lattice enthalpy

Enthalpy of formation definition

enthalpy change when 1 mole of a compound is formed from its elements under standard conditions (298 K and 100 kPa)

is enthalpy of formation exo or endo

can be endothermic or exothermic

Ionisation enthalpy definition

energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions

is ionisation energy endo or exo?

always endothermic as energy is need to overcome the attraction between an electron and the nucleus

ionisation energy equation:

Na (g) → Na⁺ (g) + e^–          ΔH_ie^ꝋ = +500 kJ mol^-1

Enthalpy change of atomisation definition

enthalpy change when 1 mole of gaseous atoms is formed from its element under standard conditions

atomisation endo or exo?

• always endothermic as energy is always required to break any bonds between the atoms in the element, to break the element into its gaseous atoms

  • Since this is always an endothermic process, the enthalpy change will always have a positivevalue

equation for atomisation:

Na (s) → Na (g)           ΔH_at^ꝋ = +108 kJ mol ^-1

Bond enthalpy definition

amount of energy required to break one mole of a specific covalent bond in the gas phase is called the bond dissociation energy

Bond enthalpy is usually treated as a bond breaking process, so it is quoted in data tables as an ____ energy change with positive values

endothermic

bond enthalpy equation:

Cl₂ (g) → 2Cl (g)    E(Cl-Cl) = +242 kJ mol ^-1

Lattice energy definition:

1 mole of an ionic compound is formed from its gaseous ions (under standard conditions)

is it endo or exo?

exothermic, as when ions are combined to form an ionic solid lattice there is an extremely large release of energy

• Since this is an exothermic process, the enthalpy change will have a negative value

• Because of the huge release in energy when the gaseous ions combine, the value will be a very large negative value

• The large negative value of ΔH_latt^ꝋ suggests that the ionic compound is much more stable than its gaseous ions

  • This is due to the strong electrostatic forces of attraction between the oppositely charged ions in the solid lattice

  • Since there are no electrostatic forces of attraction between the ions in the gas phase, the gaseous ions are less stable than the ions in the ionic lattice

  • The more exothermic the value is, the stronger the ionic bonds within the lattice are

lattice enthalpy:

Na⁺(g) + Cl^-(g) → NaCl (s)  ΔH_latt^ꝋ = -776 kJ mol ^-1

enthalpy of neutralisation definition + equation:

enthalpy change when 1 mol of water is formed in reaction between an acid and alkali under standard condition

½ H2SO4 + NaOH —> ½ Na2SO4 + H20

electron affinity def + equation:

first electron affinity= enthalpy change when each atom in one mole of geaousous atoms gains on electron to form 1 mole of geaous 1- ions

2nd= each ion in one mol of geasous 1- ions gains one electron to form on mole of geaous 2- ions

O(g) + e- —> O-(g)

electron affinity:

1st= exo

2nd= endo

hydration enthalpy:

enthalpy change when 1 mol of gaesous ions become hydrated

exo

Mg 2+ (g) + aq —> Mg 2+ (aq)

enthalpy of solu.

one mole of ionic solid dissolves in an amount of water large enough so that the dissolved ions are well seperated and do not interact with each other

exo/endo

MgCl2(s) + aq —> Mg 2+(aq) + 2Cl- (aq)

enthalpy of fusion:

enthalpy change when 1 mol of a solid is turned into a liquid

endo

Mg(s) —>Mg(l)

born-haber cycle left side:

• atoms in normal states: Na(s) + ½ Cl2

• The enthalpy of atomisation of sodium is

Na (s) → Na (g)           ΔH_at^ꝋ = +108 kJ mol ^-1

• The enthalpy of atomisation of chlorine is

½Cl₂ (g) → Cl (g)       ΔH_at^ꝋ = +121 kJ mol ^-1

right side:

• The sodium ion loses an electron, so this energy change is the first ionisation energy for sodium

Na (g) → Na⁺ (g) + e^–          ΔH_ie^ꝋ = +500 kJ mol^-1

• The change is endothermic so the direction continues upwards

• The chlorine atom gains an electron, so this is electron affinity

Cl (g) + e^– → Cl^- (g)           ΔH_ea^ꝋ = -364 kJ mol^-1

• The exothermic change means this is downwards

3:

