Voltaic (Galvanic) Cell
Voltaic (Galvanic) Cells
A voltaic cell converts chemical energy → electrical energy via a spontaneous redox reaction. The two half-cells are separated so electrons must travel through an external wire.
Core Components of Voltaic (Galvanic) Cells
- Anode: Oxidation occurs (\(\text{AN OX}\)). Negative polarity in voltaic cells.
- Cathode: Reduction occurs (\(\text{RED CAT}\)). Positive polarity in voltaic cells.
- Salt bridge: Completes the electrical circuit and maintains electrical neutrality by ion migration (e.g. \(\text{KNO}_3\) or \(\text{KCl}\)).
- Electron flow: Always from anode to cathode through the external wire.
Zn-Cu Daniell Cell Reactions
Anode (Oxidation, −): \(\text{Zn}(\text{s}) \rightarrow \text{Zn}^{2+}(\text{aq}) + 2\text{e}^-\)
Cathode (Reduction, +): \(\text{Cu}^{2+}(\text{aq}) + 2\text{e}^- \rightarrow \text{Cu}(\text{s})\)
Overall Cell Reaction: \(\text{Zn}(\text{s}) + \text{Cu}^{2+}(\text{aq}) \rightarrow \text{Zn}^{2+}(\text{aq}) + \text{Cu}(\text{s}) \quad (E^\circ_{\text{cell}} = +1.10\text{ V})\)
Zinc is more reactive than copper, so it is oxidised. The zinc electrode loses mass while the copper electrode gains mass as Cu is deposited.
Electrolytic Cells
An electrolytic cell uses electrical energy → chemical energy to force a non-spontaneous reaction to occur. An external power supply drives electrons in the opposite direction.
Core Principles of Electrolysis
- Converts electrical energy into chemical energy using an external DC power supply.
- Drives non-spontaneous redox reactions (\(\Delta G > 0, E_{\text{cell}} < 0\)).
- Anode (+): Connected to positive battery terminal → oxidation occurs.
- Cathode (−): Connected to negative battery terminal → reduction occurs.
Electrolysis of Molten NaCl
Cathode (−): \(\text{Na}^+(\text{l}) + \text{e}^- \rightarrow \text{Na}(\text{l})\) (silvery liquid metal)
Anode (+): \(2\text{Cl}^-(\text{l}) \rightarrow \text{Cl}_2(\text{g}) + 2\text{e}^-\) (pale green pungent gas)
Overall: \(2\text{NaCl}(\text{l}) \rightarrow 2\text{Na}(\text{l}) + \text{Cl}_2(\text{g})\)
Voltaic vs Electrolytic. Comparison
Voltaic vs Electrolytic Cells
| Feature | Voltaic Cell | Electrolytic Cell |
|---|---|---|
| Energy Conversion | Chemical → Electrical | Electrical → Chemical |
| Spontaneity | Spontaneous (\(\Delta G < 0, E^\circ > 0\)) | Non-spontaneous (\(\Delta G > 0, E^\circ < 0\)) |
| Anode Polarity & Reaction | Negative (−), Oxidation | Positive (+), Oxidation |
| Cathode Polarity & Reaction | Positive (+), Reduction | Negative (−), Reduction |
In BOTH types: oxidation at anode, reduction at cathode. An Ox, Red Cat.
Standard Electrode Potential E° (HL)
The standard electrode potential (E°) measures the tendency of a half-cell to undergo reduction under standard conditions (298 K, 1 mol dm⁻³, 100 kPa). All values are measured relative to the Standard Hydrogen Electrode (SHE), which is assigned E° = 0.00 V. All E° values in data booklets are written as reduction potentials.
The Standard Hydrogen Electrode (SHE)
The Standard Hydrogen Electrode (SHE) is the universal reference standard, assigned \(E^\circ = 0.00\text{ V}\) at standard conditions (\(298\text{ K}, 100\text{ kPa } \text{H}_2, 1.00\text{ mol dm}^{-3 } \text{H}^+\)):
\[2\text{H}^+(\text{aq}) + 2\text{e}^- \rightleftharpoons \text{H}_2(\text{g}) \quad (E^\circ = 0.00\text{ V})\]Interpreting Standard Electrode Potentials
- More positive \(E^\circ\): Greater tendency to be reduced (stronger oxidising agent).
