Predicting Reactions of Metals & Redox Chemistry

Definition (electron perspective)
- Chemical process involving transfer of electrons between two species
- Results in change in oxidation number of at least one element
Biological & industrial relevance
- Photosynthesis, cellular respiration
- Combustion engines, corrosion/rusting, thermite welding
Not every chemical reaction is redox
- Double‐displacement and many acid–base neutralisations ≠ electron transfer
- Hydration of oxides & simple acid/base combination to form salts usually non-redox

Represent the charge an atom appears to have when electrons are assigned by electronegativity rules
Seven guiding rules:
1. Free/uncombined element: OS = 0
2. Sum of OS values = 0 for neutral species; = ionic charge for polyatomic ions
3. Group 1 metals: +1 ; Group 2 metals: +2
4. Fluorine: always -1 in compounds
5. Hydrogen: generally +1 (except -1 in metal hydrides)
6. Oxygen: generally -2 (except peroxides -1, superoxides -\tfrac12, OF₂ +2)
7. In binary metal compounds:
- Group 17 (halogens): -1
- Group 16 (chalcogens): -2
- Group 15 (pnictogens): -3
Sum of OS values validates neutrality or ionic charge

Example 1a: \text{Fe}(s)+\text{O}2(g)\;\to\;\text{Fe}2\text{O}_3(g)
- Reactants: Fe, O₂ free elements ⇒ OS = 0
- Product: O = -2 ⇒ 3O contributes -6; Fe₂ contributes +6 ⇒ each Fe = +3
Example 1b: \text{Fe}^{2+}(aq) ⇒ OS = +2
Example 1c: \text{Ag}(s)+\text{H}2\text{S}\;\to\;\text{Ag}2\text{S}+\text{H}_2(g)
- Ag (s) = 0; H in H₂S = +1; S in H₂S = -2; H₂ = 0
- Thus Ag in Ag₂S = +1
Example 2a: \text{Na}3\text{PO}3 (sodium phosphite)
- Na = +1 (×3 = +3); O = -2 (×3 = –6)
- Let P = x; Sum = 0 ⇒ 3 + x - 6 = 0 ⇒ x = +3
Example 2b: \text{H}2\text{PO}4^{-}
- H = +1 (×2 = +2); O = -2 (×4 = –8); Charge = –1
- Let P = y; y + 2 - 8 = -1 ⇒ y = +5

OIL RIG: “Oxidation Is Loss, Reduction Is Gain” (of electrons)
Oxidising agent (OA): species that is reduced; accepts e⁻
Reducing agent (RA): species that is oxidised; donates e⁻
In disproportionation one element is simultaneously oxidised & reduced in different atoms of same species

Reaction: \text{Fe}2\text{O}3 + 2 \text{Al} \to 2\text{Fe} + \text{Al}2\text{O}3
Fe in \text{Fe}2\text{O}3: OS = +3 → 0 (reduction)
Al: 0 → +3 in \text{Al}2\text{O}3 (oxidation)
Releases intense heat – used for welding rails, incendiary devices; illustrates metal reactivity hierarchy

Standard Reduction Potential (SRP)
- Potential of a half-reaction under standard conditions 25^\circ\text{C},\;1\,\text{M},\;1\,\text{atm} measured vs. Standard Hydrogen Electrode (SHE)
- Example: \text{Cu}^{2+}+2e^-\to\text{Cu}(s) , E^\circ = +0.340\,\text{V}
Standard Oxidation Potential (SOP)
- Same magnitude, opposite sign: \text{Cu}(s)\to \text{Cu}^{2+}+2e^- , E^\circ = -0.340\,\text{V}
Relationship: E^\circ{SRP} = -E^\circ{SOP}

Reference half-cell: \text{Pt}|\text{H}_2(g,1\,\text{atm})|\text{H}^+(1\,\text{M}) , defined E^\circ = 0.00\,\text{V}
Procedure:
- Connect SHE to test half-cell via salt bridge; external circuit includes voltmeter
- Direction of electron flow indicates whether unknown half-cell acts as cathode (reduction) or anode (oxidation)
Example experimental set-up for copper (Figure 2): observed voltage +0.340\,\text{V} (Cu electrode positive vs SHE) ⇒ Cu²⁺ is being reduced

SRP table ordered from most positive (strong oxidising agents) to most negative (strong reducing agents)
Snapshot (selected values):
- \text{F}_2 + 2e^- \to 2\text{F}^- : +2.87\,\text{V} (highest)
- \text{Ag}^+ + e^- \to \text{Ag} : +0.80\,\text{V}
- 2\text{H}^+ + 2e^- \to \text{H}_2 : 0.00\,\text{V}
- \text{Zn}^{2+} + 2e^- \to \text{Zn} : -0.76\,\text{V}
- \text{Li}^+ + e^- \to \text{Li} : -3.04\,\text{V} (lowest)
Usage rules:
- Species higher in table (more positive E°) preferentially gets reduced (cathode)
- Species lower (more negative E°) preferentially gets oxidised (anode)

Metal A placed in a solution of metal B ions:
- If E^\circ{red}(\text{B}^{n+}/\text{B}) > E^\circ{red}(\text{A}^{m+}/\text{A}) → Bⁿ⁺ will be reduced & metal A will oxidise ⇒ reaction proceeds
- Example: Zn (–0.76 V) in Cu²⁺ solution (+0.34 V) ⇒ Zn oxidises, Cu²⁺ reduces ⇒ Zn dissolves & Cu metal plates

For full cell: E^\circ{cell} = E^\circ{cathode} - E^\circ_{anode}
If E^\circ{cell} > 0 ⇒ \Delta G^\circ < 0 (since \Delta G^\circ = -n F E^\circ{cell}) ⇒ spontaneous
If E^\circ_{cell} < 0 ⇒ non-spontaneous in forward direction (would proceed in reverse)

Steps (acidic medium):
1. Write separate half-reactions for oxidation & reduction
2. Balance atoms other than O & H
3. Balance O with \text{H}_2\text{O}, balance H with \text{H}^+
4. Balance charge with e⁻
5. Multiply half-reactions to equalise electrons; add together & cancel
For basic medium: add equal \text{OH}^- to neutralise \text{H}^+ → form \text{H}_2\text{O}, then cancel

Assemble two different metal/ion half-cells with salt bridge (e.g., Cu/Cu²⁺ and Zn/Zn²⁺)
Measure open-circuit voltage; compare with theoretical E^\circ_{cell}
Variables to control: temperature, ion concentration, electrode surface area
Possible extensions: concentration cells, effect of common ions, corrosion protection studies

Energy storage: batteries (alkaline, Li-ion, fuel cells) rely on predictable redox potentials
Environmental: corrosion prevention (galvanising with Zn; sacrificial anodes)
Ethical/Philosophical: sustainable sourcing of reactive metals, ecological impact of mining & disposable batteries
Biological redox chains: electron transport chain utilises sequential E° differences to synthesise ATP – illustrates universality of redox principles

Faraday constant: F = 96\,485\,\text{C}\,\text{mol}^{-1}
Relation between Gibbs energy & cell potential: \Delta G^\circ = -n F E^\circ_{cell}
Nernst equation (non-standard conditions): E = E^\circ - \frac{0.0592}{n}\log Q