Electrochemistry – NEET UG ChemistryPractice Questions & PYQs

The oxidation of manganate to permanganate involves the loss of one electron per ion, meaning the change in oxidation state or valency factor is one. The total charge required is calculated using the relation $Q = n \times z \times F$, where $n$ is the number of moles, $z$ is the valency factor, and $F$ is Faraday's constant of 96500 C per mole. Substituting 0.1 moles for $n$ and 1 for $z$ gives $Q = 0.1 \times 1 \times 96500$, resulting in a required electricity quantity of 9650 C.

Faraday's first law of electrolysis determines the mass of a substance deposited at an electrode using the relationship $W = \frac{ItE}{96500}$ where $E$ represents the equivalent mass obtained by dividing the atomic mass by the valence. Substituting the current of 10 amperes, time of 109 minutes converted to seconds, atomic mass of 59 u, and valence of 2 into this expression yields 20.0 grams of deposited cobalt metal.

The standard cell potential of a redox reaction is determined by combining the individual half-cell potentials while accounting for the reversal of oxidation processes. Reversing the manganese oxidation half-reaction changes the sign of its standard potential, and summing these values yields an overall standard cell potential of -2.69 V. Because this resulting standard cell potential is negative, the corresponding Gibbs free energy change is positive, indicating that the forward reaction is non-spontaneous.

The relationship between Gibbs free energy and cell potential is given by the equation $\Delta G^\circ = -nFE^\circ$. To find the minimum potential difference, we determine the total number of electrons transferred based on the balanced reaction where four-thirds moles of aluminium are produced, yielding a total of 4 moles of electrons transferred. Substituting the free energy change of 960000 J mol^-1 and the Faraday constant into the rearranged equation $E^\circ = -\frac{\Delta G^\circ}{nF}$, we calculate the standard potential as -2.48 V. The absolute magnitude of this value gives the minimum external voltage required, which rounds to 2.5 V.

The degree of ionization of a weak electrolyte is determined by the ratio of its molar conductance at a specific concentration to its molar conductance at infinite dilution, represented by the relation $\alpha = \frac{\Lambda_m^c}{\Lambda_m^\infty }$. Substituting the respective values into this formula gives $\alpha = \frac{9.54}{238}$. Performing this division yields 0.04008, which equals 4.008 %.

The total cell potential is determined by combining the oxidation potential of the anode and the reduction potential of the cathode using the equation $E^\circ_{\text{cell}} = E^\circ_{\text{oxidation}} + E^\circ_{\text{reduction}}$

During the operation of this button cell, zinc metal undergoes oxidation, which reverses the sign of its listed reduction potential to yield an oxidation potential of 0.76 V. Simultaneously, silver oxide undergoes reduction with a standard reduction potential of 0.34 V. Summing these two half-cell values together gives a total voltage of 1.10 V.

The Nernst equation relates the oxidation potential of a gas electrode to the standard potential, temperature, and concentrations of the participating species. For a hydrogen electrode undergoing oxidation, the potential is expressed as $E = E^\circ - \frac{0.059}{n}\log\frac{[\mathrm{H}^+]^2}{P_{\mathrm{H}_2}}$. Substituting the standard potential of zero, $n = 2$, pressure of 1 atm, and hydrogen ion concentration of $10^{-10}$ obtained from a pH of 10 results in an oxidation potential of 1.18 V.

Kohlrausch's law of independent migration of ions allows the molar conductivity of a weak electrolyte at infinite dilution to be calculated from the sum of limiting conductivities of strong electrolytes. This principle is represented by the equation $\Lambda^\circ_{\mathrm{CH}_3\mathrm{COOH}} = \Lambda^\circ_{\mathrm{CH}_3\mathrm{COONa}} + \Lambda^\circ_{\mathrm{HCl}} - \Lambda^\circ_{\mathrm{NaCl}}$ where the ionic contributions correctly account for the target acid. Substituting the respective values of 91.0, 425.9, and 126.4 into this expression results in a molar conductivity of 390.5.

Kohlrausch's Law of Independent Migration of Ions states that the limiting molar conductivity of an electrolyte is equal to the sum of the ionic conductivities of its constituent cations and anions. To determine the limiting molar conductivity of the weak electrolyte $\mathrm{NH}_4\mathrm{OH}$, strong electrolytes containing these specific ions must be combined while eliminating unwanted spectator ions. Adding the molar conductivities of $\mathrm{NH}_4\mathrm{Cl}$ and $\mathrm{NaOH}$ introduces the required $\mathrm{NH}_4^+$ and $\mathrm{OH}^-$ ions alongside extra $\mathrm{Na}^+$ and $\mathrm{Cl}^-$ ions, which are subsequently cancelled out by subtracting the molar conductivity of $\mathrm{NaCl}$.

