Energy and thermochemistry
| Standard enthalpy change (ΔHø) | The enthalpy change of a reaction carried out under standard conditions (100 KPa, 1 mol dm-3, substances in standard state) ΔH(reaction) = ∑ΔHfø(products) – ∑ΔHfø(reactants) ΔHreaction = ∑E(bonds broken) – ∑E(bonds formed) |
|---|---|
| Enthalpy change of formation | It is the enthalpy change when 1 mol of a compound is formed from its elements in their standard states under standard conditions |
| Standard enthalpy of combustion (ΔHcombø) | The enthalpy change when 1 mol of substance burns completely in O2 under standard conditions ΔHø = ∑ΔHcombø(reactants) – ∑ΔHcombø(products) |
| ΔSøreaction = ∑Sø(products) – ∑Sø(reactants) | |
| Entropy and spontaneity | ΔS(total) = ΔS(system) + ΔS(surroundings) ΔSsurroundings=-ΔHsystemT For a spontaneous process, ΔS(total) > 0 |
| Gibbs free energy (ΔG) | It is the energy associated with a chemical reaction that can be used to do work ΔG = ΔH – TΔS ΔGøreaction = ∑ΔGfø(products) – ∑ΔGfø(reactants) [standard conditions] ΔG is negative for a spontaneous process |
Energy cycles
- The first ionization energy is the minimum energy required to remove one mole of electrons from one mole of gaseous atoms
- The first electron affinity is the enthalpy change when one mole of gaseous electrons is added to one mole of gaseous atoms
- The lattice enthalpy is the enthalpy change that occurs when one mole of a solid ionic compound is separated into gaseous ions under standard conditions
- The enthalpy change of atomization ΔHøatom is the enthalpy change that occurs when one mole of gaseous atoms is formed in its standard state
Redox processes
- The more reactive metal is always reactive while the less reactive is positive. Electrons flow form the more reactive metal to the less reactive metals because metals of a higher reactivity have a greater reducing power
- Electrons always flow in the external circuit from the anode to the cathode
- The more negative the Eø (electrode potential) value for a half-cell, the more readily it is reduced (the more readily it gains electrons, meaning it is less reactive)
- Electrons always flow towards the half-cell with the highest Eø value (the less reactive metal, which is more reduced)
| Voltaic cells | Electrolytic cells |
|---|---|
| Spontaneous reactions produce electrical current Current conducted by electron flow in wire and movement of ions in salt bridge Anode = negative Cathode = positive More reactive metal loses electrons Less reactive metal gains electrons Chemical → electrical energy Exothermic | Electrical current rives non-spontaneous reactions Current conducted by electron flow in wires and movement of ions in electrolyte Anode = positive Cathode = negative Electrical → chemical energy Endothermic |
Electrolytic cells
Electrolysis of water
| Cathode | Anode |
|---|---|
| Possible reactions: Na+(aq) + e- → Na(s) 2H2O(l) + 2e- → H2(g) + 2OH-(aq) H2O is preferentially reduced and H2(g) will be discharged | Possible reactions: 4OH-(aq) → 2H2O(l) + O2(g) + 4e- 2H2O(l) → 4H+(aq) + O2(g) + 4e- OH-(aq) is preferentially oxidized and O2(g) is discharged |
| Overall balanced equation: 2H2O(l) → 2H2(g) + O2(g) Colorless gases: O2 at the anode and H2 at the cathode Ratio of gases evolved = 2:1 of H2:O2 pH at the anode decreases as H+ is released, while the pH at the cathode increases as OH- is released |