Redox- CIE iGCSE Chemistry Notes - New Syllabus
Redox for iGCSE Chemistry Notes
Core Syllabus
- Use a Roman numeral to indicate the oxidation number of an element in a compound
- Define redox reactions as involving simultaneous oxidation and reduction
- Define oxidation as gain of oxygen and reduction as loss of oxygen
- Identify redox reactions as reactions involving gain and loss of oxygen
- Identify oxidation and reduction in redox reactions
Supplement Syllabus
- Define oxidation in terms of:
(a) loss of electrons
(b) an increase in oxidation number - Define reduction in terms of:
(a) gain of electrons
(b) a decrease in oxidation number - Identify redox reactions as reactions involving gain and loss of electrons
- Identify redox reactions by changes in oxidation number using:
(a) the oxidation number of elements in their uncombined state is zero
(b) the oxidation number of a monatomic ion is the same as the charge on the ion
(c) the sum of the oxidation numbers in a compound is zero
(d) the sum of the oxidation numbers in an ion is equal to the charge on the ion - Identify redox reactions by the colour changes involved when using acidified aqueous potassium manganate(VII) or aqueous potassium iodide
- Define an oxidising agent as a substance that oxidises another substance and is itself reduced
- Define a reducing agent as a substance that reduces another substance and is itself oxidised
- Identify oxidising agents and reducing agents in redox reactions
Redox Reactions
Roman numeral to indicate the oxidation number
In chemistry, certain elements—especially transition metals—can form compounds in which the same element exists in different oxidation states (also called oxidation numbers). To clearly identify which oxidation state is present in a compound, chemists use Roman numerals in the name of the compound.
Oxidation number (oxidation state) is a number assigned to an atom in a compound or ion that reflects the number of electrons lost or gained compared to the neutral atom.
Why are Roman numerals used?
- Many elements, particularly transition metals, can form more than one type of ion with different charges (e.g. \( \text{Fe}^{2+} \) and \( \text{Fe}^{3+} \)).
- The Roman numeral in a compound name tells you the oxidation number of that specific element in that compound.
- This system helps avoid confusion and allows chemists to correctly identify the compound’s structure and chemical behavior.
Rules for using Roman numerals in compound names:
- The Roman numeral corresponds to the positive oxidation number of the metal.
- It is placed in brackets immediately after the metal’s name in the compound.
- The Roman numeral is only used when the element has more than one possible oxidation number.
- Nonmetals do not usually need Roman numerals because they usually have fixed oxidation numbers in compounds (e.g. oxygen is almost always -2).
Common oxidation numbers and corresponding Roman numerals:
| Ion / Element | Oxidation Number | Roman Numeral | Example Compound | Chemical Formula |
|---|---|---|---|---|
| Iron(II) | +2 | II | Iron(II) chloride | \( \text{FeCl}_2 \) |
| Iron(III) | +3 | III | Iron(III) chloride | \( \text{FeCl}_3 \) |
| Copper(I) | +1 | I | Copper(I) oxide | \( \text{Cu}_2\text{O} \) |
| Copper(II) | +2 | II | Copper(II) sulfate | \( \text{CuSO}_4 \) |
| Lead(IV) | +4 | IV | Lead(IV) oxide | \( \text{PbO}_2 \) |
Important Note:
The Roman numeral refers to the oxidation number, not to the number of atoms. For example, in \( \text{Fe}_2\text{O}_3 \), the name is iron(III) oxide because each iron atom has a +3 charge, even though there are two iron atoms.
Example
What is the correct name for the compound with formula \( \text{MnO}_2 \)?
▶️Answer/Explanation
Oxygen has a fixed oxidation number of -2. Since there are two oxygen atoms, the total negative charge is -4.
Therefore, manganese must have an oxidation number of +4 to balance this.
So, the compound is called manganese(IV) oxide.
Example
What is the oxidation number of iron in the compound iron(II) sulfate, \( \text{FeSO}_4 \)?
▶️Answer/Explanation
The Roman numeral (II) in the name tells us that iron has an oxidation number of +2.
Sulfate (\( \text{SO}_4^{2-} \)) is a polyatomic ion with a charge of -2, so iron must be +2 to balance the compound.
