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Class 11 Science (PCM)

Chemistry Part 2

Redox Reactions

Redox Reactions

Action PlanNotesQuestionsMock-Test
Study Action Plan (Sorted by Difficulty)
Electrode Potentials and Daniell Cell
Accuracy: 29%
Total Attempted7
Total Correct2
Total Incorrect5
Difficulty Level:
High
Redox Reactions
Accuracy: 47%
Total Attempted15
Total Correct7
Total Incorrect8
Difficulty Level:
High
Balancing Redox Reactions
Accuracy: 67%
Total Attempted9
Total Correct6
Total Incorrect3
Difficulty Level:
Medium
Titrations in Redox Reactions
Accuracy: 67%
Total Attempted6
Total Correct4
Total Incorrect2
Difficulty Level:
Medium
Standard Electrode Potentials
Accuracy: 75%
Total Attempted4
Total Correct3
Total Incorrect1
Difficulty Level:
Low
Classification of Redox Reactions
Accuracy: 80%
Total Attempted5
Total Correct4
Total Incorrect1
Difficulty Level:
Low
Oxidation Number
Accuracy: 83%
Total Attempted12
Total Correct10
Total Incorrect2
Difficulty Level:
Low
Electron Transfer and Redox Couples
Accuracy: 87%
Total Attempted15
Total Correct13
Total Incorrect2
Difficulty Level:
Low
Fractional Oxidation States
Accuracy: 100%
Total Attempted1
Total Correct1
Total Incorrect0
Difficulty Level:
Low

Learning Objectives

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Learning Objectives

  • Understand the simultaneous nature of oxidation and reduction in redox reactions, identifying oxidants and reductants.
  • Calculate oxidation numbers using established rules to determine oxidation states in compounds and ions.
  • Analyze electron transfer processes in redox reactions, defining and representing redox couples in half-reactions.
  • Balance redox equations using both the oxidation-number method and the half-reaction method, ensuring mass and charge balance in acidic and basic media.
  • Explore the concept of electrode potentials, particularly in the context of the Daniell cell, and compare oxidizing and reducing tendencies.
  • Interpret standard electrode potentials to evaluate the strength of oxidizing and reducing agents.
  • Apply titration methods in redox reactions to determine the strength of oxidants and reductants, utilizing appropriate indicators.
  • Examine the concept of fractional oxidation states in compounds, understanding their average nature and structural basis.
  • Classify redox reactions into combination, decomposition, displacement, and disproportionation reactions, including trends in metal/non-metal displacement and halogen activity.

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Revision Notes & Summary

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Chapter Notes

Redox Reactions

  • Definition: Redox reactions involve simultaneous oxidation and reduction processes.
  • Oxidation: Originally defined as the addition of oxygen to a substance.
  • Reduction: Initially described as the removal of oxygen from a compound.
  • Oxidant (Oxidising Agent): Substance that gains electrons.
  • Reductant (Reducing Agent): Substance that loses electrons.

Oxidation Number

  • Purpose: Helps in identifying oxidants and reductants in a reaction.
  • Assignment Rules:
    • Free elements have an oxidation number of zero.
    • For ions, the oxidation number equals the charge.
    • Oxygen generally has an oxidation number of -2, except in peroxides (-1) and when bonded to fluorine (+2).
    • Hydrogen is +1, except when bonded to metals in binary compounds (-1).
    • The sum of oxidation numbers in a neutral compound is zero; in a polyatomic ion, it equals the ion's charge.

Electron Transfer and Redox Couples

  • Redox Couple: Consists of the oxidised and reduced forms of a substance.
  • Half-Reactions: Redox reactions can be split into two half-reactions showing electron transfer.

Balancing Redox Reactions

  • Oxidation-Number Method:
    • Assign oxidation numbers to all elements.
    • Calculate changes in oxidation numbers to balance the equation.
  • Half-Reaction Method:
    • Separate the reaction into oxidation and reduction half-reactions.
    • Balance each half-reaction for mass and charge.
    • Combine the balanced half-reactions.

Electrode Potentials and Daniell Cell

  • Electrode Potential: Measures the tendency of a species to be reduced.
  • Daniell Cell: An electrochemical cell where zinc and copper undergo redox reactions.
  • Standard Electrode Potential (E°): Measured under standard conditions (1 M concentration, 1 atm pressure, 298 K).

