Electrochemistry MCQ Quiz - Objective Question with Answer for Electrochemistry - Download Free PDF

CONCEPT:

Calculation of Cell Potential using Nernst Equation

• The Nernst equation relates the cell potential (Ecell) to the standard cell potential (Ecello), the temperature, and the concentrations of the species involved in the electrochemical reaction.
• The Nernst equation is given by:

  Ecell = Ecello - (0.0592/n) log([Mn+]/[M])

  where:
  • Ecello is the standard electrode potential (in this case, 0.2 V)
  • n is the number of moles of electrons transferred (in this case, n = 2)
  • [Mn+] and [M] are the molar concentrations of the ions involved in the redox reaction.

EXPLANATION:
Overall reaction :–

\(\mathrm{H}_{2(\mathrm{~g})}+\mathrm{M}_{(\mathrm{aq})}^{3+} \longrightarrow \mathrm{M}_{(\mathrm{aq})}^{+}+2 \mathrm{H}_{(\mathrm{aq})}^{+} \)

\(\mathrm{E}_{\text {Cell }}=\mathrm{E}_{\text {Cathode }}^{\mathrm{o}}-\mathrm{E}_{\text {anode }}^{\mathrm{o}}-\frac{0.059}{2} \log \frac{\left[\mathrm{M}^{+}\right] \times 1^{2}}{\left[\mathrm{M}^{+3}\right] 1} \)

\(0.1115=0.2-\frac{0.059}{2} \log \frac{\left[\mathrm{M}^{+}\right]}{\left[\mathrm{M}^{+3}\right]} \)

\(3=\log \frac{\left[\mathrm{M}^{+}\right]}{\left[\mathrm{M}^{+3}\right]} \)

∴ a = 3

Therefore, the value of log([M³⁺]/[M]) is 3.

CONCEPT:

Reducing Agent Strength in Alkali Metals

• The strength of a reducing agent is determined by its ability to lose electrons easily — i.e., undergo oxidation.
• Among alkali metals (Group 1 elements), reducing strength increases as we move down the group.
• This is because the ionization energy decreases down the group, making it easier for atoms to lose electrons.
• The standard oxidation potential (E°) is used to determine reducing power — the more negative the value, the stronger the reducing agent.

EXPLANATION:

• The standard oxidation potentials of the given alkali metals are approximately:
  • Li: –3.04 V
  • Na: –2.71 V
  • K: –2.93 V
  • Rb: –2.92 V
• Although lithium has the most negative standard potential, in aqueous solution, sodium is the weakest reducing agent due to solvation effects that stabilize Na⁺ more than Li⁺.
• Thus, Na is less willing to lose its electron compared to other alkali metals under standard conditions.

Therefore, the correct answer is Na

CONCEPT:

Cell Potential Calculation from Gibbs Energy

\(\Delta G_{\text{rxn}}^{\circ} = -nFE^\circ \)

• The cell potential E^\circ for the combustion of butane is related to the standard Gibbs energy of the reaction using the equation:
  • n = Number of moles of electrons (8 electrons in this case from the stoichiometry of the reaction),
  • F = Faraday constant = 96485 C/mol,
  • E^\circ = Standard cell potential.
• The standard Gibbs energy \Delta G_{\text{rxn}}^{\circ} is calculated by summing the Gibbs energies of formation of the products and reactants.

EXPLANATION:

• The given standard Gibbs energies of formation for CO₂, H₂O, and C₄H₁₀ are used to calculate the total \( \Delta G_{\text{rxn}}^{\circ}\) for the reaction.
• Using\( \Delta G_{\text{rxn}}^{\circ} = -2743 \, \text{kJ} \), the cell potential is found to be:
• Thus, X = 105.5 .

The value of X is 105.5.

CONCEPT:

Nernst Equation for a Galvanic Cell

• The standard cell potential is calculated as:

  Ecell = E0cathode − E0anode

• The Nernst equation for the cell at 298 K is:

  Ecell = E0cell − (0.0591/n) log Q

• For the given reaction:

  ½ H2(g) + Fe3+(aq) → H+(aq) + Fe2+(aq)

  • n = 1 (1 electron transferred)
  • E0cell = E0Fe3+/Fe2+ − E0H+/H2 = 0.771 V − 0 V = 0.771 V

EXPLANATION:

\(\frac{1}{2} \mathrm{H}_{2}(\mathrm{~g})+\mathrm{Fe}^{3+}(\mathrm{aq} .) \longrightarrow \mathrm{H}^{+}(\mathrm{aq})+\mathrm{Fe}^{2+}(\mathrm{aq} .)\)

\(\mathrm{E}=\mathrm{E}^{\circ}-\frac{0.059}{1} \log \frac{\left[\mathrm{Fe}^{2+}\right]}{\left[\mathrm{Fe}^{3+}\right]}\)

⇒ \(0.712=(0.771-0)-\frac{0.059}{1} \log \frac{\left[\mathrm{Fe}^{2+}\right]}{\left[\mathrm{Fe}^{3+}\right]}\)

⇒ \(\log \frac{\left[\mathrm{Fe}^{2+}\right]}{\left[\mathrm{Fe}^{3+}\right]}=\frac{(0.771-0712)}{0.059}=1\)

⇒ \(\frac{\left[\mathrm{Fe}^{2+}\right]}{\left[\mathrm{Fe}^{3+}\right]}=10\)

Therefore, the ratio [Fe²⁺] : [Fe³⁺] is 10.

