Lattice Enthalpy

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    =         SS^q_products         -           SS^q_reactants

 

∆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!