EC2526 — Electrochemistry Tutorial
Source: EC2526 MY V3 - STDN Copy.pdf (inbox)
Note: This is a student handout with heavy image/slide content. Extractable text is limited; the summary below captures the key concepts present in the document.
Overview
Tutorial covering galvanic (voltaic) cells, cell notation, half-equations, oxidizing and reducing agents, and electromotive force (EMF).
Redox Reactions & Half-Equations
Example Reactions
i. Zinc–Cadmium $$\ce{Zn(s) + Cd^{2+}(aq) -> Zn^{2+}(aq) + Cd(s)}$$
- Anode (oxidation): $\ce{Zn(s) -> Zn^{2+}(aq) + 2e^-}$
- Cathode (reduction): $\ce{Cd^{2+}(aq) + 2e^- -> Cd(s)}$
ii. Magnesium–Scandium $$\ce{3Mg(s) + 2Sc^{3+}(aq) -> 3Mg^{2+}(aq) + 2Sc(s)}$$
- Anode: $\ce{3Mg(s) -> 3Mg^{2+}(aq) + 6e^-}$
- Cathode: $\ce{2Sc^{3+}(aq) + 6e^- -> 2Sc(s)}$
iii. Iron–Permanganate $$\ce{5Fe^{2+}(aq) + MnO4^-(aq) + 8H+(aq) -> 5Fe^{3+}(aq) + Mn^{2+}(aq) + 4H2O(l)}$$
- Anode: $\ce{5Fe^{2+}(aq) -> 5Fe^{3+}(aq) + 5e^-}$
- Cathode: $\ce{MnO4^-(aq) + 8H+(aq) + 5e^- -> Mn^{2+}(aq) + 4H2O(l)}$
Cell Notation
Cell notation follows the convention:
$$\text{Anode } | \text{ Anode ion } || \text{ Cathode ion } | \text{ Cathode }$$
Examples
| Cell | Notation |
|---|---|
| a | $\ce{Cr(s) |
| b | $\ce{Pt(s) |
| c | $\ce{Sc(s) |
Note: Platinum ($\ce{Pt}$) is used as an inert electrode when no solid metal is involved in the half-reaction.
Identifying Oxidizing & Reducing Agents
| Oxidizing Agent | Reducing Agent | |
|---|---|---|
| Description | Species that undergoes reduction | Species that undergoes oxidation |
| Electron transfer | Gains electrons | Loses electrons |
| Oxidation number | Decreases | Increases |
From Cell Examples
| Cell | Oxidizing Agent | Reducing Agent |
|---|---|---|
| a ($\ce{Cr | Pb}$) | |
| b ($\ce{Br2 | Cl2}$) | |
| c ($\ce{Sc | H2}$) |
Electromotive Force (EMF)
Galvanic Cell Basics
- Two electrodes connected by a wire and voltmeter.
- Anode: Site of oxidation; electrons leave → labeled negative terminal.
- Cathode: Site of reduction; electrons arrive → labeled positive terminal.
- Electrons always travel from anode to cathode.
- Current flows from cathode to anode (opposite to electron flow).
Spontaneity
- Positive EMF values → spontaneous redox reaction.
- Negative $\Delta G$ → releases energy to the environment.
Salt Bridge Function
- Donates anions and cations to each side to neutralize building charge.
- Without it, charge buildup would stop the reaction.
- Anode solution becomes more positive as metal ions dissolve.
- Cathode solution becomes more negative as metal ions deposit.
Relative Strength of Agents
| Oxidizing Agents | Reducing Agents | |
|---|---|---|
| Strength trend | More positive $E^\circ$ value → stronger | More negative $E^\circ$ value → stronger |
| Electrode assignment | Cathode = more positive $E^\circ$ | Anode = more negative $E^\circ$ |
Related Concepts
- Electrochemistry — General electrochemistry concepts
- FAD1018 - Basic Chemistry II — Course page