CONCEPT:

Electrochemical Reaction and Current Calculation

It = nF

• The reaction involves the reduction of dichromate ions (Cr₂O₇²⁻) in an aqueous acidic medium to Cr³⁺. The number of moles of electrons required for the reduction of 1 mole of Cr³⁺ is 3 moles, as per the balanced reaction.
• Faraday’s law relates the amount of charge passed in an electrochemical reaction to the number of moles of electrons involved and the current flowing through the circuit:
  • Where:
    • I = Current (in amperes)
    • t = Time (in seconds)
    • n = Number of moles of electrons (3 moles for Cr³⁺)
    • F = Faraday’s constant = 96,500 C/mol

EXPLANATION:

• We are given that 3 moles of electrons are required to produce 1 mole of Cr³⁺.
• The time (t) is provided as 48.25 minutes, which needs to be converted to seconds:
  • t = 48.25 × 60 = 2,895 seconds.
• We can now calculate the current (I) using the formula:

  I = (3 × 96,500) / (48.25 × 60)

  I = 100.0 A
• Therefore, the current required for the electrochemical reaction to produce 1 mole of Cr³⁺ in 48.25 minutes is 100.0 A.

Therefore, the current is 100.0 A.

Bleaching powder by its nature is an Oxidising agent.

• Stable bleaching powder is widely used as a disinfectant in water purification, as well as in the textile and pulp and paper industries.
• "Bleaching powder" is made by the action of chloride gas on calcium hydroxide.
• The reaction being essentially:
• 2Ca (OH)₂ + 2Cl₂ → Ca(OCl)₂ + CaCl₂ + 2H₂O.
• In the production of bleaching powder, slaked lime spread on the floors of large rectangular chambers of lead or concrete is exposed to chlorine gas.
• Bleaching powder, a solid combination of chlorine and slaked lime, was introduced in 1799 by Scottish chemist Charles Tennant.

The correct answer is Anodising.

Key Points

• Anodising is the process of forming a thick oxide layer of aluminum.
  • The process is called anodizing because the part to be treated forms the anode electrode of an electrolytic cell.
  • This aluminum oxide coat makes it resistant to further corrosion. 
  • It is also useful in architectural finishing.

Additional Information

• Ductility is the ability of a material to be drawn or plastically deformed without fracture.
  • The ductility of steel varies depending on the types and levels of alloying elements present.
• Galvanisation or galvanization (or galvanizing as it is most commonly called) is the process of applying a protective zinc coating to iron or steel, to prevent rusting.
  • The most common method is hot-dip galvanizing, in which steel sections are submerged in a bath of molten zinc.
• Corrosion is a natural process that converts a refined metal into a more chemically stable form such as oxide, hydroxide, carbonate, or sulfide.
  • It is the gradual destruction of materials (usually a metal) by chemical or electrochemical reactions with their environment.

• Rusting of iron is an example of a Redox reaction. 
• Redox reaction= Oxidation-Reduction reaction.In this reaction, an oxidation number of a molecule, atom, or ion changes either by gaining or losing an electron. 
• The substance which gets reduced in a chemical reaction is known as the oxidizing agent and a substance that gets oxidized in a chemical reaction is known as the reducing agent.
• Rusting is the corrosion of iron. Iron forms red-brown hydrated metal oxide (rust) in the presence of water and air. 
• Iron is oxidized to Fe²⁺ and oxygen is reduced to water. Rust keeps on forming due to the subsequent oxidation of Fe2+ by atmospheric oxygen.
• Therefore we can say that rusting is a redox reaction because oxygen acts as an oxidising agent and iron acts as a reducing agent.

Concept:

Oxidation:

It is the process in which an atom, molecule, or ion loses one or more electrons.

Reduction:

It is the process in which an atom, molecule, or ion gains one or more electrons.

Reducing agent/ Reductant:

• It is one that donates electrons to a species and thereby brings about its reduction.
• The reducing agent is oxidized by having its electrons taken away.
• Eg: In the reaction \(Mn{O_2} + 4HCl \to MnC{l_2} + C{l_2} + 2{H_2}O\), the oxidizing agent is MnO2 and the reducing agent is HCl.

