Redox Reactions - Some Important Questions

 

A. Fill in the Blanks:

1. In a redox reaction, the substance that gains electrons is called the __________.
2. The substance that loses electrons in a redox reaction is called the __________.
3. In the reaction $$2Na+Cl₂​→2NaCl, sodium (Na) is __________ (oxidized/reduced) and chlorine (Cl) is __________ (oxidized/reduced).
4. In a redox reaction, oxidation is the process of __________ (gaining/losing) electrons.
5. The oxidation state of chlorine in $$Cl₂​ is  __________ .
6. In the reaction between zinc and copper sulfate, zinc displaces copper. Here, zinc undergoes __________ (oxidation/reduction) and copper undergoes __________ (oxidation/reduction).
7. In the reaction $$Fe³⁺ + e−→Fe²⁺, iron is being __________ (oxidized/reduced).
8. The reducing agent in the reaction H₂​O₂​ → H₂​O + O₂​ is__________.
9. In the reaction 2Al + 3CuSO₄​ → 2Al₂​(SO₄​)₃​ + 3Cu$$, aluminum (Al) is __________ (oxidized/reduced) and copper (Cu) is __________ (oxidized/reduced). 
10. The oxidation number of sulfur in $$H₂​SO₄​ is __________ .​
    

   Answer:

1.     Reducer Or Oxidizing Agent

2.   Oxidizer Or Reducing Agent

3.     Oxidized, Reduced

4.     Losing

5.     Zero

6.     Oxidation, Reduction

7.     Reduced

8.     Hydrogen Peroxide

9.     Oxidized, Reduced

10.  +6

B. True/False:

1. In a redox reaction, oxidation refers to the loss of electrons and reduction refers to the gain of electrons. (True/False)
2. In the reaction 2K+Br₂​→2KBr, potassium (K) is reduced and bromine (Br) is oxidized. (True/False)
3. In the reaction 2Na+Cl₂​→2NaCl, both sodium and chlorine undergo oxidation. (True/False)
4. The oxidation number of oxygen in water (H₂O) is -2. (True/False)
5. In a redox reaction, the substance that loses electrons is called the oxidizing agent. (True/False)
6. In the reaction 2Cu+O₂​→2CuO, oxygen is the reducing agent. (True/False)
7. The process of reducing an element involves a decrease in its oxidation state. (True/False)
8. In the reaction 4Fe+3O₂​→2Fe₂​O₃​, iron is reduced to iron oxide. (True/False)
9. When hydrogen is added to a substance in a reaction, the substance undergoes oxidation. (True/False)
10. In the reaction 2FeCl₃​+3Zn→3ZnCl₂​+2Fe, zinc is reduced and iron(III) chloride is oxidized. (True/False)

$\mathrm{}$

Answer:

1.     True

2.     False [Explanation: Potassium (K) undergoes oxidation (loses electrons), and bromine (Br) undergoes reduction (gains electrons)].

3.     False [Explanation: Sodium undergoes oxidation (loses electrons), and chlorine undergoes reduction (gains electrons)]

4.     True

5.     False [Explanation: The substance that loses electrons is called the reducing agent, while the oxidizing agent gains electrons.]

6.     False [Explanation: Oxygen is the oxidizing agent because it gains electrons; copper is the reducing agent because it loses electrons.]

7.     True

8.     False [Explanation: Iron is oxidized (its oxidation state increases) to form iron(III) oxide, while oxygen is reduced (its oxidation state decreases).]

9.     False [Explanation: Adding hydrogen usually results in reduction (the substance gains electrons).]

10.  False [Explanation: Zinc is oxidized (loses electrons), and iron(III) chloride is reduced (gains electrons).]

C. Some Important Questions

Question 1: Define oxidation and reduction in terms of electron transfer. Give an example of each.

Answer: Oxidation is the process where a species loses electrons, resulting in an increase in its oxidation state.

Example: The reaction of magnesium metal with oxygen gas to form magnesium oxide.

2Mg(s) + O2(g) → 2MgO(s)

In this reaction, magnesium loses electrons and its oxidation state increases from 0 to +2. Therefore, magnesium is oxidized.

 

Reduction is the process where a species gains electrons, resulting in a decrease in its oxidation

state.

Example: The reaction of copper(II) oxide with hydrogen gas to form copper and water.

CuO(s) + H2(g) → Cu(s) + H2O(l)

In this reaction, copper(II) oxide gains electrons and its oxidation state decreases from +2 to 0. Therefore, copper(II) oxide is reduced.

