Oxidation And Reduction Reactions MCQ Quiz - Objective Question with Answer for Oxidation And Reduction Reactions - Download Free PDF

Concept:

Disproportionation reaction - A type of redox reaction in which the same element or species is oxidised and reduced simultaneously to give two different products is called disproportionation reaction.

• These are also called dismutation reaction.
• An element undergoing disproportionation reaction must show minimum 3 different oxidation state.
• The element should be less stable in a particular oxidation state.

Explanation:

(i)

• This is only a redox reaction.
• In this reaction Cl gets reduced to Cl^- and I gets oxidized to I₂.

(ii)

• This is a disproportionation reaction.
• In this reaction Cl is simultaneously reduced and oxidised to Cl^- and ClO^- respectively.

(iii) 

• This is only a redox reaction.
• In this reaction, H(+1) gets reduced to H(0) and Mg(0) gets oxidized to Mg(+2).

(iv)

• This is the decomposition reaction of hydrogen peroxide also a disproportionation reaction.
• Here oxygen shows disproportionation.
• In peroxide 'O' is in -1 oxidation state, gets oxidised into 0 oxidation state in 'O₂' and reduced to -2 oxidation state in H₂O.

Conclusion:

Therefore only (ii) and (iv) are the disproportionation reaction.

Correct answer: 1) 

Concept:

• The titrations involving redox reaction are called redox titrations. 
• A reaction, which involves simultaneous oxidation and reduction, is called a redox reaction.
• A titration is a technique where a solution of known concentration is used to determine the concentration of an unknown solution.
• Typically, the titrant (the know solution) is added from a burette to a known quantity of the analyte (the unknown solution) until the reaction is complete.
• Knowing the volume of titrant added allows the determination of the concentration of the unknown.
• Often, an indicator is used to usually signal the end of the reaction, the endpoint.
• Oxalic acid acts as a reducing agent, it can be titrated against potassium permanganate in the acidic medium.

Explanation:

• Oxalic acid can be oxidized to carbon dioxide by a suitable oxidizing agent like aqueous potassium permanganate.
• KMnO4 is an oxidizing agent which works in an acidic medium more strongly than an alkaline medium.
• So, for quantitative analysis potassium permanganate is generally used in an acidic medium only. 
• The reaction will be faster in aqueous HCl because chlorine ion present HCl get oxidised into chlorine gas by KMnO₄.
• So, the amount of KMnO₄ will be more as it oxidises both oxalic acid and chlorine ion when HCl is present.

Conclusion:

Thus, the reaction will be faster with aq.HCl

The correct answer is Carbon Monoxide.

Key Points

• In a blast furnace, carbon monoxide (CO) is the primary reducing agent used to convert iron ore (Fe₂O₃) into metallic iron (Fe).
• Carbon monoxide is generated inside the blast furnace by the reaction of coke (carbon) with oxygen at high temperatures.
• The reduction reaction involves carbon monoxide reducing iron oxide as:
  Fe₂O₃ + 3CO → 2Fe + 3CO₂.
• This process is crucial for the extraction of iron in the steel-making industry, where the molten iron is further processed into steel.
• The use of carbon monoxide ensures efficient reduction and separation of iron from impurities in the ore.

Additional Information

• Blast Furnace:
  • A blast furnace is a large, vertical furnace used for smelting to produce industrial metals, primarily iron.
  • It operates at high temperatures, typically over 1,500°C, and uses raw materials like iron ore, coke, and limestone.
• Role of Coke:
  • Coke, derived from coal, acts both as a fuel and as a reducing agent in the blast furnace.
  • It reacts with oxygen to produce carbon monoxide, which is essential for the reduction process.
• Reduction Process:
  • The reduction process removes oxygen from iron oxides (hematite and magnetite) to obtain metallic iron.
  • Carbon monoxide plays a key role in this step, facilitating the conversion of Fe₂O₃ into Fe.
• Slag Formation:
  • Limestone (CaCO₃) is added to the furnace to act as a flux, combining with impurities like silica to form slag.
  • The slag, being less dense than molten iron, floats on top and is removed, leaving behind pure iron.
• Environmental Impact:
  • The blast furnace process produces carbon dioxide (CO₂) as a byproduct, contributing to greenhouse gas emissions.
  • Efforts are ongoing to develop more sustainable methods to reduce emissions during iron production.

Option 1 is correct.

Key Points

• Electrons are lost in an oxidation reaction. 
• Oxidation simply means a gain of oxygen and an oxidizing agent is a substance that oxidizes something.
• The reduction is the gain of electrons. A reducing agent reduces something in a reaction.
• As when one molecule gets reduced, another molecule gets oxidized. Most of the time, Oxidation and Reduction take place simultaneously and such a reaction is called a redox reaction.
• Redox reaction example- In a reaction between hydrogen and fluorine, hydrogen is oxidized and fluorine is reduced.