• The enthalpy of formation of sodium chloride is added at the bottom of the diagram

Na(s) + ½Cl₂ (g) → NaCl (s)            ΔH_f^ꝋ = -411 kJ mol ^-1

• This is an exothermic change for sodium chloride so the arrow points downwards

• Enthalpy of formation can be exothermic or endothermic, so you may need to show it above the elements ( and displaced to the right) for a endothermic change

• The final change is lattice enthalpy, which is usually shown a formation. For sodium chloride the equation is

Na⁺(g) + Cl^-(g) → NaCl (s)  ΔH_latt^ꝋ 

how to work out overall change:

standard enthalpy change of solution (ΔH_sol) is the enthalpy change when ____

1 mole of an ionic substance dissolves in sufficient water to form an infinitely dilute solution

standard enthalpy change of hydration (ΔH_hyd) is the enthalpy change when _______

1 mole of a specified gaseous ion dissolves in sufficient water to form an infinitely dilute solution

• When an ionic solid dissolves in water, positive and negative ions are formed

• Water is a polar molecule with a δ- oxygen (O) atom and δ+ hydrogen (H) atoms which will form ion-dipole attractions with the ions present in the solution

• The oxygen atom in water will be attracted to the positive ions and the hydrogen atoms will be attracted to the negative ions

• According to Hess’s law, the enthalpy change for both routes is the same, such that:

ΔHhydꝋ = ΔHlattꝋ + ΔHsolꝋ

What is the affect of increasing ionic radius on lattice energy?

The lattice energy becomes less exothermic as the ionic radius of the ions increases

why?

• because the charge on the ions is more spread out over the ion when the ions are larger

• The ions are also further apart from each other in the lattice

  • The attraction between ions is between the centres of the ions involved, so the bigger the ions the bigger the distance between the centre of the ions

• Therefore, the electrostatic forces of attraction between the oppositely charged ions in the lattice are weaker

• For example, the lattice energy of caesium fluoride (CsF) is less exothermic than the lattice energy of potassium fluoride (KF)

  • Since both compounds contain a fluoride (F^-) ion, the difference in lattice energy must be due to the caesium (Cs⁺) ion in CsF and potassium (K⁺) ion in KF

  • Potassium is a Group 1 and Period 4 element

  • Caesium is a Group 1 and Period 6 element

  • This means that the Cs⁺ ion is larger than the K⁺ ion

  • There are weaker electrostatic forces of attraction between the Cs⁺ and F^- ions compared to K⁺ and F^- ions

  • As a result, the lattice energy of CsF is less exothermic than that of KF

What happens to lattice energy as ionic charge increases?

lattice energy gets more exothermic as the ionic charge of the ions increases

why?

• greater the ionic charge, the higher the charge density

• This results in stronger electrostatic attraction between the oppositely charged ions in the lattice

• As a result, the lattice energy is more exothermic

• For example, the lattice energy of calcium oxide (CaO) is more exothermic than the lattice energy of potassium chloride (KCl)

  • Calcium oxide is an ionic compound which consists of calcium (Ca²⁺) and oxide (O^2-) ions

  • Potassium chloride is formed from potassium (K⁺) and chloride (Cl^-) ions

  • The ions in calcium oxide have a greater ionic charge than the ions in potassium chloride

  • This means that the electrostatic forces of attraction are stronger between the Ca²⁺ and O^2-compared to the forces between K⁺ and Cl^-

  • Therefore, the lattice energy of calcium oxide is more exothermic, as more energy is released upon its formation from its gaseous ions

  • Ca²⁺ and O^2- are also smaller ions than K⁺ and Cl^-, so this also adds to the value for the lattice energy being more exothermic

what 2 key factors affect lattice enthalpy?

charge and radius of the ions that make up the crystalline lattice

what factors affect the enthalpy of hydration?

• amount that the ions are attracted to the water molecules

• The factors which affect this attraction are the ionic charge and radius

If ionic radius decreases what happens to enthalpy of hydration?

