Nootan solutions for Chemistry Part 1 and 2 [English] Class 12 ISC chapter 3 - Electrochemistry [Latest edition]

Define a galvanic cell.

Define reference electrode.

Define standard electrode potential.

What is understood by a normal hydrogen electrode?

Give the significance of the normal hydrogen electrode.

Define standard electrode potential.

What do you understand by electrochemical series?

How does the electrochemical series help in predicting whether a given reaction is feasible in a given direction or not?

It is impossible to measure the potential of a single electrode. Comment.

What is standard hydrogen electrode?

Give the reaction that occurs at standard hydrogen electrode when it acts as positive electrode in an electrochemical cell.

What is the effect of change in concentration on the electrode potential of a given half-cell?

What is the effect of change in temperature on the electrode potential of a given half cell?

Write the Nernst equation and explain the terms involved.

State the values of `E_(Ca//Ca^(2+))^circ, E_(Ni//Ni^(2+))^circ` and `E_(Cl^-//1/2Cl_2)^circ` if  `E_(Ca^(2+)//Ca)^circ` = −2.87 volts, `E_(Ni^(2+)//Ni)^circ` = −0.25 volts and `E_(1/2Cl_2//Cl^-)^circ` = +1.36 volts.

A potential difference of 0.25 volts is observed when an electrode system M/M²⁺ (1 mol L⁻¹) is connected to a standard hydrogen electrode. If the direction of the flow of current is from hydrogen electrode to the metal electrode, calculate the standard potential of the electrode.

The standard reduction potential of the electrode X/X⁺ (1 mol L⁻¹) is +0.80 volts. Predict the direction of the flow of electrons when it is connected to a standard hydrogen electrode.

Predict whether the following reaction is feasible under standard conditions or not: 

\[\ce{Zn_{(s)} + Cu{^{2+}_{(aq)}} -> Zn{^{2+}_{(aq)}} + Cu_{(s)}}\]

Given that `E_(Zn^(2+)//Zn)^circ` = −0.76 volts and `E_(Cu^(2+)//Cu)^circ` = +0.34 volts.

On the basis of electrochemical series, explain why Fe can displace hydrogen from dilute HCl but silver cannot.

On the basis of electrochemical series, explain why HCl cannot be stored in an aluminium vessel.

On the basis of electrochemical series, explain why Au₂O₃ decomposes on heating.

On the basis of electrochemical series, explain why a coating of copper is deposited on an iron nail when placed in an aqueous solution of copper sulphate.

Can we store copper sulphate solution in zinc vessel? Give suitable explanation. 

Given: \[\ce{E_{Cu^{2+}/Cu}}\] = +0.34 V, \[\ce{E_{Zn^{2+}/Zn}}\] = −0.76 V, \[\ce{E_{Ag^{+}/Ag}}\] = +0.80 V

Can we store copper sulphate solution in silver vessel? Give suitable explanation. 

Given: \[\ce{E{^{\circ}_{Cu^{2+}/Cu}}}\] = +0.34 V, \[\ce{E{^{\circ}_{Zn^{2+}/Zn}}}\] = −0.76 V, \[\ce{E{^{\circ}_{Ag^{+}/Ag}}}\] = +0.80 V.

The two half cell reactions and their oxidation potentials are

1. \[\ce{Pb_{(s)} - 2e- -> Pb{^{2+}_{(aq)}}, E^{\circ}_{Pb/Pb^{2+}}}\] = +0.13 V
2. \[\ce{Ag_{(s)} - e- -> Ag{^+_{(aq)}}, E^{\circ}_{Ag/Ag^{+}}}\] = −0.80 V

Write the cell reaction and calculate the cell emf.

A cell is prepared by dipping a copper rod in 1 M copper sulphate solution and a zinc rod in 1 M zinc sulphate solution. The standard reduction potentials of copper and zinc are 0.34 V and −0.76 V respectively.

1. What will be the cell reaction?
2. What will be the standard emf of the cell?
3. Which electrode will be positive?
4. How can the cell be represented?

Calculate at 25°C the electrode potential of Mg²⁺/Mg electrode in which the concentration of Mg²⁺ ions is 0.1 M. Given \[\ce{E^{\circ}_{Mg^{2+}/Mg}}\] = −2.36 V, R = 8.314 JK⁻¹, F = 96500 coulombs mol⁻¹.

Why does the blue colour of copper sulphate solution get discharged when zinc rod is dipped into it?

Given: \[\ce{E^{\circ}_{Cu^{2+}/Cu}}\] = +0.34 V and \[\ce{E^{\circ}_{Zn^{2+}/Zn}}\] = −0.76 V.

Using the standard electrode potentials, predict if the reaction between the following is feasible:

\[\ce{Fe^{3+}_{ (aq)}}\] and \[\ce{I^-_{ (aq)}}\]

Using the standard electrode potentials, predict if the reaction between the following is feasible:

\[\ce{Ag^+_{ (aq)}}\] and Cu_(s)

Using the standard electrode potentials, predict if the reaction between the following is feasible:

\[\ce{Fe^{3+}_{ (aq)}}\] and \[\ce{Br^-_{ (aq)}}\]

Using the standard electrode potential, predict if the reaction between the following is feasible:

Ag_(s) and \[\ce{Fe^{3+}_{ (aq)}}\]

Using the standard electrode potential, predict if the reaction between the following is feasible:

Br_2(aq) and \[\ce{Fe^{2+}_{ (aq)}}\]

Calculate the electrode potentials of the following half cell at 25°C. 

Pt, Cl₂ (10 atm) | HCl (0.1 M)

Given: \[\ce{E^{\circ}_{\frac{1}{2}Cl_2/Cl^-}}\] = +1.36 V and \[\ce{E{^{\circ}_{H^+/\frac{1}{2}H_2}}}\] = 0.00 V.

Calculate the electrode potentials of the following half cell at 25°C.

Pt, H₂ (5 atm) | HCl (0.5 M)

Given: \[\ce{E^{\circ}_{\frac{1}{2}Cl_2/Cl^-}}\] = +1.36 V and \[\ce{E{^{\circ}_{H^+/\frac{1}{2}H_2}}}\] = 0.00 V.

Calculate the electrode potentials of the following half cells at 298 K.

Ag_(s) | AgNO₃ (0.1 M)

Given: \[\ce{E^{\circ}_{Ag^+/Ag}}\] = +0.80 V, \[\ce{E^{\circ}_{{Co^{2+}/Co}}}\] = −0.28 V

Calculate the electrode potentials of the following half cells at 298 K.

Co_(s) | Co²⁺ (0.01 M)

Given: \[\ce{E^{\circ}_{Ag^+/Ag}}\] = +0.80 V, \[\ce{E^{\circ}_{{Co^{2+}/Co}}}\] = −0.28 V

Write Nernst equation for the following cell reaction:

\[\ce{2 A + 3 B -> C + 2 D}\]

How can the cell potential be used in predicting the feasibility of a cell reaction?

Calculate the EMF of the cell at 298 K:

Pt, Br₂ | Br⁻ (0.01 M) || H⁺ (0.03 M) | H₂ (1 atm), Pt

Given: \[\ce{E^{\circ}_{\frac{1}{2}Br/Br^-}}\] = +1.08 V.

Calculate the standard electrode potential of the Ni²⁺/Ni electrode, if the cell potential of the cell is 0.59 V.

Ni | Ni²⁺ (0.01 M) || Cu²⁺ (0.1 M) | Cu

Given: \[\ce{E^{\circ}_{Cu{^{2+}/Cu}}}\] = +0.34 V

Calculate the emf of the following cell at 298 K:

Mg | Mg²⁺ (0.001 M) || Cu²⁺ (0.0001 M) | Cu

Given: \[\ce{E^{\circ}_{{Cu^{2+}/{Cu}}}}\] = 0.337 V, \[\ce{E^{\circ}_{{Mg^{2+}/{Mg}}}}\] = −2.37 V, F = 96500 C mol⁻¹.

Calculate the EMF of the following cell at 298 K:

Mg | Mg²⁺ (0.1 M) || Ag⁺ (0.1 M) | Ag

\[\ce{E^{\circ}_{{Mg^{2+}/{Mg}}}}\] = −2.37 V, \[\ce{E^{\circ}_{{Ag^{+}/{Ag}}}}\] = +0.80 V, R = 8.31 J K⁻¹, F = 96500 C mol⁻¹.

Consider a cell composed of two half cells:

1. \[\ce{Cu_{(s)} | Cu{^{2+}_{(aq)}}}\] and
2. \[\ce{Ag_{(s)} | Ag{^+_{(aq)}}}\].

Calculate: (a) the standard cell potential, and (b) the cell potential when [Cu²⁺] is 2M and [Ag⁺] is 0.05 M.

Given: \[\ce{E^{\circ}_{{Cu^{2+}/{Cu}}}}\] = +0.33 V, \[\ce{E^{\circ}_{{Ag^{+}/{Ag}}}}\] = +0.80 V, R = 8.31 J K⁻¹ mol⁻¹, F = 96500 C mol⁻¹.

Consider a cell composed of the following two half cells:

1. \[\ce{Mg_{(s)} | Mg{^{2+}_{(aq)}}}\], and
2. \[\ce{Ag_{(s)} | Ag{^{+}_{(aq)}}}\]

The EMF of the cell is 2.96 V, [Mg²⁺] = 0.130 M and [Ag⁺] = 1.0 × 10⁻⁴ M. Calculate the standard EMF of the cell. (R = 8.31 J K⁻¹ mol⁻¹, F = 96500 C mol⁻¹)

Calculate the cell potential at 298 K: 

\[\ce{Zn_{(s)} | Zn^{2+} (0.1 M) || Sn^{2+} (0.001 M) | Sn_{(s)}}\]

(Given: \[\ce{E{^{\circ}_{Zn^{2+}/Zn}}}\] = −0.76 V, \[\ce{E{^{\circ}_{Sn^{2+}/Sn}}}\] = −0.14 V, Gas constant, R = 8.314 J K⁻¹ mol⁻¹, Faraday constant, F = 96500 C mol⁻¹)

The standard potentials are given as: 

\[\ce{E{^{\circ}_{Cu^{2+}/Cu}}}\] = +0.34 V, and \[\ce{E{^{\circ}_{Ag^{+}/Ag}}}\] = +0.80 V.

Calculate the cell potential (E) for the cell containing 0.100 M Ag⁺ and 4.00 M Cu²⁺ at 298 K.

Predict whether the following reaction will occur spontaneously at 298 K: 

\[\ce{Co_{(s)} + Fe{^{2+}_{(aq)}} -> Co{^{2+}_{(aq)}} + Fe_{(s)}}\]

Given: \[\ce{E^{\circ}_{{Co^{2+}/{Co}}}}\] = −0.28 V and \[\ce{E^{\circ}_{{Fe^{2+}/{Fe}}}}\] = −0.44 V.

