalent bonding. The
electrons in the negative ions are pulled towards the positive ions.
The smaller the positive ion, and the higher its charge, the more polarising
it is. Larger negative ions are polarised more easily.
Be able to calculate enthalpies of solution for ionic compounds from lattice enthalpies
and enthalpies of hydration
Lattice dissociation enthalpy + enthalpy of hydration = enthalpy of
solution
, When a solid ionic lattice dissolves in water:
1. The bonds between the ions break - DISSOCIATION
2. Bonds between the ions and water are made - HYDRATION
Be able to use mean bond enthalpies to calculate an approximate value of ΔH for other
reactions
Bond breaking is endothermic
Bond making is exothermic
Bond enthalpy:
The energy when one mole of bonds are broken with all species in
gaseous state.
AB (g) A (g) + B (g)
Mean bond enthalpy:
The average of several values of the bond dissociation enthalpy for a
given type of bond, taken from a range of compounds.
In cycle for bond enthalpies, arrows go DOWN.
Or calculation:
∑b(reactants) - ∑b(products)
Be able to explain why values from mean bond enthalpy calculations differ from those
determined from enthalpy cycles
Values from mean bond enthalpy calculations differ from those
determined from enthalpy cycles because experimental data is specific to
the compound you are testing. Whereas, a mean bond enthalpy is an
average taken from a range of different compounds.
Free-energy change (ΔG) and entropy change (ΔS)
Understand that ΔH, whilst important, is not sufficient to explain spontaneous change
(eg. spontaneous endothermic reactions)
A spontaneous reactions is one which takes place of its own accord.
Many spontaneous reactions are exothermic. You normally have to supply
energy to make an endothermic reaction happen, but you can get
spontaneous endothermic reactions. Negative ΔH (exo) is a factor in
whether a reaction is spontaneous, but it does not explain why a number
of endothermic reactions are spontaneous.
This will happen when the entropy increases so much that the reaction will
happen by itself.
A reaction will not happen unless the total entropy change is positive.
Understand that the concept of increasing disorder (entropy change ΔS) accounts for the
above deficiency, illustrated by physical change (eg. melting, evaporation) and chemical
change (eg. dissolution, evolution of CO2 from hydrogencarbonates with acid)
Inorganic Chemistry
3.5.1 – Thermodynamics (4 pages)
Enthalpy change (ΔH)
Be able to define and apply the terms enthalpy of formation, ionisation enthalpy,
enthalpy of atomisation of an element and of a compound, bond dissociation enthalpy,
electron affinity, lattice enthalpy (defined as either lattice dissociation or lattice
formation), enthalpy of hydration and enthalpy of solution
Enthalpy of formation:
The enthalpy change when one mole of a compound is formed from its
constituent elements under standard conditions, all reactants and
products in their standard states.
Ionisation enthalpy:
First: The standard enthalpy change when one mole of gaseous atoms is
converted into a mole of gaseous ions each with a single positive charge.
Second: The standard enthalpy change when one mole of gaseous
unipositive ions is converted into a mole of gaseous ions each with a
double positive charge.
Enthalpy of atomisation of an element:
Enthalpy change when one mole of gaseous atoms is formed from the
element in its standard state.
Enthalpy of atomisation of a compound:
Enthalpy change when one mole of a compound in its standard state is
converted into its free gaseous atoms.
Bond dissociation enthalpy:
Standard enthalpy change when one mole of gaseous molecules each
breaks a covalent bond to form two free radicals (averaged over a range
of compounds).
Electron affinity:
First: The standard enthalpy change when a mole of gaseous atoms is
converted to a mole of gaseous ions, each with a single negative charge.
Second: The standard enthalpy change when a mole of electrons is added
to a mole of gaseous ions each with a single negative charge to form ions
each with a double negative charge.
Lattice dissociation enthalpy:
The standard enthalpy change when one mole of solid ionic compound is
separated into its free gaseous ions.
