Lecture Notes on Redox Equilibria and Electrochemical Cells

Electrochemical Cells

• Consist of two half cells, each with a metal electrode and a solution of a compound containing that metal (e.g., Cu and CuSO₄).
• A salt bridge, usually potassium nitrate, connects the two half cells, allowing charge conduction by free-moving ions.
• A wire isn't used as it would create its own electrode system.
• Voltage is generated due to different tendencies of metals to oxidize or reduce, measured using a high-resistance voltmeter.

Function of a Voltmeter

• High resistance prevents current flow, allowing measurement of maximum potential difference (E).
• If current flows, reactions occur at each electrode until reactants are depleted and voltage drops to zero.

Cell Diagrams

• Represented as: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s).
• Solid vertical line indicates phase boundary (solid and electrolyte), and double line represents the salt bridge.
• More positive half cell written on the right, if possible.

Systems Without Metals

• Use a platinum electrode if no metal acts as an electrode.
• Platinum conducts and is unreactive (e.g., Fe²⁺/Fe³⁺ system).

Measuring Electrode Potential

• Potential difference measured between two connected half cells.
• Standard Hydrogen Electrode (SHE) used as a reference with 0 Volts.

Standard Conditions

• Hydrogen electrode conditions: 100 kPa H₂ gas, 1 mol dm⁻³ H⁺ ions, 298K.
• Other standard electrodes calibrated against SHE (secondary standards).

Calculating EMF (Electromotive Force)

• Ecell calculated using standard electrode potentials: Ecell = Erhs - Elhs.
• Positive Ecell indicates spontaneous reaction.

Using Electrode Potentials

• Predict direction of spontaneous redox reactions.
• More negative half cell oxidizes; more positive reduces.

Effect of Conditions on Ecell

• Ecell is affected by concentration, temperature, and reaction conditions (Le Chatelier's principle).

Types of Cells

• Electrochemical cells used as energy sources can be non-rechargeable, rechargeable, or fuel cells.

Non-Rechargeable Cells

• Primary cells like Dry Cells: irreversible reactions.

Rechargeable Cells

• Include lead-acid, nickel-cadmium, and lithium-ion cells used in electronics.
• Reactions reversible upon recharging.

Fuel Cells

• Use fuel (e.g., hydrogen) with oxygen to produce voltage.
• Maintain constant voltage by continuous supply of reactants.
• Advantages: less pollution, higher efficiency.
• Limitations: cost, storage and transport of hydrogen, limited lifetime.

Ethanol Fuel Cells

• Renewable and less risky compared to hydrogen, produce fewer pollutants.
• Ethanol can be produced abundantly and in a carbon-neutral way.