The half reactions are:

1. \[\ce{Fe^3+ + e- -> Fe^2+}\]; E° = 0.76 V
2. \[\ce{Ag+ + e- -> Ag}\]; E° = 0.80 V

Calculate K_c for the following reaction at 25°C:

\[\ce{Ag+ + Fe^2+ -> Fe^3+ + Ag}\] (F = 96500 C mol⁻¹)

Calculate the emf of the cell:

Mg_(s) | Mg²⁺ (0.1 M) || Cu²⁺ (1 × 10⁻³ M) | Cu_(s)

Given \[\ce{E^{\circ}_{{Cu^{2+}/{Cu}}}}\] = +0.34 V, \[\ce{E^{\circ}_{{Mg^{2+}/{Mg}}}}\] = −2.37 V

What are electrolytic conductors?

Define the following term:

Specific conductance

Define the term equivalent conductance.

Define the term cell constant.

Define “Molar conductivity”.

What is the SI unit of molar conductivity?

Define “Molar conductivity”.

Mention the effect of temperature on molar conductivity.

Define the term equivalent conductance.

Define “Molar conductivity”.

What are the physical significances of equivalent conductivity?

The resistance of 0.01 M NaCl solution in a conductivity cell was found to be 210 Ω. The specific conductance of this solution is 4.5 × 10⁻³ S cm⁻¹. What is the cell constant of the cell?

The resistance of 0.01 N NaCl solution is 200 Ω at 25°C. Cell constant of the conductivity cell is unity. Calculate the equivalent conductance of the solution.

The specific conductance of `N/50` solution of KCl at 25°C is 0.002765 mho cm⁻¹. If the resistance of a cell containing this solution is 400 ohms, find out the cell constant.

Calculate the equivalent conductance of 1 M H₂SO₄ solution, if its conductivity is 2.6 × 10⁻² ohm⁻¹ cm⁻¹.

The resistance of a decinormal solution of a salt occupying a volume between two platinum electrodes 1.80 cm apart and 5.4 cm² in area was found to be 32 ohms.

Calculate the specific and equivalent conductances of the solution.

A conductivity cell has its electrodes 1 cm apart and each electrode has a cross section of 2 cm². When filled with `N/50` solution of an electrolyte MX, the cell shows a resistance of 166.5 ohms. Calculate the equivalent conductivity of MX at the given concentration.

A potential difference of 20 volts applied to the ends of a column of `N/10` AgNO₃ solution, 4 cm in diameter and 12 cm in length, gives a current of 1.198 amperes. Calculate the specific and equivalent conductivity of the solution.

A conductivity cell when filled with 0.05 M solution of KCl records a resistance of 410.5 ohm at 25°C. When filled with CaCl₂ solution (11 g in 500 mL), it records a resistance of 990 ohms. If specific conductance of 0.05 M KCl is 0.00189 ohm⁻¹ cm⁻¹, calculate

1. cell constant; 
2. specific conductance of CaCl₂;
3. equivalent conductance of CaCl₂ solution;
4. molar conductance of CaCl₂ solution.

A cell with N/50 KCl solution showed a resistance of 550 ohms at 25°C. The specific conductivity of N/50 KCl at 25°C is 0.002768 ohm⁻¹,cm⁻¹. The cell filled with N/10 ZnSO₄ solution at 25°C shows a resistance of 72.18 ohms. Calculate the cell constant and molar conductivity of ZnSO₄ solution.

How does molar conductance of a strong electrolyte vary with its concentration in solution?

How does molar conductivity of a weak electrolyte vary with its concentration in solution?

What is the effect of decreasing concentration on the molar conductivity of a weak electrolyte?

State and explain Kohlrausch’s law.

How can the degree of dissociation of acetic acid be calculated from its molar conductance data?

At 298 K, the molar conductivities at infinite dilution of sodium propionate (CH₃CH₂COONa), HCl and NaCl are 85.9, 426.1 and 126 ohm⁻¹ cm² mol⁻¹ respectively. Calculate the molar conductivity of propionic acid at infinite dilution.

The molar conductivity of a weak monobasic acid at infinite dilution is 387 ohm⁻¹ cm² mol⁻¹. The 0.02 M solution of the same acid has a specific conductivity of 3.3 × 10⁻⁴ ohm⁻¹ cm² mol⁻¹. Calculate the degree of dissociation of the acid at the given concentration.

The molar conductivity of CH₃COOH at infinite dilution is found to be 387 ohm⁻¹ cm² mol⁻¹ but at a dilution of 1 g mol in 1000 litres it is 55 ohm⁻¹ cm² mol⁻¹. What is the percent dissociation of the acid at the given dilution?

Calculate the molar conductivity at infinite dilution for CH₃COOH if the molar conductivity at infinite dilution for NaCl, HCl and CH₃COONa are 126.45, 426.16 and 91.0 ohm⁻¹ cm² mol⁻¹ respectively.

The molar conductance at infinite dilution of Al₂(SO₄)₃ is 858 ohm⁻¹ cm² mol⁻¹. Calculate the molar conductance at infinite dilution of Al³⁺ ion if that of \[\ce{SO^{2-}_{4}}\] ion is 160 ohm⁻¹ cm² mol⁻¹.

What is electrolysis?

Explain the mechanism of electrolysis by taking a suitable example.

Define the following term:

Faraday

Define the following term:

Electrochemical equivalent

Explain the criteria for product formation during electrolysis taking the example of aqueous sodium chloride.

State Faraday’s Laws of electrolysis.

Predict the products of electrolysis of an aqueous solution of CuBr₂ using inert electrodes.

Given: \[\ce{E^{\circ}_{{Cu^{2+}/{Cu}}}}\] = 0.34 V, \[\ce{E^{\circ}_{{H_{2}O}/H_{2}}}\], OH⁻ = −0.83 V, \[\ce{E^\circ_{Br/Br^-}}\] = 1.08 V and \[\ce{E^{\circ}_{\frac{1}{2} O_{2}, H{^{+}}/{H_{2}O}}}\] = 1.23 V.

How many coulombs of electricity are required for reduction of 1 mole of \[\ce{Cr2O^{2-}_{ 7}}\] to Cr³⁺?

How much electricity is required in coulomb for the oxidation of 1 mol of H₂O to O₂?

How many coulombs of electricity are required for reduction of 1 mole of Sn⁴⁺ to Sn²⁺?

Calculate the number of coulombs required to deposit 5.4 g of Al when the electrode reaction is Al₃ + 3e^– + Al.
(Given: Atomic mass of Al = 27 g mol⁻¹, F = 96500 C mol⁻¹)

Calculate the quantity of electricity that would be required to reduce 12.3 g of nitrobenzene to aniline, if the current efficiency is 50%.

How many molecules of chlorine should be deposited from sodium chloride in 1 minute by a current of 300 mA?

A constant current was passed through a solution of \[\ce{AuCl^-_4}\] between gold electrodes. After a period of 14 minutes, the cathode recorded an increase in mass of 1.808 g. Calculate the quantity of current passed in solution. What was the magnitude of current? (Atomic mass of gold = 197.0)

A current of 4 amperes was passed for 2 hours through a solution of copper sulphate when 5.0 g of copper was deposited. Calculate the current efficiency.

Two voltameters containing CuSO₄ and acidulated water respectively are connected in series and the same current is passed for some time. 0.3177 g copper is deposited in the first one. Calculate the weight of liberated hydrogen in the second. (At. wt. of Cu = 63.5 and H = 1.008)

An aqueous solution of NaCl on electrolysis gives H_2(g), Cl_2(g) and NaOH according to the reaction

\[\ce{2Cl^-_{ (aq)} + 2H2O -> 2OH^-_{ (aq)} + H2_{(g)} + Cl2_{(g)}}\].

A direct current of 25 amperes with a current efficiency of 62% is passed through 20 litres of NaCl solution (20% by weight).

1. Write down the reactions taking place at the electrodes.
2. How long will it take to produce 1 kg of Cl₂?
3. What will be the molarity of solution with respect to OH⁻? Assume no loss in volume due to evaporation.

Explain, why electrolysis of aqueous solution of NaCl gives H₂ at cathode and Cl₂ at anode. Write overall reaction.

[Given: \[\ce{E^{\circ}_{{Na^{+}/{Na}}}}\] = 2.71 V, \[\ce{E^{\circ}_{Cl_2/2Cl^-}}\] = 1.36 V and \[\ce{\frac{1}{2} O2_{(g)} + 2H+_{ (aq)} + 2e- -> H2O_{(l)}}\]; E° = 1.23 V]

Silver is electrodeposited on a metallic vessel of surface area 800 cm² by passing current of 0.2 amp for 3 hours. Calculate the thickness of silver deposited. 

(Density of silver = 10.47 g cm⁻³, atomic mass of silver = 107.92 amu)

Predict the products of electrolysis obtained at the electrodes in the case when the electrodes used are of platinum:

An aqueous solution of AgNO₃.

Predict the product of electrolysis in the following:

A dilute solution of H₂SO₄ with platinum electrodes.

How much charge is required for the following reduction? 

1 mol of Cu²⁺ to Cu.

Write the differences between the primary cell and the secondary cell.

What do you understand by the term fuel-cell?

Give electrode reactions of hydrogen-oxygen fuel cell.

What is a mercury cell?

Give the electrode reactions of the mercury cell.

What is corrosion?

Give a brief account of the mechanism of corrosion.

What is corrosion?

How is the cathodic protection of iron different from its galvanisation?

What is galvanisation?

In what respect is galvanisation better than tin coating used for the protection of iron from corrosion?

Write two applications of fuel cells.

Describe the limitations of fuel cells.

Explain why tin protects copper more efficiently as compared to iron?

Discuss the factors which govern the rate of corrosion of a metal.

Give a brief account of the mechanism of corrosion.

Give an example of a fuel cell and write the anode and cathode reactions for it.

How does a fuel cell operate?

Write the cell reactions which occur in lead storage battery when the battery is in use.

Write the cell reactions which occur in a lead storage battery when the battery is on charging.

Differentiate between conductors and insulators.

What do you understand by non-electrolytes?

What is the effect of temperature on metallic conduction?

What type of species are responsible for electrolytic conduction?

Can electric current be passed through a solid ionic compound?

How does electrolytic conduction vary with temperature?

What is meant by the term ‘electrolyte’?

Give two examples of electrolytes.

Distinguish between strong electrolyte and weak electrolyte.

Define the degree of dissociation.

How does solute-solute interaction affect electrolytic conduction?

What is the effect of solute-solvent interaction on electrolytic conduction?

State Ostwald’s dilution law.

Define the following term:

Specific conductance

What are the units of specific conductivity?

How is equivalent conductivity related to specific conductivity?

Define the term equivalent conductance.

