Redox Reactions & Simple Galvanic Cells

Redox Reactions in Society

• Redox (reduction–oxidation) reactions underpin many technologies, especially batteries that power mobile phones, laptops, calculators, etc.

• Rechargeable batteries rely on repeatable redox processes; without them, devices would need constant connection to mains electricity.

• In galvanic/voltaic cells, the chemical (redox) energy is directly converted to electrical energy—practical evidence of electron transfer.

Simple Galvanic Cells

• A spontaneous metal–non-metal reaction is typically exothermic (heat-releasing).

• By physically separating the oxidation and reduction half-reactions into two half-cells, the spontaneous redox process can be harnessed as an electric current instead of heat.

• Core apparatus (see Fig. 12.3.1):

  • Two beakers (half-cells) each containing an electrode immersed in an electrolyte.

  • External wire connecting electrodes; may include a galvanometer/light globe to detect or utilise electron flow.

  • Salt bridge (filter paper or U-tube filled with inert electrolyte such as \text{KNO}_3) to maintain charge balance by ion migration.

Evidence for Electron Transfer (Zn | Cu Example)

• Direct mixing: \text{Zn(s)}+\text{Cu}^{2+}(\text{aq})\to\text{Zn}^{2+}(\text{aq})+\text{Cu(s)} releases heat.

• Separated in a cell:

  • Oxidation at Zn electrode: \text{Zn(s)}\to\text{Zn}^{2+}(\text{aq})+2e^{-}

  • Reduction at Cu electrode: \text{Cu}^{2+}(\text{aq})+2e^{-}\to\text{Cu(s)}

  • Electrons forced through the wire, detected by a galvanometer (needle deflects ➔ evidence of flow).

  • Negative ions in the salt bridge migrate toward the Cu half-cell (same direction as electrons); positive ions migrate toward the Zn half-cell, completing the circuit.

Components & Terminology

• Half-cell = electrode + electrolyte that together host one redox half-reaction.

  • Electrode: solid conductor (metal strip, graphite) providing electron pathway.

  • Electrolyte: ionic solution conducting via ion motion.

• In any galvanic cell:

  • Strongest reducing agent (lowest in reactivity series) is oxidised.

  • Oxidation site = Anode (negative in a galvanic cell).

  • Reduction site = Cathode (positive in a galvanic cell).

  • Memory aid: “AnOx / RedCat” (Anode–Oxidation, Reduction–Cathode).

Typical Metal | Metal-Ion Half-Cells (Cu│Cu²⁺ and Ag│Ag⁺)

• Denoted as \text{Cu(s)}/\text{Cu}^{2+}!(\text{aq}) and \text{Ag(s)}/\text{Ag}^{+}!(\text{aq}).

• When connected (Fig. 12.3.3):

  • Copper is the stronger reducing agent than silver (from reactivity series) ➔ Cu oxidises.

  • Half-equations:

  • \text{Cu(s)} \to \text{Cu}^{2+}!(\text{aq}) + 2e^- (anode, negative electrode)

  • \text{Ag}^{+}!(\text{aq}) + e^- \to \text{Ag(s)} (cathode, positive electrode)

  • Electrons flow Cu → Ag through the wire; ions in salt bridge move to balance charges, lighting an external bulb.

Worked Example 12.3.1 (Zn │Zn²⁺ & Pb │Pb²⁺)

• Reactivity series: Zn is lower (stronger reducing agent) than Pb.

• (a) Reducing agent: Zn(s); Oxidising agent: \text{Pb}^{2+}!(\text{aq}).

• (b) Half-reactions:

  • Anode: \text{Zn(s)} \to \text{Zn}^{2+}!(\text{aq}) + 2e^-

  • Cathode: \text{Pb}^{2+}!(\text{aq}) + 2e^- \to \text{Pb(s)}

• (c) Electron flow: Zn electrode → Pb electrode via external circuit.

• (d) Negative electrode: Zn strip; Positive electrode: Pb strip.

• (e) Anode: Zn; Cathode: Pb.

• Diagrammatic tips:

  • Always show both beakers, electrodes, salt bridge, wire with indicator, and arrow for “electron flow.”

Try-Yourself Example (Cu │Cu²⁺ & Ni │Ni²⁺)

• Predict using reactivity series (Ni is stronger reducing agent than Cu):

  • Ni oxidises (anode, negative).

  • \text{Ni(s)} \to \text{Ni}^{2+}!(\text{aq}) + 2e^-

  • \text{Cu}^{2+}!(\text{aq}) + 2e^- \to \text{Cu(s)} (cathode, positive).

  • Electrons flow Ni → Cu.

Historical Context: Earliest Batteries

• Luigi Galvani (1737–1798):

  • Observed frog-leg muscle contractions when two dissimilar metals contacted a nerve during thunderstorms.

  • Coined idea of “animal electricity.”

• Alessandro Volta (1745–1827):

  • Disagreed with “animal electricity”; argued metals created the emf, frog was a conductor.

  • 1792–1800: experimented with metal disks (Zn & Cu or Zn & Ag) separated by salt-soaked cloth; detected weak current on tongue.

  • Built the first practical battery—the “voltaic pile”—and “crown of cups” series cell (multiple galvanic cells in series for higher voltage).

  • Unit of electromotive force, the volt (V), named after him (1881).

• Case-study prompts:

  1. In Galvani’s frog-on-rail setup: Zn (if present) or iron acts as anode; Cu hook is cathode (Cu less reactive than Fe).

  2. Modern measurement of twitching: use high-speed camera with motion-tracking software or force transducer attached to muscle.

  3. Possible experiment variables: Independent – metal pair used; Controlled – temperature, frog tissue freshness, electrode spacing.

Galvanic Cells in Modern Society

• “Battery” (strictly) = series of cells; often a single cell is colloquially called a battery.

• Cell types:

  • Primary cell: disposable, non-rechargeable (e.g. zinc-carbon dry cell, alkaline cell).
    • Dry cell (1866, Leclanché): acidic electrolyte; sand/paste immobilised.
    • Alkaline cell: \text{KOH} or other basic electrolyte ➔ longer shelf-life & higher current.

  • Secondary cell: rechargeable (Li-ion, Ni-MH, lead-acid, etc.); key to electric vehicles, grid storage.

• Societal importance: facilitates portable electronics, renewable-energy uptake, reduced fossil-fuel reliance.

Recycling & Environmental / Ethical Implications

• July 2019: Victoria (AU) banned e-waste—including batteries—from landfill.

• Reasons to recycle:

  • Recover non-renewable metals (Zn, Ni, Li, Co, etc.).

  • Prevent leaching of toxic metals (Pb, Hg) into soil & waterways.

  • Ethical sourcing: metals like Co often mined under poor labor conditions—recycling lessens demand.

• Practical tip: Dedicated battery-recycling bins available at municipal waste transfer stations.

Key Numerical / Scientific Facts & Mnemonics

• In galvanic cells: Electrode potentials drive electron flow from lower reduction potential (stronger reducing agent) to higher reduction potential (stronger oxidising agent).

• Memory aids:

  • “LEO the lion says GER” – Loss of Electrons = Oxidation, Gain of Electrons = Reduction.

  • “AnOx / RedCat” – Anode Oxidation, Reduction Cathode.

• Salt bridge role: completes internal circuit; prevents charge build-up; typically contains inert ions (e.g. \text{K}^+ & \text{NO}_3^-) that do not react in either half-cell.

• Sign convention (galvanic): Anode = negative (electron source); Cathode = positive (electron sink).

• Electrons & negative ions travel in same direction within overall loop; positive ions migrate opposite direction—maintains electroneutrality.