What Do Oxidation and Reduction Mean?
Electron Transfer: The Core Idea
- Many reactions involve the transfer of electrons from one species to another.
- These are called redox reactions (short for reduction–oxidation).
- We use the acronym: OIL RIG
- Oxidation Is Loss (of electrons)
- Reduction Is Gain (of electrons)
- Which means:
- Oxidation = loss of electrons
- Reduction = gain of electrons
- These two processes always happen together: if one species loses electrons, another species must gain them.
Note
- There is no oxidation without reduction and no reduction without oxidation.
- Redox reactions always involve paired electron transfers.
Oxidation – Loss of Electrons
- When a particle (atom, ion or molecule) loses electrons, it is oxidised.
- We sometimes also describe oxidation as:
- gain of oxygen (e.g. iron forming iron oxide), or
- loss of hydrogen (e.g. hydrocarbons burning in oxygen),
- but the electron definition is the most general and works for all redox reactions.
Example
Aluminium metal forming aluminium ions
$$\mathrm{Al}(\mathrm{~s}) \rightarrow \mathrm{Al}^{3+}(\mathrm{aq})+3 e^{-}$$
- The aluminium atom loses 3 electrons.
- It goes from Al to Al³⁺.
- Therefore, aluminium is oxidised.
Reduction – Gain of Electrons
- When a particle gains electrons, it is reduced.
- Other older ways to describe reduction (often used in simple examples):
- loss of oxygen (e.g. copper(II) oxide → copper),
- gain of hydrogen (e.g. hydrogen added to an organic molecule).
- Again, the electron definition is the most important: Reduction = gain of electrons.
Example
Aluminium ion gaining electrons
$$\mathrm{Al}^{3+}(\mathrm{aq})+3 e^{-} \rightarrow \mathrm{Al}(\mathrm{~s})$$
- The aluminium ion gains 3 electrons.
- It goes from Al³⁺ to Al.
- Therefore, aluminium ion is reduced.
Oxidising Agents and Reducing Agents
In a redox reaction:
- The species that gets oxidised (loses electrons) is the reducing agent
– it reduces something else by giving it electrons. - The species that gets reduced (gains electrons) is the oxidising agent
– it oxidises something else by taking electrons.
Tip
- Oxidising agent – causes oxidation of another species, is reduced itself.
- Reducing agent – causes reduction of another species, is oxidised itself.
How Do We Identify What Is Oxidised and What Is Reduced?
There are two main methods you can use:
- Track electrons in half-equations (best when they’re given).
- Look at oxidation states (useful when equations are more complex).
Note
At MYP level, you can focus mainly on electron transfer and some simple oxidation state changes.
Using Half-Equations
A half-equation shows either oxidation or reduction on its own.
- If electrons appear on the right → oxidation (electrons lost).
- If electrons appear on the left → reduction (electrons gained).
Example
Magnesium and copper(II) ions
Overall reaction: $$\mathrm{Mg}(\mathrm{~s})+\mathrm{Cu}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Mg}^{2+}(\mathrm{aq})+\mathrm{Cu}(\mathrm{~s})$$
Half-equations:
- Oxidation (magnesium): $$\operatorname{Mg}(\mathrm{s}) \rightarrow \mathrm{Mg}^{2+}(\mathrm{aq})+2 e^{-}$$
- Reduction (copper(II) ion): $$\mathrm{Cu}^{2+}(\mathrm{aq})+2 e^{-} \rightarrow \mathrm{Cu}(\mathrm{~s})$$
Thus:
- Mg is oxidised (loses electrons) → Mg is the reducing agent.
- Cu²⁺ is reduced (gains electrons) → Cu²⁺ is the oxidising agent.
Using Oxidation States (Numbers)
Each element can be assigned an oxidation state (oxidation number).
- When the oxidation state increases → the element is oxidised.
- When the oxidation state decreases → the element is reduced.
Example
Rusting of iron
Simplified reaction:
$$4 \mathrm{Fe}(\mathrm{~s})+3 \mathrm{O}_2(\mathrm{~g}) \rightarrow 2 \mathrm{Fe}_2 \mathrm{O}_3(\mathrm{~s})$$
Oxidation states:
- Fe: 0 in Fe(s) → +3 in Fe₂O₃ → oxidation (0 to +3)
- O: 0 in O₂(g) → −2 in Fe₂O₃ → reduction (0 to −2)
Thus:
- Iron is oxidised (loses electrons).