Mention the units of equivalent conductivity.

What is the relationship between molar conductivity and specific conductivity?

Define “Molar conductivity”.

What is the SI unit of molar conductivity?

What is cell constant?

Define molar conductivity at infinite dilution.

Why is the molar conductivity of a weak electrolyte much lower than that of a strong electrolyte at moderate concentrations?

What is specific resistance?

How would you represent a Daniell cell?

What is the polarity of anode in a galvanic cell and why?

What do you understand by a half cell?

How would you represent a half cell in which H⁺ ions of conc. 0.2 M get reduced into hydrogen gas at atmospheric pressure?

What is the direction of flow of electrons in the outer circuit in a galvanic cell?

Define electrode potential.

What is a null electrode?

Name the electrode and the potential when the electrode reaction involves reduction.

What do you understand by standard electrode potential?

What is the reference electrode generally used in the measurement of standard electrode potential of an electrode assembly?

What is the function of platinised platinum in the standard hydrogen electrode?

Represent a standard hydrogen electrode.

What is the direction of flow of electrons in the outer circuit when a standard hydrogen electrode is connected to a copper electrode?

How does the tendency to gain electrons vary when one moves from the top to the bottom in the electrochemical series?

It is impossible to measure the potential of a single electrode. Comment.

Predict whether F₂ and Na will react together:

\[\ce{E^{\circ}_{F/F^-}}\] = +2.87 V, \[\ce{E^{\circ}_{{Na^{+}/{Na}}}}\] = −2.71 V

Which substance can be used to oxidise fluorides to fluorine?

How does the reducing power of elements vary on moving down the electrochemical series?

Can copper displace zinc from its salt solution?

\[\ce{E^{\circ}_{{Zn^{2+}/{Zn}}}}\] = −0.76 V, \[\ce{E^{\circ}_{{Cu^{2+}/{Cu}}}}\] = +0.34 V

Can nickel displace hydrogen from hydrochloric acid?

\[\ce{E^{\circ}_{{Ni^{2+}/{Ni}}}}\] = −0.25 V

Write the Nernst equation for a reduction electrode.

Define emf of a galvanic cell.

What is the convention for the emf of a cell?

Mention the relationship between the standard cell potential and standard electrode potentials of the two electrodes.

Write the Nernst equation for a cell having the cell reaction. 

\[\ce{aA + bB -> cC + dD}\].

The standard reduction potential for \[\ce{Zn{^{2+}_{(aq)}} | Zn_{(s)}}\] half cell is −0.76 V. Write the reactions occurring at the electrodes when coupled with NHE.

Define electrolysis.

Why does an aqueous solution of sodium chloride on electrolysis give H₂ at cathode instead of sodium?

What is overvoltage?

Why do Cl⁻ ions get oxidised at anode in preference to H₂O during the electrolysis of an aqueous solution of sodium chloride?

Predict the products of electrolysis of a dilute aqueous solution of sulphuric acid using inert electrodes.

State Faraday’s Laws of electrolysis.

What is electrochemical equivalent?

How can electrochemical equivalent be calculated for a particular substance?

What is the quantity of electricity needed to liberate one equivalent of a substance?

Calculate the electric charge carried by one electron on the basis of Faraday’s constant.

How much charge is required to reduce 8 g of H⁺ ions to H₂ gas?

How can you increase the reduction potential of an electrode?

Electrolysis of KBr_(aq) gives Br₂ at anode but that of KF_(aq) does not give F₂. Give reason.

Define electrochemistry.

Discuss the scope of electrochemistry.

What is metallic conduction?

How does metallic conduction differ from electrolytic conduction?

Discuss the factors which affect electrolytic dissociation.

What is specific resistance?

What do you understand by specific conductance of a solution?

Derive the units for specific resistance of a solution.

Derive the units for conductance of a solution.

Derive the units of specific conductivities of an electrolytic solution and write expressions which show their mutual relationship.

Define the term equivalent conductance.

How is equivalent conductivity related to specific conductivity?

Derive the units of equivalent conductivity of a solution.

Define “Molar conductivity”.

What is the SI unit of molar conductivity?

What is the relationship between molar conductivity and specific conductivity?

Derive a relationship between molar conductivity and equivalent conductivity.

What is conductivity cell?

For what purpose is a conductivity cell used?

What is wheatstone bridge?

What is the wheatstone bridge’s significance in the determination of the conductivity of a solution?

What is cell constant?

How is cell constant determined?

Can a direct current be used for the measurement of electrical conductivity of a solution? If not, why?

How does the molar conductivity for the solution of a strong electrolyte vary with concentration? Explain graphically.

What is Debye-Huckel-Onsagar equation and what is its significance?

Explain why the molar conductivity of a strong electrolytic solution increases only slightly when concentration is decreased.

What do you understand by molar conductivity at infinite dilution?

How does the molar conductivity for the solution of a strong electrolyte vary with concentration? Explain graphically.

How does the molar conductivity of a weak electrolytic solution vary with concentration?

The molar conductivity at infinite dilution of a strong electrolyte can be obtained by extrapolating the curve Λ_m vs `sqrt c`, but the same cannot be done for a weak electrolyte. Explain why.

Why does the molar conductivity of a weak electrolytic solution increase sharply at very low concentrations?

How does the molar conductivity of a weak electrolytic solution vary with concentration?

What is the effect of dilution on the specific conductivity of a solution?

What is the effect of dilution on the equivalent conductivity of a solution?

What is the effect of dilution on the molar conductivity of a solution?

What is a galvanic cell?

How can a galvanic cell be constructed?

What do you understand by a half cell?

What are the half cells present in a Daniell cell?

How would you represent a Daniell cell?

How would you represent a Daniell cell?

Explain why the anode is of negative polarity in a galvanic cell.

Explain why the cathode is of positive polarity in a galvanic cell.

What type of reactions do occur at the two electrodes at a galvanic cell and what is the nature of the net cell reaction? Explain with an example.

What is electrode potential?

How does electrode potential come into existence?

What is a salt bridge?

What are the functions of a salt bridge in a galvanic cell?

Write the ion electron equations for the half cell reactions and overall cell reaction for the cell.

\[\ce{Pt, H2_{(g)} | HCl_{(aq)} || Fe^{3+}, Fe^{2+} | Pt}\]

Depict the electrochemical cell in which the cell reaction is

\[\ce{2Cr_{(s)} + 3Cu{^{2+}_{(aq)}} -> 2Cr{^{3+}_{(aq)}} + 3Cu_{(s)}}\]

1. mark the anode and the cathode;
2. show movement of electrons and flow of current;
3. write the half cell reactions.

Write the symbolic representation of the electrochemical cell with the cell reaction.

\[\ce{Zn_{(s)} + 2Ag{^+_{(aq)}} -> Zn{^{2+}_{(aq)}} + 2Ag_{(s)}}\]

Indicate the oxidation electrode and the direction of movement of electrons and write the electrode reactions.

What is standard hydrogen electrode?

What is standard hydrogen electrode significance?

Define electrode potential.

It is impossible to measure the potential of a single electrode. Comment.

How are standard electrode potentials measured?

When a Zn²⁺/Zn electrode is connected to a SHE, the electrons flow from zinc electrode to hydrogen electrode but the flow of electrons is in opposite direction when a Cu²⁺/Cu electrode is connected to a SHE. Explain giving reasons.

How does the tendency to gain electrons vary when one moves from the top to the bottom in the electrochemical series?

How does the following property vary on moving down the electrochemical series?

Reducing power of the reduced form.

How does the following property vary on moving down the electrochemical series?

Oxidising power of the oxidised form.

Justify the statement:

If \[\ce{E^{\circ}_{{M^{n+}/{M}}}}\] is positive, Mⁿ⁺ can be reduced to M by H₂ gas under standard conditions.

Comment on the statement:

If \[\ce{E^{\circ}_{{M^{n+}/{M}}}}\] is negative, H⁺ ions cannot be reduced to H_2(g) by the metal M.

On what factor does the reactivity of a metal depend and why?

Explain, why can zinc displace Cu from an aqueous solution of CuSO₄ but Cu cannot displace zinc from aqueous ZnSO₄ solution?

What is meant by reduction electrode potentials of zinc and copper being −0.76 V and + 0.34 V respectively? Can an aqueous solution of CuSO₄ be stored in a zinc vessel? Answer with reason.

What do you understand by electrochemical series?

How does the electrochemical series help in predicting whether a given reaction is feasible in a given direction or not?

What do you understand by electrochemical series?

How does the electrochemical series help in calculating the emf of a standard cell?

Give reason why the blue colour of copper sulphate solution is discharged slowly when an iron rod is dipped in it.

Given: \[\ce{E^{\circ}_{{Cu^{2+}/{Cu}}} = 0.34 V}\], \[\ce{E^{\circ}_{{Fe^{2+}/{Fe}}} = -0.44 V}\].

Write the equation which correlates the electrode potential of an electrode assembly with the concentration of ions and the temperature. What is the name of the equation and what is its significance?

What do you understand by EMF of a galvanic cell?

How can EMF of a galvanic cell be calculated if the cell is in standard condition?

Write the Nernst equation for the cell potential of a galvanic cell.

Apply the Nernst equation to a Daniell cell.

Write the Nernst equation for a cell having the cell reaction. 

\[\ce{aA + bB -> cC + dD}\].

What is an electrolytic cell?

How does electrolytic cell differ from a galvanic cell?

Define electrolysis.

Discuss the electrolysis mechanism by taking the example of fused sodium chloride.

What is overvoltage?

How does overvoltage affect the electrolysis of an aqueous solution of sodium chloride?

Predict the products of electrolysis obtained at the electrodes in the case when the electrodes used are of platinum:

An aqueous solution of AgNO₃.

Predict the product of electrolysis in the following:

An aqueous solution of AgNO₃ with silver electrodes.

What are the criteria for product formation during electrolysis?

Why does the anode dissolve when an aqueous solution of CuSO₄ is subjected to electrolysis using copper electrodes?

What are the products of electrolysis of an aqueous solution of dilute sulphuric acid? Explain with reasons.

State Faraday’s Laws of electrolysis.

How much charge is necessary to liberate one equivalent of a substance and why? Calculate the value of this charge.

All the energy released from the reaction \[\ce{X -> Y}\], Δ_rG° = −193 KJ mol⁻¹ is used for oxidizing \[\ce{M+ -> M^{3+} + 2e−}\], E° = −0.25 V.

Under standard conditions, the number of moles of M⁺ oxidized when one mole of X is converted to Y is [F = 96500 C mol⁻¹].

What are electrolytic conductors?

How are electrolytic conductors classified?

Define the following term:

Specific conductance

Define the term equivalent conductance.

Define “Molar conductivity”.

Derive the units of specific conductivities of an electrolytic solution and write expressions which show their mutual relationship.

