Redox Reaction Class 11 Chemistry Notes

Key Definitions

Oxidation

The process involving the loss of one or more electrons by an atom, ion, or molecule, leading to an increase in its oxidation number. Classically, it was defined as addition of oxygen or removal of hydrogen.

Reduction

The process involving the gain of one or more electrons by an atom, ion, or molecule, leading to a decrease in its oxidation number. Classically, it was defined as addition of hydrogen or removal of oxygen.

Redox Reaction

A chemical reaction where both oxidation and reduction occur simultaneously. The total number of electrons lost during oxidation equals the total number of electrons gained during reduction.

Oxidizing Agent (Oxidant)

A substance that causes oxidation in another substance by accepting electrons from it, and itself gets reduced. It facilitates the removal of electrons.

Reducing Agent (Reductant)

A substance that causes reduction in another substance by donating electrons to it, and itself gets oxidized. It facilitates the addition of electrons.

Oxidation Number (Oxidation State)

A hypothetical charge assigned to an atom in a molecule or ion, assuming that all bonds are ionic. It represents the number of electrons lost or gained by an atom in a compound.

Understanding Redox Reactions: A Deeper Dive

Redox reactions are at the heart of much of chemistry, from metabolic processes in living organisms to industrial manufacturing. Historically, oxidation was described as the addition of oxygen or removal of hydrogen, and reduction as the removal of oxygen or addition of hydrogen. For example, the rusting of iron (Fe → Fe₂O₃) is an oxidation process where iron gains oxygen. The reduction of copper oxide by hydrogen (CuO + H₂ → Cu + H₂O) involves CuO losing oxygen.

However, these definitions were limited to reactions involving oxygen or hydrogen. The modern and more comprehensive definition relies on electron transfer. According to this view, oxidation is the loss of electrons, and reduction is the gain of electrons. This electronic concept applies to a much broader range of reactions. For instance, in the reaction 2Na + Cl₂ → 2NaCl, Sodium (Na) loses electrons to become Na⁺ (oxidation), and Chlorine (Cl₂) gains electrons to become Cl⁻ (reduction). This simultaneous occurrence of electron loss and gain is what defines a redox reaction.

The concept of oxidation number (or oxidation state) is a powerful tool to track electron shifts and identify redox processes. It's a hypothetical charge an atom would have if all bonds were purely ionic. An increase in oxidation number signifies oxidation (loss of electrons), while a decrease signifies reduction (gain of electrons). Understanding how to assign and interpret oxidation numbers is crucial for balancing redox equations and predicting reaction outcomes. The ability to identify oxidizing and reducing agents is also key: the substance that gets reduced is the oxidizing agent, and the substance that gets oxidized is the reducing agent.

Rules for Assigning Oxidation Numbers

Oxidation vs. Reduction

Worked Example: Oxidation Number Calculation

• Example 1: Find the oxidation number of Cr in K₂Cr₂O₇. Let the oxidation number of Cr be 'x'. K is a Group 1 metal, so its oxidation number is +1. Oxygen usually has an oxidation number of -2. For a neutral molecule, the sum of oxidation numbers is 0. 2(+1) + 2(x) + 7(-2) = 0 2 + 2x - 14 = 0 2x - 12 = 0 2x = 12 x = +6 Thus, the oxidation number of Cr in K₂Cr₂O₇ is +6.

Balancing Redox Reactions (Ion-Electron Method)

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Exam Tip: Common Pitfalls in Redox Reactions

When balancing redox reactions, always check both mass and charge balance in the final equation. A common mistake is to forget to balance charges after balancing atoms. Pay close attention to the medium (acidic or basic) as it dictates how you add H⁺/OH⁻ and H₂O. For oxidation number calculations, remember the exceptions for oxygen (peroxides, superoxides, OF₂) and hydrogen (metal hydrides). Practice assigning oxidation numbers for complex ions to avoid errors. Showing clear steps for balancing will fetch you partial marks even if the final answer has a minor error.

Key Points to Remember

• Redox reactions involve simultaneous oxidation (loss of electrons, increase in oxidation number) and reduction (gain of electrons, decrease in oxidation number).
• The substance oxidized is the reducing agent; the substance reduced is the oxidizing agent.
• Oxidation number is a hypothetical charge, crucial for tracking electron transfer and balancing equations.
• Rules for assigning oxidation numbers must be memorized, especially exceptions for H and O.
• Balancing redox reactions can be done by the oxidation number method or the ion-electron (half-reaction) method.
• In acidic medium, H⁺ and H₂O are used for balancing H and O; in basic medium, OH⁻ and H₂O are used.
• Disproportionation reactions are special redox reactions where the same element is simultaneously oxidized and reduced.
• Redox reactions are fundamental in electrochemistry, biological processes, and industrial applications.

Practice Questions with Solutions

• Q: Define an oxidizing agent. A: An oxidizing agent is a substance that causes the oxidation of another substance by accepting electrons from it, and itself gets reduced.
• Q: What is the oxidation number of Sulfur in H₂SO₄? A: Let S be 'x'. 2(+1) + x + 4(-2) = 0 => 2 + x - 8 = 0 => x = +6. So, S is +6.
• Q: In the reaction Zn + CuSO₄ → ZnSO₄ + Cu, identify the species undergoing oxidation and reduction. A: Zn goes from 0 to +2 (oxidation). Cu in CuSO₄ goes from +2 to 0 (reduction).
• Q: What is a disproportionation reaction? Give an example. A: A disproportionation reaction is a type of redox reaction where the same element in a single reactant is simultaneously oxidized and reduced. Example: 2H₂O₂ → 2H₂O + O₂ (Oxygen in H₂O₂ is -1, becomes -2 in H₂O and 0 in O₂).

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