Reactivity 3.2 Flashcards

Loss of electrons (OIL). Also: an increase in oxidation state.

Gain of electrons (RIG). Also: a decrease in oxidation state.

A substance that accepts electrons and causes another substance to be oxidised. The oxidising agent itself is reduced.

A substance that donates electrons and causes another substance to be reduced. The reducing agent itself is oxidised.

1. Loss of electrons. 2. Increase in oxidation state. 3. Gain of oxygen or loss of hydrogen.

−2. Exceptions: peroxides (−1, e.g. H₂O₂), OF₂ (+2), and as the free element O₂ (0).

+1. Exception: in metal hydrides (e.g. NaH), it is −1.

−1 (the most electronegative element).

Zero. In a polyatomic ion, the sum equals the overall ion charge.

Assign oxidation states to all elements on both sides. If any element’s oxidation number changes, it is a redox reaction.

A redox reaction where a single element is simultaneously oxidised AND reduced. Example: Cl₂ + H₂O → HClO + HCl (Cl goes from 0 to +1 and −1).

1. Balance atoms other than O and H. 2. Balance O with H₂O. 3. Balance H with H⁺. 4. Balance charge with e⁻. 5. Combine, cancelling electrons.

Mg → Mg²⁺ + 2e⁻

Cu²⁺ + 2e⁻ → Cu

The negative electrode (−). Oxidation occurs here. Electrons flow FROM the anode through the external circuit.

The positive electrode (+). Reduction occurs here. Metal ions gain electrons and deposit as solid metal.

Completes the circuit by allowing ion flow between half-cells, maintaining electrical neutrality. Contains unreactive ions like K⁺ and NO₃⁻.

Voltaic: chemical → electrical energy (spontaneous). Electrolytic: electrical → chemical energy (non-spontaneous, requires external power).

Positive (+). This is opposite to a voltaic cell where the anode is negative (−). But in both cells: An Ox, Red Cat.

Cathode: Na metal (Na⁺ + e⁻ → Na). Anode: Cl₂ gas (2Cl⁻ → Cl₂ + 2e⁻).

A more reactive metal displays a less reactive metal from its compound. E.g. Zn + CuSO₄ → ZnSO₄ + Cu.

No. Copper is BELOW zinc in the reactivity series, so it cannot displace zinc. No reaction occurs.

Metals ABOVE hydrogen in the reactivity series (e.g. Mg, Zn, Fe). They produce a salt + hydrogen gas.

By electrolysis of their molten compounds, because carbon is not reactive enough to reduce their oxides.

Reduction with carbon monoxide in a blast furnace: Fe₂O₃ + 3CO → 2Fe + 3CO₂.

Both oxygen AND water must be present. Saltwater accelerates rusting by increasing water conductivity.

Coating iron with a MORE reactive metal (e.g. Zinc/galvanisation). The zinc oxidises preferentially, protecting the iron even if scratched.

1. Barrier protection (paint, oil, grease, plastic). 2. Sacrificial protection / galvanisation (coating with a more reactive metal like zinc).

Using electrolysis to coat an object with a thin metal layer. The object is the cathode; the plating metal is the anode; the electrolyte contains ions of the plating metal.

E°(cell) = E°(cathode) − E°(anode). If E°(cell) > 0, the reaction is spontaneous.

A stronger tendency to be oxidised → a stronger reducing agent (more reactive metal).

A stronger tendency to be reduced → a stronger oxidising agent.

The reference electrode assigned E° = 0.00 V. All other standard electrode potentials are measured relative to it.

ΔG° = −nFE°(cell), where n = moles of electrons, F = 96 485 C mol⁻¹. Positive E° → negative ΔG° → spontaneous.

NO. E° is an intensive property. When multiplying half-equations to equalise electrons, never multiply the E° value.

When the metal has a more negative reduction potential than water (−0.83 V). Water is reduced instead, evolving H₂ at the cathode.