Standard Electrode Potentials

  • Table of Standard Potentials: Lists electrode potentials for various redox couples.
  • Interpretation:
    • Positive E° indicates a strong oxidising agent.
    • Negative E° indicates a strong reducing agent.

Titrations in Redox Reactions

  • Redox Titration: Used to determine the concentration of an oxidant or reductant.
  • Indicators: Substances that change color at the endpoint of the titration.

Fractional Oxidation States

  • Concept: Occurs when an element in a compound exhibits more than one oxidation state.
  • Example: C3O2C_3O_2 where carbon has an average oxidation state of 4/3.

Classification of Redox Reactions

  • Combination Reactions: Two or more elements combine to form a compound.
  • Decomposition Reactions: A compound breaks down into two or more components.
  • Displacement Reactions: An element displaces another in a compound.
  • Disproportionation Reactions: An element in one oxidation state is simultaneously oxidised and reduced.

Example Problems

  • Balancing Redox Reactions: Use the oxidation number or half-reaction method to balance equations.
  • Electrode Potentials: Use standard electrode potentials to predict the feasibility of redox reactions.

Exercises

  • Assign oxidation numbers, balance redox reactions, and identify oxidising/reducing agents in given reactions.

Summary

  • Redox reactions are fundamental to chemistry, involving electron transfer processes.
  • Understanding oxidation numbers and electrode potentials is crucial for analyzing redox reactions.
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Practice Test – MCQs, True/False

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Multiple Choice Questions

A.

A negative E∘E^\circ indicates a stronger reducing agent than the H+/H2H^+/H_2 couple.

B.

A positive E∘E^\circ indicates a stronger reducing agent than the H+/H2H^+/H_2 couple.

C.

A negative E∘E^\circ indicates a weaker reducing agent than the H+/H2H^+/H_2 couple.

D.

The E∘E^\circ value does not indicate the strength of a reducing agent.
Correct Answer: A

Solution:

A negative standard electrode potential (E∘E^\circ) means the redox couple is a stronger reducing agent than the H+/H2H^+/H_2 couple.

Chapter Concept:

Electrode Potentials and Daniell Cell

A.

Fe3+(aq)+I−(aq)→Fe2+(aq)+12I2(s)Fe^{3+}(aq) + I^-(aq) \rightarrow Fe^{2+}(aq) + \frac{1}{2}I_2(s)

B.

Ag+(aq)+Cu(s)→Ag(s)+Cu2+(aq)Ag^+(aq) + Cu(s) \rightarrow Ag(s) + Cu^{2+}(aq)

C.

Fe3+(aq)+Cu(s)→Fe2+(aq)+Cu2+(aq)Fe^{3+}(aq) + Cu(s) \rightarrow Fe^{2+}(aq) + Cu^{2+}(aq)

D.

Ag(s)+Fe3+(aq)→Ag+(aq)+Fe2+(aq)Ag(s) + Fe^{3+}(aq) \rightarrow Ag^+(aq) + Fe^{2+}(aq)
Correct Answer: B

Solution:

The reaction Ag+(aq)+Cu(s)→Ag(s)+Cu2+(aq)Ag^+(aq) + Cu(s) \rightarrow Ag(s) + Cu^{2+}(aq) is feasible as the standard electrode potential for Ag+/AgAg^+/Ag is higher than that of Cu2+/CuCu^{2+}/Cu, indicating Ag+Ag^+ can oxidize CuCu.

Chapter Concept:

Electrode Potentials and Daniell Cell

A.

Sodium (NaNa)

B.

Chlorine (Cl2Cl_2)

C.

Sodium chloride (NaClNaCl)

D.

None of the above
Correct Answer: B

Solution:

Chlorine is reduced because it gains electrons from sodium, forming Cl−Cl^- ions.

Chapter Concept:

Oxidation Number

A.

To allow the direct flow of electrons between electrodes.

B.

To prevent the flow of ions between the two half-cells.

C.

To maintain electrical neutrality by allowing the flow of ions.

D.

To increase the potential difference between the electrodes.
Correct Answer: C

Solution:

The salt bridge maintains electrical neutrality by allowing the flow of ions between the two half-cells, preventing the solutions from mixing while completing the electrical circuit.

Chapter Concept:

Electrode Potentials and Daniell Cell

A.

Zinc (Zn\text{Zn})

B.

Copper ion (Cu2+\text{Cu}^{2+})

C.

Zinc ion (Zn2+\text{Zn}^{2+})

D.

Copper (Cu\text{Cu})
Correct Answer: A

Solution:

Zinc is the reducing agent as it donates electrons to copper ions, reducing them to copper metal.