- Oxygen is reduced (gains electrons).
Exam technique
- Assign oxidation states.
- Compare before and after.
- If the number goes up, it’s oxidation; if it goes down, it’s reduction.
Redox Reactions in Everyday Life
Redox reactions are everywhere: in rusting (introduced in the example earlier), batteries, bleaching, combustion, photosynthesis and even cellular respiration.
Redox Reactions in Batteries
Batteries are practical devices that convert chemical energy into electrical energy via redox reactions.
In a simple cell:
- The anode is where oxidation happens (electrons are released).
- The cathode is where reduction happens (electrons are accepted).
- Electrons flow through the external circuit, powering devices.
Example
Zn/MnO₂ Dry Cell (Alkaline/Leclanché-type idea)
In a typical zinc–manganese dioxide dry cell:
- Anode (negative electrode): zinc metal
- Zinc is oxidised: $$\mathrm{Zn}(\mathrm{~s}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+2 e^{-}$$
- Cathode (positive electrode): manganese dioxide (MnO₂) mixture
- MnO₂ is reduced (simplified): $$\mathrm{MnO}_2+\mathrm{H}^{+}+e^{-} \rightarrow \mathrm{MnO}(\mathrm{OH}) \quad \text { (simplified) }$$
Electrons move from Zn (anode) → through the circuit → to MnO₂ (cathode).
Fuel Cells (e.g. Hydrogen Fuel Cells)
- A fuel cell continuously converts energy from a fuel (such as hydrogen) into electricity through redox reactions.
- In a hydrogen fuel cell:
- Anode: $$\mathrm{H}_2 \rightarrow 2 \mathrm{H}^{+}+2 e^{-} \quad \text { (hydrogen is oxidised) }$$
- Cathode: $$\frac{1}{2} \mathrm{O}_2+2 \mathrm{H}^{+}+2 e^{-} \rightarrow \mathrm{H}_2 \mathrm{O} \text { (oxygen is reduced) }$$
- Overall reaction: $$\mathrm{H}_2+\frac{1}{2} \mathrm{O}_2 \rightarrow \mathrm{H}_2 \mathrm{O}$$
- Electrons released at the anode travel through the circuit (doing work, e.g. driving a motor).
- Fuel cells are important in discussions about clean energy and hydrogen-powered vehicles.
Photosynthesis
- Photosynthesis is a biological redox process that occurs in chloroplasts of plants.
- Overall equation: $$6 \mathrm{CO}_2+6 \mathrm{H}_2 \mathrm{O} \xrightarrow{\text { light }} \mathrm{C}_6 \mathrm{H}_{12} \mathrm{O}_6+6 \mathrm{O}_2$$
- Carbon dioxide is reduced to glucose.
- Water is oxidised to oxygen.
- Electrons move (via complex pathways) from water to carbon dioxide, storing energy in glucose.
Bleaching
- Many bleaching agents (e.g. sodium hypochlorite in household bleach, hydrogen peroxide) work by oxidizing coloured substances.
- The coloured molecules (often organic dyes) are oxidized, changing their structure so they no longer absorb visible light → they appear colourless.
- The bleaching agent itself is reduced in the process.
Example
Sodium hypochlorite (NaClO) in bleach acts as an oxidizing agent, removing colour from stains on fabrics or paper.
Other Everyday Redox Processes
- Combustion (burning fuels like methane, petrol) – fuel is oxidized, releasing energy.
- Respiration – glucose is oxidized in our cells to release energy, with oxygen being reduced.
- Metal refining, electroplating, and corrosion protection all rely on controlled redox reactions.
Self review
- What is the difference between oxidation and reduction in terms of electron transfer? How are they related in a redox reaction?
- For the reaction $$\mathrm{Zn}(\mathrm{~s})+\mathrm{Cu}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+\mathrm{Cu}(\mathrm{~s})$$
- Which species is oxidized and which is reduced?
- Which is the reducing agent and which is the oxidizing agent?
