Problem 1 - A Complete Resource Book in Chemistry for JEE Main

Question

Which of the following is a redox reaction? (a) \(\mathrm{NaCl}+\mathrm{KNO}_{3} \longrightarrow \mathrm{NaNO}_{3}+\mathrm{KCl}\) (b) \(\mathrm{CaC}_{2} \mathrm{O}_{4}+2 \mathrm{HCl} \longrightarrow \mathrm{CaCl}_{2}+\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) (c) \(\mathrm{Mg}(\mathrm{OH})_{2}+2 \mathrm{NH}_{4} \mathrm{Cl} \longrightarrow \mathrm{MgCl}_{2}+2 \mathrm{NH}_{4} \mathrm{OH}\) (d) \(\mathrm{Zn}+2 \mathrm{AgCN} \longrightarrow 2 \mathrm{Ag}+\mathrm{Zn}(\mathrm{CN})_{2}\)

Step-by-Step Solution

Answer

Reaction (d) is a redox reaction.

Step 1: Understanding Redox Reactions

A redox reaction is a chemical reaction in which there is a change in the oxidation states of the substances involved. This involves the transfer of electrons between two species: one gains electrons (reduced), and the other loses electrons (oxidized). We'll identify the changes in oxidation states to determine which reaction is a redox reaction.

Step 2: Examining Reaction (a)

(a) \( \mathrm{NaCl} + \mathrm{KNO}_{3} \longrightarrow \mathrm{NaNO}_{3} + \mathrm{KCl} \)Examine the oxidation states:- Na in NaCl: +1- Cl in NaCl: -1- K in KNO₃: +1- N in KNO₃: +5- O in KNO₃: -2In the products:- Na in NaNO₃: +1- K in KCl: +1- Cl in KCl: -1- N in NaNO₃: +5- O in NaNO₃: -2No change in oxidation states occurs, so this is not a redox reaction.

Step 3: Examining Reaction (b)

(b) \( \mathrm{CaC}_{2}\mathrm{O}_{4} + 2 \mathrm{HCl} \longrightarrow \mathrm{CaCl}_{2} + \mathrm{H}_{2}\mathrm{C}_{2}\mathrm{O}_{4} \)Examine the oxidation states:- Ca in CaC₂O₄: +2- C in oxalate \((\mathrm{C}_{2}\mathrm{O}_{4})\): +3- H in HCl: +1- Cl in HCl: -1In the products:- Ca in CaCl₂: +2- H in H₂C₂O₄: +1- C in C₂O₄ (oxalic acid): +3- Cl in CaCl₂: -1No change in oxidation states occurs, so this is not a redox reaction.

Step 4: Examining Reaction (c)

(c) \( \mathrm{Mg(OH)}_{2} + 2 \mathrm{NH}_{4}\mathrm{Cl} \longrightarrow \mathrm{MgCl}_{2} + 2 \mathrm{NH}_{4}\mathrm{OH} \)Examine the oxidation states:- Mg in Mg(OH)₂: +2- O in OH: -2- H in OH/NH₄: +1- N in NH₄: -3- Cl in NH₄Cl: -1In the products:- Mg in MgCl₂: +2- Cl in MgCl₂: -1- N in NH₄OH: -3- H in NH₄OH: +1- O in NH₄OH: -2No change in oxidation states occurs, so this is not a redox reaction.

Step 5: Examining Reaction (d)

(d) \( \mathrm{Zn} + 2 \mathrm{AgCN} \longrightarrow 2 \mathrm{Ag} + \mathrm{Zn(CN)}_{2} \)Examine the oxidation states:- Zn in elemental form: 0- Ag in AgCN: +1- C in CN: +2- N in CN: -3In the products:- Zn in Zn(CN)₂: +2- Ag in elemental form: 0- C in CN: +2- N in CN: -3Zn goes from 0 to +2, indicating it is oxidized. Ag goes from +1 to 0, indicating it is reduced. There is a change in oxidation states, making this a redox reaction.

Key Concepts

Oxidation States

Oxidation states, also known as oxidation numbers, are a way to keep track of electrons during chemical reactions, especially redox reactions. They help us determine how electrons are transferred between atoms. Typically, the oxidation state of an element in its natural state is zero. For compounds, we assign oxidation states based on a set of rules, such as:

• The oxidation state of oxygen is usually -2.
• The oxidation state of hydrogen is usually +1.
• For monoatomic ions, the oxidation state equals the ion's charge.

By assigning these numbers, we can observe changes in oxidation states from reactants to products. If the oxidation number of an element increases, it means it has lost electrons and is oxidized. Conversely, if the oxidation number decreases, it has gained electrons and is reduced.
This concept is central to identifying redox reactions, as they involve changes in oxidation states. For example, in reaction (d) from the exercise, Zn changes from an oxidation state of 0 to +2, indicating it is oxidized, and Ag changes from +1 to 0, indicating it is reduced.

Electron Transfer

In redox reactions, the transfer of electrons is what defines the process. The species that loses electrons gets oxidized, while the species that gains electrons gets reduced. This electron transfer is crucial for the conservation of charge in the reaction.
Imagine electrons as tiny particles being passed from one atom to another. When an atom loses an electron, it becomes positively charged because it has more protons than electrons. We call this process oxidation. On the other hand, when an atom gains electrons, it becomes negatively charged, which is known as reduction. These changes can be tracked using oxidation states, as previously discussed.

• Example of oxidation: Zinc (Zn) loses electrons, going from 0 to +2 oxidation state.
• Example of reduction: Silver (Ag) gains electrons, going from +1 to 0 oxidation state.

In reaction (d), when Zn is oxidized, it loses electrons, which are then gained by Ag, causing Ag to be reduced. This consistent movement and balance of electrons is what makes a redox reaction a key component of many chemical processes.

Chemical Reactions

Chemical reactions are transformations where reactants are converted into products. In the context of redox reactions, these involve the transfer of electrons between substances, leading to the rearrangement of atoms and bonds.
Redox reactions play a vital role in processes such as combustion, metabolism, and even in the functioning of batteries. It's essential to distinguish between reactions that are redox-based and those that are not, by observing changes in oxidation states.
Consider how in the exercise, we have various reactions presented. To identify the redox reaction, we examine each reaction for electron movement:

• Reactions (a), (b), and (c) do not involve any change in oxidation states, indicating they aren't redox reactions.
• Reaction (d), however, involves a clear change in oxidation states, with Zn being oxidized and Ag being reduced, marking it as a redox reaction.

Understanding chemical reactions, especially redox reactions, deepens our comprehension of various processes occurring in our daily lives and enables us to predict the behavior of different substances in various environments.