Lecture Notes: Redox and Electrode Potentials for OCR-A Chemistry

Introduction

• Presenter: Chris Harris from Allery Chemistry
• Topic: Redox and Electrode Potentials, specifically for OCR-A Chemistry
• Allery Chemistry offers free YouTube videos for UK exam boards
• Important to practice exam technique in addition to content knowledge.
• Slides available for purchase for study support.

Redox Reactions

• Redox: Reduction and Oxidation processes involving electron transfer.
• Acronym: OIL RIG - Oxidation Is Loss, Reduction Is Gain of electrons.
• Example: Calcium oxidized in the presence of oxygen forming calcium ion (Ca²⁺), and oxygen is reduced forming oxide ion (O²⁻).
• Agents:
  • Reducing agents lose electrons and are oxidized.
  • Oxidizing agents gain electrons and are reduced.

Balancing Half Equations

• Half equations show separate reduction and oxidation processes.
• Balancing Rules:
  1. Write down species before and after the reaction.
  2. Balance all atoms except for oxygen and hydrogen.
  3. Balance oxygens with water (H₂O).
  4. Balance hydrogens with H⁺ ions.
  5. Balance charges with electrons.
• Example: Conversion of Fe²⁺ to Fe³⁺.

Combining Half Equations

• Combine half equations to form a full ionic equation.
• Ensure electrons are canceled out between equations.

Redox Titrations

• Acid-Base Titration versus Redox Titration:
  • Focus on electron transfer rather than proton transfer.
• Transition metals useful due to variable oxidation states.
• Use of color change without indicators due to transition metal properties.
• Example: Acidified KMnO₄ as oxidizing agent.

Titration Calculations

• Calculate concentration of a reducing/oxidizing agent using titration.
• Example Assessment: Manganate titration with Fe²⁺.
• Steps involve balancing equations, calculating moles, and determining concentrations.

Electrochemistry and Half Cells

• Half Cells: Part of an electrochemical cell, metal in its ion solution or inert platinum in mixed ion solutions.
• Electrochemical Cells: Consist of two half cells, a wire and a salt bridge.
• Electrons flow from more reactive to less reactive metal.
• Example: Zinc and copper half-cells, reactions and observations.

Standard Electrode Potentials

• Each half cell has an E⁰ value indicating electron donation ease.
• Calculated under standard conditions (298K, 100kPa, 1 mol/dm³ concentration).
• Standard Hydrogen Electrode (SHE): Reference for measuring electrode potentials.

Electrochemical Series

• Table of half-cell reactions with E⁰ values.
• Strong oxidizing agents have higher E⁰ values; reducing agents have lower.
• Used in calculations to predict reaction outcomes.

Calculations with Electrode Potentials

• Calculation Formula: E⁰(cell) = E⁰(reduced) - E⁰(oxidized)
• Example Problems: Determining cell potentials with given electrode potentials.

Cell Notation

• Simplified representation of electrochemical cells.
• Most negative half cell on the left; double line for salt bridge.

Feasibility Predictions

• Use E⁰ to predict if reactions proceed under standard conditions.
• Positive E⁰ indicates feasible reactions.
• Real-world conditions (temperature, concentration) can influence outcomes.

Applications: Cells and Batteries

• Electrochemical cells used in batteries (e.g., lithium-ion for mobile devices).
• Fuel cells (e.g., hydrogen-oxygen) offer efficient energy but require constant fuel supply.
• Pros and cons of fuel cells include efficiency, environmental impact, and safety concerns.

Conclusion: Redox and electrode potentials are critical in understanding chemical reactions and practical applications such as batteries and fuel cells. Thorough understanding and practice with examples and calculations are essential for mastery in OCR-A Chemistry.