L1 L2 Electrochemistry

Lectures 1-2 on Electrochemistry for FAD1018 Basic Chemistry II. Source file: L1 L2 EC - STDN Copy.pdf (82 pages, PowerPoint slides).

[!note] Direct Image Processing Content reconstructed from direct visual processing of all 82 slide images.

Learning Outcomes

  • Define electrochemistry and differentiate between galvanic and electrolytic cells
  • Write balanced redox half-reactions and identify anode/cathode processes
  • Represent electrochemical cells using standard cell notation
  • Define standard electrode potential and standard cell potential
  • Calculate cell potential using standard reduction potentials

1. Introduction to Electrochemistry

Electrochemistry is the study of chemical processes that cause electrons to move — involving the interconversion of electrical and chemical energy.

Two Types of Electrochemical Cells

Feature Galvanic (Voltaic) Cell Electrolytic Cell
Energy conversion Chemical → Electrical Electrical → Chemical
Spontaneity Spontaneous reaction Non-spontaneous (forced)
Anode Negative electrode Positive electrode
Cathode Positive electrode Negative electrode
Example Battery Electroplating
  • Galvanic cell: Naturally occurs; spontaneous redox reaction generates electricity.
  • Electrolytic cell: Electricity drives a reaction that would not normally happen; pulls out electrons at anode, pushes in electrons at cathode.
[Zn+2]

[Cu+2]

2. Redox Reactions in Electrochemical Cells

Electrochemical cells rely on reduction-oxidation (redox) reactions:

  • Oxidation (at anode): loss of electrons
  • Reduction (at cathode): gain of electrons

Mnemonic: OIL RIG — Oxidation Is Loss, Reduction Is Gain (of electrons)

Example: Zn-Cu Galvanic Cell

  • Anode (Zn): $\text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2e^-$ — oxidation
  • Cathode (Cu): $\text{Cu}^{2+}(aq) + 2e^- \rightarrow \text{Cu}(s)$ — reduction
  • Overall: $\text{Zn}(s) + \text{Cu}^{2+}(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{Cu}(s)$

How a Galvanic Cell Works

  1. Zn electrode erodes as Zn oxidises to Zn²⁺
  2. Cu²⁺ deposits on Cu electrode as it reduces to Cu
  3. Anode solution becomes positively charged (Zn²⁺ accumulation)
  4. Cathode solution becomes negatively charged (Cu²⁺ depletion)
  5. Salt bridge neutralises charge buildup with ion flow (anions → anode, cations → cathode)

3. Cell Notation (Cell Diagram)

Standard notation for representing electrochemical cells:

$$ \text{Anode} \mid \text{Anode electrolyte} \parallel \text{Cathode electrolyte} \mid \text{Cathode} $$

  • Single vertical line (|): phase boundary
  • Double vertical line (||): salt bridge
  • Inert electrodes (Pt, graphite): used when no solid metal participates

Examples:

  • $\text{Zn}(s) \mid \text{Zn}^{2+}(aq) \parallel \text{Cu}^{2+}(aq) \mid \text{Cu}(s)$
  • $\text{Pt}(s) \mid \text{Br}^-(aq), \text{Br}_2(l) \parallel \text{Cl}_2(g), \text{Cl}^-(aq) \mid \text{Pt}(s)$
BrBr

ClCl

[Cl-]

[Sc+3]

[H+]

[H][H]

4. Standard Electrode Potential (E°)

The potential of a half-cell under standard conditions (1 M solutions, 1 atm gases, 25°C).

Standard Hydrogen Electrode (SHE)

Reference electrode with $E° = 0.00$ V: $$2\text{H}^+(aq) + 2e^- \rightarrow \text{H}_2(g)$$

Relative Strength of Agents

  • Oxidising agent: species that is reduced (gains electrons); stronger when $E°$ is more positive
  • Reducing agent: species that is oxidised (loses electrons); stronger when $E°$ is more negative
  • Anode = more negative $E°$ (better reducing agent)
  • Cathode = more positive $E°$ (better oxidising agent)

5. Standard Cell Potential (E°cell)

$$ E°{cell} = E°{cathode} - E°_{anode} $$

  • Positive $E°_{cell}$ → spontaneous reaction
  • Negative $E°_{cell}$ → non-spontaneous reaction

Worked Example: Zinc Electrode Potential

Zinc has a negative electrode potential but can produce a positive cell potential: $$E_{cell} = E°{cathode} - E°{anode} = 0\text{ V} - (-0.76\text{ V}) = +0.76\text{ V}$$

This relates to two different ways of measuring electrochemical activity.

6. Worked Examples from Lecture

Example 1: Identifying Cathode and Anode

Given half-equations, identify which is cathode and which is anode by referring to standard reduction potentials.

Strategy:

  1. Look up $E°$ values for each half-reaction
  2. The half-reaction with more positive $E°$ is the cathode (reduction)
  3. The half-reaction with more negative $E°$ is the anode (oxidation)
  4. Calculate $E°{cell} = E°{cathode} - E°_{anode}$

Example 2: Cell Potential Calculation

Given $E°{\text{Mg}^{2+}/\text{Mg}} = -2.37$ V and $E°{\text{Ca}^{2+}/\text{Ca}} = -2.76$ V:

  • Mg²⁺/Mg is more positive → cathode
  • Ca²⁺/Ca is more negative → anode
  • $E°_{cell} = (-2.37) - (-2.76) = +0.39$ V

Key Concepts

  • Electrochemistry — Concept page
  • Galvanic Cell — Spontaneous electrochemical cell
  • Electrolytic Cell — Non-spontaneous electrochemical cell
  • Standard Reduction Potential — Half-cell potential
  • Redox Reactions — Oxidation-reduction fundamentals
  • Cell Notation — Standard cell diagram representation

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