Derive the units of equivalent conductivity of a solution.

How is equivalent conductivity related to specific conductivity?

What is the SI unit of molar conductivity?

Derive a relationship between molar conductivity and equivalent conductivity.

What is the relationship between molar conductivity and specific conductivity?

What is cell constant?

How is cell constant determined?

What is the significance of the cell constant in the measurement of conductivity of an unknown electrolytic solution?

Discuss the method used to measure the electrolytic conductivity of a solution.

How would you determine the equivalent and molar conductivities of a given solution?

How does molar conductance of a strong electrolyte vary with its concentration in solution?

How does the molar conductivity of a weak electrolytic solution vary with concentration?

How does the molar conductivity for the solution of a strong electrolyte vary with concentration? Explain graphically.

What do you understand by molar conductivity at infinite dilution?

How does the molar conductivity of a weak electrolytic solution vary with concentration?

State and explain Kohlrausch’s law.

How would you determine the molar conductivity at infinite dilution of NH₄OH with the help of Kohlrausch’s law?

What is a galvanic cell?

Describe the construction and working of a Daniell cell.

Draw a labelled diagram of Daniell cell and show

1. electrodes working as anode and cathode and their polarity;
2. direction of flow of electrons and direction of flow of conventional current;
3. half cell reactions and the net cell reaction.

What is electrode potential?

How does electrode potential come into existence?

Explain with examples the term oxidation potential.

Explain with examples the term reduction potential.

What do you understand by standard electrode potential?

How are standard electrode potentials measured?

What do you understand by electrochemical series?

What are the important features of the electrochemical series?

How does the electrochemical series help in comparing the relative reducing or oxidising powers of different elements?

How does the electrochemical series help in predicting the displacement of hydrogen by metals from dilute acids?

On the basis of electrochemical series, how would you predict displacement of metals from salt solutions?

On the basis of electrochemical series, how would you predict feasibility of a redox reaction?

On the basis of electrochemical series, how would you predict thermal stability of oxides of different metals?

What is Nernst equation for the potential of an electrode?

Write the Nernst equation for a cell having the cell reaction. 

\[\ce{aA + bB -> cC + dD}\].

What do you understand by EMF of a galvanic cell?

Write the IUPAC convention for calculating the standard emf of a cell.

Write the Nernst equation for the cell potential of a galvanic cell.

What is the significance of electrode potentials in predicting the products of electrolysis of an aqueous solution of an electrolyte? Explain with a suitable example.

Discuss the electrolysis of an aqueous solution of sodium chloride and explain the products formed at the two electrodes.

Mention the criteria to predict the formation of products during electrolysis of an aqueous solution of an electrolyte. Apply these criteria to the electrolysis of an aqueous solution of copper sulphate using inert electrodes.

State Faraday’s Laws of electrolysis.

How are Faraday’s laws of electrolysis helpful in exploring the quantitative aspect of electrolysis?

The molar conductivities at infinite dilution of barium chloride, sulphuric acid and hydrochloric acid are 280, 860 and 426 S cm² mol⁻¹ respectively. The molar conductivity at infinite dilution of barium sulphate is ______ S cm² mol⁻¹.

(Round off to the nearest integer)

A KCl solution of conductivity 0.14 S m⁻¹ shows a resistance of 4.19 Ω in a conductivity cell. If the same cell is filled with an HCl solution, the resistance drops to 1.03 Ω. The conductivity of the HCl solution is ______ × 10⁻² S m⁻¹

(Round off to the nearest integer)

At 298 K, the limiting molar conductivity of a weak mono basic acid is 4 × 10² S cm² mol⁻¹. At 298 K, for an aqueous solution of the acid, the degree of dissociation is ‘α’ and the molar conductivity is y × 10² S cm² mol⁻¹. At 298 K, upon 20 times dilution with water, the molar conductivity of the solution becomes 3y × 10² S cm² mol⁻¹.

1. The value of ‘α’ is ______.
2. The value of ‘y’ is ______.

The cell potential for the following cell:

\[\ce{Pt | H2_{(g)} | H{^+_{(aq)}} || Cu^+2 (0.01 M) | Cu}\] is 0.576 at 298 K.

The pH of the solution is ______. (Nearest integer)

The cell potential for the given cell at 298 K.

\[\ce{Pt | H2_{(g)} (1 bar) | H{^+_{(aq)}} || Cu{^{+2}_{(aq)}} | Cu_{(s)}}\] is 0.31 V.

The pH of the acidic solution is found to be 3 whereas the concentration of Cu⁺² is 10^−x M. The value of x is ______ (Nearest integer).

(Given, \[\ce{E^{\circ}_{{Cu^{2+}/{Cu}}} = 0.34 V}\] and \[\ce{\frac{2.303 RT}{F} = 0.06 V}\]).

A solution of Fe₂(SO₄)₃ is electrolysed for ‘x’ min with a current of 1.5 A to deposit 0.3482 g of Fe. The value of ‘x’ is ______. (Nearest integer)

(Given, 1 F = 96500 C mol⁻¹, Atomic mass of Fe = 56)

\[\ce{Cu_{(s)} + Sn{^{+2}} (0.001 M) -> Cu^{+2} (0.01 M) + Sn_{(s)}}\]

The Gibbs free energy change for the above reaction at 298 K is x × 10⁻¹ KJ mol⁻¹. The value of x is ______ (Nearest integer).

(Given \[\ce{E^{\circ}_{{Cu^{2+}/{Cu}}} = 0.34 V}\], \[\ce{E^{\circ}_{{Sn^{2+}/{Sn}}} = -0.14 V}\], F = 96500 C mol⁻¹)

For the given reactions

\[\ce{Sn^{+2} + 2e- -> Sn}\]

\[\ce{Sn^{+4} + 2e- -> Sn}\]

The electrode potentials are \[\ce{E^{\circ}_{{Sn^{+2}/{Sn}}}}\] = −0.14 V and \[\ce{E^{\circ}_{{Sn^{+4}/{Sn}}}}\] = 0.010 V. The magnitude of standard electrode potential for \[\ce{Sn^{+4}/Sn^{+2}}\] i.e., \[\ce{E^{\circ}_{{Sn^{+4}/{Sn^{+2}}}}}\] is ______ × 10⁻² V (Nearest integer).

The reduction potential (E° in V) of \[\ce{MnO^-_{4(aq)}/Mn_{(s)}}\] is ______.

Given \[\ce{E^{\circ}_{MnO^-_{4(aq)}/MnO_{2(s)}}}\] = 1.68 V

\[\ce{E^{\circ}_{MnO_{2(s)}/Mn^{+2}_{(aq)}}}\] = 1.21 V

\[\ce{E^{\circ}_{Mn^{+2}_{(aq)}/Mn_{(s)}}}\] = −1.03 V

Molten sodium chloride conducts electricity due to the presence of ______.

When one Faraday of electric current is passed, the mass deposited is equal to ______.

A current liberated 0.50 g of hydrogen in 2 hours. How many grams of copper can be liberated by the same current flowing for the same time in CuSO₄ solution?

One Faraday of charge was passed through the electrolytic cells placed in series containing solutions of Ag⁺, Ni²⁺ and Cr³⁺ respectively. The amounts of Ag (At. mass = 108), Ni (At. mass = 59) and Cr (At. mass = 52) deposited will be:

Which one is the correct equation that represents the first law of electrolysis?

When electricity is passed through a solution of AlCl₃, 13.5 g of Al is discharged, the amount of charge passed is ______.

The cathodic reaction in the electrolysis of dilute H₂SO₄ with platinum electrodes is ______.

The reaction, \[\ce{\frac{1}{2} H2_{(g)} + AgCl_{(s)} -> H^+_{ (aq)} + Cl^-_{ (aq)} + Ag_{(s)}}\], occurs in the galvanic cell

The standard oxidation potentials, \[\ce{E{^{\circ}_{oxi}}}\] for the half cell reactions are as follows:

\[\ce{Zn -> Zn^2+ + 2e-}\]; \[\ce{E{^{\circ}_{oxi}}}\] = +0.76 V

\[\ce{Fe -> Fe^2+ + 2e-}\]; \[\ce{E{^{\circ}_{oxi}}}\] = +0.41 V

The emf of the cell, \[\ce{Fe^2+ + Zn -> Zn^2 + Fe}\], is ______.

If a spoon of copper metal is placed in a solution of ferrous sulphate, ______.

Which one will liberate Br₂ from KBr?

The standard reduction potentials at 25°C for the following half cell reactions are given against each:

\[\ce{Zn{^{2+}_{(aq)}} + 2e- -> Zn_{(s)}}\] = −0.762 V

\[\ce{Cr{^{3+}_{(aq)}} + 3e- -> Cr_{(s)}}\] = −0.740 V

\[\ce{2H+ + 2e- -> H2_{(g)}}\] = 0.000 V

\[\ce{Fe^3+ + e- -> Fe^2+}\] = 0.770 V

Which is the strongest reducing agent?

Which is not true for a standard hydrogen electrode?

The standard electrode potentials of four elements A, B, C, D are −3.05, 1.66, −0.40 and 0.80 volts respectively. The highest chemical activity will be shown by ______.

Which one of the following will increase the voltage of the cell whose cell reaction is given below?

\[\ce{Sn + 2Ag+ -> Sn^2+ + 2Ag}\]

When electric current is passed through a cell having an electrolyte, the positive ions move towards the cathode and the negative ions towards the anode. If the cathode is pulled out of the solution ______.

If a salt bridge is removed from the two half cells, the voltage ______.

1 Faraday of electricity will liberate 1 g atom of the metal from ______.

The metal that cannot be obtained by electrolysis of the aqueous solution of its salts is ______.

The standard reduction electrode potential values of the elements A, B and C are +0.68, −2.50 and −0.50 V respectively. The order of their reducing power is ______.

How many electrons carry one coulomb of charge?

If \[\ce{E^{\circ}_{{Ag^{+}/{Ag}}}}\] = +0.80 V and \[\ce{E^{\circ}_{{Sn^{2+}/{Sn}}}}\] = −0.14 V, the standard emf of the cell \[\ce{Sn | Sn^2+ ||  Ag+ | Ag}\], will be ______.

The hydrogen electrode is dipped in a solution of pH = 3 at 25°C. The reduction potential of the electrode would be ______.

The standard EMF for the cell reaction,
\[\ce{Zn + Cu^{2+} -> Zn^{2+} + Cu}\], 
is 1.10 V at 25°C. The EMF of the cell reaction, when 0.1 M Cu²⁺ and 0.1 M Zn²⁺ solutions are used at 25°C is:

The number of Faradays required to reduce one mol of Cu²⁺ to metallic copper is ______.

A standard hydrogen electrode has zero electrode potential because ______.