Correct answer: 1)

Concept:

• Disproportionation reaction: Disproportionation reaction is a special type of redox reaction in which an element in one oxidation state is simultaneously oxidized and reduced.
• Disproportionation reaction, also sometimes called dismutation reaction.

Explanation:

• A disproportionation reaction is a type of redox reaction involving the simultaneous reduction and oxidation of atoms of the same element from one oxidation state (OS) to two different oxidation states. 

• In this reaction, N is both oxidized as well as reduced since the oxidation number of N increases from +4 in \(NO_{3}^{-}\) to +5 in NO2​ and decreases from +4 in NO to +3 in \(NO_{2}^{-}\)
• Thus this is an example of a disproportionation reaction. 

​Conclusion:

Thus, 2NO2 + 2OH– → \(\rm NO^-_2\) + \(\rm NO^-_3\) + H2O is a disproportionation reaction.

Option 4 is correct.

Key Points

• Electrons are lost in an oxidation reaction. 
• Oxidation simply means a gain of oxygen and an oxidizing agent is a substance that oxidizes something.
• The reduction is the gain of electrons. A reducing agent reduces something in a reaction.
• As when one molecule gets reduced, another molecule gets oxidized. Most of the time, Oxidation and Reduction take place simultaneously and such a reaction is called a redox reaction.
• Redox reaction example- In a reaction between hydrogen and fluorine, hydrogen is oxidized and fluorine is reduced.

The correct answer is  Rust and iron have the same composition.

Key Points

• A reddish-brown deposit called rust, forms over a piece of iron when it is exposed to moist air for some time. Rust is hydrated iron (III) oxide (Fe2O3.xH2O).
• When iron metal is exposed to air for a long time, it oxidises and forms a reddish-brown colour iron oxide on its surface.
• The iron rusting formula is 4Fe + 3O2 + 6H2O → 4Fe(OH)3.

Additional Information

• The slow conversion of iron into its hydrated oxide, in the presence of moisture and air is called rusting, whereas the hydrated oxide of iron is called rust.
• When some grease or oil is applied to the surface of an iron object, then air and moisture cannot come in contact with it and hence rusting is prevented.
• The most common method of preventing the rusting of iron (or corrosion of iron) is to coat its surface with paint.
• Galvanization is the process of coating a metal with a zinc layer for protection. It is a widely used technique for protecting iron from rusting.

The correct answer is Option 2. 

Key Points Oxidation is the loss of hydrogen. Reduction is the gain of hydrogen.

• Reduction Reaction is always accompanied by an oxidation reaction where the reactant loses one or more electron.
• Reduction-oxidation reactions are often called redox equations.
• A reduction reaction is only one half of a redox reaction. The other half is the oxidation reaction.​

Additional Information

•  Oxidation Reaction refers to a reaction in which either the addition of Oxygen takes place or the removal of Hydrogen takes place.
• Rusting is an example of an oxidation reaction.
  • The iron reacts with water and oxygen to form hydrated iron oxide, which we see as rust.
    iron + water + oxygen → hydrated iron oxide
• Decomposition is the process of breakdown of the complex organic matter into a simpler inorganic matter like carbon dioxide, water, and nutrients.  The fungi, bacteria, and flagellates initiate the process of decomposition and are known as decomposers.

Explanation:

• Originally, the term oxidation was used to describe the addition of oxygen to an element or a compound.
• Because of the presence of dioxygen in the atmosphere (~20%), many elements combine with it and this is the principal reason why they commonly occur on the earth in the form of their oxides.
• The following reactions represent oxidation processes according to the limited definition of oxidation:
  1. 2 Mg (s) + O₂ (g) → 2 MgO (s) 
  2. S (s) + O₂ (g) → SO₂ (g)
• To summarise, the term “oxidation” is defined as the addition of oxygen /electronegative element to a substance or removal of hydrogen/ electropositive element from a substance.

Concept:

• The more positive the value of reduction potential, the greater will be its ability to be reduced.
• If the difference between the reduction potentials of any two elements is minimal then they will not undergo redox reduction because both will have the same potential to be reduced.
• The species which are strong oxidants have greater tendency to be reduced.
• Those species which have less positive value of electrode potential are strong reducing agents.
• The strongest reducing agent means that it has the greater tendency to be oxidized.

EXPLANATION:

For a reaction to be possible, the value of the reduction potential of a chemical reaction must be positive.

