Are these practice questions for Redox Reactions based on the CBSE pattern?

Yes. Every practice question is modeled on the CBSE Class 11 - Science Chemistry exam pattern — including MCQs, assertion-reason questions, and short-answer formats that frequently appear in CBSE papers.

Are the practice questions for Redox Reactions free?

Yes. Every chapter includes free practice questions for Class 11 - Science students to get you started.

What types of questions are included?

The practice set for Redox Reactions mainly includes MCQs and CBSE-style short-answer questions, with instant feedback on each answer. For full-length textbook practice, use the NCERT Solutions page.

Are the practice questions suitable for board exam preparation?

Yes. The questions test the same concepts and depth that CBSE asks in board exams. After finishing, also solve the NCERT and NCERT Exemplar exercises for Redox Reactions for full coverage.

How many practice questions are there for Redox Reactions?

The number scales with chapter length and concept density. Each chapter typically has 15-50 practice questions across MCQ and short-answer formats.

Redox Reactions Class 11 Chemistry Practice Questions with Answers

easy

Identify the oxidizing agent in the following reaction:

\mathrm{Zn(s)} + \mathrm{Cu^{2+}(aq)} \rightarrow \mathrm{Zn^{2+}(aq)} + \mathrm{Cu(s)}

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easy

Analyze the reaction

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

and identify the species undergoing oxidation.

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easy

Explain 'reduction' according to the classical idea. Provide one example for the addition of hydrogen and one for the removal of oxygen.

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easy

Define the term 'reductant' or 'reducing agent'.

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easy

List any three rules used for the calculation of the oxidation number of an element in a compound or ion.

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easy

Justify why fluorine, the most powerful oxidizing agent, exhibits only a -1 oxidation state in its compounds and cannot undergo disproportionation.

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easy

Calculate the oxidation number of Phosphorus in Magnesium Phosphate,

Mg_3(PO_4)_2

.

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easy

What is the oxidation number of an element when it is in its free or uncombined state? Give two examples.

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medium

Justify the classification of the thermal decomposition of potassium chlorate (

\text{KClO}_3

) as an intramolecular redox reaction.

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medium

Analyze the following reaction and identify the substance oxidized, substance reduced, oxidizing agent, and reducing agent.

I_2(s) + 10HNO_3(aq) \rightarrow 2HIO_3(aq) + 10NO_2(g) + 4H_2O(l)

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medium

Design a simple galvanic cell using cobalt and nickel electrodes. Given the standard electrode potentials:

\text{Co}^{2+}(\text{aq}) + 2\text{e}^- \rightarrow \text{Co}(\text{s})

,

E^\ominus = -0.28

V

\text{Ni}^{2+}(\text{aq}) + 2\text{e}^- \rightarrow \text{Ni}(\text{s})

,

E^\ominus = -0.25

V

Justify your choice of anode and cathode, write the half-reactions, and formulate the overall cell reaction.

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medium

List and describe the five main steps involved in balancing a redox reaction using the oxidation number method in an acidic solution.

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medium

Compare the reactions of zinc metal with dilute hydrochloric acid and with cold water. Explain the difference in reactivity and classify the type of redox reaction in both cases.

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medium

What is a redox couple?

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medium

Define oxidation and reduction in terms of electron transfer. Identify the species being oxidized and reduced in the reaction:

2\mathrm{Na(s)} + \mathrm{Cl_2(g)} \rightarrow 2\mathrm{NaCl(s)}

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medium

Describe the primary function of a salt bridge in an electrochemical cell like the Daniell cell.

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medium

Summarize the key differences between the classical concept and the electron transfer concept of redox reactions. Provide one example reaction that can be explained by both concepts.

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medium

Demonstrate why the decomposition of silver carbonate,

Ag_2CO_3(s) \rightarrow Ag_2O(s) + CO_2(g)

, is not a redox reaction, whereas the decomposition of silver oxide,

2Ag_2O(s) \rightarrow 4Ag(s) + O_2(g)

, is a redox reaction.

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medium

A strip of nickel metal is placed in a 1M solution of silver nitrate. Analyze the reaction that occurs and write the balanced net ionic equation. Identify which metal is more reactive based on this observation.

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medium

Evaluate the statement: 'The concept of fractional oxidation number, as in tetrathionate ion (

\text{S}_4\text{O}_6^{2-}

), is physically meaningless because electrons cannot be transferred in fractions.' Propose the correct interpretation.

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medium

Identify the type of redox reaction for each of the following: (i)

\mathrm{C(s)} + \mathrm{O_2(g)} \rightarrow \mathrm{CO_2(g)}

(ii)

2\mathrm{H_2O(l)} \rightarrow 2\mathrm{H_2(g)} + \mathrm{O_2(g)}

(iii)

\mathrm{Zn(s)} + \mathrm{CuSO_4(aq)} \rightarrow \mathrm{ZnSO_4(aq)} + \mathrm{Cu(s)}

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medium

Propose an experimental procedure using chlorine water (

\text{Cl}_2(\text{aq})

) and an organic solvent like

\text{CCl}_4

to distinguish between aqueous solutions of sodium bromide (NaBr) and sodium iodide (NaI). Justify the underlying redox principles.

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medium

Calculate the oxidation number of the central carbon atom (marked with *) in the structure of carbon suboxide (

O=\stackrel{+2}{C}=\stackrel{*}{C}=\stackrel{+2}{C}=O

).

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medium

Identify the reducing agent in the reaction:

3CuO(s) + 2NH_3(g) \rightarrow 3Cu(s) + N_2(g) + 3H_2O(l)

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medium

Describe the difference between an oxidant and a reductant in a redox reaction.

