Electro Chemistry: One Mark Questions, Book Back Answers and Notes

About Electro Chemistry

Hard ~120 min study

A torch cell turns a chemical reaction into electric current. Electrolysis of alumina dissolved in molten cryolite does the reverse, using current to force a reaction that would not happen on its own. Electrochemistry studies both directions of this conversion.

The chapter first looks at how solutions carry current. In a metal, electrons move; in an electrolyte, ions move. By measuring resistance you can find specific and molar conductivity and see how they change on dilution. The second half turns to cells: how a spontaneous redox reaction is split into two half-cells, how the cell's emf is measured against the standard hydrogen electrode, and how the Nernst equation corrects that emf for real concentrations.

The numerical problems follow a few set patterns: calculate molar conductivity from κ and concentration, use Kohlrausch's law to find \(\Lambda_{m}^{0}\) of acetic acid, find the emf of a cell from E° values, apply the Nernst equation, and calculate the mass deposited by a given current in a given time. The link \(\Delta G = -nFE_{cell}\) also appears often, since it tells you whether a cell reaction is spontaneous.

What you'll learn

• Distinguish metallic from electrolytic conductors and explain how temperature affects each.
• Calculate molar and equivalent conductance from resistance, cell constant and concentration.
• Use Kohlrausch's law to find the limiting molar conductivity of a weak electrolyte.
• Write cell notation for a galvanic cell using IUPAC conventions.
• Predict whether a redox reaction is feasible from E° values and the electrochemical series.
• Solve problems with the Nernst equation for non-standard concentrations.
• Write the reactions at each electrode of a lead storage battery during discharge and recharge.

Before you start

• Oxidation, reduction and balancing redox reactions.
• Electrical units: volt, ampere, ohm and coulomb.
• Chemical equilibrium and the law of mass action.

Electro Chemistry explained

Conductance, Cells and Electrolysis

Metallic and Electrolytic Conduction

Metals conduct by the flow of electrons, and their conductance falls as temperature rises because vibrating metal ions get in the way. Electrolytes conduct by the movement of ions, and their conductance rises with temperature because the solvent becomes less viscous and ion–ion attraction weakens. Resistance \(R\) is measured in ohms, and its reciprocal, conductance \(C\), is measured in ohm^-1 or siemens (S).

Specific and Molar Conductivity

Specific conductivity \(\kappa\) is the conductance of a solution placed between two electrodes of 1 m² area kept 1 m apart. To compare electrolytes, we use molar conductivity \(\Lambda_{m}\) = κ/C (with suitable unit conversion), the conductance of all the ions from one mole of electrolyte. On dilution κ decreases, because there are fewer ions per unit volume, but \(\Lambda_{m}\) increases. For a strong electrolyte such as KCl, \(\Lambda_{m}\) decreases linearly as √C increases (the Debye–Huckel–Onsager relation), so a graph of \(\Lambda_{m}\) against √C is a straight line that can be extended to zero concentration. A weak electrolyte such as acetic acid shows a steep rise in \(\Lambda_{m}\) only at very high dilution, so its limiting value cannot be read from the graph.

Kohlrausch's Law

At infinite dilution each ion contributes a fixed amount to the molar conductivity, whatever ion it is paired with. So \(\Lambda_{m}^{0}\)(CH₃COOH) = \(\Lambda_{m}^{0}\)(CH₃COONa) + \(\Lambda_{m}^{0}\)(HCl) − \(\Lambda_{m}^{0}\)(NaCl). The law gives \(\Lambda_{m}^{0}\) for weak electrolytes, the degree of dissociation α = \(\Lambda_{m}\)/\(\Lambda_{m}^{0}\), the dissociation constant of weak acids and the solubility product of sparingly soluble salts like AgCl.

Galvanic and Electrolytic Cells

A galvanic cell such as the Daniell cell produces electricity from a spontaneous reaction. Oxidation occurs at the anode, which is negative, and reduction at the cathode, which is positive. A salt bridge of KCl or KNO₃ in agar keeps both half-cells electrically neutral. An electrolytic cell uses an external source to drive a non-spontaneous reaction. Oxidation is still at the anode, but now the anode is the positive electrode.