Hence, oxidation number of iron = +2.
Example
Why is the compound \( \text{Cu}_2\text{O} \) called copper(I) oxide?
▶️Answer/Explanation
Each oxygen has an oxidation number of -2. Since there is one oxygen atom, total negative charge = -2.
There are two copper atoms, so the total positive charge must be +2, meaning each copper atom has +1.
Therefore, this is copper(I) oxide.
Redox Reactions
Redox reactions are one of the most important types of chemical reactions in chemistry. The term redox comes from the two processes that happen together during such a reaction: reduction and oxidation.
A redox reaction is a chemical reaction in which oxidation and reduction occur at the same time. It can be defined in three ways:
| Type | Oxidation | Reduction |
|---|---|---|
| Oxygen transfer | Gain of oxygen | Loss of oxygen |
| Electron transfer | Loss of electrons | Gain of electrons |
| Oxidation number | Increase in oxidation number | Decrease in oxidation number |
Oxidation and Reduction
Oxidation is the gain of oxygen by a substance during a chemical reaction.
Reduction is the loss of oxygen from a substance during a chemical reaction.
This definition applies particularly well to reactions involving metal oxides and gases like hydrogen or carbon monoxide.
Key points:
- If a substance gains oxygen, it is said to be oxidised
- If a substance loses oxygen, it is said to be reduced
- These changes happen together in redox reactions
In such reactions, there is always a substance that donates oxygen (the oxidising agent) and one that receives it (the substance being oxidised).
Therefore, a redox reaction involves one substance being oxidised (gaining oxygen) and another being reduced (losing oxygen).
How to identify oxidation and reduction:
- Compare the reactants and products
- Check if a substance has gained or lost oxygen
- The substance that gains oxygen → oxidised
- The substance that loses oxygen → reduced
- If both oxidation and reduction occur in the same reaction, it is a redox reaction
Note: You may also observe a change in colour or the formation of a new substance (like a metal or gas) as clues for oxidation or reduction.
Example
Is the following a redox reaction? Identify what is oxidised and what is reduced:
\( \text{CuO}_{(s)} + \text{H}_2{}_{(g)} \rightarrow \text{Cu}_{(s)} + \text{H}_2\text{O}_{(l)} \)
▶️Answer/Explanation
Copper(II) oxide loses oxygen and becomes copper → reduction.
Hydrogen gains oxygen and becomes water → oxidation.
So, CuO is reduced and H2 is oxidised.
Both oxidation and reduction occur → this is a redox reaction.
Example
Identify the oxidised and reduced substances in:
\( \text{Zn}_{(s)} + \text{CuO}_{(s)} \rightarrow \text{ZnO}_{(s)} + \text{Cu}_{(s)} \)
▶️Answer/Explanation
Zn → ZnO: Zinc gains oxygen → oxidised.
CuO → Cu: Copper(II) oxide loses oxygen → reduced.
So, zinc is oxidised and copper(II) oxide is reduced.
Example
State what is being oxidised and what is being reduced in the following reaction:
\( 2\text{Mg}_{(s)} + \text{O}_2{}_{(g)} \rightarrow 2\text{MgO}_{(s)} \)
▶️Answer/Explanation
Oxidation: Magnesium (Mg) is oxidised.
In the reactants, magnesium is in its elemental form (oxidation number = 0). In the product \( \text{MgO} \), magnesium has an oxidation number of +2.
So, magnesium loses electrons (oxidation is loss of electrons).
Reduction: Oxygen (O2) is reduced.
In the reactants, oxygen is in its elemental form (oxidation number = 0). In the product \( \text{MgO} \), each oxygen atom has an oxidation number of -2.
So, oxygen gains electrons (reduction is gain of electrons).
Redox Reactions involving gain and loss of electrons
Oxidation (Electron-Based):
Oxidation is the loss of electrons from an atom, ion, or molecule during a chemical reaction.
When a substance loses electrons, it becomes more positively charged (its oxidation state increases).
Oxidation Number:
- Oxidation is also defined as an increase in oxidation number.
- The oxidation number is a value that represents the effective charge of an atom in a compound.
- When an atom’s oxidation number increases, it means it has lost electrons.
