Electrochemistry
Electrochemistry is the story of chemical reactions powered by electricity and electricity generated by chemical reactions.
Start with the simplest version: this lesson is about Electrochemistry. If you can explain the core idea to a friend using everyday language, examples, and one clear reason why it matters, you have moved from memorising to understanding.
Electrochemistry is the story of chemical reactions powered by electricity and electricity generated by chemical reactions. It's the science behind batteries, corrosion, electroplating, and fuel cells. When electrons move from one substance to another, they carry energy that can light bulbs, charge phones, or drive chemical transformations. This chapter explores how redox reactions occur in electrochemical cells and how we measure their spontaneity and power.
The Electrochemical Cell Story
Imagine two metals sitting in salt water, connected by a wire. Electrons begin flowing through the wire, creating an electric current. This is an electrochemical cell—a device where chemical reactions produce or consume electricity. There are two main types, like two sides of the same coin.
Galvanic (Voltaic) Cells: Chemistry Powers Electricity
A galvanic cell is a battery. A spontaneous redox reaction pushes electrons through a circuit, creating useful current. Your smartphone battery is a galvanic cell—inside, zinc and carbon electrodes sit in acidic paste, driving electrons toward your screen.
How It Works: The Two-Electrode Drama
The Anode (negative terminal): Where oxidation occurs—an element loses electrons and becomes oxidized. These electrons travel through the external circuit (powering devices) and eventually reach the other electrode.
The Cathode (positive terminal): Where reduction occurs—another element gains those electrons and becomes reduced. Simultaneously, ions move through an internal conductor (electrolyte) to complete the circuit and balance charge.
Salt Bridge: A pathway inside the cell allowing ions to flow between electrodes, maintaining electrical neutrality as one side accumulates positive ions and the other accumulates negative ions.
Example: In a simple Zn-Cu cell:
- Anode: Zn → Zn²⁺ + 2e⁻ (oxidation, zinc corrodes)
- Cathode: Cu²⁺ + 2e⁻ → Cu (reduction, copper deposits)
- Result: Electrons flow from zinc to copper through the external circuit, creating current.
Electrolytic Cells: Electricity Powers Chemistry
An electrolytic cell is the opposite. Instead of spontaneous reactions generating electricity, we apply external electrical voltage to force non-spontaneous reactions. This is how aluminum is extracted from ore, how silver is purified, and how electrolysis of water produces hydrogen and oxygen gases.
Key Differences from Galvanic Cells
Polarity reverses: The positive electrode is now the anode (oxidation site), the negative electrode is the cathode (reduction site).
External power required: We must apply voltage greater than the cell's opposing EMF to make the reaction go.
Non-spontaneous reactions occur: We can drive reactions "backward" that wouldn't happen naturally.
Example: Electroplating uses this principle. A negative voltage on a jewelry item (cathode) forces metal ions from solution to deposit as a protective coat, while a positive electrode (anode) of pure metal dissolves and replenishes the ions.
Cell Potential and EMF: The Voltage of Chemistry
The cell potential (E) measures the driving force of a redox reaction. Standard potentials—measured under standard conditions (1 M concentration, 1 atm pressure, 25°C)—tell us which element "wants" to be oxidized or reduced.
Standard Reduction Potentials are measured relative to the hydrogen electrode (set at 0 V as reference). Metals like lithium (high negative potential) are easily oxidized; halogens like fluorine (high positive potential) are easily reduced.
Calculating Cell EMF
E°cell = E°cathode - E°anode
If E°cell is positive, the reaction is spontaneous. If negative, it's non-spontaneous without external power.
Example: A Zn-Cu cell has E°Zn = -0.76 V and E°Cu = +0.34 V, so E°cell = 0.34 - (-0.76) = +1.10 V. This positive value confirms zinc and copper naturally form a galvanic cell.
The Nernst Equation: Real-World Adjustments
The relationship between cell potential and Gibbs free energy is ΔG° = -nFE°cell, where n is electrons transferred and F is Faraday's constant (96,500 C/mol).
But real conditions rarely match standard conditions. The Nernst equation adjusts for non-standard concentrations:
E = E° - (RT/nF) ln Q
At 25°C, this simplifies to: E = E° - (0.059/n) log Q
where Q is the reaction quotient. As a reaction proceeds and Q increases, the cell potential decreases until equilibrium (E = 0).
Conductivity and Ionic Solutions
Conductivity measures how easily electricity flows through a solution. It depends on:
Concentration: More ions mean more charge carriers, so conductivity increases initially.
Temperature: Higher temperature increases ion mobility.
Nature of ions: Small, highly charged ions conduct better than large, singly charged ones.
Molar conductivity (Λm) accounts for concentration differences: the conductivity divided by molar concentration. As concentration increases, molar conductivity decreases because ions shield each other's electric fields. For strong electrolytes at infinite dilution, molar conductivity reaches a maximum, and we can add individual ionic conductivities to predict total conductivity.
Applications: The Electrochemical Revolution
Batteries and Fuel Cells: Galvanic cells power everything from watches to electric vehicles. Hydrogen fuel cells generate electricity with only water as exhaust—representing clean energy's future.
Electroplating: Creating protective coatings of chrome, gold, or silver on objects, or refining metals to extreme purity (like copper wires for electronics).
Corrosion and Protection: Preventing steel from rusting using sacrificial anodes (like zinc blocks on ships) that oxidize preferentially.
Electrolysis: Producing chemicals like chlorine gas and sodium hydroxide from salt water; obtaining metals like aluminum and magnesium from ores.
Socratic Questions
- Why does a galvanic cell work as long as there's a salt bridge, but stops generating current if the salt bridge is removed?
- In an electrolytic cell, why do we apply an external voltage greater than the theoretical cell potential, and what happens to that "excess" energy?
- If you place a zinc strip and a copper strip in saltwater without connecting them with a wire, will a reaction occur? What changes if you connect them?
- A battery's voltage drops as it's used. According to the Nernst equation, why does this happen in terms of concentration changes?
- How could you design an experiment to prove that cell potential depends on ionic concentrations, not just on which metals are used?
Related Topics
Solutions - Ionic concentrations and colligative properties Chemical Kinetics - Reaction rates and activation energy Redox Reactions - Electron transfer mechanisms