This is a draft cheat sheet. It is a work in progress and is not finished yet.
Key Terms and Definitions
Term |
Definition |
Entropy (∆S) |
A measure of the disorder in a system. Units: J K⁻¹ mol⁻¹ |
Standard Entropy Change (∆S⁰) |
∆S⁰ = ΣS⁰(products) – ΣS⁰(reactants) |
Spontaneous Process |
A process where ∆G < 0; can occur without external input. |
Gibbs Free Energy (∆G) |
The energy available to do work: ∆G = ∆H – T∆S |
Lattice Enthalpy of Formation (∆H⁰LE) |
Enthalpy change when 1 mole of an ionic solid forms from its gaseous ions (exothermic). |
Lattice Enthalpy of Dissociation |
Enthalpy change when 1 mole of ionic solid dissociates into gaseous ions (endothermic). |
Atomisation Enthalpy (∆H⁰at) |
Enthalpy change to form 1 mole of gaseous atoms from the element. |
Electron Affinity (∆H⁰ea) |
Enthalpy change when 1 mol of electrons is added to 1 mole of gaseous atoms to form 1⁻ ions. |
Hydration Enthalpy (∆H⁰hyd) |
Enthalpy change when gaseous ions dissolve in water to form aqueous ions. |
Solution Enthalpy (∆H⁰sol) |
Enthalpy change when 1 mol of solute dissolves to infinite dilution. |
Perfect Ionic Model |
Assumes ions are spherical and charge is evenly distributed with no covalent character. |
Entropy Effect
Change |
Entropy Effect |
Melting / Boiling |
Entropy increases significantly |
Dissolving ionic lattice |
Entropy increases |
Increase in gas moles in products |
Entropy increases |
Cooling a gas |
Entropy decreases |
Formation of solid from gas |
Entropy decreases significantly |
Gibbs Free Energy vs. Temperature
Equation: ∆G = –∆S·T + ∆H (y = mx + c)
Slope = –∆S
Y-intercept = ∆H
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Key Formulas
✅ Overall Entropy of a system ΔStotal = ΔSsystem + ΔSsurroundings
✅ Entropy of surroundings ΔSsurroundings = - ΔH/T
✅ Entropy Change of a reaction within a system
ΔS∘system = ∑S∘products − ∑S∘reactants
If entropy change is +ve, products more disordered than reactants (natural direction of change)
& if entropy change is -ve, reactants more disordered than products.
✅ Gibbs Free Energy ΔG=ΔH−TΔSsystem
- ∆G: Gibbs free energy (kJ·mol⁻¹)
- ∆H: Enthalpy change (kJ·mol⁻¹)
- T: Temperature (K)
- ∆S: Entropy change (J·K⁻¹·mol⁻¹; convert to kJ by dividing by 1000)
✅ Feasibility Conditions
A reaction is feasible when: ΔG < 0
To find minimum temperature for spontaneity: ΔG = 0 ⇒ T = ΔH/ΔSsystem
✅ Born-Haber Cycle Equation
For lattice formation enthalpy:
ΔH∘f= ∑enthalpies (atomisation, ionisation, electron affinity, etc.) + ΔHLE
✅ Hydration & Solution Enthalpy (Hess's Cycle) Δ𝐻∘sol = ∑Δ𝐻∘ − Δ𝐻∘LE |
Factors Affecting Lattice & Hydration Enthalpies
Factor |
Effect on ∆H (more negative) |
Smaller ionic radius |
Increases attraction → More negative ∆H |
Higher ionic charge |
Increases attraction → More negative ∆H |
Higher ionic charge Increases attraction → More negative ∆H Greater hydration energy |
Stronger interactions with water → More negative ∆H |
Kinetics vs. Thermodynamics
A reaction may be thermodynamically feasible (∆G < 0) but may not occur due to:
- High activation energy
- Slow reaction rate due to kinetic barriers |
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