Chapter Concept:

Redox Reactions

A.

Br2Br_2

B.

BrO3−BrO_3^-

C.

Br−Br^-

D.

Br2OBr_2O
Correct Answer: B

Solution:

Br−Br^- is oxidized to BrO3−BrO_3^- in the presence of MnO4−MnO_4^- in basic medium.

Chapter Concept:

Titrations in Redox Reactions

A.

Copper rod

B.

Zinc rod

C.

Copper sulfate solution

D.

Zinc sulfate solution
Correct Answer: B

Solution:

In the Daniell cell, the zinc rod acts as the anode where oxidation occurs.

Chapter Concept:

Electrode Potentials and Daniell Cell

A.

SO2SO_2

B.

SO3SO_3

C.

Na2S4O6Na_2S_4O_6

D.

H2SO4H_2SO_4
Correct Answer: C

Solution:

In Na2S4O6Na_2S_4O_6, the average oxidation state of sulfur is 2.52.5, indicating a fractional oxidation state due to the presence of sulfur atoms in different oxidation states.

Chapter Concept:

Fractional Oxidation States

A.

Fluorine (F2F_2)

B.

Chlorine (Cl2Cl_2)

C.

Bromine (Br2Br_2)

D.

Iodine (I2I_2)
Correct Answer: A

Solution:

Fluorine (F2F_2) is the strongest oxidizing agent as it has the highest standard electrode potential.

Chapter Concept:

Electron Transfer and Redox Couples

A.

Zn2++2e−→ZnZn^{2+} + 2e^- \rightarrow Zn

B.

Cu2++2e−→CuCu^{2+} + 2e^- \rightarrow Cu

C.

Zn→Zn2++2e−Zn \rightarrow Zn^{2+} + 2e^-

D.

Cu→Cu2++2e−Cu \rightarrow Cu^{2+} + 2e^-
Correct Answer: B

Solution:

In a Daniell cell, the reduction half-reaction occurs at the copper electrode where Cu2+Cu^{2+} ions gain electrons to form copper metal: Cu2++2e−→CuCu^{2+} + 2e^- \rightarrow Cu.

Chapter Concept:

Electrode Potentials and Daniell Cell

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Redox Reactions

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True or False

Correct Answer: False

Solution:

The reaction is a redox reaction because zinc is oxidized to Zn2+Zn^{2+} and copper ions are reduced to copper metal.

Chapter Concept :

Electron Transfer and Redox Couples

Correct Answer: True

Solution:

Fractional oxidation states are average values because electrons are never shared or transferred in fractions. The structural parameters reveal that the element is present in different oxidation states.

Chapter Concept :

Fractional Oxidation States

Correct Answer: True

Solution:

The permanganate ion (MnO4−MnO_4^-) acts as a self-indicator in redox titrations due to its intense purple color, which disappears when it is reduced, indicating the endpoint of the titration.

Chapter Concept :

Titrations in Redox Reactions

Correct Answer: True

Solution:

In peroxides, the oxidation number of oxygen is -1, and in superoxides, it is -1/2.

Chapter Concept :

Redox Reactions

Correct Answer: True

Solution:

The standard electrode potential (E°) is a measure of the tendency of a redox couple to gain electrons and thus act as an oxidizing agent. A higher positive E° indicates a stronger oxidizing ability.

Chapter Concept :

Electrode Potentials and Daniell Cell

Correct Answer: False

Solution:

In this reaction, sodium is oxidized as it loses electrons to form Na+\text{Na}^+ ions.

Chapter Concept :

Redox Reactions

Correct Answer: False

Solution:

In the reaction 2Mg+O2→2MgO2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}, magnesium is oxidized due to the addition of oxygen.

Chapter Concept :

Redox Reactions

Correct Answer: False

Solution:

The oxidation number of an element can be a fraction in certain compounds, such as C3O2C_3O_2, where the average oxidation state of carbon is 43\frac{4}{3}. This occurs because the element is present in different oxidation states within the compound.

Chapter Concept :

Oxidation Number

Correct Answer: False

Solution:

In peroxides, the oxidation number of oxygen is -1, not -2.

Chapter Concept :

Redox Reactions

Correct Answer: True

Solution:

The sum of the oxidation numbers in a polyatomic ion is equal to the charge on the ion, as per the rules for assigning oxidation numbers.

Chapter Concept :

Balancing Redox Reactions

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