Has been given,EΘ value:

Fe3+/Fe2+ = + 0.77; I2 /I- = + 0.54;

Cu2+/Cu = + 0.34; Ag+ /Ag = + 0.80 V

Option 1: Fe3+  and I-

2Fe3+ +2I-→2Fe2+ + I2

\(E_{Cell}^{\circ }= E_{Fe^{3+}/Fe^{2+}}^{\circ }-E_{I_{2}/I^{-}}^{\circ } \)

=0.77-(0.54)= +0.23 V

Thus, the value is positive. Therefore, the reaction is possible.

Option 2:  Ag+ and Cu

Cu +2Ag+ → Cu2+ +2Ag

\(E_{Cell}^{\circ }= E_{Ag^{+}/Ag}^{\circ }-E_{Cu^{2+}/Cu}^{\circ } \)

= 0.80- (0.34)= +0.46 V

Thus, the value is positive. Therefore, the reaction is possible.

Option 3: Fe³⁺ and Cu

2Fe3+ +Cu → 2Fe2+ + Cu2+

\(E_{Cell}^{\circ }= E_{Fe^{3+}/Fe^{2+}}^{\circ }-E_{Cu^{2+}/Cu}^{\circ } ​​\)

=0.77-(0.34)= 0.43 V

Thus, the value is positive. Therefore, the reaction is possible.

Option 4 : Ag and Fe³⁺

Ag +Fe3+ → Ag+ +Fe2+

\(E_{Cell}^{\circ }= E_{Fe^{3+}/Fe^{2+}}^{\circ }-E_{Ag^{+}/Ag}^{\circ } \)

=0.77- (-0.80)= -0.03 V

Since, the value is negative. Hence, the reaction is not possible.

conclusion:

Hence, the value in option 4 is negative. Therefore, the reaction is not possible

Correct option: 4)

Correct answer: 4)

Concept:

• Disproportionation reaction: Disproportionation reaction is a special type of redox reaction in which an element in one oxidation state is simultaneously oxidized and reduced.
• Disproportionation reaction, also sometimes called dismutation reaction.

Explanation:

• A disproportionation reaction is a type of redox reaction involving the simultaneous reduction and oxidation of atoms of the same element from one oxidation state (OS) to two different oxidation states. 

• In this reaction, N is both oxidized as well as reduced since the oxidation number of N increases from +4 in \(NO_{3}^{-}\) to +5 in NO2​ and decreases from +4 in NO to +3 in \(NO_{2}^{-}\)
• Thus this is an example of a disproportionation reaction. 

​Conclusion:

Thus, 2NO2 + 2OH– → \(\rm NO^-_2\) + \(\rm NO^-_3\) + H2O is a disproportionation reaction.

Concept:

Disproportionation reaction - 

• The reaction in which one element is simultaneously oxidized or reduced is called the disproportionate reaction.
• In a disproportionation reaction, a single substance gives two products one is oxidized and the other is reduced.
• The common example of a disproportionation reaction is - H₂O₂  → H₂O + O₂ 

Explanation:

→ In redox reaction the compound which undergoes oxidation and reaction both is able to show disproportionation tendency.

The given elements are from group 17 of the periodic table and are called halogens.

The general valence shell configuration for halogens is - ns² np⁵.

Thu, total valence electrons in all halogens are seven.

 disproportionation tendency in halogens - 

• All halogens except F show a variable oxidation state between -1 to +7.
• 'F' being small in size and the most electronegative element in the periodic table have no tendency to lose its electron.
• Oxidation state F is -1 only.
• Thus, F can only accept one electron and can undergo reduction only.
• The oxidation of F is not possible.
• Except for the F, all other elements Cl, Br, and I can be easily oxidized and reduced or can lose and accept electrons easily.

Thus, F (Fluorine) does not show the disproportionation tendency.

Conclusion:

Therefore, Fluorine (F)  does not show a disproportionation tendency.

Hence, the correct answer is option 3.

Concept:

Oxidation Reduction in terms of oxidation number: 

• Oxidation: An increase in the oxidation number of an element in a given substance.
• Reduction: A decrease in the oxidation number of the element in a given substance.
• Oxidizing agent(Oxidant): A reagent that can increase the oxidation number of an element in a given substance.
• Reducing agent (Reductants): A reagent which lowers the oxidation number of an element in a given substance.
• Redox reactions: Reactions which involve a change in the oxidation number of the interacting species.