- More negative \(E^\circ\): Greater tendency to be oxidised (stronger reducing agent).
Standard Cell Potential Formula
If \(E^\circ_{\text{cell}} > 0\), the reaction is spontaneous under standard conditions.
Calculating E°cell for Zn-Cu Cell
\(E^\circ(\text{Cu}^{2+}/\text{Cu}) = +0.34\text{ V}\), \(E^\circ(\text{Zn}^{2+}/\text{Zn}) = -0.76\text{ V}\)
\[E^\circ_{\text{cell}} = +0.34 - (-0.76) = +1.10\text{ V}\]Relationship: ΔG° and E°cell
- \(n\) = number of moles of electrons transferred in balanced equation.
- \(F\) = Faraday constant = \(96500\text{ C mol}^{-1}\).
- \(E^\circ_{\text{cell}} > 0 \iff \Delta G^\circ < 0\) (spontaneous).
Electroplating
Electroplating uses electrolysis to coat an object with a thin layer of metal. The object to be plated is the cathode; the plating metal is the anode. The electrolyte contains ions of the plating metal.
Electrolysis of Aqueous Solutions (HL)
When electrolyzing an aqueous solution, there is competition between the solute ions and water molecules at each electrode.
Selective Discharge at Cathode
The species with the most positive \(E^\circ\) is reduced first at the cathode:
- In dilute \(\text{NaCl}(\text{aq})\): \(\text{H}_2\text{O}\) (\(E^\circ = -0.83\text{ V}\)) is reduced instead of \(\text{Na}^+\) (\(E^\circ = -2.71\text{ V}\)) → \(\text{H}_2(\text{g})\) produced.
Selective Discharge at Anode
The species with the least positive (most negative) \(E^\circ\) is oxidised first at the anode:
- In concentrated \(\text{NaCl}(\text{aq})\): high \([\text{Cl}^-]\) causes \(\text{Cl}^-\) to discharge as \(\text{Cl}_2(\text{g})\) despite \(E^\circ\) considerations (concentration effect).
Electrolysis of Brine (Conc. NaCl)
Cathode (−): \(2\text{H}_2\text{O}(\text{l}) + 2\text{e}^- \rightarrow \text{H}_2(\text{g}) + 2\text{OH}^-(\text{aq})\)
Anode (+): \(2\text{Cl}^-(\text{aq}) \rightarrow \text{Cl}_2(\text{g}) + 2\text{e}^-\)
Residual Solution: \(\text{Na}^+\) and \(\text{OH}^-\) remain → alkaline \(\text{NaOH}(\text{aq})\) solution forms.
Primary and Secondary Cells
The IB syllabus requires you to distinguish between primary cells, secondary cells, and fuel cells.
Primary (Disposable) Cells
Non-rechargeable electrochemical cells where the redox reaction is irreversible. As reactants are consumed, battery voltage drops until flat (e.g. alkaline zinc-manganese dioxide batteries).
Secondary (Rechargeable) Cells
Rechargeable batteries where applying an external electrical current reverses the cell redox reaction (e.g. lithium-ion, lead-acid car batteries).
Fuel Cells (H₂/O₂)
Galvanic cells supplied continuously with external fuel (e.g. \(\text{H}_2\)) and oxidant (e.g. \(\text{O}_2\)), producing water and electrical energy with zero direct greenhouse gas emissions.
| Primary | Secondary | Fuel Cell | |
|---|---|---|---|
| Rechargeable? | No | Yes | N/A (continuous fuel) |
| Reaction type | Irreversible | Reversible | Continuous |
| Example | Alkaline battery | Li-ion battery | H₂/O₂ fuel cell |
Salt Bridge Electrolyte Selection
The salt bridge must contain inert ions (\(\text{KNO}_3\)) that will not precipitate with cell electrolytes (unlike \(\text{Cl}^-\) which would precipitate \(\text{Ag}^+\) or \(\text{Pb}^{2+}\)).
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