Enthalpy change
ΔlatticeH - lattice enthalpy of formation
Enthalpy change when 1 mole of a solid ionic compound is formed from its gaseous ions under
standard conditions (exothermic so negative sign)
For example - Ca2+(g) + 2Cl-(g) - CaCl2(s)
Ionic compounds form giant ionic lattices. When gaseous ions combine to make a giant ionic lattice,
energy is released (lattice energy).
The standard lattice enthalpy is a measure of ionic bond strength. The more negative the lattice
energy, the stronger the bond.
Smaller the size of the ion, the more negative the lattice energy is as they have a higher charge
density.
Larger ionic charge results in a more negative lattice energy since the ionic bond (electrostatic
attraction) would be higher.
ΔatH - enthalpy change of atomisation
Enthalpy change when 1 mole of gaseous atoms is made from an element in its standard state
For example - 0.5F2(g) - F(g)
Standard enthalpy change of atomisation of an element is the enthalpy change measured at 298K and
100kPa when one mole of gaseous atoms is formed from an element in its standard state.
Δea1H - 1st electron af nity
Enthalpy change when 1 mole of gaseous 1- ions are made from 1 mole of gaseous atoms
O(g) - O-(g) (Adding electrons)
Δea2H - 2nd electron af nity
Enthalpy change when 1 mole of gaseous 2- ions are made from 1 mole of gaseous 1- ions
O-(g) - O2-(g)
Ionic bonding
Definition of ionic bonding
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, Electrostatic attraction between oppositely charged ions
What affects strength of ionic bonding?
Size of charge on the ion - bigger the charge on an ion the stronger the electrostatic attraction
between the ions. More energy is required to overcome these forces so they have high melting and
boiling points
Ionic radii - smaller the ion, the stronger the electrostatic attraction between ions. Smaller ions can
pack together more closely and more energy is required to overcome the stronger forces. The
melting and boiling points are higher.
Smaller the ion and the higher the charge the stronger the electrostatic attraction and the higher
the melting point (high charge density)
Theoretical and experimental
la ice enthalpies
Theoretical and experimental values of lattice enthalpies can be different depending on how “purely
ionic” the compound is.
Theoretical lattice enthalpies can be calculated from data assuming a perfectly ionic model
Perfect ionic model
Ions that are perfectly spherical (no distortion)
Charge is evenly distributed in this sphere (point charges)
Ions are in contact with each other
We can carry out an experiment to work out the lattice enthalpy but could nd this is different to the
theoretical lattice enthalpy. This tells us that the compound being experimented on doesn’t follow the
perfectly ionic model and has some covalent characteristics.
The positive ion distorts the charge distribution in the negative ion. The positive ion polarises the
negative ion. The more polarisation/distortion we get, the more covalent character there will be.
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ΔlatticeH - lattice enthalpy of formation
Enthalpy change when 1 mole of a solid ionic compound is formed from its gaseous ions under
standard conditions (exothermic so negative sign)
For example - Ca2+(g) + 2Cl-(g) - CaCl2(s)
Ionic compounds form giant ionic lattices. When gaseous ions combine to make a giant ionic lattice,
energy is released (lattice energy).
The standard lattice enthalpy is a measure of ionic bond strength. The more negative the lattice
energy, the stronger the bond.
Smaller the size of the ion, the more negative the lattice energy is as they have a higher charge
density.
Larger ionic charge results in a more negative lattice energy since the ionic bond (electrostatic
attraction) would be higher.
ΔatH - enthalpy change of atomisation
Enthalpy change when 1 mole of gaseous atoms is made from an element in its standard state
For example - 0.5F2(g) - F(g)
Standard enthalpy change of atomisation of an element is the enthalpy change measured at 298K and
100kPa when one mole of gaseous atoms is formed from an element in its standard state.
Δea1H - 1st electron af nity
Enthalpy change when 1 mole of gaseous 1- ions are made from 1 mole of gaseous atoms
O(g) - O-(g) (Adding electrons)
Δea2H - 2nd electron af nity
Enthalpy change when 1 mole of gaseous 2- ions are made from 1 mole of gaseous 1- ions
O-(g) - O2-(g)
Ionic bonding
Definition of ionic bonding
fifi
, Electrostatic attraction between oppositely charged ions
What affects strength of ionic bonding?
Size of charge on the ion - bigger the charge on an ion the stronger the electrostatic attraction
between the ions. More energy is required to overcome these forces so they have high melting and
boiling points
Ionic radii - smaller the ion, the stronger the electrostatic attraction between ions. Smaller ions can
pack together more closely and more energy is required to overcome the stronger forces. The
melting and boiling points are higher.
Smaller the ion and the higher the charge the stronger the electrostatic attraction and the higher
the melting point (high charge density)
Theoretical and experimental
la ice enthalpies
Theoretical and experimental values of lattice enthalpies can be different depending on how “purely
ionic” the compound is.
Theoretical lattice enthalpies can be calculated from data assuming a perfectly ionic model
Perfect ionic model
Ions that are perfectly spherical (no distortion)
Charge is evenly distributed in this sphere (point charges)
Ions are in contact with each other
We can carry out an experiment to work out the lattice enthalpy but could nd this is different to the
theoretical lattice enthalpy. This tells us that the compound being experimented on doesn’t follow the
perfectly ionic model and has some covalent characteristics.
The positive ion distorts the charge distribution in the negative ion. The positive ion polarises the
negative ion. The more polarisation/distortion we get, the more covalent character there will be.
tt fi
