Chapter 23 (Part 1) Energetics Flashcards
Lattice Enthalpy. (2)
- a measure of the strength of forces between the ions in an ionic solid
- the greater the lattice enthalpy, the stronger the forces
Lattice Formation Enthalpy
- enthalpy change when one mole of an ionic lattice is formed from its isolated gaseous ions
- highly exothermic
- strong electrostatic attraction btw oppositely-charged ions
- always negative
Lattice Dissociation Enthalpy
- enthalpy change when one mole of an ionic lattice dissociates into isolated gaseous ions
- highly endothermic
- strong electrostatic attraction btw oppositely-charged ions
- always positive
Factors affecting lattice enthalpy
- charge density
- ionic radius
Charge Density
- the strength of the bond between the ions of opposite charge in an ionic compound
- depends on charges on the ion & the distance btw the centres of ions when they pack to form a crystal
- greater charge = greater force = lattice energies more exothermic
Ionic Radius
- smaller ionic radius = greater attraction between ions because of greater charge density
- increased lattice enthalpies, increased melting points
Electron affinity
- energy change when 1 mole of electron is added to 1 mole of gaseous atom to form a negatively-charged ion
- X (g) + e- –> X- (g)
Why is 1st EA exothermic & 2nd, 3rd etc. EA endothermic?
2nd EA onwards:
- electron repels the anion
- energy absorbed to overcome repulsion between electron & anion
Steps for Simple Hess Cycle.
- Write out balanced H f & H LE.
- Atomise.
- Ionise. Metals must always be ionised before non-metals because electrons are needed for the non-metal.
- Ions come together to form solid.
- refer to class discussion
Steps for constructing Born Haber Cycle.
- Big arrow upwards for energy (E).
- Arrow upwards = endothermic / enthalpy change +ve
- Arrow downwards = exothermic / enthalpy change -ve
Enthalpy of atomisation.
- Enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state, under standard conditions.
Bond enthalpy (ΔHBe)
- Energy required to break a particular covalent bond in 1 mole of molecule in a gaseous state
- 1/2 Cl2 (g) –> Cl (g)
Carbonate anion dot-&-cross
- Refer to class discussion
- 2 single, 1 double
Nitrate anion dot-&-cross
- Refer to class discussion
- 1 double, 1 single, 1 dative
Explain delocalisation on carbonate & nitrate ions.
- pi bond on C=O & lone pairs overlaps with delocalised pi system
-anion is polarised due to attraction towards G2 cation (think of dipole) - under heating, weakest (polarised) bond breaks
Thermal stability of G2 carbonates & nitrates.
Going down G2:
- ionic radius increase
- charge density decrease
- anion is less polarised
- more heat needed to decompose
- thermal stability increase
Enthalpy change of solution (ΔH Sol)
- Enthalpy change when 1 mole of solute is dissolved in water to form infinitely dilute solution under standard conditions.
- Soluble compounds = exothermic, -ve
- Insoluble compounds = endothermic, +ve
Enthalpy change of hydration (ΔH Hyd)
- Enthalpy change when 1 mole of gaseous ion dissolves in water to give infinitely dilute solution.
- Mg2+ (g) –> Mg2+ (aq)
- exothermic (ALWAYS)
Variation of G2 sulfates down the group.
- ΔH LE & ΔH hyd become less exothermic
- ionic radius increases, charge density decreases
- ΔH hyd decreases more than ΔH LE
- ΔH Sol becomes more endothermic (more energy required)
- Solubility of G2 sulfates decreases
Variation of G2 hydroxides down the group.
- ΔH LE & ΔH hyd become less exothermic
- ionic radius increases, charge density decreases
- ΔH LE decreases more than ΔH Hyd
- ΔH Sol becomes more exothermic
- Solubility of G2 hydroxides increases
ΔH Sol = ?
- ΔH LE (formation) + ΔH Hyd