Oxidizing agent/ Oxidizer:

• It is defined as a substance that removes electrons from another reactant in a redox chemical reaction.
• The oxidizing agent is reduced by taking electrons onto itself.
• In the reaction \(2Mg + {O_2} \to 2MgO\), the oxidizing agent is O2 and the reducing agent is Mg.

Explanation:

• As the electronegativity of a species increases, its tendency to pull electrons also increases and eventually behaves as a stronger oxidizing agent.
• Greater the reduction potential, greater will be the tendency of oxidizing agent to get reduced easily and hence acts as a stronger oxidizing agent.

H₂O₂:

• Hydrogen peroxide acts as an oxidizing agent and reducing agent.
• Eg: of reaction in which H₂O₂ acts as an oxidizing agent is as follows: \({H_2}{O_2} + 2{H^ + }_{(aq)} + 2{e^ - } \to 2{H_2}{O_{(l)}};\,{E^0} = + 1.77V\)

O₃:

• Ozone is a strong oxidizing agent.
• It can easily lose nascent oxygen.
• An example of a reaction in which O₃ acting as an oxidizing agent is: \({O_3} + 2{H^ + }_{(aq)} + 2{e^ - } \to {O_2} + 2O{H^ - };\,{E^0}/SRP = + 2.07V\)

K₂Cr₂O₇:

• It acts as an oxidizing agent in an acidic medium and can oxidize reducing agents like ferrous sulfate, nitrite, sulfite, etc.
• In these reactions in acid solution, dichromate is reduced to green Cr³⁺ ions: \(C{r_2}O_7^{2 - } + 14{H^ + }_{(aq)} + 6{e^ - } \to 2C{r^{3 + }} + 7{H_2}O;\,{E^0}/SRP = + 1.33V\)

KMnO₄:​

• It dissolves in water to give an intensely purple solution.
• KMnO₄ is stronger than any other oxidizing agent because it contains Mn in its highest oxidation state +7.
• Elements become more electronegative as the oxidation state of their atoms is increased.
• Their reaction is as follows: \(MnO_4^ - + 8{H^ + }_{(aq)} + 5{e^ - } \to M{n^{2 + }} + 4{H_2}O;\,{E^0}/SRP = + 1.51V\)

The increasing order of reduction potential of the given oxidizing agent is: \({E^o}C{r_2}O_7^{2 - }/C{r^{3 + }}( + 1.33V) < {E^o}Mn{O_4}^ - /M{n^{2 + }}( + 1.51V) < {O_3}/{O_2}( + 1.77V) < {H_2}{O_2}/{H_2}O( + 2.07V)\)

Hence, the increasing order of power of the oxidizing agent is : \({K_2}C{r_2}{O_7} < KMn{O_4} < {H_2}{O_2} < {O_3}\)

Therefore, ozone is the strongest oxidizing agent.

The correct answer is AlCl3 + 3H2O → Al(OH)3 + 3HCl.

Explanation:

• Redox reactions involve the oxidation and reduction of reacting species simultaneously. Thus, a change in oxidation state determines whether or not a reaction is a redox.
• An oxidizing agent (also oxidant) is the element or compound that accepts an electron from another species in an oxidation-reduction (redox) reaction. Because the oxidizing agent is gaining electrons (and thus is frequently referred to as an electron acceptor), it is said to be reduced.
• During a redox reaction, a reductant is a reactant that donates electrons to other reactants.

AlCl3 + 3H2O → Al(OH)3 + 3HCl is not an example of a redox reaction because in this reaction oxidation and reduction do not take place.

Additional Information

2NaH → 2Na + H₂ 
          oxidation-reduction  

4Fe + 3O₂ → 2Fe₂O₃
               oxidation reduction
     
CuSO₄ + Zn → Cu + ZnSO₄
                    reduction  oxidation

Thus, the above 3 reactions are an example of redox reactions because in these reactions oxidation and reduction take place simultaneously.

Explanation:

Kohlrausch's law states that the equivalent conductivity of an electrolyte at infinite dilution is equal to the sum of the conductances of the anions and cations. If a salt is dissolved in water, the conductivity of the solution is the sum of the conductances of the anions and cations.