Connection Between the Two Definitions:
- Loss of electrons → increase in oxidation number → oxidation
Visualizing Oxidation:
- Electron loss happens from the outer shell of atoms.
- This is common with metals, which tend to form positive ions by losing electrons.
Common Examples of Oxidation:
- Magnesium burning in air: \( \text{Mg} \rightarrow \text{Mg}^{2+} + 2e^- \)
- Iron rusting: \( \text{Fe} \rightarrow \text{Fe}^{3+} + 3e^- \)
- Hydrogen gas turning into H⁺ ions: \( \text{H}_2 \rightarrow 2\text{H}^+ + 2e^- \)
Example
Magnesium reacts with oxygen to form magnesium oxide:
\( 2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO} \)
Identify what is being oxidised and how.
▶️Answer/Explanation
Magnesium atoms lose two electrons each → \( \text{Mg} \rightarrow \text{Mg}^{2+} + 2e^- \)
This is oxidation (loss of electrons and increase in oxidation number from 0 to +2).
Example
In the reaction \( \text{Fe}^{2+} \rightarrow \text{Fe}^{3+} + e^- \), identify what is happening to iron.
▶️Answer/Explanation
Iron(II) ion loses one electron and becomes Iron(III) ion.
This is oxidation (loss of electrons and increase in oxidation number from +2 to +3).
Reduction (Electron-Based):
Reduction is the gain of electrons by an atom, ion, or molecule during a chemical reaction.
This causes the atom to become more negatively charged (or less positive).
Reduction (Oxidation Number):
- Reduction is also defined as a decrease in oxidation number.
- When an atom’s oxidation number decreases, it means it has gained electrons.
Connection Between the Two Definitions:
- Gain of electrons → decrease in oxidation number → reduction
Visualizing Reduction:
- Non-metals commonly gain electrons during reactions to form negative ions.
- Reduction is commonly observed during displacement or extraction reactions.
Common Examples of Reduction:
- Oxygen being removed from copper(II) oxide: \( \text{CuO} + \text{H}_2 \rightarrow \text{Cu} + \text{H}_2\text{O} \)
- Chlorine gaining electrons: \( \text{Cl}_2 + 2e^- \rightarrow 2\text{Cl}^- \)
- Iron(III) ion gaining electrons: \( \text{Fe}^{3+} + e^- \rightarrow \text{Fe}^{2+} \)
Example
Chlorine gas reacts to form chloride ions:
\( \text{Cl}_2 + 2e^- \rightarrow 2\text{Cl}^- \)
Identify what is being reduced and how.
▶️Answer/Explanation
Each chlorine atom gains one electron to form chloride ion.
This is reduction (gain of electrons and oxidation number decreases from 0 to -1).
Example
In the reaction: \( \text{Fe}^{3+} + e^- \rightarrow \text{Fe}^{2+} \), determine what is happening to the iron ion.
▶️Answer/Explanation
The iron(III) ion gains one electron and becomes iron(II) ion.
This is reduction (gain of electrons and oxidation number decreases from +3 to +2).
Redox Reactions (Electron Transfer):
A redox (reduction–oxidation) reaction is one in which:
- Oxidation is the loss of electrons
- Reduction is the gain of electrons
In every redox reaction, electrons are transferred from one substance to another.
Key Features of Redox Reactions:
- One substance loses electrons (gets oxidised)
- Another substance gains those electrons (gets reduced)
- These processes always happen together in the same reaction
Recognizing Redox Reactions:
- Look for electron transfer (half-equations help)
- Observe changes in oxidation number (if available)
- Follow color changes or gas evolution in some redox reactions
Common Situations Where Redox Reactions Occur:
- Metal + acid → salt + hydrogen gas
- Displacement reactions
- Electrolysis (both oxidation and reduction occur at electrodes)
Example
Zinc reacts with dilute hydrochloric acid:
\( \text{Zn} + 2\text{H}^+ \rightarrow \text{Zn}^{2+} + \text{H}_2 \)
Identify the oxidation and reduction processes.
▶️Answer/Explanation
Zinc loses 2 electrons to form \( \text{Zn}^{2+} \) → oxidation
Hydrogen ions gain 2 electrons to form \( \text{H}_2 \) → reduction
This is a redox reaction: zinc is oxidised and hydrogen is reduced.
