Redox Reactions — Study Notes

1. Introduction

Many of the most important chemical reactions in nature and industry are oxidation-reduction (redox) reactions. Combustion, corrosion, photosynthesis, respiration, the reactions in batteries and cells, and numerous industrial processes such as the extraction of metals all involve the transfer of electrons. This chapter develops the systematic framework for understanding and balancing such reactions.

Redox chemistry is fundamentally about the transfer of electrons between species. An oxidising agent gains electrons and is itself reduced, while a reducing agent loses electrons and is itself oxidised. The oxidation number concept provides a bookkeeping tool that lets us track electrons in even the most complex reactions, including those that do not obviously involve electron transfer.

Beyond classification, this chapter teaches the tools used to balance redox equations in acidic and basic media, and introduces the quantitative relationships between redox reactions and electrochemistry. These ideas connect to electrolysis, cells and the extraction of metals, forming a bridge between chemistry and its applications in energy storage and metallurgy.

2. Classical and Electronic Concepts of Oxidation and Reduction

In the classical definition, oxidation is the addition of oxygen or removal of hydrogen, while reduction is the removal of oxygen or addition of hydrogen. For example, the conversion of copper to copper oxide is an oxidation, while the conversion of copper oxide back to copper is a reduction. Similarly, hydrogen sulphide to sulphur is an oxidation because hydrogen is removed.

The modern electronic concept defines oxidation as the loss of electrons and reduction as the gain of electrons. These always occur together: one species cannot be oxidised without another being reduced. A redox reaction is therefore the sum of an oxidation half-reaction and a reduction half-reaction.

An oxidising agent (oxidant) is a substance that accepts electrons from another species and is itself reduced in the process. A reducing agent (reductant) is a substance that donates electrons and is itself oxidised. For example, in the reaction between copper and silver ions, copper is the reducing agent and silver ions are the oxidising agent.

3. Oxidation Number and Oxidation State

The oxidation number is a formal charge assigned to an atom in a compound, which helps us keep track of electron transfer. Some elements can have more than one oxidation state, such as iron in +2 and +3 states and sulphur in many states from -2 to +6.

The rules for assigning oxidation numbers are essential. The oxidation number of an element in its elemental form is zero. For a monatomic ion, it equals the charge on the ion. In most compounds, hydrogen has +1 and oxygen has -2, except in peroxides where oxygen is -1 and in metal hydrides where hydrogen is -1. The sum of the oxidation numbers of all atoms in a neutral molecule is zero, and in a polyatomic ion it equals the charge on the ion.

Using these rules, the oxidation number of an atom can be determined even in complex compounds. For example, in K2Cr2O7, chromium has an oxidation number of +6, and in KMnO4, manganese is +7.

4. Types of Redox Reactions

Redox reactions can be classified into several types. Combination reactions involve two substances combining, such as the formation of hydrogen chloride from hydrogen and chlorine. Decomposition reactions are the reverse, such as the decomposition of hydrogen peroxide into water and oxygen.

Displacement reactions involve one element displacing another from a compound. In metal displacement, a more reactive metal displaces a less reactive one, and in non-metal displacement, a more reactive halogen displaces a less reactive one. The reactivity series determines which displacement reactions occur spontaneously.

Disproportionation reactions are a special type in which a single substance is simultaneously oxidised and reduced. For example, in the decomposition of hydrogen peroxide, oxygen in the peroxide is both oxidised to O2 and reduced to H2O. In these reactions, the same element appears in two different products with different oxidation states.

5. Balancing of Redox Reactions: The Oxidation Number Method

Redox equations can be balanced using the oxidation number method. The steps are to write the skeleton equation, identify the atoms whose oxidation numbers change, and balance the increase and decrease in oxidation number by adjusting coefficients so that the total increase equals the total decrease.

For example, to balance the reaction of potassium permanganate with iron(II) in acidic medium, we note that Mn changes from +7 to +2 (a decrease of 5) while Fe changes from +2 to +3 (an increase of 1). Therefore, one MnO4- ion requires five Fe2+ ions, and the coefficients are adjusted accordingly.

The oxidation number method is straightforward but must be followed carefully: identify the oxidised and reduced species, compute the changes in oxidation numbers, and equalise the total increase and decrease before balancing the remaining atoms and charges.