Electrode Potential and the Nernst Equation

Electrode potentials are measured against the standard hydrogen electrode, whose potential is taken as zero. Arranging standard reduction potentials in order gives the electrochemical series. For a cell, E°_cell = E°_cathode − E°_anode. The Nernst equation gives the emf at other concentrations: at 298 K, E_cell = E°_cell − (0.0591/n) log Q, where Q is the reaction quotient. The emf is related to free energy by \(\Delta G = -nFE_{cell}\), so a positive E_cell means a negative ΔG and a spontaneous reaction.

Faraday's Laws of Electrolysis

The mass of substance liberated at an electrode is proportional to the charge passed, m = ZIt, where Z is the electrochemical equivalent. When the same charge passes through different electrolytes, the masses liberated are proportional to their equivalent masses. One faraday, taken as 96500 C, is the charge on one mole of electrons.

Batteries and Fuel Cells

Primary cells such as the Leclanche dry cell and the mercury button cell cannot be recharged. Secondary cells such as the lead storage battery and the lithium-ion battery can, because their reactions reverse when current is passed in the opposite direction. In a lead storage battery, PbSO₄ forms at both electrodes during discharge and sulphuric acid is used up. The hydrogen-oxygen fuel cell converts the energy of H₂ and O₂ directly into electricity with water as the product, and it was used in space missions.

Corrosion and Its Prevention

Rusting of iron is an electrochemical process. At anodic spots iron is oxidised to Fe²⁺; at cathodic spots oxygen is reduced in the presence of water. Fe²⁺ is further oxidised to hydrated ferric oxide, rust. Corrosion is prevented by coating with paint or oil, by galvanising (coating iron with zinc), by sacrificial protection with a more reactive metal such as magnesium, and by alloying.

Common mistakes to avoid

• Mixing up electrode signs. Oxidation is always at the anode, but the anode is negative in a galvanic cell and positive in an electrolytic cell.
• Leaving time in minutes or hours in Faraday's law problems. Charge Q = It needs time in seconds.
• Taking the wrong value of n in the Nernst equation. Count the electrons in the balanced overall reaction, not in one half-reaction written separately.
• Multiplying E° values by stoichiometric coefficients when adding half-reactions. Electrode potential is an intensive property, so it does not change when a half-reaction is multiplied.

Test yourself on these with the practice test, then check the worked reasoning in the solved MCQs.

Frequently asked questions

Why does specific conductivity decrease on dilution?

Specific conductivity is the conductance of a fixed volume of solution, so it depends on how many ions are in that volume. Dilution spreads the ions through more solvent, leaving fewer ions per cm3, and κ falls. Molar conductivity rises at the same time because it counts all the ions from one mole, however much volume they occupy.

What is the difference between a primary and a secondary cell?

In a primary cell, like the Leclanche dry cell, the reaction cannot be reversed, so the cell is thrown away once the reactants are used. In a secondary cell, like the lead storage battery, passing current from an outside source reverses the reaction and restores the electrodes, so the cell can be recharged many times.

How does temperature affect conductance of an electrolyte?

It increases it, which is the opposite of metals. A hotter solvent is less viscous and the attraction between oppositely charged ions weakens, so ions move faster towards the electrodes and the solution conducts better.

Can we measure the potential of a single electrode?

No. A voltmeter needs two electrodes to complete the circuit, so only a potential difference can be measured. We therefore pair each electrode with the standard hydrogen electrode, which is given a potential of zero, and report the measured emf as that electrode's potential.

What is a salt bridge and why is it used?

It is a U-tube filled with a gel of agar and an inert electrolyte such as KCl or KNO3, joining the two half-cells. Its ions move into each compartment to balance the charge building up there. Without it, one side would become positively charged and the other negatively charged, and the current would stop almost at once.

Why is zinc used to galvanise iron?

Zinc has a lower (more negative) standard reduction potential than iron, so it is oxidised more easily. It corrodes in place of iron, acting as a sacrificial anode. Even if the zinc coating is scratched, the exposed iron stays protected as long as zinc remains nearby.