The standard reduction potential values of three metallic cations, X, Y, Z are 0.52, −3.03 and −1.18 V respectively. The order of reducing power of the corresponding metals is ______.

A gas X at 1 atm is bubbled through a solution containing a mixture of 1 MY⁻ and 1 MZ⁻ at 25°C. If the reduction potential of Z > Y > X, then ______.

For the electrochemical cell:

\[\ce{M | M+ || X- | X}\]; 

\[\ce{E^{\circ}_{{M^{+}/{M}}}}\] = 0.44 V,

\[\ce{E^{\circ}_{X/X^-}}\] = 0.33 V

Which of the following is TRUE for this data?

At 25°C, the standard e.m.f. of a cell having reaction involving two electron change is found to be 0.295 V. The equilibrium constant of the reaction is ______.

Saturated solution of KNO₃ is used to make salt bridge because:

The correct order of equivalent conductance at infinite dilution of LiCl, NaCl and KCl is ______.

For a cell given below:

\[\ce{Ag | Ag+ || Cu^2+ | Cu}\]

\[\ce{Ag+ + e- -> Ag}\]; E° = x

\[\ce{Cu^2+ + 2e- -> Cu}\]; E° = y

\[\ce{E{^{\circ}_{cell}}}\] is ______.

Conductivity (siemens, S) is directly proportional to area of the vessel and the concentration of the solution in it and is inversely proportional to the length of the vessel, then constant of proportionality is expressed in ______.

When during electrolysis of a solution of AgNO₃, 9650 coulombs of charge pass through the electroplating bath, the mass of silver deposited on the cathode will be ______.

For the redox reaction, 

\[\ce{Zn_{(s)} + Cu^2+ (0.1 M) -> Zn^2+ (1M) + Cu_{(s)}}\],

taking place in a cell, \[\ce{E{^{\circ}_{cell}}}\] is 1.10 volt. E_cell for the cell will be: \[\ce{(2.303 \frac{RT}{F} = 0.0591)}\]

For a cell reaction in involving a two electron change, the Stanford emf of the cell is found to be 0.295 V at 25°C. The equilibrium constant of the reaction at 25°C will be:

Standard reduction electrode potentials of three metals A, B and C are +0.5 V, −3.0 V and −1.2 V respectively. The reducing power of these metals is ______.

In a hydrogen-oxygen fuel cell, combustion of hydrogen occurs to ______.

Consider the following E° values:

\[\ce{E^{\circ}_{{Fe^{3+}/{Fe^{2+}}}}}\] = + 0.77 V

\[\ce{E^{\circ}_{{Sn^{2+}/{Sn}}}}\] = −0.14 V.

Under standard conditions the potential for the reaction:

\[\ce{Sn_{(s)} +  2Fe^{3+}_{ (aq)} -> 2Fe^{2+}_{ (aq)} + Sn^{2+}_{ (aq)}}\] is:

The standard emf of a cell, involving one electron charge is found to be 0.591 V at 25°C. The equilibrium constant of the reaction is ______. (F = 96,500 C mol⁻¹)

The limiting molar conductivities Λ^° for NaCl, KBr and KCl are 126, 152 and 150 S cm² mol^–1 respectively. The limiting molar conductivity Λ^° for NaBr is ______.

In a cell that utilises the reaction \[\ce{Zn_{(s)} + 2H{^{+}_{(aq)}} -> Zn{^{2+}_{(aq)}} + H2_{(g)}}\] addition of H₂SO₄ to cathode compartment will ______.

The \[\ce{E^{\circ}_{M^{3+}/M^{2+}}}\] values for Cr, Mn, Fe and Co are −0.41, +1.57, +0.77 and +1.97 V respectively. For which one of these metals the change in oxidation state from +2 to +3 is easiest?

Aluminium oxide may be electrolysed at 1000°C to furnish aluminium metal.

(Atomic mass = 27 amu; 1 Faraday = 96,500 coulombs)

The cathode reaction is:

\[\ce{Al^3+ + 3e- -> Al^0}\]

To prepare 5.12 kg of aluminium metal by this method would require:

Given the data at 25°C,

\[\ce{Ag + I- -> AgI + e-}\]; E° = – 0.152 V

\[\ce{Ag -> Ag+ + e-}\]; E° = – 0.800 V.

What is the value of K_sp for AgI?

\[\ce{\left(2.303 \frac{RT}{F} = 0.059 V\right)}\]

Resistance of a conductivity cell filled with a solution of an electrolyte of concentration 0.1 M is 100 Ω. The conductivity of this solution is 1.29 S m⁻¹. Resistance of the same cell when filled with 0.02 M of the same solution is 520 Ω. The molar conductivity of 0.02 M solution of the electrolyte will be ______.

The cell, \[\ce{Zn | Zn^{2+} (1 M) || Cu^{2+} (1 M) | Cu}\], (\[\ce{E^{\circ}_{cell} = 1.10 V}\]), was allowed to be completely discharged at 298 K. The relative concentration of Zn²⁺ to \[\ce{Cu^{2+}\left(\frac{[Zn^{2+}]}{[Cu{2+}]}\right)}\] is:

The equivalent conductances of two strong electrolytes at infinite dilution in H₂O (where ions move freely through a solution) at 25°C are given below:

\[\ce{\Lambda^{\circ}_{CH_3COONa}}\] = 91.0 S cm²/equiv.

\[\ce{\Lambda^{\circ}_{HCl}}\] = 426.2 S cm²/equiv.

What additional information/quantity does one need to calculate Λ° of an aqueous solution of acetic acid?

Given \[\ce{E^{\circ}_{Cr^{3+}/Cr}}\] = −0.72 V, \[\ce{E^{\circ}_{Fe^{2+}/Fe}}\] = −0.42 V. The potential for the cell \[\ce{Cr | Cr^{3+} (0.1 M) || Fe^{2+} (0.01 M) | Fe}\] is:

For the reduction of silver ions with copper metal, the standard cell potential was found to be +0.46 V at 25°C. The value of standard Gibbs energy, ΔG° will be ______. (F = 96500 C mol⁻¹)

Standard electrode potential of three metals X, Y and Z are –1.2 V, +0.5 V and –3.0 V, respectively. The reducing power of these metals will be ______.

If the \[\ce{E^{\circ}_{cell}}\] for a given reaction has a negative value, then which of the following gives the correct relationships for the values of ΔG° and K_eq?

The electrode potentials for \[\ce{Cu{^{2+}_{(aq)}} + e- -> Cu{^+_{(aq)}}}\] and \[\ce{Cu{^{+}_{(aq)}} + e- -> Cu_{(s)}}\] are +0.15 V and +0.50 V respectively. The value of  \[\ce{E^{\circ}_{{Cu^{2+}/{Cu}}}}\] will be:

A buffer solution is prepared in which the concentration of NH₃ is 0.30 M and the concentration of \[\ce{NH^+_4}\] is 0.20 M. If the equilibrium constant, K_b for NH₃ equals 1.8 × 10⁻⁵, what is the pH of this solution? (log 2.7 = 0.43)

Standard electrode potential for Sn⁴⁺/Sn²⁺ couple is +0.15 V and that for the Cr³⁺/Cr couple is −0.74. These two couples in their standard state are connected to make a cell. The cell potential will be ______.

\[\ce{\Lambda^{\circ}_m(NH_4OH)}\] [i.e., \[\ce{\Lambda^{\circ}_m(NH4OH)}\]] is equal to ______.

When 0.1 mol \[\ce{MnO^{2-}_4}\] is oxidised, the quantity of electricity required to completely oxidise \[\ce{MnO^{2-}_4}\] to \[\ce{MnO^-_4}\] is ______.

The weight of silver (at. wt. = 108) displaced by a quantity of electricity which displaces 5600 mL of O₂ at STP will be ______.

Resistance of 0.2 M solution of an electrolyte is 50 Ω. The specific conductance of the solution is 1.4 S m⁻¹. The resistance of 0.5 M solution of the same electrolyte is 280 Ω. The molar conductivity of 0.5 M solution of the electrolyte in S mol⁻¹ is ______.

The equivalent conductance of NaCl at concentration C and at infinite dilution are λ_c and λ_∞, respectively. The correct relationship between λ_c and λ_∞ is given as:

The metal that cannot be obtained by the electrolysis of an aqueous solution of its salts is ______.

Given below are the half-cell reactions:

\[\ce{Mn^2+ + 2e- -> Mn}\], E° = −1.18 V

\[\ce{2(Mn^3+ + e- -> Mn^2+)}\], E° = +1.51 V

The E° for \[\ce{3Mn^2+ -> Mn + 2Mn^3+}\] will be ______.

The electrochemical equivalent is the amount of substance which gets deposited from its solution on passing electrical charge equal to ______.

Two Faraday of electricity is passed through a solution of CuSO₄. The mass of copper deposited at the cathode is ______. (at. mass of Cu = 63.5 amu)

Galvanization is applying a coating of ______.

The pressure of H₂ required to make the potential of H₂-electrode zero in pure water at 298 K is ______.

In the electrochemical cell, Zn | ZnSO₄ (0.01 M) || CuSO₄ (1.0 M) | Cu, the emf of this Daniell cell is E₁. When the concentration of ZnSO₄ is changed to 1.0 M and that of CuSO₄ changed to 0.01 M, the emf changes to E₂. From the following, which one is the relationship between E₁ and E₂?

`("Given", (RT)/F = 0.059)`

Given: \[\ce{E^{\circ}_{Cl_2/Cl^-}}\] = 1.36 V, \[\ce{E^{\circ}_{Cr^{3+}/Cr}}\] = −0.74 V

\[\ce{E^{\circ}_{Cr_2O^{2-}_7/Cr^{3+}}}\] = 1.33 V, \[\ce{E^{\circ}_{MnO^-_4/Mn^{2+}}}\] = 1.51 V

Among the following, the strongest reducing agent is:

How long (approximate) should water be electrolysed by passing through 100 amperes current so that the oxygen is released can completely burn 27.66 g of diborane? 

(Atomic weight of B = 10.8 u)

Consider the change in the oxidation state of Bromine corresponding to different emf values as shown in the expression below:

\[\ce{BrO^-_4 ->[1.82 V] BrO^-_3 ->[1.5 V] HBrO ->[1.595 V] Br2 ->[1.0652 V] Br^-}\]

Then the species undergoing disproportionation is:

Consider the statements S₁ and S₂.

S₁: Conductivity always increases with decrease in the concentration of electrolyte.

S₂: Molar conductivity always increases with decrease in the concentration of electrolyte.

The correct option among the following is:

Which one of the following graphs between molar conductivity (Λ_m) versus `sqrt C` is correct?

