Maharashtra Board Class 11 Chemistry Chapter 6 Redox Reactions PDF Download

Redox Reactions

6.1 Introduction

Redox is short for "oxidation and reduction". Many things we see happen every day involve redox reactions. For example, when we breathe, when iron rusts, and when fuel burns, all these are redox reactions.

Why does a cut apple turn brown when exposed to air? Why does an old car bumper change colour? Why do new batteries become useless after some days? These are all redox reactions happening around us.

6.1.1 Classical Ideas Of Redox Reactions

Long ago, scientists thought oxidation meant adding oxygen to something.

For example, when carbon burns, it joins with oxygen:

\[C(s) + O_2(g) \rightarrow CO_2(g)\]

When magnesium burns, it also joins with oxygen:

\[2Mg + O_2 \rightarrow 2MgO(s)\]

In these reactions, carbon and magnesium are oxidized because they join with oxygen.

But sometimes, something loses oxygen. This is called reduction. For example:

\[2Fe_2O_3 + 3C(s) \rightarrow 4Fe(s) + 3CO_2(g)\]

Here, iron loses oxygen, so it is reduced.

Sometimes hydrogen is removed. This is also oxidation. For example:

\[2H_2S(g) + O_2(g) \rightarrow 2S(s) + 2H_2O(l)\]

Here, sulfur loses hydrogen and is oxidized.

Now we know that oxidation does not always mean adding oxygen. It can also mean:

Joining with an element that pulls electrons away. For example:

\[Mg(s) + F_2(g) \rightarrow MgF_2(s)\]

\[Mg(s) + S(s) \rightarrow MgS(s)\]

These are also oxidation reactions even though oxygen is not involved.

Oxidation also means losing an element that pushes electrons toward us. For example:

\[Hg_2Cl_2(s) \rightarrow HgCl_2(s) + Hg(s)\]

Teacher's Note

When we cut an apple and it turns brown, that is oxidation happening in real life. The apple is losing electrons to oxygen in the air, just like how metals rust.

Exam Trick

Remember: Oxidation is not always about oxygen! It can mean losing electrons or joining with elements that pull electrons away.

Points to Remember

Oxidation means joining with oxygen or losing electrons.
Reduction means losing oxygen or gaining electrons.
These reactions happen all around us every day.
Old ideas about oxidation have been expanded to include electron transfer.

Now let us think about reduction. Reduction means:

Removing oxygen. For example:

\[2HgO(s) \rightarrow 2Hg(l) + O_2(g)\]

Removing an element that pulls electrons away. For example:

\[2FeCl_3 + H_2(g) \rightarrow 2FeCl_2(aq) + 2HCl\]

Adding hydrogen. For example:

\[CH_2 = CH_2(g) + H_2(g) \rightarrow CH_3 - CH_3(g)\]

Adding an element that pushes electrons toward us. For example:

\[2HgCl_2(aq) + SnCl_2(aq) \rightarrow Hg_2Cl_2(s) + SnCl_4(aq)\]

An oxidizing agent is a substance that causes another substance to lose electrons. It takes electrons from the other substance.

A reducing agent is a substance that causes another substance to gain electrons. It gives electrons to the other substance.

In the reaction above, HgCl₂ is reduced and SnCl₂ is oxidized. So this is a redox reaction because both oxidation and reduction happen at the same time.

Teacher's Note

Think of an oxidizing agent like a person who takes your money, and a reducing agent like a person who gives you money. In India, when old bumpers rust, that is oxidation happening because the iron is losing electrons to oxygen.

Exam Trick

Remember: Oxidizing agent = takes electrons. Reducing agent = gives electrons. Just like taking and giving!

Points to Remember

Oxidizing agents take electrons from other substances.
Reducing agents give electrons to other substances.
When both happen together, it is a redox reaction.
Oxidation and reduction always happen together in a redox reaction.

