Lattice Enthalpy
The enthalpy change when ONE MOLE of a solid ionic lattice is separated into its GASEOUS IONS under standard conditions.
- There is a huge release in energy when two
ions of opposite charge combine to form a solid. (Bonds are being made)
- Lattice enthalpy indicates the strength of
an ionic bond.
- It is a highly EXOthermic reaction as it
has strong electrostatic attraction between two oppositely charged ions.
- The larger the negative lattice enthalpy
shows, the larger the electrostatic forces of attraction there are between the oppositely charged ions in the lattice.
Example:
Na⁺(g) + Cl⁻(g) → NaCl(s)
- Lattice enthalpy cannot be measured
directly because it is impossible to form one mole of an ionic lattice from
gaseous ions.
- As you can’t calculate lattice enthalpy
directly, you use Hess’ Law. This is called a Born-Haber Cycle.
Factors that affect
Lattice Enthalpy
①
Ionic Size
· Lattice Enthalpy becomes less exothermic and less
negative as the size of the negative ions increases.
· The charge density decreases
· The attraction gets weaker between the ions, hence the
weaker ionic bonding.
② Ionic Charge
Also
the charge density increases.
Enthalpy change of
solution
The
standard enthalpy change of solution is the enthalpy change that takes place
when ONE MOLE of a COMPOUND is completely DISSOLVED in WATER under standard
conditions.
Example - This process can be endothermic or exothermic:
KCl(s) + aq → K+(aq)
+ Cl-(aq)
When
an ionic solid dissolves in water, two processes take place:
1. Breakdown of the ionic lattice into gaseous ions.
- Overcoming the attractive forces between the oppositely charged
ions requires energy (it is an endothermic reaction – opposite to lattice
enthalpy)
KCl(s) + aq
→ K+(g)
+ Cl-(g)
2. Hydration
Definition: The standard enthalpy change of hydration is the enthalpy change when ONE MOLE of ISOLATED GASEOUS IONS is dissolved in water forming ONE MOLE of AQUEOUS IONS under standard conditions.
Example: K+(g) + aq → K+(aq)
Cl-(g) + aq → Cl-(aq)
- Hydration involves the GASEOUS IONS bonding with H2O molecules.
- The +vely charged ions will be attracted to the slightly negative oxygen atoms in water.
- The –vely charged ions will be attracted to the slightly positive hydrogen atoms in H2O
- It is an EXOthermic process as it is MAKING BONDS.
Q: Why is the enthalpy change of hydration of magnesium ions
more exothermic than the enthalpy change of hydration of calcium ions. Explain
why. (2)
A: The
Mg2+ ion is smaller than the Ca2+ ion. This means Mg2+
has a stronger attraction to H2O.
Q: Explain the differences in the lattice enthalpies of
magnesium fluoride (-2957), sodium fluoride (-918) and sodium chloride (-780).
(3)
A: * Cl- is larger than F- in ionic size. This makes NaCl less exothermic.
* Mg2+ has a greater charge than Na+, this makes Mg2+ more exothermic. Also Mg2+ is smaller than Na+. This means Mg2+ is more exothermic as there is greater attraction.
* F- has greater attraction for Na+ ions and Mg2+ has greater attraction for F- ions.
Born
Haber Cycles
Enthalpy of formation = Sum of all the other
enthalpy values
Steps:
1. Standard enthalpy change of formation, ∆Hf:
2. The standard enthalpy change of atomisation, ∆Hat
3. First Ionisation energy, ∆HI1:
4. (Possible) Second Ionisation energy, ∆HI2:
5. First Electron affinity, ∆HEA1:
6. Second Electron affinity, ∆HEA2:
Example of a born Haber Cycle for Sodium Oxide:
Entropy
Entropy (s) is a measure of the degree of
disorder/randomness in a chemical system.
The value for s is ALWAYS POSITIVE
- All substances possess some degree
of disorder as particles are always in constant motion
- There is a tendency for entropy to
increase, i.e. for things to become more disordered.
- At 0K, perfect crystals
have zero entropy.
Entropy,
S (unlike enthalpy, H) of a substance can be measured directly. It is measured
in J K-1 mol-1
DS = SSqproducts
-
SSqreactants
∆S can
be +ve or –Ve, whereas S is always +ve.
Free Energy
v Spontaneous reactions
occur on their own and move to products with a lower energy and more stability.
Exothermic Reactions
- Many occur
spontaneously at room temperature
- The enthalpy content of
the chemical system decreases during the reaction as the energy is released to
the surroundings.
- This increases
stability.
Endothermic Reactions
- Can also take place spontaneously at room temperature
- The enthalpy content of the chemical system increases during the reaction as energy is being taken in from surroundings.
- This must also increase stability.
- But this is dependent on entropy.
Gibbs
Free Energy
Gibbs
free energy (∆G) is a thermodynamic function that relates enthalpy, entropy and
temperature to spontaneity.
- Spontaneous reactions (feasible reactions at a given temperature) occur if the value for ∆G is NEGATIVE!
