Redox Reactions (OCR A-Level Chemistry A): Revision Notes

Redox Reactions

Introduction to redox reactions

Redox reactions are fundamental chemical processes involving the transfer of electrons between chemical species. The term "redox" combines two complementary processes: reduction and oxidation, which always occur together in these reactions. Understanding redox chemistry is essential for explaining a wide range of chemical phenomena, from simple displacement reactions to complex electrochemical cells.

Defining reduction and oxidation

Electron transfer definitions

When we consider electron movement, redox processes can be defined as follows:

• Reduction is the gain of electrons by a chemical species
• Oxidation is the loss of electrons by a chemical species

A helpful mnemonic to remember this is OILRIG:

• Oxidation Is Loss (of electrons)
• Reduction Is Gain (of electrons)

Oxidation number definitions

Alternatively, redox processes can be described in terms of oxidation number changes:

• Reduction is a decrease in oxidation number
• Oxidation is an increase in oxidation number

Rules for assigning oxidation numbers

To use the oxidation number method effectively, you need to know the standard rules for assigning oxidation numbers to atoms in different chemical environments. These rules follow a specific hierarchy:

Oxidising agents and reducing agents

In every redox reaction, there must be both an oxidising agent and a reducing agent. These terms describe the roles that different species play in facilitating electron transfer.

The oxidising agent

An oxidising agent is a substance that causes oxidation in another species. To do this, the oxidising agent must accept electrons from the species being oxidised. In accepting these electrons, the oxidising agent itself undergoes reduction.

The reducing agent

A reducing agent is a substance that causes reduction in another species. To do this, the reducing agent must donate electrons to the species being reduced. In donating these electrons, the reducing agent itself undergoes oxidation.

Worked example: identifying oxidising and reducing agents

This reaction is visible as the copper wire becomes coated with silver metal deposits, while the solution develops a blue colour from $\text{Cu}^{2+}$ ions.

Constructing redox equations from half-equations

Half-equations are a powerful tool for representing the separate reduction and oxidation processes that occur in a redox reaction. By combining appropriately balanced half-equations, you can construct the overall redox equation.

The balancing procedure

The key principle when combining half-equations is ensuring that the number of electrons lost during oxidation exactly equals the number of electrons gained during reduction. This maintains charge balance in the overall equation.

Step 1: Balance the electrons

Examine both half-equations and identify how many electrons are involved in each. Find the lowest common multiple of these numbers, then multiply each half-equation by the appropriate factor to equalise electron numbers.

Step 2: Add and cancel electrons

Add the two balanced half-equations together. The electrons should appear equally on both sides, allowing you to cancel them completely.

Step 3: Cancel common species

Check if any species appear on both sides of the equation. If so, cancel the appropriate number to simplify the equation.

Worked example: combining half-equations

Using oxidation numbers to write equations

An alternative method for constructing redox equations uses oxidation number changes directly. This approach is particularly useful when half-equations aren't provided, or when dealing with complex reactions involving multiple elements.

The oxidation number balancing procedure

The fundamental principle is that the total increase in oxidation number during oxidation must equal the total decrease in oxidation number during reduction.

Step 1: Write the unbalanced equation

Start by writing out all the reactants and products, even if you're unsure of some coefficients.

Step 2: Assign oxidation numbers

Determine the oxidation number of each element in all species. Identify which elements undergo oxidation number changes.

Step 3: Balance the species undergoing redox

Only balance the species containing elements that change oxidation number. Use the oxidation number changes to determine the correct stoichiometric ratios.

Step 4: Balance remaining atoms

Finally, balance any atoms that didn't change oxidation number (often hydrogen and oxygen).

Worked example: balancing using oxidation numbers

Writing half-equations

Sometimes you'll need to construct half-equations rather than just combine them. This skill is essential for understanding electrode processes and electrochemical cells.

Procedure for writing half-equations

Step 1: Assign oxidation numbers

Determine the oxidation numbers of the species being reduced or oxidised, and calculate the change.

Step 2: Balance the electrons

Add electrons to the appropriate side of the equation. Remember:

• For reduction (decrease in oxidation number), electrons go on the left (reactant side)
• For oxidation (increase in oxidation number), electrons go on the right (product side)

Step 3: Balance remaining atoms

Balance any other atoms present, then predict any additional species that must be formed.

Worked example: writing a half-equation

Predicting products of redox reactions

In exam questions and practical situations, you may not be given all the species involved in a redox reaction. You'll need to predict missing reactants or products based on your understanding of redox chemistry.

Key considerations when predicting products

1. Water is commonly formed

In aqueous redox reactions, water ($\text{H}_2\text{O}$) is frequently produced, particularly when reactions involve oxygen-containing species.

2. Hydrogen ions or hydroxide ions

Depending on whether the reaction occurs in acidic or alkaline conditions:

• Acidic conditions: $\text{H}^+$ ions are likely to be involved
• Alkaline conditions: $\text{OH}^-$ ions are likely to be involved

3. Charge must balance

As a final check, ensure that the total charge on the left side of your equation equals the total charge on the right side. This is a fundamental requirement for all balanced chemical equations.

4. Use oxidation numbers to predict half-equations

If you need to predict a missing half-equation, use oxidation numbers to identify what species makes chemical sense. The change in oxidation number will tell you how many electrons must be involved.

Example of predicting products

Suppose you're told that hydrogen iodide reacts with concentrated sulfuric acid, and you need to predict all products. You might know that iodide can be oxidised to iodine, and that sulfuric acid can be reduced. Using oxidation numbers:

• Iodine in HI: -1 → 0 (in $\text{I}_2$), so oxidation occurs
• Sulfur in $\text{H}_2\text{SO}_4$: +6 → ? (reduction must occur)

Common reduction products of sulfuric acid include $\text{H}_2\text{S}$ (where S is -2), showing sulfur has been reduced. Water would also be formed. This systematic thinking allows you to predict products logically rather than guessing.

Remember!