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medium

Balance the following redox reaction in an acidic medium using the half-reaction method:

MnO_4^-(aq) + C_2O_4^{2-}(aq) \rightarrow Mn^{2+}(aq) + CO_2(g)

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medium

Using Stock notation, represent the compound Manganese(VII) oxide.

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medium

Hydrogen peroxide (

\text{H}_2\text{O}_2

) can act as both an oxidizing and a reducing agent. Formulate two separate balanced net ionic equations to justify this dual behavior. In the first, it oxidizes

\text{Fe}^{2+}

to

\text{Fe}^{3+}

in an acidic solution. In the second, it reduces

\text{MnO}_4^-

to

\text{Mn}^{2+}

in an acidic solution. Critique why

\text{H}_2\text{O}_2

acts as a reducing agent in one case and an oxidizing agent in the other.

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medium

Formulate a balanced net ionic equation for the redox titration of ferrous ions (

\text{Fe}^{2+}

) with dichromate ions (

\text{Cr}_2\text{O}_7^{2-}

) in an acidic medium, resulting in ferric ions (

\text{Fe}^{3+}

) and chromic ions (

\text{Cr}^{3+}

). A student suggests that since the dichromate solution is orange and the chromic ion solution is green, the endpoint can be detected without an indicator. Critique this suggestion and propose a more reliable method for endpoint detection in this specific titration.

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medium

Critique the classical definition of oxidation (addition of oxygen/electronegative element) using the reaction

2\text{Na(s)} + \text{H}_2(\text{g}) \rightarrow 2\text{NaH(s)}

. Justify why the electron transfer concept is a superior model for this specific case.

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medium

Evaluate the two methods for balancing redox reactions (Oxidation Number vs. Half-Reaction). Propose a specific type of reaction where the half-reaction method is clearly superior and justify your choice.

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medium

Name the specific type of redox reaction in which an element in one oxidation state is simultaneously oxidized and reduced.

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medium

Balance the following redox reaction using the oxidation number method in a basic medium. Also, identify the oxidant and reductant.

Cr(OH)_3(s) + IO_3^-(aq) \rightarrow I^-(aq) + CrO_4^{2-}(aq)

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hard

Design a series of experiments to identify the contents of three unlabelled beakers containing aqueous solutions of silver nitrate (

\text{AgNO}_3

), zinc sulfate (

\text{ZnSO}_4

), and magnesium sulfate (

\text{MgSO}_4

). You are provided with strips of copper (Cu) and zinc (Zn) metal. Justify your entire procedure based on the principles of competitive electron transfer and the relative positions of these metals in the electrochemical series (

E^\ominus_{\text{Mg}^{2+}/\text{Mg}} < E^\ominus_{\text{Zn}^{2+}/\text{Zn}} < E^\ominus_{\text{Cu}^{2+}/\text{Cu}} < E^\ominus_{\text{Ag}^{+}/\text{Ag}}

).

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hard

A reaction involves permanganate ion (

MnO_4^−

) reacting with hydrogen peroxide (

H_2O_2

) in an acidic solution to produce manganese(II) ion (

Mn^{2+}

) and oxygen gas (

O_2

). (a) Write the balanced half-reactions for oxidation and reduction. (b) Write the overall balanced net ionic equation. (c) Analyze the role of hydrogen peroxide in this reaction and compare it with its role in the reaction

2H_2O_2(aq) \rightarrow 2H_2O(l) + O_2(g)

.

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hard

A student attempts to balance the following reaction in a basic medium using the half-reaction method and presents the final incorrect answer. Critique the student's final equation, identify the likely errors in their process, and formulate the correctly balanced equation.

Student's Incorrect Equation:

2\text{MnO}_4^-(\text{aq}) + 3\text{I}^-(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightarrow 2\text{MnO}_2(\text{s}) + \text{I}_3^-(\text{aq}) + 2\text{OH}^-(\text{aq})

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hard

Explain the paradox of fractional oxidation numbers. Describe the actual oxidation states of the sulfur atoms in the tetrathionate ion,

\mathrm{S_4O_6^{2-}}

, and show how the fractional value is obtained.

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hard

Balance the following equation using the ion-electron (half-reaction) method in an acidic medium:

Zn(s) + NO_3^-(aq) \rightarrow Zn^{2+}(aq) + NH_4^+(aq)

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hard

Calculate the oxidation number of iron in

Fe_3O_4

. Explain why this compound gives a fractional average oxidation number and what the actual oxidation states of the iron atoms are, given that it is a mixed oxide of

FeO

and

Fe_2O_3

.

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hard

Chlorine gas is passed through a hot concentrated solution of sodium hydroxide. The reaction results in the formation of sodium chloride and sodium chlorate. Write the balanced chemical equation for this disproportionation reaction.

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hard

Propose a reason why in the reaction of

\text{Pb}_3\text{O}_4

with HCl,

\text{Pb(IV)}

is reduced to

\text{Pb(II)}

, but with

\text{HNO}_3

, it remains as

\text{PbO}_2

.

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hard

The oxoanions of chlorine are

\text{ClO}^-

,

\text{ClO}_2^-

,

\text{ClO}_3^-

, and

\text{ClO}_4^-

. Evaluate their ability to undergo disproportionation. For each species that can disproportionate, formulate a balanced chemical equation for the reaction. Justify why one of them cannot.

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hard

The reaction of white phosphorus (

P_4

) with aqueous NaOH is a disproportionation reaction. Formulate the balanced ionic equation for this reaction, where the products are phosphine gas (

\text{PH}_3

) and the hypophosphite ion (

\text{H}_2\text{PO}_2^-

). Justify its classification as disproportionation by assigning oxidation numbers.

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hard

Explain what is meant by a disproportionation reaction. Use the decomposition of hydrogen peroxide as an example to illustrate your explanation.

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Redox ReactionsClass 11 Chemistry Practice Questions