Example
Iron reacts with copper(II) sulfate solution:
\( \text{Fe} + \text{Cu}^{2+} \rightarrow \text{Fe}^{2+} + \text{Cu} \)
Is this a redox reaction?
▶️Answer/Explanation
Yes. Iron loses electrons → oxidised to \( \text{Fe}^{2+} \)
Copper(II) ions gain electrons → reduced to copper metal
Since both oxidation and reduction occur, it is a redox reaction.
Redox Reactions involving changes in Oxidation Number and Colour
What is Oxidation Number?
The oxidation number (or oxidation state) of an atom in a compound is a number that represents the charge the atom would have if electrons were transferred completely, rather than shared.
General Rules for Assigning Oxidation Numbers:
- Uncombined elements (like \( \text{Na} \), \( \text{Cl}_2 \), \( \text{O}_2 \)) have an oxidation number of 0.
- A monatomic ion (like \( \text{Fe}^{3+} \) or \( \text{Cl}^- \)) has an oxidation number equal to its charge.
- The sum of the oxidation numbers in a compound must equal 0.
- The sum of the oxidation numbers in a polyatomic ion equals the charge of the ion.
Typical Oxidation Numbers for Common Elements:
- Group 1 metals: +1
- Group 2 metals: +2
- Hydrogen: +1 (except in metal hydrides where it is -1)
- Oxygen: -2 (except in peroxides where it is -1)
- Halogens: usually -1 unless combined with oxygen or more electronegative elements
How This Helps in Identifying Redox Reactions:
A redox reaction involves a change in oxidation numbers:
- If an element’s oxidation number increases → oxidation
- If an element’s oxidation number decreases → reduction
Example
Identify the oxidised and reduced species in the following reaction using oxidation numbers:
\( \text{Mg} + \text{Cl}_2 \rightarrow \text{MgCl}_2 \)
▶️Answer/Explanation
Mg (uncombined) = 0
Cl2 (uncombined) = 0
In \( \text{MgCl}_2 \):
Mg = +2, each Cl = -1
So, Mg is oxidised (0 → +2) and Cl is reduced (0 → -1)
Therefore, this is a redox reaction.
Example
Determine if the reaction below is redox:
\( \text{H}_2\text{O}_2 \rightarrow \text{H}_2\text{O} + \text{O}_2 \)
▶️Answer/Explanation
In \( \text{H}_2\text{O}_2 \): O has an oxidation number of -1
In \( \text{H}_2\text{O} \): O is -2
In \( \text{O}_2 \): O is 0
One oxygen atom is reduced (-1 → -2), the other is oxidised (-1 → 0)
Since both oxidation and reduction occur, this is a redox reaction.
Redox Reactions involving Colour Changes
1. Acidified Potassium Manganate(VII), \( \text{KMnO}_4 \)
It acts as a strong oxidising agent. When added to a reducing agent in acidic solution, it undergoes the following change:
\( \text{MnO}_4^- \) (purple) → \( \text{Mn}^{2+} \) (colourless or very pale pink)
This colour change is used to detect the presence of a reducing agent. The disappearance of the purple colour indicates that reduction has occurred.
2. Potassium Iodide, \( \text{KI} \)
Potassium iodide is a reducing agent. When it reacts with an oxidising agent, iodide ions are oxidised to iodine:
\( 2\text{I}^- \rightarrow \text{I}_2 \)
The colour change observed is:
\( \text{I}^- \) (colourless) → \( \text{I}_2 \) (brown solution)
This brown colour indicates that oxidation of iodide has taken place and thus a redox reaction has occurred.
Why Colour Changes Are Useful in Redox Identification:
Certain reagents undergo visible colour changes when they are oxidised or reduced. These changes provide direct evidence that electrons have been transferred, confirming a redox reaction.
Example
What colour change would you observe when acidified potassium manganate(VII) is added to an iron(II) sulfate solution?
▶️Answer/Explanation
The \( \text{Fe}^{2+} \) ions are oxidised to \( \text{Fe}^{3+} \) and the purple \( \text{MnO}_4^- \) ions are reduced to pale pink \( \text{Mn}^{2+} \) ions.