ΔH_hyd^ꝋ becomes more exothermic with decreasing ionic radii

why?

• Smaller ions have a greater charge density resulting in stronger ion-dipole attractions between the water molecules and the ions in the solution

• Therefore, more energy is released when they become hydrated and ΔH_hyd^ꝋ becomes more exothermic

e.g.

• For example, the ΔH_hyd^ꝋ of magnesium sulfate (MgSO₄) is more exothermic than the ΔH_hyd^ꝋ of barium sulfate (BaSO₄)

  • Since both compounds contain a sulfate (SO₄^2-) ion, the difference in ΔH_hyd^ꝋ must be due to the magnesium (Mg²⁺) ion in MgSO₄ and barium (Ba²⁺) ion in BaSO₄

  • Magnesium is a Group 2 and Period 3 element

  • Barium is a Group 2 and Period 6 element

  • This means that the Mg²⁺ ion is smaller than the Ba²⁺ ion

  • The attraction is therefore much stronger for the Mg²⁺ ion

  • As a result, the standard enthalpy of hydration of MgSO₄ is more exothermic than that of BaSO₄

Ionionic charge what is the effect on enthalpy of hydration?

ΔH_hyd^ꝋ is more exothermic for ions with larger ionic charges

why?

• Ions with large ionic charges have a greater charge density resulting in stronger ion-dipole attractions between the water molecules and the ions in the solution

• Therefore, more energy is released when they become hydrated and ΔH_hyd^ꝋbecomes more exothermic

e.g.

ΔH_hyd^ꝋ of calcium oxide (CaO) is more exothermic than the ΔH_hyd^ꝋof potassium chloride (KCl)

• Calcium oxide is an ionic compound that consists of calcium (Ca²⁺) and oxide (O^2-) ions

• Potassium chloride is formed from potassium (K⁺) and chloride (Cl^-) ions

• Both of the ions in calcium oxide have a greater ionic charge than the ions in potassium chloride

• This means that the attractions are stronger between the water molecules and Ca²⁺ and O^2-ions upon hydration of CaO

• The attractions are weaker between the water molecules and K⁺ and Cl^- ions upon hydration of KCl

• Therefore, the ΔH_hyd^ꝋ of calcium oxide is more exothermic as more energy is released upon its hydration

The entropy (S) of a given system _____

is the number of possible arrangements of the particles and their energy in a given system

• In other words, it is a measure of how disordered or chaotic a system is

When a system becomes more disordered, its entropy will _____

increase

An increase in entropy means that the system becomes _______

energetically more stable

during the thermal decomposition of calcium carbonate (CaCO₃): CaCO₃(s) → CaO(s) + CO₂(g)

what happens to the entropy + why?

increase

• a gas molecule (CO₂) is formed

• The CO₂ gas molecule is more disordered than the solid reactant (CaCO₃), as it is constantly moving around

• As a result, the system has become more disordered and there is an increase in entropy

the system with the higher entropy will be ______

energetically favourable (as the energy of the system is more spread out when it is in a disordered state)

feasibility takes no account of the _______

rate of reaction and states only what is possible, not what actually happens. A feasible reaction might be incredibly slow, such as the rusting of iron

What is the unit for entropy?

J K^-1 mol^-1

equation to calculate the standard entropy change of a system is:

ΔS^ꝋ = ΣS_products^ꝋ - ΣS_reactants^ꝋ

where Σ = sum of

entropies (S) of the reactants and products

Gibbs equation:

ΔGꝋ = ΔHreactionꝋ - TΔSsystemꝋ

When ΔG^ꝋ is negative, the reaction is ______ and likely to occur

feasible

other equation to work out gibbs:

ΔG^ꝋ = ΣΔG_products^ꝋ - ΣΔG_reactants^ꝋ

While ∆G can be used to determine the feasibility of a reaction, it does not take into account the ______ of the reaction i.e. rate of reaction

• There might be a large energy barrier (E_a) which the reacting species have to overcome before a reaction can occur

• Some reactions are feasible since ∆G is negative, but kinetically not feasible since it just occurs too slowly

• Such reactions are feasible but very slow

kinetics