Chapter 7: Redox Reactions Short Questions | chemistry Class 11 CBSE

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Chapter 7: Redox Reactions Short Questions | chemistry Class 11 CBSE | ByteNotes.in

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Question 1

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Define oxidation and reduction in terms of oxygen and hydrogen transfer.

Oxidation is the addition of oxygen or an electronegative element to a substance, or removal of hydrogen or an electropositive element from a substance.
Reduction is the removal of oxygen or an electronegative element from a substance, or addition of hydrogen or an electropositive element to a substance.
Both processes always occur simultaneously; the substance that gets oxidised is the reducing agent and the one that gets reduced is the oxidising agent.

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Question 2

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Define oxidation and reduction in terms of electron transfer.

Oxidation is the loss of one or more electrons by any species in a chemical reaction.
Reduction is the gain of one or more electrons by any species in a chemical reaction.
An oxidising agent is an electron acceptor, while a reducing agent is an electron donor.

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Question 3

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State the rules for assigning oxidation number to oxygen and hydrogen in their compounds.

The oxidation number of oxygen is -2 in most compounds; exceptions are peroxides (e.g., H2O2) where it is -1, superoxides (e.g., KO2) where it is -1/2, and in OF2 where it is +2.
The oxidation number of hydrogen is +1 in most compounds, but is -1 when bonded to metals in binary compounds such as LiH, NaH, and CaH2.
Fluorine always has an oxidation number of -1 in all its compounds because it is the most electronegative element.

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Question 4

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What is Stock notation? Give two examples.

Stock notation, devised by Alfred Stock, expresses the oxidation number of a metal using a Roman numeral in parentheses immediately after the symbol of the metal in the molecular formula.
Aurous chloride (AuCl) is written as Au(I)Cl and auric chloride (AuCl3) is written as Au(III)Cl3 in Stock notation.
Similarly, stannous chloride (SnCl2) is written as Sn(II)Cl2 and stannic chloride (SnCl4) is written as Sn(IV)Cl4.

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Question 5

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What is a disproportionation reaction? Give one example from the chapter.

A disproportionation reaction is a special type of redox reaction in which an element present in a single intermediate oxidation state is simultaneously oxidised and reduced in the same reaction.
The reacting substance must contain an element that can exist in at least three oxidation states; the element in the reacting form is in the intermediate state.
The decomposition of hydrogen peroxide, 2H2O2 → 2H2O + O2, is a classic example where oxygen in the -1 state is simultaneously reduced to -2 (in H2O) and oxidised to 0 (in O2).

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Question 6

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Calculate the oxidation number of chromium in K2Cr2O7.

In K2Cr2O7, potassium has oxidation number +1 (alkali metal) and oxygen has oxidation number -2 (standard rule).
Applying the rule that the sum of all oxidation numbers in a neutral compound equals zero: 2(+1) + 2x + 7(-2) = 0, giving 2 + 2x - 14 = 0.
Solving: 2x = 12, therefore x = +6; the oxidation number of chromium in K2Cr2O7 is +6.

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Question 7

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Explain why the reaction 2Na(s) + H2(g) → 2NaH(s) is a redox reaction.

In NaH, the compound is ionic and can be represented as Na+H-, indicating that hydrogen has an oxidation number of -1.
The half-reaction for sodium shows oxidation: 2Na(s) → 2Na+(g) + 2e- (sodium loses electrons, oxidation number changes from 0 to +1).
The half-reaction for hydrogen shows reduction: H2(g) + 2e- → 2H-(g) (hydrogen gains electrons, oxidation number changes from 0 to -1); therefore, the reaction involves both oxidation and reduction, making it a redox reaction.

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Question 8

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What is a combination redox reaction? Give one example with oxidation numbers.

A combination reaction is of the type A + B → C, where at least one of the reactants A or B must be in elemental form for it to qualify as a redox reaction.
The combustion of carbon: C(s) + O2(g) → CO2(g) is a combination redox reaction where carbon is oxidised from 0 to +4 and oxygen is reduced from 0 to -2.
All combustion reactions using elemental dioxygen, as well as reactions like 3Mg(s) + N2(g) → Mg3N2(s), are examples of combination redox reactions.

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Question 9

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Why does fluorine not show disproportionation in alkaline medium unlike other halogens?

Fluorine is the most electronegative element in the periodic table and cannot exhibit any positive oxidation state.
Disproportionation requires an element to exist in at least three oxidation states (a lower, intermediate, and higher state), but fluorine cannot achieve a positive oxidation state.
Therefore, when fluorine reacts with alkali, the reaction 2F2(g) + 2OH-(aq) → 2F-(aq) + OF2(g) + H2O(l) occurs instead of a disproportionation reaction.

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Question 10

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What is a redox couple? Give an example from the Daniell cell.

A redox couple consists of the oxidised and reduced forms of a substance that together participate in either an oxidation half-reaction or a reduction half-reaction.
The two forms are separated by a vertical line (slash) representing the interface, with the oxidised form written before the reduced form.
In the Daniell cell, the two redox couples are Zn2+/Zn (at the anode, where zinc is oxidised) and Cu2+/Cu (at the cathode, where copper ions are reduced).

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Question 11

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Define standard electrode potential. What is the standard electrode potential of the hydrogen electrode?

Standard electrode potential (E°) is the potential of an electrode measured at 298 K when all ionic species involved in the electrode reaction are at unit concentration (1 mol L-1), and any gas is at 1 atmospheric pressure.
By convention, the standard electrode potential of the hydrogen electrode (H+/H2 couple) is assigned a value of exactly 0.00 volts.
A negative E° value for a redox couple indicates it is a stronger reducing agent than H+/H2, while a positive E° value indicates it is a weaker reducing agent than H+/H2.

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Question 12

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In a Daniell cell, what is the role of the salt bridge?

The salt bridge is a U-tube containing an electrolyte solution such as potassium chloride or ammonium nitrate, usually solidified with agar-agar to form a jelly-like substance.
It provides electrical contact between the two half-cell solutions, allowing ions to migrate and complete the electrical circuit inside the cell.
At the same time, the salt bridge prevents the two solutions from mixing with each other, which is essential for maintaining the potential difference between the electrodes.

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Question 13

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Identify the oxidising agent and reducing agent in the reaction: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s).

In this reaction, zinc (Zn) loses electrons and is oxidised from oxidation state 0 to +2, forming Zn2+ ions; therefore zinc is the reducing agent (reductant).
Copper ions (Cu2+) gain electrons and are reduced from oxidation state +2 to 0, depositing as metallic copper; therefore Cu2+ is the oxidising agent (oxidant).
The equilibrium for this reaction greatly favours the products, indicating that zinc has a much greater tendency to lose electrons compared to copper.

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Chapter 7: Redox Reactions | Short Questions

Define oxidation and reduction in terms of oxygen and hydrogen transfer.

Define oxidation and reduction in terms of electron transfer.

State the rules for assigning oxidation number to oxygen and hydrogen in their compounds.

What is Stock notation? Give two examples.

What is a disproportionation reaction? Give one example from the chapter.

Calculate the oxidation number of chromium in K2Cr2O7.

Explain why the reaction 2Na(s) + H2(g) → 2NaH(s) is a redox reaction.

What is a combination redox reaction? Give one example with oxidation numbers.

Why does fluorine not show disproportionation in alkaline medium unlike other halogens?

What is a redox couple? Give an example from the Daniell cell.

Define standard electrode potential. What is the standard electrode potential of the hydrogen electrode?

In a Daniell cell, what is the role of the salt bridge?

Identify the oxidising agent and reducing agent in the reaction: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s).

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