Observation: The purple colour of the manganate(VII) disappears.
This confirms a redox reaction.
Example
What happens when potassium iodide solution is added to chlorine water?
▶️Answer/Explanation
Chlorine is a stronger oxidising agent than iodine. It oxidises iodide ions to iodine:
\( \text{Cl}_2 + 2\text{I}^- \rightarrow 2\text{Cl}^- + \text{I}_2 \)
Observation: The solution turns brown due to the formation of iodine.
This indicates that a redox reaction has taken place.
Oxidising Agents and Reducing Agents
Oxidising agent
An oxidising agent is a substance that causes another substance to undergo oxidation (lose electrons) and in the process, it itself is reduced (gains electrons).
Key ideas:
- The oxidising agent removes electrons from another substance.
- It accepts those electrons, so its own oxidation number decreases (reduction).
General process:
Oxidising agent: gains electrons → reduced
Substance oxidised: loses electrons → oxidation
Common oxidising agents:
- Acidified potassium manganate(VII) \( (\text{KMnO}_4) \)
- Acidified potassium dichromate(VI) \( (\text{K}_2\text{Cr}_2\text{O}_7) \)
- Chlorine \( (\text{Cl}_2) \)
- Hydrogen peroxide \( (\text{H}_2\text{O}_2) \)
Example
In acidic solution:
\( \text{MnO}_4^- + 8\text{H}^+ + 5\text{Fe}^{2+} \rightarrow \text{Mn}^{2+} + 5\text{Fe}^{3+} + 4\text{H}_2\text{O} \)
Which substance is the oxidising agent?
▶️Answer/Explanation
\( \text{Fe}^{2+} \) loses electrons → oxidised to \( \text{Fe}^{3+} \)
\( \text{MnO}_4^- \) gains electrons → reduced to \( \text{Mn}^{2+} \)
So \( \text{MnO}_4^- \) is the oxidising agent as it causes the oxidation of iron(II) ions and is reduced itself.
Reducing agent
A reducing agent is a substance that causes another substance to be reduced (gain electrons) and in the process, it itself is oxidised (loses electrons).
Key ideas:
- The reducing agent donates electrons to another substance.
- As it gives away electrons, its own oxidation number increases (oxidation).
General process:
Reducing agent: loses electrons → oxidised
Substance reduced: gains electrons → reduction
Common reducing agents:
- Carbon (C)
- Hydrogen gas \( (\text{H}_2) \)
- Metals like zinc and iron
- Sulfur dioxide \( (\text{SO}_2) \)
Identifying oxidising agents and reducing agents in redox reactions
- Look for the species that gains electrons → it is the oxidising agent.
- Look for the species that loses electrons → it is the reducing agent.
Clues to identify:
- Oxidation numbers: If a substance’s oxidation number decreases, it is reduced → it’s an oxidising agent.
- Electron transfer: Electron gain means reduction, and the substance that enables this is the oxidising agent.
- Color change: In some redox reactions, oxidising agents cause visible color changes (e.g. purple \( \text{MnO}_4^- \) to colorless).
Example
In the reaction:
\( \text{CuO} + \text{H}_2 \rightarrow \text{Cu} + \text{H}_2\text{O} \)
Identify the reducing agent.
▶️Answer/Explanation
Copper(II) oxide is reduced to copper.
Hydrogen is oxidised to water.
So, hydrogen is the reducing agent because it donates electrons to copper(II) oxide and is itself oxidised.
Example
Identify the oxidising and reducing agents in the reaction:
\( \text{Zn} + \text{Cu}^{2+} \rightarrow \text{Zn}^{2+} + \text{Cu} \)
▶️Answer/Explanation
Zinc loses electrons and is oxidised → it is the reducing agent.
Copper(II) ions gain electrons and are reduced → they are the oxidising agent.
Example
Identify the oxidising and reducing agents in:
\( \text{H}_2 + \text{Cl}_2 \rightarrow 2\text{HCl} \)
▶️Answer/Explanation
Hydrogen is oxidised (loses electrons) → it is the reducing agent.
Chlorine is reduced (gains electrons) → it is the oxidising agent.