Module 5: 5 - Enthalpy and Entropy
What does enthalpy change of solution depend on?
Lattice enthalpy of the solid Enthalpy of hydration for the ions
Effect of ionic charge on lattice enthalpy
- Greater charge over same size = greater charge density. - Greater force of attraction between ions - Stronger ionic bond - Lattice enthalpy is more exothermic.
Effect of ionic size on lattice enthalpy
- Ionic radius is smaller - Greater charge density - Greater force of attraction between ions - Stronger ionic bond - Lattice enthalpy is more exothermic
CaO vs. MgS
Ca²⁺ ions and Mg²⁺ ions have the same charge but Mg²⁺ ions are smaller and so have a greater charge density. O²⁻ and S²⁻ have the same charge but O²⁻ ions are smaller and so have a greater charge density. It is difficult to predict the magnitude of lattice enthalpy between the two without relative values of the electron charge density to see the difference of effects.
Enthalpy of hydration of ions
Depends on size and charge of the ions More exothermic if the ion is small and highly charge (high charge density) This is because the ions will attract the water molecules more and form a stronger bond.
Increasing entropy
During changes of state that give more entropy: solid → liquid → gas When a solid lattice dissolves When a gas is evolved Number of gas molecules increase
Born-Haber Cycle
Lattice enthalpy cannot be measured directly and must be measured using known energy changes in an energy cycle. The indirect determination of lattice enthalpy needs a Born-Haber cycle.
Factors affecting lattice enthalpy
Lattice enthalpy values give a measure of ionic bond strength. The more exothermic the lattice enthalpy values, the 'stronger' the ionic bond. There is a greater force of attraction between the ions. Two factors influence this : ionic size ionic charge
NaCl vs. MgCl₂
Mg²⁺ ions are smaller than Na⁺ ions, and Mg²⁺ ions have a greater charge than Na⁺ ions. This means that Mg²⁺ ions have a greater charge density so there is a greater force of attraction between Mg²⁺ and Cl⁻ ions, so a stronger ionic bond is formed. The lattice enthalpy of MgCl₂ is more exothermic than NaCl.
Energy cycle for enthalpy change of solution
NB: Enthalpy change of solution has to be calculated first because it could be exothermic or endothermic ΔH𝑙𝑎𝑡𝑡 + ΔH𝑠 = ΔH𝘩𝑦𝑑 of ions
Lattice enthalpy of the solid
Opposite of the lattice enthalpy (lattice dissociation enthalpy) Breaking lattice apart Endothermic process
Spontaneous reactions
Reaction that happens of its own accord, without external help. Exothermic reactions are usually spontaneous because they go from higher to lower enthalpy.
Entropy change, Δ𝑠
The difference in the entropy of the products compared to the reactants. Can be positive or negative.
Enthalpy change of solution, ΔH𝑠
The enthalpy change when one mole of a compound is completely dissolved in water under standard conditions. 1) Breakdown of lattice into its gaseous ions 2) Surround the ions with water molecules - hydrate
Lattice Enthalpy
The enthalpy change when one mole of an ionic solid lattice is formed from its gaseous ions under standard conditions. Lattice enthalpy involves ionic bond formation from separate gaseous ions. It is an exothermic change and the enthalpy change will always be negative.
Enthalpy change of hydration, ΔH𝘩𝑦𝑑
The enthalpy change when one mole of gaseous ions is dissolved in water to form one mole of aqueous ions under standard conditions. 1) Forming new bonds with the water molecules so it is an exothermic reaction.
Predicting melting points
The magnitude of lattice enthalpy is a good indicator of melting point of an ionic compound. Very exothermic lattice enthalpies = high melting point
Free energy
The overall change in energy in a chemical reaction is called the free energy change, ΔG and is made up of two types of energy: - Enthalpy change ΔH - Entropy change at the temperature of the reaction TΔS
Entropy, 𝑠
The quantitative measure of the disorder of a system. Entropy is sometimes considered as 'the amount of randomness' in a system. All substances possess some degree of disorder because particles are always in constant motion so S is always a positive number.
Standard entropy
The standard entropy of a substance is the entropy of one mole of a substance under standard conditions. Measured in J K⁻¹ mol⁻¹ Always positive
Successive electron affinities
When an anion has a greater charge than 1-, successive electron affinities are required. Second electron affinities are endothermic. A second electron is gained by a negative ion which repels the electron away. So energy must be put in to force the negatively charged electron onto the negative ion.
Feasibility
Whether a reaction is able to happen and is energetically feasibly.
Gibbs Free Energy Equation
ΔG = ΔH - TΔS For a reaction to be feasible, ΔG < 0 NB Temp in K Remember to convert everything to KJmol⁻¹
Standard enthalpy change of atomisation
ΔH𝑎𝑡 is the enthalpy change that takes place for the formation of one mole of gaseous atoms from its element in its standard state in standard conditions. It is always endothermic because bonds are broken to form gaseous atoms.
First electron affinity
ΔH𝑒𝑎 is the enthalpy change that takes place when one electron is added to each atom in one mole of gaseous atoms to form one mole of gaseous 1- ions. First electron affinities are exothermic because the electron being added is attracted in towards the nucleus.
First Ionisation energy
ΔH𝑖𝑒 is the enthalpy change required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions. Ionisation energies are endothermic because energy is required to overcome the attraction between a negative electron and the positive nucleus.
Standard enthalpy change of formation
ΔH𝘧 is the enthalpy change that takes place when one mole of a compound is formed from its elements under standard conditions, with all reactants and products in their standard states.
Calculating entropy change
ΔS = Σ𝐬 (products) - Σ𝐬 (reactants)