Lattice formation enthalpy:
The standard enthalpy change when one mole of solid ionic compound is
formed from its free gaseous ions.
Enthalpy of hydration:
The standard enthalpy change when water molecules surround one mole
of gaseous ions.
Enthalpy of solution:
,The standard enthalpy change when one mole of an ionic solid dissolves
completely in sufficient water to form a solution in which the ions are far
enough apart not to interact with each other.
Be able to construct Born-Haber cycles to calculate lattice enthalpies from experimental
data. Be able to compare lattice enthalpies from Born-Haber cycles with those from
calculations based on a perfect ionic model to provide evidence for covalent character in
ionic compounds
A Born-Haber cycle is a cycle that includes all the enthalpy changes in the
formation of an ionic compound.
The lattice enthalpy of a compound is an indication of the strength of the
ionic bonding. The greater the magnitude of the lattice enthalpy, the
stronger the bonding. Compounds with smaller ions have stronger
attractions and so greater lattice enthalpy.
Lattice enthalpies can be calculated either by:
1. A Born-Haber cycle (which uses experimentally measured values).
2. Theoretical calculation based on the charge and size of positive and
negative ions.
The closer the experimentally measured value, using a Born-Haber cycle,
to the theoretical value, the purer the ionic bonding. If the difference is
large, then the compound has some covalent character. The more
polarising the positive ion, the more the negative ion is distorted in shape
and the more covalent character.
Theoretical values assume Perfect Ionic Bonding: a mathematical
calculation of the lattice formation enthalpy of a compound, which
assumes that the positive and negative ions are perfectly spherical and
that there is no covalent character in the compound (ie. only ionic bonding
present – ions are unpolarised).
Experimental values assume there is partial covalent bonding. The
electrons in the negative ions are pulled towards the positive ions.
The smaller the positive ion, and the higher its charge, the more polarising
it is. Larger negative ions are polarised more easily.
Be able to calculate enthalpies of solution for ionic compounds from lattice enthalpies
and enthalpies of hydration
Lattice dissociation enthalpy + enthalpy of hydration = enthalpy of
solution
, When a solid ionic lattice dissolves in water:
1. The bonds between the ions break - DISSOCIATION
2. Bonds between the ions and water are made - HYDRATION
Be able to use mean bond enthalpies to calculate an approximate value of ΔH for other
reactions
Bond breaking is endothermic
Bond making is exothermic
Bond enthalpy:
The energy when one mole of bonds are broken with all species in
gaseous state.
AB (g) A (g) + B (g)
Mean bond enthalpy:
The average of several values of the bond dissociation enthalpy for a
given type of bond, taken from a range of compounds.
In cycle for bond enthalpies, arrows go DOWN.
Or calculation:
∑b(reactants) - ∑b(products)
Be able to explain why values from mean bond enthalpy calculations differ from those
determined from enthalpy cycles
Values from mean bond enthalpy calculations differ from those
determined from enthalpy cycles because experimental data is specific to
the compound you are testing. Whereas, a mean bond enthalpy is an
average taken from a range of different compounds.
Free-energy change (ΔG) and entropy change (ΔS)
Understand that ΔH, whilst important, is not sufficient to explain spontaneous change
(eg. spontaneous endothermic reactions)
A spontaneous reactions is one which takes place of its own accord.
Many spontaneous reactions are exothermic. You normally have to supply
energy to make an endothermic reaction happen, but you can get
spontaneous endothermic reactions. Negative ΔH (exo) is a factor in
whether a reaction is spontaneous, but it does not explain why a number
of endothermic reactions are spontaneous.
This will happen when the entropy increases so much that the reaction will
happen by itself.
A reaction will not happen unless the total entropy change is positive.
Understand that the concept of increasing disorder (entropy change ΔS) accounts for the
above deficiency, illustrated by physical change (eg. melting, evaporation) and chemical
change (eg. dissolution, evolution of CO2 from hydrogencarbonates with acid)