6.1.2 Redox Reaction In Terms Of Electron Transfer

We can describe redox reactions by counting electrons that move.

\[Mg + O_2 \rightarrow Mg^{2+} + O^{2-}\]

\[Mg + F_2 \rightarrow Mg^{2+} + 2F^-\]

When magnesium is oxidized to MgO, the neutral magnesium atom loses 2 electrons to become Mg²⁺. The oxygen gains these 2 electrons to become O²⁻.

Each step where electrons are lost or gained is called a half reaction.

For example:

\[Fe(s) \rightarrow Fe^{2+}(aq) + 2e^-\]

\[Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)\]

\[Fe(s) + Cu^{2+}(aq) \rightarrow Fe^{2+}(aq) + Cu(s)\]

In the first half reaction, iron loses electrons. In the second half reaction, copper gains electrons. When we add them together, we get a redox reaction.

Fe acts as a reducing agent because it gives electrons. Cu²⁺ acts as an oxidizing agent because it takes electrons.

Teacher's Note

Displacement reactions, like when one metal pushes another out, are also redox reactions. For example, when zinc pushes copper out of copper sulfate solution.

Exam Trick

Remember: In half reactions, one loses electrons (oxidation) and one gains electrons (reduction). Add them and you get a redox reaction.

Points to Remember

Half reactions show electrons being lost or gained.
Oxidation is loss of electrons.
Reduction is gain of electrons.
Redox reactions have two half reactions added together.

6.2 Oxidation Number

We need a way to talk about redox reactions that works for all types of compounds, not just those with ionic bonds.

For example:

\[2H_2(g) + O_2(g) \rightarrow 2H_2O(l)\]

\[H_2(g) + Cl_2(g) \rightarrow 2HCl(g)\]

In the first reaction, hydrogen joins with oxygen. So it is oxidation. In the second reaction, an element that pulls electrons is added to hydrogen. So it is also oxidation.

But these products are not ionic. They are covalent. In covalent bonds, the electrons are not completely transferred. They just shift a little bit toward one atom.

To handle this, we use something called oxidation number. We pretend that the electrons in a covalent bond belong completely to the atom that pulls them harder.

The oxidation number of an element in a compound is the number of electrical charges it would have if all electrons shifted completely to the atom that pulls them harder.

6.2.1 Rules To Assign Oxidation Number

Rule 1: In free elements, the oxidation number is zero. For example, in H₂, Cl₂, O₃, and S₈, each atom has oxidation number zero.

Rule 2: In an ion with one atom, the oxidation number equals the charge of the ion. Alkali metals like sodium have oxidation number +1 in all compounds. Alkaline earth metals like calcium have oxidation number +2 in all compounds. Aluminum has oxidation number +3 in all compounds.

Rule 3: Oxygen usually has oxidation number -2 in all compounds. But in peroxides like H₂O₂, it has -1. In superoxides, it has -1/2. In OF₂, oxygen has +2.

Rule 4: Hydrogen usually has oxidation number +1. But when it is bonded to metals, it has -1.

Rule 5: Fluorine always has oxidation number -1. Other halogens like chlorine, bromine, and iodine usually have -1 in their compounds with nonmetals. But in compounds with oxygen, they have +1.

Rule 6: In a neutral molecule, the sum of all oxidation numbers is zero.

Rule 7: In an ion, the sum of all oxidation numbers equals the charge of the ion.

Rule 8: When there are two or more atoms of the same element in a molecule, the oxidation number of each atom is the average.

Teacher's Note

Oxidation numbers help us track electrons in all types of reactions, even covalent ones. Just like how we count runs in cricket, we count oxidation number changes to find redox reactions.

Exam Trick

Remember: Oxidation number of an element can be positive, negative, whole number, or a fraction. It is just a counting tool to find redox reactions.

Points to Remember

Free elements always have oxidation number zero.
Oxygen is usually -2, except in peroxides and superoxides.
In molecules, all oxidation numbers add up to zero.
In ions, all oxidation numbers add up to the ion charge.

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