Calculation:

According to Kohlrausch’s law, the molar conductivity of HA at infinite dilution is given as,

\({\Lambda_m}^\circ \left( {{\rm{HA}}} \right) = \left[ {{\Lambda _m}^\circ \left( {{{\rm{H}}^ + }} \right) + {\Lambda _m}^\circ \left( {{\rm{C}}{{\rm{l}}^ - }} \right)} \right] + \left[ {{\Lambda _m}^\circ \left( {{\rm{N}}{{\rm{a}}^ + }} \right) + {\Lambda _m}^\circ \left( {{A^ - }} \right)} \right] - \left[ {{\Lambda _m}^\circ \left( {{\rm{N}}{{\rm{a}}^ + }} \right) + {\Lambda _m}^\circ \left( {{\rm{C}}{{\rm{l}}^ - }} \right)} \right.\) ]

= 425.9 + 100.5 – 126.4

= 400 S cm² mol^-1

Also, molar conductivity (Λm°) at given concentration is given as,

\(\Lambda_m = \frac{{1000 \times k}}{M}\)

Given, k = conductivity ⟹ 5 × 10^-5 S cm^-1

M = Molarity ⟹ 0.001 M

\({\therefore \Lambda_m} = \frac{{1000 \times 5 \times {{10}^{ - 5}}{\rm{\;Sc}}{{\rm{m}}^{ - 1}}}}{{{{10}^{ - 3}}{\rm{M}}}}\)

= 50 S cm² mol^-1

Therefore, degree of dissociation (α), of HA is,

\(\alpha = \frac{{{\Lambda_m}}}{{{\Lambda_m}^\circ {\rm{\;}}}} = \frac{{50\;S\;c{m^2}mo{l^{ - 1}}}}{{400\;S\;c{m^2}\;mo{l^{ - 1}}}} = 0.125\)

Option 2 is the correct answer: Decomposition reactions are always endothermic in nature.

• In a decomposition reaction, a chemical compound is broken into its constituent components.
• The process takes place through the breaking of bonds between the constituent atoms of the compound.
• The reaction in which heat or light energy is absorbed during the process, are called endothermic reactions.
• In decomposition reactions, the energy is required to break the chemical bonds. therefore they are aways endothermic in nature.
• Combustion reaction - Combustion means reaction with oxygen, hence combustion reactions are generally oxidation reactions (Exothermic).
• Displacement reaction - One constituent component is replaced by some other component (Spontaneous and exothermic).
• Combination reaction - Two or more elements or compounds combine. New bonds are formed and energy is released (Exothermic).

Concept:

• The conductivity is depending on the number of ions present in the unit volume of solution. The decreasing order of the electrical conductivity of the aqueous solution is based on the acid strength.
• The more the acid strength more will be the dissociation of acid into ion and more will be the conductivity.
• Acid strength refers to the tendency of an acid, symbolized by the chemical formula HA, to dissociate into a proton H⁺, and an anion, A⁻. The dissociation of a strong acid in solution is effectively complete, except in its most concentrated solutions.
• Electrical conductivity is the measure of the amount of electrical current a material can carry or its ability to carry a current. Electrical conductivity is also known as specific conductance.

The Order of acidic strength is HCOOH (formic acid) > C₆ H₅ COOH (benzoic acid) > CH₃COOH (Acetic acid).

The correct option is a Double Displacement reaction.

Key PointsDouble Displacement Reaction

• The chemical reaction is in which one component of both of the reacting molecules gets exchanged to form the products.
• In other words, the reaction in which two different atoms or groups of atoms are replaced by other atoms or groups of atoms.
• CuSO₄ + H₂S  → CuS ↓ + H₂SO₄
• In the above reaction on passing hydrogen sulphide gas through an aqueous solution of copper sulphate, a black precipitate of copper sulphide is formed.
• Downward Arrow(↓) indicates the formation of a precipitate.
• Two compounds exchange their ions and one of the products formed is insoluble which is precipitate.
• These reactions usually occur in ionic compounds when dissolved in water.
• These reactions are fast reactions and take place within a fraction of a second.

Additional Information Addition reaction

• An addition reaction is also known as a combination reaction.
• The reaction in which two or more substances (elements are compounds) simply combine to form a new substance.

Decomposition reaction

• A reaction in which a single compound breaks down to produce two or simpler substances is called a decomposition reaction.
• It is opposite to the combination reaction.
• A decomposition reaction is of three types:
• thermal decomposition, electrolytic decomposition and photochemical decomposition reaction.

Displacement reaction

• Those reactions in which more reactive elements can displace less reactive elements from a compound are called displacement reactions.

2e^- + 2H⁺ (aq) → H₂(g)

\(\mathrm{E}=\mathrm{E}^{\circ}-\frac{0.059}{\mathrm{n}} \log \frac{\mathrm{P}_{\mathrm{H}_2}}{\left[\mathrm{H}^{+}\right]^2} \)

\(0=0-\frac{0.059}{2} \log \frac{\mathrm{P}_{\mathrm{H}_2}}{\left(10^{-7}\right)^2} \)

\(\log \frac{\mathrm{P}_{\mathrm{H}_2}}{\left(10^{-7}\right)^2}=0 \)

\(\frac{\mathrm{P}_{\mathrm{H}_2}}{10^{-14}}=1 \)

\(\mathrm{P}_{\mathrm{H}_2}=10^{-14} \text { bar } \)