Thermodynamics
17.1 Enthalpy Change:
17.1 Enthalpy change
- The lattice enthalpy of formation ∆! 𝐻∅ is the standard
enthalpy change when 1 mole of a solid ionic compound is 17.2 Born-Haber cycles
formed from its gaseous ions. 17.3 More enthalpy changes
- Lattice enthalpies are always exothermic, and they cannot be
determined directly, they must be determined from other data. 17.4 Why do chemical reactions take place
- The lattice enthalpy of dissociation is the standard enthalpy change when 1 mole of solid ionic
compound is dissociated into its gaseous ions.
- The standard enthalpy change of atomisation ∆#$ 𝐻∅ is the enthalpy which accompanies the formation
of 1 mole of gaseous atoms from the element in its standard state under standard conditions
- For covalent elements (e.g., Cl2) the enthalpy change is ½ of the bond enthalpy of the element.
- The first ionisation energy ∆%& 𝐻∅ is the standard enthalpy change when one mole of gaseous atoms is
converted into a mole of gaseous ions each with a single positive charge.
- The electron affinity ∆'(& 𝐻∅ is the standard enthalpy change when a mole of gaseous atoms is
converted to a mole of gaseous ions, each with a single negative charge.
- The first is always negative, and all subsequent are positive.
- The standard entropy 𝑆 ∅ is the entropy of one mole of the substance under standard conditions.
- The standard enthalpy of solution ∆)*+ 𝐻∅ is the standard enthalpy change when one mole of a solute
dissolves completely in sufficient solvent to form a solution in which the molecules or ions are far
enough apart not to interact with each other.
- The standard enthalpy change of hydration ∆,-. 𝐻∅ is the standard enthalpy change when 1 mole of
gaseous ions are converted into aqueous ions.
- They are all exothermic.
17.2 Born-Haber Cycles:
- An energy scale is drawn on the Y-axis, every step must be thoroughly labelled, with state symbols.
- An endothermic reactions will point up, exothermic points down.
- The larger the ionic radius the less
∆#$ 𝐻∅ Na = +108 kJ mol-1 exothermic the lattice enthalpy of
∆#$ 𝐻∅ Cl = + 122 kJ mol-1 formation. Larger ions have the charge
∆%& 𝐻∅ = + 496 kJ mol-1 spread out over a larger volume. This
∅
∆'(& 𝐻 = -349 kJ mol -1 means there is a lower charge density.
∅
∆/ 𝐻 = -411 kJ mol -1 Therefore, there is a lower electrostatic
attraction and so a less exothermic
lattice enthalpy.
- The higher the charge on the ion the
more exothermic the lattice enthalpy, as
there is a stronger electrostatic
attraction.
17.1 Enthalpy Change:
17.1 Enthalpy change
- The lattice enthalpy of formation ∆! 𝐻∅ is the standard
enthalpy change when 1 mole of a solid ionic compound is 17.2 Born-Haber cycles
formed from its gaseous ions. 17.3 More enthalpy changes
- Lattice enthalpies are always exothermic, and they cannot be
determined directly, they must be determined from other data. 17.4 Why do chemical reactions take place
- The lattice enthalpy of dissociation is the standard enthalpy change when 1 mole of solid ionic
compound is dissociated into its gaseous ions.
- The standard enthalpy change of atomisation ∆#$ 𝐻∅ is the enthalpy which accompanies the formation
of 1 mole of gaseous atoms from the element in its standard state under standard conditions
- For covalent elements (e.g., Cl2) the enthalpy change is ½ of the bond enthalpy of the element.
- The first ionisation energy ∆%& 𝐻∅ is the standard enthalpy change when one mole of gaseous atoms is
converted into a mole of gaseous ions each with a single positive charge.
- The electron affinity ∆'(& 𝐻∅ is the standard enthalpy change when a mole of gaseous atoms is
converted to a mole of gaseous ions, each with a single negative charge.
- The first is always negative, and all subsequent are positive.
- The standard entropy 𝑆 ∅ is the entropy of one mole of the substance under standard conditions.
- The standard enthalpy of solution ∆)*+ 𝐻∅ is the standard enthalpy change when one mole of a solute
dissolves completely in sufficient solvent to form a solution in which the molecules or ions are far
enough apart not to interact with each other.
- The standard enthalpy change of hydration ∆,-. 𝐻∅ is the standard enthalpy change when 1 mole of
gaseous ions are converted into aqueous ions.
- They are all exothermic.
17.2 Born-Haber Cycles:
- An energy scale is drawn on the Y-axis, every step must be thoroughly labelled, with state symbols.
- An endothermic reactions will point up, exothermic points down.
- The larger the ionic radius the less
∆#$ 𝐻∅ Na = +108 kJ mol-1 exothermic the lattice enthalpy of
∆#$ 𝐻∅ Cl = + 122 kJ mol-1 formation. Larger ions have the charge
∆%& 𝐻∅ = + 496 kJ mol-1 spread out over a larger volume. This
∅
∆'(& 𝐻 = -349 kJ mol -1 means there is a lower charge density.
∅
∆/ 𝐻 = -411 kJ mol -1 Therefore, there is a lower electrostatic
attraction and so a less exothermic
lattice enthalpy.
- The higher the charge on the ion the
more exothermic the lattice enthalpy, as
there is a stronger electrostatic
attraction.
