LATTICE Flashcards
Lattice energy
is the enthalpy change when 1 mole of an ionic compound is formed from
its gaseous ions under standard conditions
The lattice energy is always
exothermic; the more exothermic the lattice energy, the
stronger the ionic bonding in the lattice
The standard enthalpy change of atomisation,
∆HꝊ
at, is the enthalpy change when 1 mole
of gaseous atoms is formed from its element under standard conditions; is endothermic
The first electron affinity, ∆HꝊ
ea1,
, is the enthalpy change when 1 mole of electrons is
added to 1 mole of gaseous atoms to form 1 mole of gaseous 1- ions under standard
conditions; is exothermic
The second electron affinity, ∆HꝊ
ea2,
is the enthalpy change when 1 mole of electrons is
added to 1 mole of gaseous 1– ions to form 1 mole of gaseous 2– ions under standard
conditions; is endothermic – so are the 3rd electron affinities
Mg2+ requires
1st and 2nd ionisation energy to be calculated
Two chloride ions in MgCl2,
hence the values of ∆HꝊ
at & ∆HꝊ
ea1 should be multiplied
by 2
Lattice energy arises from the
electrostatic force of attraction of oppositely charged ions
when the crystalline lattice is formed
As the size of the ion
increases, the lattice energy becomes less exothermic, e.g. the lattice
energy gets less exothermic as the size of the anion increases from F- to I-
why does as the size of the ion increases
the lattice energy becomes less exothermic
Due to the decrease in charge density with the same ionic charge, as the same
charge is spread over a larger volume, resulting in weaker electrostatic forces of
attraction in the ionic lattice, e.g. NaF has a less exothermic lattice energy than LiF
The lattice energy becomes more exothermic (stronger ionic bonds formed) as the
ionic
charge increases (higher charge density), e.g. LiF < MgO:
The positive charge on the cation in an ionic lattice may attract
the electrons in the anion
towards it, resulting to distortion of the electron cloud of the anion, causing it to no longer
be spherical (ion polarisation)
polarising power of the cation
the ability of a cation to attract electrons and distort an
anion
The degree of polarisation on an anion depends on:
The degree of polarisation on an anion depends on:
And more polarised if:
The cation is small
The cation has a charge of 2+ or 3+
The anion is large
The anion has a charge of 2- or 3-
Many ionic bonding have some
covalent character due to ion polarisation
The Group 2 carbonates decompose to
their oxides and carbon dioxide on heating:
The further down the group,
the higher temperature required to decompose the carbonate
their relative stabilities increases down the group:
Ion polarisation of carbonates:
The ionic carbonate ion has large ionic radius, hence easily polarised (given a small
highly charged cation)
Group 2 cations’ ionic radius increase down the group:
The smaller the ionic radius of the cation, the better the polarising power, hence
degree of polarisation of carbonate ion by Group 2 cation follows:
The greater the polarisation, the easier it is to weaken a C – O bond in the
carbonate and form the CO2 and oxide on heating
Thermal decomposition of Group 2 nitrates has a similar pattern, decompose to form
nitrogen dioxide, oxygen and the oxide:
The enthalpy change of solution, ∆HꝊ
sol,
is the energy when 1 mole of an ionic solid
dissolve in sufficient water to form a very dilute solution; can be exothermic and
endothermic
Ion-dipole bonds are formed when
ionic solid dissolves in water:
The enthalpy change of hydration, ∆HꝊ
hyd,
is the enthalpy change when 1 mole of a
gaseous ion dissolves in sufficient water to form a very dilute solution; is exothermic
enthaly change for hydration is more exothermic when
More exothermic for ions with the same charge but smaller ionic radii, e.g. ∆HꝊ
hyd is
more exothermic for Li+ than for Na+
More exothermic for ions with the same radii but a larger charge, e.g. ∆HꝊ
hyd is
more exothermic for Mg2+ than for Li+.
The solubility of Group 2 sulfates decreases as the
radius of the metal ion increases
Change in hydration enthalpy down the group:
Smaller ions (with same charge) have greater enthalpy changes of hydration
So the enthalpy change of hydration decreases (gets less exothermic)
following:
hydration enthalpy Decrease is large down the group, depending entirely on the increase in size
of the cation, as the anion is unchanged
Lattice energy is inversely proportional to the sum of the radii (cation &
anion)
Sulfate ion much larger than group 2 cations, hence it contributes a greater
part to the change in lattice energy down the group
Hence decrease in lattice energy is small down the group, determined more
by the size of the sulfate ions than the size of the cations
Change in lattice energy down the group:
Smaller ions form greater lattice energy
So the lattice energy decreases following:
The lattice energy of the sulfates decreases by
relatively smaller values
The enthalpy change of hydration decreases by
relatively larger values down the
group
∆HꝊ
sol becomes more
endothermic down the group
Solubility of Group 2 sulfates decreases
down the group
The higher the positive value of ∆HꝊ
sol
he less soluble the salt
Entropy
a measure of the ‘disorder’ of a system, and that a system becomes more stable
when its energy is spread out in a more disordered state
Standard molar entropy is
the entropy when one mole of substance in its standard state
The values of all molar entropies are
positive
Gases generally have much
higher entropy values than liquids – which have higher entropy
values than solids; hence the more gas molecules present, the greater the number of ways
of arranging them, hence higher entropy:
Increase in entropy of the system due to the greater number of moles of gas
molecules in the products (5 molecules) than in the reactants (2 molecules); there
are two different product molecules and only one type of reactant molecule,
contributing to a greater disorder, increasing the stability (energetically) of the system
Decrease in entropy, hence the reactants are more stable than the product
Simpler substances with fewer atoms have lower entropy values than more complex
substances (more no. of atoms)
For similar substances, the harder substance has lower entropy value
Gradual increase in entropy as the temperature of the substance is increased
highest to lowest entropy
g>l>s
For an exothermic reaction, energy released to
the surroundings, and causes translation
and rotation of molecules in the surroundings – increasing its arrangements – hence there
is likely to be an increase in entropy and increase in the chance for chemical change to
occur spontaneously
For an endothermic reaction, energy absorbed from
the surroundings, decreasing its
arrangements, hence there is likely to be a decrease in entropy and decrease in the chance
for a chemical reaction to occur spontaneously
When the total enthalpy change is positive,
the reaction will occur spontaneously,
reaction is feasible
When total enthalpy change is negative,
he reaction is not likely to occur (not feasible)
entropy formula
Entropy change of the surroundings is given by:
total enthalpy change formula
During exothermic reactions
the enthalpy change plays a bigger role than the entropy
change of the system
kJmol-1 to JKmol-1
*1000
gibbs free energy formula
An exothermic reaction causes ∆Hreaction to be negative
If the value of ∆Ssystem is positive, the reaction will be spontaneous, ∆G is
negative
If the value of ∆Ssystem is negative in low temperatures, ∆G is negative
If the value of ∆Ssystem is negative in high, ∆G is positive
An endothermic reaction causes ∆Hreaction to be positive
If the value of ∆Ssystem is negative, ∆G is positive
If the value of ∆Ssystem is positive in high temperatures, ∆G is negative
If the value of ∆Ssystem is positive in low temperatures, ∆G is positive
The electron affinity is a measure of the
attraction between the incoming electron and the
nucleus - the stronger the attraction, the more energy is released – hence electron affinity
increases upward for the groups and from left to right across periods of a periodic table
