Grade 11 ↓
Redox reactions
Redox reactions, also called reduction-oxidation reactions, are a family of reactions that involve the transfer of electrons between chemicals. These reactions are central to many biological and industrial processes, so understanding them is an essential part of chemistry.
What are redox reactions?
Redox reactions involve two major processes: oxidation and reduction.
- Oxidation is the process in which a substance loses electrons.
- Reduction is the process in which a substance gains electrons.
A mnemonic to help remember this is OIL RIG which stands for:
- Oxidation Is Loose
- Deduction is a benefit
Basic example of a redox reaction
A simple example of a redox reaction is the reaction between hydrogen and oxygen to form water:
2 H2 + O2 → 2 H2O
In this response:
- Hydrogen (
H2
) is oxidized as it loses electrons. The oxidation state of hydrogen increases from 0 to +1. - Oxygen (
O2
) is reduced as it gains electrons. The oxidation state of oxygen decreases from 0 to -2.
The diagram above shows hydrogen and oxygen forming water through a redox reaction. Hydrogen atoms transfer electrons to oxygen atoms, forming water molecules.
Oxidation number
To fully understand redox reactions, we need to talk about oxidation numbers. Oxidation numbers are a way of keeping track of electrons during a chemical reaction. They can be thought of as imaginary charges assigned to atoms in molecules or ions.
Here are some basic rules for determining oxidation numbers:
- The oxidation number of an atom in its elemental form is always zero. For example, the oxidation number of O2, H2, N2, etc. is zero.
- For a simple (monatomic) ion, the oxidation number is equal to the charge on the ion. For example, the oxidation number of Na+ is +1, and the oxidation number of Cl- is -1.
- The oxidation number of oxygen is usually -2, except when bonded to peroxides or fluorine.
- The oxidation number of hydrogen is generally +1 when it combines with non-metals, and -1 when it combines with metals.
- In a neutral compound the sum of the oxidation numbers must be zero.
- The sum of the oxidation numbers in a polyatomic ion must be equal to the charge of the ion.
Let's apply these rules to find the oxidation number of water (H2O):
- The oxidation number of oxygen is usually -2.
- The oxidation number of each hydrogen is +1.
- The sum of oxidation numbers in H2O = 2(+1) + (-2) = 0, which satisfies the rule of neutral compounds.
Identifying redox reactions
Identifying redox reactions involves checking for changes in the oxidation numbers of the substances in the reaction. If the oxidation numbers change, the reaction is redox.
Consider the following reaction:
4Zn + CuSO4 → ZnSO4 + Cu
To determine if this is a redox reaction, calculate the oxidation number:
- In Zn: Oxidation number of Zn = 0, ZnSO4: Zn = +2 (in compound form).
- In CuSO4: Cu = +2, elemental Cu: Cu = 0.
Information:
- Zn goes from 0 to +2 (loses electrons, i.e. is oxidized).
- Cu changes from +2 to 0 (gains electrons, i.e., is reduced).
The diagram shows the transfer of electrons from zinc to copper, forming zinc sulfate and copper metal. Zinc oxidation forms zinc ions and in parallel copper gains electrons, turning it into native copper.
Balancing redox reactions
Balancing redox reactions ensures that the number of electrons lost in oxidation is equal to the number of electrons gained in reduction. Here's how to balance using the half-reaction method:
Steps to balancing redox reactions
- Split the unbalanced redox equation into two half-reactions: one for oxidation and one for reduction.
- Balance each half-reaction for mass, then charge using electrons.
- Multiply the half-reactions by appropriate integers so that the electrons in both half-reactions are equal.
- Recombine the balanced half-reactions, making sure that the electrons cancel out.
- Verify the atomic and charge balance in the last equation.
Let's balance the redox reaction of iron with chlorine:
2Fe + Cl2 → FeCl3
Balancing process
Step 1: Split into half-reactions:
Oxidation: Fe → Fe3+ + 3e- Reduction: Cl2 + 2e- → 2Cl-
Step 2: Balance atoms other than O and H, then balance oxygen and hydrogen if present (not needed here).
Step 3: Balance the charges by adding electrons:
- The oxidation is already balanced for the charge.
- 3 electrons required for reduction: Cl2 + 6e- → 2Cl2
Step 4: Multiply the 2 reductions by appropriate integers to make the electrons equal to the oxidations:
Oxidation: Fe → Fe3+ + 3e- Reduction: Cl2 + 6e- → 2Cl- (same electron gain required)
Step 5: Add and remove electrons:
2Fe + 3/2Cl2 → FeCl3
The balanced reaction ensures equality in electron transfer and atom list.
Common redox reactions in everyday life
Redox reactions are all around us and are part of everyday life. Here are some examples:
Combustion
Combustion is a redox reaction in which fuel reacts with oxygen to release energy. For example, the burning of methane gas:
CH4 + 2 O2 → CO2 + 2 H2O
Batteries
Batteries operate on redox reactions, driven by chemical differences to transfer electrons from the anode to the cathode. In a simple zinc-carbon battery:
Zn → Zn2+ + 2e- 2MnO2 + 2e- + 2NH4Cl → Mn2O3 + H2O + 2NH3
War
Rusting of iron is a redox process in which iron is oxidized by oxygen and water to form iron oxide.
4 Fe + 3 O2 + 6 H2O → 4 Fe(OH)3
Conclusion
Redox reactions are fundamental to understanding both basic and complex chemical processes. From the burning of fuels to biological processes and energy storage in batteries, redox reactions facilitate essential transformations involving electron transfer. Understanding the basics of oxidation and reduction, as well as efficiently determining oxidation numbers, is critical to exploring the wide applications and implications of redox chemistry.