Following limiting molar conductivities are given as:

\[\ce{\lambda^{\circ}_{m} (H2SO4)}\] = x S cm² mol⁻¹

\[\ce{\lambda{^{\circ}_{m}} (K2SO4)}\] = y S cm² mol⁻¹

\[\ce{\lambda{^{\circ}_{m}} (CH3COOK)}\] = z S cm² mol⁻¹

\[\ce{\lambda^{\circ}_{m}}\] (in S cm² mol⁻¹) for CH₃COOH will be:

The standard electrode potential (E°) values of Al³⁺/Al, Ag⁺/Ag, K⁺/K and Cr³⁺/Cr are −1.66 V, 0.80 V, −2.93 V and −0.74 V, respectively. The correct decreasing order of reducing power of the metal is ______.

For the cell reaction:

\[\ce{2Fe^{3+}_{ (aq)} + 2l^-_{ (aq)} -> 2Fe^{2+}_{ (aq)} + l2_{(aq)}}\]

\[\ce{E{^{\circ}_{cell}}}\] = 0.24 V at 298 K. The standard Gibbs energy (Δ_rG°) of the cell reaction is:

[Given that Faraday constant, F = 96500 C mol⁻¹]

For a cell involving one electron \[\ce{E^{\circ}_{cell}}\] = 0.59 V at 298 K, the equilibrium constant for the cell reaction is:

[Given that \[\ce{\frac{2.303 RT}{F}}\] = 0.059 V at T = 298 K]

On electrolysis of dilute sulphuric acid using platinum electrodes, the product obtained at the anode will be ______.

The number of Faradays (F) required to produce 20 g of calcium from molten CaCl₂ (Atomic mass of Ca = 40 g mol⁻¹) is ______.

Let C_NaCl and \[\ce{C{_{BaSO_4}}}\] be the conductances (in S) measured for saturated aqueous solutions of NaCl and BaSO₄ respectively, at a temperature T. Which of the following is false?

250 mL of a waste solution obtained from the workshop of a goldsmith contains 0.1 M AgNO₃ and 0.1 M AuCl. The solution was electrolyzed at 2 V by passing a current of 1 A for 15 minutes. The metal/metals electrodeposited will be ______.

(\[\ce{E^{\circ}_{Ag^+/Ag}}\] = 0.80 V, \[\ce{E^{\circ}_{Au^{2+}/Au}}\] = 1.69 V)

The molar conductance of NaCl, HCl and CH₃COONa at infinite dilution are 126.45, 426.16 and 91.0 S cm² mol⁻¹ respectively. The molar conductance of CH₃COOH at infinite dilution is ______.

Choose the right option for your answer.

The pK_b of dimethylamine and pK_a of acetic acid are 3.27 and 4.77 respectively at T(K). The correct option for the pH of dimethylammonium acetate solution is ______.

The molar conductivity of 0.007 M acetic acid is 20 S cm² mol⁻¹. What is the dissociation constant of acetic acid? Choose the correct option.

\[\begin{array}{cc}
\end{array}\]\[\begin{bmatrix}
\ce{\Lambda^{\circ}_{H^+} = 350 S cm^2 mol^{-1}}\\
\ce{\Lambda^{\circ}_{CH_3COO^-} = 50 S cm^2 mol^{-1}}
\end{bmatrix}\]

Some standard electrode potentials at 298 K are given below:

To a solution containing 0.001 M of X⁺² and 0.1 M of Y⁺², the metal rods X and Y are inserted (at 298 K) and connected by a conducting wire. This resulted in dissolution of X. The correct combination(s) of X and Y respectively is (are):

(Given R = 8.314 J K⁻¹ mol⁻¹, F = 96500 (mol⁻¹)

At 298 K, the standard electrode potentials of Cu⁺²/Cu, Zn⁺²/Zn, Fe⁺²/Fe and Ag⁺/Ag are 0.34 V, −0.76 V, −0.44 V and −0.80 V respectively.

On the basis of standard electrode potential, predict which of the following reaction cannot occur?

Given below are half cell reactions:

\[\ce{MnO^-_4 + 8H+ + 5e- -> Mn^{+2} + 4H2O}\]

\[\ce{E^{\circ}_{Mn^{+2}/MnO^-_4}}\] = −1.510 V

\[\ce{\frac{1}{2} O2 + 2H+ + 2e- -> H2O}\]

\[\ce{E^{\circ}_{O_2/H_2O}}\] = +1.223 V

Will the permanganate ion \[\ce{MnO^-_4}\] liberate O₂ from water in presence of an acid?

Find the emf of the cell in which the following reaction takes place at 298 K.

\[\ce{Ni_{(s)} + 2Ag+ (0.001 M) -> Ni^{+2} (0.001 M) + 2Ag_{(s)}}\]

(Given that \[\ce{E{^{\circ}_{cell}}}\] = 1.05 V, \[\ce{\frac{2.303 RT}{F}}\] = 0.059 at 298 K)

The `((del E)/(del T))_P`of different types of half cells are as follows:

(where E is the electromotive force)

Which of the above half cells would be preferred to be used as reference electrode?

The correct order of reduction potentials of the following pairs is:

1. Cl₂/Cl
2. I₂/I⁻
3. Ag⁺/Ag
4. Na⁺/Na
5. Li⁺/Li

Choose the correct answer from the options given below:

The conduction of current by a conductor is either due to flow of ______ or due to the movement of ______.

The electrolytic conduction ______ with a rise in temperature.

The fraction of total number of moles which undergoes dissociation is called ______.

The ions remain ______ in solution.

Higher the interaction between solute and solvent particles ______ is the salvation of ions and ______ is the current conducted by ions.

The properties of an electrolytic solution are the properties of ______ present in the solution.

The degree of dissociation of a weak electrolyte is ______ proportional to the ______ of its molar concentration.

According to Ohm’s law, the potential difference across a conductor is ______ proportional to the ______ flowing through it.

The unit of conductance is ______ which is also expressed as ______.

Equivalent conductivity is the conducting power of all the ions furnished by one ______ of an electrolyte present in a definite ______ of the solution.

D.C. cannot be used to measure the electrical conductivity of an electrolytic solution because it causes ______ effect.

Conductivity water has a very ______ conductance.

According to Debye-Huckel-Onsagar equation, `Lambda_m = Lambda_m^infty -` ______.

The molar conductivity of a weak electrolytic solution is much ______ than that of a strong electrolytic solution.

The salt bridge used in a galvanic cell completes the ______ and maintains the ______ of the solutions present in the two ______.

While representing a galvanic cell, the oxidation electrode is always written at the ______ while the reduction electrode at the ______.

The electrode potential of a standard hydrogen electrode is ______.

When a SHE is connected to a Zn²⁺/Zn electrode, the electrons flow from ______ electrode to the ______ electrode and the current flows in the ______ direction.

The electrode systems having positive values of standard electrode potentials act as ______ when connected to a standard hydrogen electrode.

F₂ is the ______ oxidising agent because it possesses the ______ value of standard electrode potential.

A metal can displace any other metal placed ______ it in the electrochemical series from its salt solution.

A cell reaction is feasible when the EMF of the cell is ______.

The cell which converts ______ energy into ______ energy is called an electrolytic cell.

The electrolysis of fused sodium chloride gives ______ at the cathode and ______ at the anode, whereas the products obtained on electrolysis of an aqueous solution of sodium chloride are ______ at the cathode and ______ at the anode.

During the electrolysis of an aqueous solution of sodium chloride, the chloride ions get oxidised at anode in preference to water due to ______.

If the cation produced from the electrolyte has higher reduction potential than that of water, the ______ will get reduced at the cathode.

The electrochemical equivalent is the amount of substance liberated by a current of ______ ampere passed for ______ second.

The quantity of electricity required to liberate one equivalent of any substance is ______ Faraday.

Reduction of 1 mole of \[\ce{MnO^-_4}\] to Mn²⁺ requires ______ coulombs of charge.

The cell potential of a dry cell is approximately ______ V to ______ V.

Assertion: When a Zn | Zn²⁺ electrode is connected to a standard hydrogen electrode, the electrons flow from zinc electrode to the hydrogen electrode.

Reason: The standard reduction potential of zinc electrode is lower than that of hydrogen electrode.

Assertion: The specific conductivity of an electrolytic solution decreases with increase in dilution.

Reason: Specific conductivity is the conductance due to the ions present in 1 cm³ of the solution.

Assertion: The electrode potential of a standard hydrogen electrode is equal to zero.

Reason: The value has been obtained on the basis of experimental observations.

Assertion: Copper can displace zinc from the solution of zinc sulphate.

Reason: Copper is placed below zinc in the electrochemical series.

Assertion: When \[\ce{E{^{\circ}_{cell}}}\] is positive, the cell reaction is feasible and proceeds spontaneously.

Reason: When \[\ce{E{^{\circ}_{cell}}}\] is positive, ΔG° is negative which makes the cell reaction feasible.

When a conductivity cell was filled with 0.02 M KCl, it had a resistance of 82.4 ohm at 25°C and when filled with 0.005 N K₂SO₄ it had a resistance of 326 ohm. Calculate

1. cell constant
2. specific conductivity, and
3. equivalent conductivity of 0.005 N K₂SO₄ solution.

The specific conductance of 0.02 M KCl is 0.002768 ohm⁻¹ cm⁻¹.

The resistance of a 0.5 M solution of an electrolyte in a cell was found to be 45 ohm. Calculate the molar conductivity of the solution if the electrodes in the cell are 2.2 cm apart and have an area of 3.8 cm².

During electrolysis, 4828.4 mg of iodine requires 3671.3 coulombs of electricity for its deposition at the anode. Calculate the value of Faraday’s constant.

Calculate the volume of Cl₂ at STP produced during electrolysis of MgCl₂ which produces 6.50 g of Mg. (At. wt. of Mg = 24.3)

A current of 1.70 amp is passed through 500 mL of 0.16 M solution of ZnSO₄ for 230 seconds with a current efficiency of 90%. Find the molarity of Zn²⁺ after the deposition of zinc. Assume that the volume of the solution remains constant during electrolysis.

The reaction \[\ce{Zn_{(s)} + Co^{2+} -> Co_{(s)} + Zn^{2+}}\] occurs in a cell. Write the electrode reactions and compute the standard emf of the cell. 

Given that \[\ce{E^{\circ}_{{Zn/{Zn}^{2+}}}}\] = +0.76 V and \[\ce{E^{\circ}_{{Co/{Co}^{2+}}}}\] = 0.2 V.

There are two hydrogen electrodes in a cell. The anode is in contact with the hydrogen ion concentration of 10⁻⁶ M. The emf of the cell at 25°C is 0.118 volt. Calculate the hydrogen ion concentration at cathode.

How many grams of silver could be plated out on a serving tray by electrolysis of a solution containing silver in +1 oxidation state for a period of 8.0 hours at a current of 8.46 amperes? What is the area of the tray if the thickness of the silver plating is 0.00254 cm? Density of silver is 10.5 g cm⁻³.