6. Balancing of Redox Reactions: The Half-Reaction Method

The half-reaction method (ion-electron method) is more powerful and widely used. It treats the oxidation and reduction as separate half-reactions, balances each, and then combines them. The steps are to write the two half-reactions, balance all atoms except hydrogen and oxygen, balance oxygen by adding water and hydrogen by adding H+, and finally balance charge by adding electrons.

For acidic medium, the steps above apply directly. For basic medium, the same steps are followed and then OH- ions are added to both sides to neutralise the H+ ions, converting the solution to basic conditions. Each half-reaction must be balanced independently before they are combined.

The number of electrons lost in oxidation must equal the number gained in reduction. Once the electrons are balanced, the two half-reactions are added together and the common species are cancelled. This method works systematically for the most complex redox reactions.

7. Redox Reactions and Electrode Processes

Redox reactions are the basis of electrochemistry. A galvanic cell converts chemical energy into electrical energy through spontaneous redox reactions. At the anode, oxidation occurs, and at the cathode, reduction occurs. The electrons flow through the external circuit from the anode to the cathode.

An electrolytic cell uses electrical energy to drive non-spontaneous redox reactions. The electrode at which oxidation occurs is always called the anode, and the electrode at which reduction occurs is the cathode. In electrolysis, anions migrate to the anode and cations to the cathode.

The redox potential of a couple measures its tendency to gain or lose electrons. The standard electrode potential is measured against the standard hydrogen electrode, which is assigned a potential of zero. A more positive reduction potential indicates a stronger oxidising agent, and a more negative one a stronger reducing agent. The cell potential is the difference between the cathode and anode potentials:

$$E_{\text{cell}} = E_{\text{cathode}} - E_{\text{anode}}$$

Quick Revision Tables

Table 1: Oxidation Number Rules

Table 2: Common Oxidising and Reducing Agents

Table 3: Common Oxidation States

Mind Map

Important Diagrams (SVG)

Diagram 1: Galvanic Cell

Diagram 2: Oxidation Number Calculation Examples

Common Mistakes

1. Confusing the oxidising and reducing agents; the oxidant is reduced (gains electrons) and the reductant is oxidised (loses electrons).
2. Forgetting that oxygen in peroxides has an oxidation number of -1, not -2.
3. Writing hydrogen as +1 in metal hydrides like NaH; hydrogen is -1 in such compounds.
4. Attempting to balance atoms before balancing electrons in the half-reaction method; the charge must be balanced with electrons first.
5. Forgetting that in disproportionation reactions the same species is both oxidised and reduced.
6. Failing to convert the acidic-medium balancing steps to basic medium by adding OH- to neutralise H+.
7. Using E cell = E anode - E cathode instead of E cell = E cathode - E anode; the standard cell potential always equals cathode minus anode potential.

Exam Tips

1. Memorise the oxidation number rules completely, especially the exceptions for peroxides (-1) and hydrides (-1 for H).
2. Practise computing oxidation numbers in compounds like K2Cr2O7, KMnO4, H2O2 and Fe3O4 before exams.
3. In the half-reaction method, remember the sequence: balance atoms (except H and O), then water for O, then H+ for H, then electrons for charge.
4. Know that acidic KMnO4 reduces Mn +7 to Mn +2 (pink to colourless), and acidic K2Cr2O7 reduces Cr +6 to Cr +3 (orange to green); these colour changes are classic exam clues.
5. Remember that the standard hydrogen electrode has a potential of zero by definition.
6. For cell potential calculations, always write E cell = E cathode - E anode, using the larger reduction potential as the cathode.
7. Recognise disproportionation examples like H2O2 -> H2O + O2 and 2Cu+ -> Cu + Cu2+.

Conclusion

Redox reactions unify a vast range of chemical phenomena under the simple idea of electron transfer. The oxidation number concept provides the bookkeeping tool to track electrons, and the oxidation number and half-reaction methods allow any redox equation to be balanced systematically, whether in acidic or basic medium. Redox chemistry extends into electrochemistry, where galvanic cells convert chemical energy into electricity and electrolytic cells do the reverse. The standard electrode potentials provide a quantitative measure of oxidising and reducing strength and predict the direction of spontaneous change. Mastery of redox concepts is essential for understanding metals extraction, corrosion, batteries, biological energy conversion and the many analytical applications that rely on oxidation-reduction chemistry.