Explanation:

Given reactions:

2\(\rm S_2O^{2-}_3\) + I2 → \(\rm S_4O^{2-}_6\) + 2I–

 \(\rm S_2O^{2-}_3\) + 2Br2 + 5H2O → \(\rm 2SO^{2-}_4\) + 2Br– + 10 H+ 

In reaction 1 when thiosulphate gets oxidised with iodine,  the oxidation number of sulphur in \(\rm S_2O^{2-}_3\) in reactant side is +2 and the oxidation number of sulphur In \(\rm S_4O^{2-}_6\) in product, side is 2.5.

In reaction 2 when thiosulphate gets oxidised with bromine,  the oxidation number of sulphur in \(\rm S_2O^{2-}_3\) in reactant side is +2 and the oxidation number of sulphur in \(\rm 2SO^{2-}_4\) in the product, side is +6.

Thus, Br₂​ is a stronger oxidizing agent than I₂​, as it oxidizes sulphur from a lower oxidation state(+2) to a higher oxidation state(+6). Iodine is a weaker oxidizing agent and is not able to oxidize Sulphur to a higher oxidation state.

Correct answer: 1)

Conclusion:

Thus, Bromine oxidises sulphur to a higher oxidation state and therefore it is a stronger oxidizing agent than iodine. Additional Information

Correct answer: 4)

Concept:

• Redox reactions involve the oxidation and reduction of reacting species at the same time.
• An oxidation-reduction (redox) reaction is a type of chemical reaction that involves the transfer of electrons between two reacting species.
• An oxidation-reduction reaction is a chemical reaction in which the change in the oxidation number of a molecule, atom, or ion takes place by gaining or losing an electron.
• Redox reactions are very common and play a vital role in the basic functions of life, including photosynthesis, respiration, combustion, and corrosion or rusting.

Explanation:

• Option 1:
  • CuO + H2 → Cu + H2O
  • Reactant side:
  • The oxidation number of Cu=+2
  • The oxidation number of O=-2
  • The oxidation number of H=0
  • On the product side:
  • The oxidation number of Cu=0
  • The oxidation number of O=-2
  • The oxidation number of H=+1
  • So, here, oxidation of hydrogen (0 to +1)and reduction of copper (+2 to 0) takes place. Therefore, it is a redox reaction. 
• Option 2:
  • Fe2O3 + 3CO → 2Fe + 3CO2
  • Reactant side:
  • The oxidation number of Fe=+3
  • The oxidation number of O=-2
  • The oxidation number of C=+2
  • On the product side:
  • The oxidation number of Fe=0
  • The oxidation number of O=-2
  • The oxidation number of C=+4
  • So, here, oxidation of carbon (+2 to +4) and reduction of iron(+3 to 0) takes place. Therefore, it is a redox reaction. 
• Option 3:
  • 2K + F2 → 2KF
  • Reactant side:
  • The oxidation number of K=0
  • The oxidation number of F=0
  • On the product side:
  • The oxidation number of Fe=0
  • The oxidation number of K=+1
  • The oxidation number of F=-1
  • So, here, oxidation of potassium (0 to +1)and reduction of fluorine(0 to -1) takes place. Therefore, it is a redox reaction. 
• Option 4:
  • BaCl2 + H2SO4 → BaSO4 + 2HCl is not a redox reaction as it does not involve any change in oxidation number.
  • It is not a redox reaction.
  • It is an example of a double displacement reaction.

Conclusion:

Therefore, option 4: BaCl2 + H2SO4 → BaSO4 + 2HCl is not an example of redox reaction.

Additional Information 

Concept:

Anode - 

• Anode is a positive electrode.
• Normally oxidation occurs at the anode in an electrochemical reaction i.e. it will lose electrons.
• Thus, anode act as a stronger reducing agent than hydrogen when connected to standard hydrogen electrode.
• Therefore, EΘ value for anode will be less than zero or negative because it is having negative electrode potential.
• For example, Al/Al3+  EΘ = –1.66 will act as anode due to having -ve electrode potential.

Explanation:

→ The electrode which has more negative standard reduction potential as compare to standard hydrogen electrode. As they are stronger reducing agent than H₂ gas.

Therefore, 

• Al/Al3+  EΘ = –1.66
• Fe/Fe2+  EΘ = – 0.44

will act as anode as they both have negative reduction potential i.e. EΘ< 0.

→ The electrode which has more positive standard reduction potential as compared to standard hydrogen electrode.

Therefore, 

• Cu/Cu2+  EΘ = + 0.34
• F2 (g)/2F– (aq)  EΘ = + 2.87

will act as cathode as they both have positive reduction potential i.e. EΘ> 0.

Conclusion:

Hence, 

• Al/Al3+  EΘ = –1.66
• Fe/Fe2+  EΘ = – 0.44

will act as anodes, when connected to Standard Hydrogen Electrodes.

\therefore, the correct answers are option 1 and option 2.