 

Question 2: Balance the following redox reaction using the half-reaction method:

Answer: Step 1: Write the half-reactions

Oxidation half-reaction: Fe²⁺(aq) → Fe³⁺(aq)          

Reduction half-reaction: MnO^4-(aq) → Mn²⁺(aq)

 

Step 2: Balance the atoms other than O and H

Oxidation half-reaction: Fe²⁺(aq) → Fe³⁺(aq)           

Reduction half-reaction: MnO^4-(aq) → Mn²⁺(aq)

Step 3: Balance oxygen atoms by adding H2O

Oxidation half-reaction: Fe²⁺(aq) → Fe³⁺(aq)

Reduction half-reaction: MnO^4-(aq) → Mn²⁺(aq) + 4H₂O(l)

 

Step 4: Balance hydrogen atoms by adding H+

Oxidation half-reaction: Fe²⁺(aq) → Fe³⁺(aq)

Reduction half-reaction: MnO^4-(aq) + 8H⁺(aq) → Mn²⁺(aq) + 4H₂O(l)

 

Step 5: Balance charges by adding electrons

Oxidation half-reaction: Fe²⁺(aq) → Fe³⁺(aq) + e-

Reduction half-reaction: MnO^4-(aq) + 8H⁺(aq) + 5e- → Mn²⁺(aq) + 4H₂O(l)

 

Step 6: Multiply the half-reactions to make the number of electrons equal

Oxidation half-reaction: 5Fe²⁺(aq) → 5Fe³⁺(aq) + 5e-

 Reduction half-reaction: MnO^4-(aq) + 8H⁺(aq) + 5e- → Mn²⁺(aq) + 4H₂O(l)

 

Step 7: Add the half-reactions and cancel out the electrons

MnO^4-(aq) + 8H⁺(aq) + 5Fe²⁺(aq)  → Mn²⁺(aq) + 4H₂O(l) + 5Fe³⁺(aq)

 

Question 3: What is the role of a reducing agent in a redox reaction? Give an example of a reducing agent.

Answer: A reducing agent is a species that donates electrons to another species in a redox reaction. In doing so, it reduces the other species and itself gets oxidized. A reducing agent is also known as a reductant.

Example: In the reaction of zinc with hydrochloric acid, zinc acts as a reducing agent. Zinc donates electrons to hydrogen ions, reducing them to hydrogen gas, while zinc itself gets oxidized to zinc ions.

Question 4: What is the difference between a strong oxidizing agent and a weak oxidizing agent? Give examples.

Answer: An oxidizing agent is a species that accepts electrons from another species in a redox reaction. Strong oxidizing agents have a high affinity for electrons and readily accept electrons from other species. Weak oxidizing agents, on the other hand, have a lower affinity for electrons and are less likely to accept electrons.

Example of a Strong Oxidizing Agent: Potassium permanganate (KMnO₄)

Example of a Weak Oxidizing Agent:  Iodine (I₂)

 

Question 5: Explain how redox reactions are involved in the process of rusting. What is the role of oxygen in this process?

Answer: Rusting is a chemical process that involves the oxidation of iron in the presence of oxygen and water. The process begins when iron metal reacts with oxygen from the air to form iron(II) oxide (FeO).

2Fe(s) + O₂(g) → 2FeO(s)

This iron(II) oxide can then further react with oxygen and water to form hydrated iron(III) oxide, commonly known as rust (Fe₂O₃.xH₂O).

4FeO(s) + O₂(g) + 2xH₂O(l) → 2Fe₂O₃.xH₂O(s)

Oxygen acts as the oxidizing agent in this process, accepting electrons from iron and causing its oxidation.

 

Question 6: What is a disproportionation reaction? Give an example.

Answer: A disproportionation reaction is a type of redox reaction where the same species undergoes both oxidation and reduction. In other words, the same element is present in both the oxidized and reduced forms of the products.

Example: The reaction of chlorine with hydroxide ions (OH-) in an alkaline solution.

Cl₂(g) + 2OH^-(aq) → ClO-(aq) + Cl^-(aq) + H₂O(l)

In this reaction, chlorine is both oxidized to hypochlorite (ClO-) and reduced to chloride (Cl).

 

Question 7: How are redox reactions used in electrochemical cells? Explain with an example.

Answer: Redox reactions are the foundation of electrochemical cells, which convert chemical energy into electrical energy or vice versa. Electrochemical cells consist of two electrodes, an anode and a cathode, immersed in an electrolyte solution. At the anode, oxidation takes place, releasing electrons that flow through an external circuit to the cathode, where reduction occurs.

Example: A simple galvanic cell (voltaic cell) can be constructed using a zinc electrode and a copper electrode immersed in solutions of zinc sulfate and copper sulfate, respectively.

In this cell, zinc gets oxidized at the anode, releasing electrons that flow to the cathode, where copper ions get reduced. This flow of electrons generates an electric current.

 

Question 8: What are the applications of redox reactions in everyday life?

Answer: Redox reactions are ubiquitous in our daily lives. They are essential for various processes, including:

Combustion: The burning of fuels like wood, propane, or natural gas involves redox reactions where the fuel is oxidized and oxygen is reduced, releasing energy as heat and light.

Batteries: Batteries use redox reactions to store and release electrical energy. The flow of electrons in a battery is driven by redox reactions occurring at the electrodes.

Corrosion: The rusting of iron or the tarnishing of silver are examples of corrosion, a process involving redox reactions. In these cases, metals are oxidized, leading to their deterioration.

Photosynthesis: Plants use redox reactions to convert sunlight, water, and carbon dioxide into glucose and oxygen during photosynthesis. This process is crucial for life on Earth.

Respiration: Living organisms utilize redox reactions in respiration, a process where glucose is oxidized to release energy in the form of ATP. This process is vital for all living cells.