Calculate the quantity of electricity required to reduce 6.15 g of nitrobenzene to aniline if the current efficiency is 68 per cent. If potential drop across the cell is 7.0 volts, calculate the energy consumed in the process.

Calculate the pH of the following half cell solutions (at 25°C):

\[\ce{Pt, H2 (1 atm) | HCl; E_{H/H^+}}\] = 0.25 V

Calculate the pH of the following half cell solutions (at 25°C):

\[\ce{Pt, H2 (1 atm) | H2SO; E_{H/H^+}}\] = 0.3 V

Arrange the following metals in the order in which they displace each other from the solution of their salts.

Al, Cu, Fe, Mg and Zn.

Given the standard electrode potentials,

K⁺/K = −2.93 V, Ag⁺/Ag = 0.80 V,

Hg²⁺/Hg = 0.79 V

Mg²⁺/Mg = −2.37 V, Cr³⁺/Cr = −0.74 V

Arrange these metals in their increasing order of reducing power.

Depict the galvanic cell in which the reaction takes place.

\[\ce{Zn_{(s)} + 2Ag+_{ (aq)} -> Zn^{2+}_{ (aq)} + 2Ag_{(s)}}\]

Further show:

1. Which of the electrode is negatively charged?
2. The carriers of the current in the cell.
3. Individual reaction at each electrode.

Calculate the standard cell potential of a galvanic cell in which the following reaction takes place:

\[\ce{2Cr_{(s)} + 3Cd{^{2+}_{(aq)}} -> 2Cr{^{3+}_{(aq)}} + 3Cd}\]

Calculate the Δ_rG° and equilibrium constant of the reaction.

Calculate the standard cell potential of a galvanic cell in which the following reaction takes place:

\[\ce{Fe{^{2+}_{(aq)}} + Ag{^{+}_{(aq)}} -> Fe{^{3+}_{(aq)}} + Ag_{(s)}}\]

Calculate the Δ_rG° and equilibrium constant of the reaction.

Write the Nernst equation and emf of the following cell at 298 K:

Mg_(s) | Mg²⁺ (0.001 M) || Cu²⁺ (0.0001 M) | Cu_(s)

Write the Nernst equation and emf of the following cell at 298 K:

Fe_(s) | Fe²⁺ (0.001 M) || H⁺ (1 M) | H_2(g) (1 bar) | Pt_(s)

Write the Nernst equation and emf of the following cell at 298 K:

\[\ce{Sn_{(s)} | Sn^{2+} (0.050 M) || H^+ (0.020 M) | H2_{(g)} (1 bar) | Pt_{(s)}}\]

Write the Nernst equation and emf of the following cell at 298 K:

Pt_(s) | Br⁻ (0.010 M) | Br_2(l) || H⁺ (0.030 M) | H_2(g) (1 bar) | Pt_(s)

In the button cells widely used in watches and other devices, the following reaction takes place:

\[\ce{Zn_{(s)} + Ag2O_{(s)} + H2O_{(l)} -> Zn{^{2+}_{(aq)}} + 2Ag_{(s)} + 2OH^-_{ (aq)}}\]

Determine Δ_rG° and E° for the reaction.

Define conductivity for the solution of an electrolyte.

Define “Molar conductivity”.

Discuss the variation of conductivity and molar conductivity with concentration.

The conductivity of 0.20 M solution of KCl at 298 K is 0.0248 S cm⁻¹. Calculate its molar conductivity.

The resistance of a conductivity cell containing 0.001 M KCl solution at 298 K is 1500 Ω. What is the cell constant if the conductivity of 0.001 M KCl solution at 298 K is 0.146 × 10⁻³ S cm⁻¹.

The conductivity of sodium chloride at 298 K has been determined at different concentrations and the results are given below:

Calculate ∧_m for all concentrations and draw a plot between ∧_m and c^1/2. Find the value of `∧_m^0`.

Conductivity of 0.00241 M acetic acid is 7.896 × 10⁻⁵ S cm⁻¹. Calculate its molar conductivity and if `∧_m^0` for acetic acid is 390.5 S cm² mol⁻¹, what is its dissociation constant?

How much charge is required for the following reduction:

1 mol of Al³⁺ to Al?

How much charge is required for the following reduction? 

1 mol of Cu²⁺ to Cu.

How much charge is required for the following reduction:

1 mol of \[\ce{MnO^-_4}\] to Mn²⁺?

How much electricity in terms of Faraday is required to produce 20 g of Ca from molten CaCl₂?

(Given: Molar mass of Calcium is 40 g mol⁻¹.)

How much electricity in terms of Faraday is required to produce 40.0 g of Al from molten Al₂O₃?

(Given: Molar mass of Aluminium is 27 g mol⁻¹.)

How much electricity is required in coulomb for the oxidation of 1 mol of H₂O to O₂?

How much electricity is required in coulomb for the oxidation of 1 mol of FeO to Fe₂O₃?

A solution of Ni(NO₃)₂ is electrolysed between platinum electrodes using a current of 5 amperes for 20 minutes. What mass of Ni is deposited at the cathode?

Three electrolytic cells A, B, C containing solutions of ZnSO₄, AgNO₃ and CuSO₄, respectively, are connected in series. A steady current of 1.5 amperes was passed through them until 1.45 g of silver deposited at the cathode of cell B. How long did the current flow? What mass of copper and zinc were deposited?

Using the standard electrode potentials, predict if the reaction between the following is feasible:

\[\ce{Fe^{3+}_{ (aq)}}\] and \[\ce{I^-_{ (aq)}}\]

Using the standard electrode potentials, predict if the reaction between the following is feasible: 

\[\ce{Ag{^{+}_{(aq)}}}\] and Cu_(aq)

Using the standard electrode potentials, predict if the reaction between the following is feasible:

\[\ce{Fe^{3+}_{ (aq)}}\] and \[\ce{Br^-_{ (aq)}}\]

Using the standard electrode potential, predict if the reaction between the following is feasible:

Ag_(s) and \[\ce{Fe^{3+}_{ (aq)}}\]

Using the standard electrode potential, predict if the reaction between the following is feasible:

Br_2(aq) and \[\ce{Fe^{2+}_{ (aq)}}\]

Predict the product of electrolysis in the following:

An aqueous solution of AgNO₃ with silver electrodes.

Predict the product of electrolysis in the following:

An aqueous solution of AgNO₃ with platinum electrodes.

Predict the product of electrolysis in the following:

A dilute solution of H₂SO₄ with platinum electrodes.

Predict the product of electrolysis in the following:

An aqueous solution of CuCl₂ with platinum electrodes.

Correct the following statement:

Copper displaces both zinc and silver from their solutions.

A conductivity cell has a cell constant of 0.5 cm⁻¹. This cell when filled with 0.01 M sodium chloride solution, has a resistance of 384 ohms at 25°C. Calculate the equivalent conductivity of 0.01 M sodium chloride at 25°C.

For the cell, Cu | Cu²⁺ (0.13 M) || Ag⁺ (0.01 M) | Ag,

1. Calculate the reduction potential of each electrode if the standard reduction potential for copper and silver electrodes are 0.34 V and 0.80 V respectively.
2. Calculate the emf of the cell.
3. Write the cell reaction.
4. Is this cell reaction spontaneous? Why?

Correct the following statement:

For a strong electrolyte, the plot of equivalent conductance versus concentration of the solution is a straight line.

A current of 4 amperes is passed through a molten solution for 45 minutes. 2.977 g of metal is deposited. Calculate the charge carried by the metal cation if its atomic mass is 106.4 g mol⁻¹.

As concentration of an electrolyte solution increases, its ______ conductance increases but ______ conductance decreases.

A cell is constructed by dipping a zinc rod in 0.1 M zinc nitrate solution and a lead rod in 0.2 M lead nitrate solution.

\[\ce{E^{\circ}_{Pb^{2+}/Pb}}\] = −0.13 V and \[\ce{E^{\circ}_{Zn^{2+}/Zn}}\] = −0.76 V

1. Write the spontaneous cell reaction.
2. Calculate standard emf and emf of the cell.

The number of Faradays which will be required to reduce 4 gram equivalents of Cu²⁺ to Cu metal is ______.

For the cell:

Zn | Zn²⁺ (a = 1) || Cu²⁺ (a = 1) | Cu

Given that \[\ce{E^{\circ}_{Zn/Zn^{2+}}}\] = 0.761 volt; \[\ce{E^{\circ}_{Cu^{2+}/Cu}}\] = +0.339 volt

1. Write the cell reaction.
2. Calculate the emf and free energy change at 298 K involved in the cell.

[Faraday’s constant, F = 96500 coulomb mol⁻¹]

Calculate the equivalent conductivity of 1 M H₂SO₄, whose specific conductivity is 26 × 10⁻² ohm⁻¹ cm⁻¹.

A current of 10 A is passed for 80 min. and 27 seconds through a cell containing dilute sulphuric acid.

1. How many moles of oxygen gas will be liberated at the anode?
2. Calculate the amount of zinc deposited at the cathode when another cell containing ZnSO₄ solution is connected in series (Zn = 65).

Calculate:

E_cell at 25°C for the reaction:

\[\ce{Zn + Cu{^{2+}} (0.20 M) -> Zn{^{2+}} (0.50 M) + Cu}\]

Given: \[\ce{E^{\circ}_{({Zn^{2+}/{Zn}})}}\] = −0.76 volt, \[\ce{E^{\circ}_{({Cu^{2+}/{Cu}})}}\] = 0.34 volt

Give a reason for the following:

Specific conductance decreases with dilution, whereas equivalent conductance increases with dilution.

When zinc granule is dipped in copper sulphate solution, copper is precipitated because ______.

In a galvanic cell, the movement of electrons in the external circuit is from ______ to ______.

Zinc displaces hydrogen from acid solution (\[\ce{E^{\circ}_{Zn^{2+}/Zn}}\] = −0.76 V). Explain.

A 0.05 M NaOH solution offered a resistance of 31.6 ohms in a conductivity cell. If the cell constant of the conductivity cell is 0.378 cm⁻¹, determine the molar conductivity of sodium hydroxide solution at this temperature.

Mention any two factors affecting the electrode potential of a metal.

For the following cell, calculate the emf:

\[\ce{Al | Al^{3+} (0.01 M) || Fe^{2+} (0.02 M) | Fe}\]

Given: \[\ce{E^{\circ}_{{Al^{3+}/{Al}}}}\] = −1.66 V, \[\ce{E^{\circ}_{{Fe^{2+}/{Fe}}}}\] = −0.44 V

When zinc granule is dipped in copper sulphate solution, copper is precipitated because ______.

State Faraday’s first law of electrolysis.

How many electrons will flow when a current of 5 amperes is passed through a solution for 200 seconds?

Consider the reaction:

\[\ce{2Ag+ + Cd -> 2Ag + Cd^{2+}}\]

The reduction potentials of Ag⁺/Ag and Cd²⁺/Cd are +0.80 volt and −0.40 volt respectively.

1. Give the cell representation.
2. What is the standard cell emf, E°?
3. What will be the emf of the cell if concentration of Cd²⁺ is 0.1 M and Ag⁺ is 0.2 M?
4. Will the cell work spontaneously for the condition given in (c) above?

Zinc can displace ______ from CuSO₄ solution, but cannot displace ______ from MgSO₄ solution.

The quantity of electricity required to deposit 1.15 g of sodium from molten NaCl (Na = 23, Cl = 35.5) is ______.

The conductivity of 0.2 M KCl solution is 3 × 10⁻² ohm⁻¹ cm⁻¹. Calculate its molar conductance.

Magnesium displaces hydrogen from dilute acid solution because ______.

Calculate the number of coulombs required to deposit 20.25 g of aluminium (at. mass = 27) from a solution containing Al⁺³.

What do you understand by specific conductance of a solution?

What are the units of specific conductivity?

How is equivalent conductivity related to specific conductivity?

2.5 amperes of current is passed through copper sulphate solution for 30 minutes. Calculate the number of copper atoms deposited at the cathode (Cu = 63.54).

Four metals W, X, Y and Z have the following values of E_red:

E_red

W = −0.140 V 

X = −2.93 V 

Y = +0.80 V

Z = +1.50 V

Arrange them in the increasing order of reducing power.

A current liberates 0.50 g of hydrogen in 2 hours. The weight of copper (at. wt. = 63.5) deposited at the same time by the same current through copper sulphate solution is ______.

What happens when a nickel rod is dipped into a copper sulphate solution? Justify your answer.

[\[\ce{E{^{\circ}_{Ni^{+2}/Ni}}}\] = − 0.25 V and \[\ce{E{^{\circ}_{Cu^{+2}/Cu}}}\] = +0.34 V]

What is standard hydrogen electrode?

0.05 M NaOH solution offered resistance of 31.6 ohms in a conductivity cell at 298 K. If the cell constant of the cell is 0.367 cm⁻¹, calculate the molar conductivity of the NaOH solution.

0.3605 g of a metal is deposited on the electrode by passing 1.2 amperes of current for 15 minutes through its salt solution. The atomic weight of the metal is 96. What is the valency of the metal?

How many hours does it take to reduce 3 moles of Fe³⁺ to Fe²⁺ with 2.0 A current intensity?

The specific conductance of a 0.01 M solution of acetic acid at 298 K is 1.65 × 10⁻⁴ ohm⁻¹ cm⁻¹. The molar conductance at infinite dilution for H⁺ ion and CH₃COO⁻ ion is 349.1 ohm⁻¹ cm² mol⁻¹ and 40.9 ohm⁻¹ cm² mol⁻¹ respectively.

Calculate:

1. Molar conductance of the solution.
2. The degree of dissociation of CH₃COOH.
3. A dissociation constant for acetic acid.

Calculate the e.m.f. of the following cell reaction at 298 K:

\[\ce{Mg_{(s)} + Cu{^{2+}} (0.0001 M) -> Mg{^{2+}} (0.001M) + Cu_{(s)}}\]

The standard potential (E°) of the cell is 2.71 V.

Consider the following cell reaction at 298 K:

\[\ce{2Ag+ + Cd -> 2Ag + Cd^{2+}}\]

The standard reduction potentials (E°) for Ag⁺/Ag and Cd²⁺/Cd are 0.80 V and −0.40 V respectively.

1. Write the cell representation.
2. What will be the emf of the cell if concentration of Cd²⁺ is 0.1 M and that of Ag⁺ is 0.2 M?
3. Will the cell work spontaneously for the condition given in (b) above?

A 0.05 M NH₄OH solution offers the resistance of 50 ohms to a conductivity cell at 298 K. If the cell constant is 0.50 cm⁻¹ and molar conductance of NH₄OH at infinite dilution is 471.4 ohm⁻¹ cm² mol⁻¹, calculate:

1. Specific conductance
2. Molar conductance
3. Degree of dissociation

On dilution of a solution, its specific conductance ______ while its equivalent conductance ______.

More ______ the standard reduction potential of a substance, the ______ is its ability to displace hydrogen from acids.

For a spontaneous reaction ΔG° and \[\ce{E{^{\circ}_{cell}}}\] will be respectively:

Specific conductivity of 0.20 M solution of KCl at 298 K is 0.025 S cm⁻¹. Calculate its molar conductivity.

The following electrochemical cell is set up at 298 K:

\[\ce{Zn | Zn^{2+}_{ (aq)} (1 M) || Cu^{2+}_{ (aq)} (1 M) | Cu}\]

Given: \[\ce{E^{\circ}_{Zn^{2+}/Zn}}\] = −0.761 V, \[\ce{E^{\circ}_{Cu^{2+}/Cu}}\] = +0.339 V

Write the cell reaction.

The following electrochemical cell is set up at 298 K: 

\[\ce{Zn | Zn^{2+}_{ (aq)} (1 M) || Cu^{2+}_{ (aq)} (1 M) | Cu}\]

Given: \[\ce{E^{\circ}_{Zn^{2+}/Zn}}\] = −0.761 V, \[\ce{E^{\circ}_{Cu^{2+}/Cu}}\] = +0.339 V

Calculate the emf and free energy change at 298 K.

What is the effect of temperature on the ionic product of water (K_w)?

What happens to the ionic product of water (K_w) if some acid is added to it?

Name the law or principle to which the following observations confirm:

When 9650 coulombs of electricity is passed through a solution of copper sulphate, 3.175 g of copper is deposited on the cathode (at. wt. of Cu = 63.5).

A galvanic cell converts ______ energy into ______ energy.

Calculate the emf and ΔG° for the cell reaction at 25°C:

\[\ce{\underset{(0.1 M)}{Zn_{(s)} | Zn^{2+}_{ (aq)}} || \underset{(0.01 M)}{Cd^{2+}_{ (aq)} | Cd_{(s)}}}\]

Given \[\ce{E^{\circ}_{Zn^{2+}/Zn}}\] = −0.763 and \[\ce{E^{\circ}_{Cd^{2+}/Cd}}\] = −0.403

Define the term equivalent conductance.

Define the following term:

Corrosion of metals

The specific conductivity of a solution containing 5 g of anhydrous BaCl₂ (mol. wt. = 208) in 1000 cm³ of a solution is found to be 0.0058 ohm⁻¹ cm⁻¹. Calculate the molar and equivalent conductivity of the solution.

What do you understand by electrochemical series?

How is the electrochemical series useful in predicting whether a metal can liberate hydrogen from acid or not?

Calculate the mass of silver deposited at cathode when a current of 2 amperes is passed through a solution of AgNO₃ for 15 minutes.

(at. wt. of Ag = 108, 1 F = 96500 C)

Calculate the emf and ΔG for the cell reaction at 298 K:

Mg_(s) | Mg²⁺ (0.1 M) || Cu²⁺ (0.01 M) | Cu_(s)

Given: \[\ce{E^{\circ}_{cell}}\] = 2.71 V

1 F = 96500 C

Define the following term:

Specific conductance

Define the following term:

Kohlrausch’s Law.

The resistance of a conductivity cell containing 0.001 M KCl solution at 298 K is 1500 ohm. What is the cell constant and molar conductivity of 0.001 M KCl solution, if the conductivity of this solution is 0.146 × 10⁻³ ohm⁻¹ cm⁻¹ at 298 K?

Calculate the molar conductivity at infinite dilution for CH₃COOH if the molar conductivity at infinite dilution for NaCl, HCl and CH₃COONa are 126.45, 426.16 and 91.0 ohm⁻¹ cm² mol⁻¹ respectively.

Calculate the emf and ΔG for the given cell at 25°C:

Cr_(s) | Cr³⁺ (0.1 M) || Fe²⁺ (0.01 M) || Fe_(s)

Given: \[\ce{E^{\circ}_{Cr^{3+}/Cr}}\] = −0.74 V, \[\ce{E^{\circ}_{Fe^{2+}/Fe}}\] = −0.44 V

(1 F = 96500 C mol⁻¹, R = 8.314 JK⁻¹ mol⁻¹)

Calculate the degree of dissociation (α) of acetic acid if its molar conductivity (Λ_m) is 39.05 S cm² mol⁻¹.

(Given \[\ce{\lambda^{\circ}_{(H^+)}}\] = 349.6 S cm² mol⁻¹ and \[\ce{\lambda^{\circ}_{(CH_3COO^-)}}\] = 40.95 S cm² mol⁻¹)

Which of the following statements is true for electrochemical cell?

Electrochemical equivalent (z) is the amount of substance which gets deposited from its solutions on passing electrical charge equal to solutions on passing electrical charge equal to ______.

The standard electrode potential of Cu⁺²/Cu is +0.34 V and that of Cr⁺³/Cr is −0.74 V. These two electrodes are connected in their standard state to make an electrochemical cell. What will be the standard electrode potential (E°) of this cell?

The ionic conductance at infinite dilution for Ba⁺² and Cl⁻ ions are 127 ohm⁻¹ cm² mol⁻¹ and 76 ohm⁻¹ cm² mol⁻¹ respectively. What will be the molar conductance of BaCl₂ at infinite dilution?

A conductivity cell is filled with 0.05 M NaOH solution offering a resistance of 31.6 ohm. If the cell constant of the cell is 0.347 cm⁻¹, calculate the following:

The value of specific conductance:

A conductivity cell is filled with 0.05 M NaOH solution offering a resistance of 31.6 ohm. If the cell constant of the cell is 0.347 cm⁻¹, calculate the following:

The value of molar conductance:

The standard electrode potential for Sn⁺⁴/Sn⁺² couple is +0.15 V and that for the Cr⁺³/Cr couple is −0.74 V. These two couples in their standard states are connected to make a cell (1 Faraday = 96500 mol⁻¹)

What will be the value of \[\ce{E{^{\circ}_{cell}}}\]?

The standard electrode potential for Sn⁺⁴/Sn⁺² couple is +0.15 V and that for the Cr⁺³/Cr couple is −0.74 V. These two couples in their standard states are connected to make a cell. (1 Faraday = 96500 mol⁻¹)

What will be the value of standard Gibbs energy (ΔG°)?

A solution containing 2 g of anhydrous barium chloride in 400 cm³ of water has a specific conductivity of 0.0058 S cm⁻¹. (at. wt. of Ba = 137, Cl = 35.5)

What is the molarity of the above solution?

A solution containing 2 g of anhydrous barium chloride in 400 cm³ of water has a specific conductivity of 0.0058 S cm⁻¹. (at. wt. of Ba = 137, Cl = 35.5)

What is the molar conductivity of the above solution?