The strength of an oxidizing agent is directly proportional to its standard reduction potential, meaning species with more positive values readily accept electrons and act as powerful oxidants. Conversely, the strength of a reducing agent increases as the reduction potential becomes more negative, making species with low or negative potentials eager to donate electrons. Comparing the given values, fluorine possesses the highest positive reduction potential and acts as the strongest oxidizing agent, whereas the iodide ion corresponds to the most easily oxidized species and functions as the strongest reducing agent.

A redox reaction is thermodynamically favorable and spontaneous when the species with the higher standard reduction potential undergoes reduction while the species with the lower reduction potential undergoes oxidation. Comparing the given standard reduction potentials, the iron(III) to iron(II) couple has a higher potential of +0.77 V compared to the iodine to iodide couple at +0.536 V. Therefore, iron(III) ions will spontaneously undergo reduction to form iron(II) ions, and consequently, iodide ions must be oxidized to molecular iodine to balance the electron transfer.

The thermodynamic feasibility of an electrochemical cell is linked to standard cell potential via the equation $\Delta G^0 = -nFE^0_{\mathrm{cell}}$. When the standard cell potential is negative, the resulting Gibbs free energy change becomes positive, indicating a non-spontaneous process under standard conditions. Furthermore, the equilibrium constant relates to standard Gibbs free energy through $\Delta G^0 = -RT \ln K_{\mathrm{eq}}$, which dictates that a positive free energy change corresponds to an equilibrium constant less than one.

Reducing power is the inherent tendency of a chemical species to lose electrons and undergo oxidation, which increases as the standard reduction potential becomes more negative. A lower electrode potential indicates a stronger reducing agent because the species readily donates electrons to other substances in a reaction. Comparing the given values, metal Z has a potential of -3.0 V, metal X has +1.2 V, and metal Y has +0.5 V. Ordering these values from the most negative to the most positive potential naturally establishes the decreasing trend in reducing capability.

The standard cell potential of a galvanic cell is calculated by subtracting the standard reduction potential of the anode from the standard reduction potential of the cathode, represented by the formula $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$. The species with the more negative reduction potential acts as the anode and undergoes oxidation, while the species with the more positive reduction potential acts as the cathode and undergoes reduction. Assigning the chromium couple at -0.74 V as the anode and the tin couple at +0.15 V as the cathode, we substitute these values into the formula. Subtracting -0.74 V from +0.15 V results in a positive cell potential of 0.89 V.

Gibbs free energy is an extensive property, meaning the free energy change for an overall reaction is the sum of the free energy changes of its individual steps, expressed through the relation $\Delta G^\circ = -nFE^\circ$. Because free energy is additive while electrode potentials are not directly additive, we combine the individual step equations into the overall two-electron reduction of copper(II) to solid copper. Substituting the individual potentials and electron counts into the combined relation yields the equation $E^\circ = \frac{n_1E_1^\circ + n_2E_2^\circ}{n}$. Evaluating this expression with the given values of 0.15 V and 0.50 V for single-electron steps divided by a total of 2 electrons gives the final standard electrode potential.

The relationship linking standard Gibbs free energy change to standard cell potential is expressed through the equation $\Delta G^\circ = -nFE^\circ_{\text{cell}}$. In this redox process, two electrons are transferred between the copper and silver species, establishing that $n = 2$. Multiplying this electron count by the Faraday constant and the cell potential of 0.46 V gives a total work value of -88780 J. Converting this energy from joules to kilojoules produces -88.78 kJ, which rounds appropriately to -89.0 kJ.

Strong electrolytes undergo complete dissociation into their constituent ions across all concentrations, meaning that dilution does not alter the total number of charge carriers present in solution. However, adding more solvent physically separates the ions further apart from one another, which significantly diminishes the electrostatic forces of attraction holding them back. This reduction in interionic hindrance allows the ions to migrate much more freely under an applied electrical field, leading directly to higher ionic mobility and a corresponding rise in equivalent conductance.

Kohlrausch's law of independent migration of ions establishes that the equivalent conductance of any electrolyte at infinite dilution is equal to the sum of the equivalent ionic conductances of its constituent cation and anion. Because equivalent conductance measures the conducting power of one gram equivalent of an electrolyte, stoichiometric coefficients are already accounted for and do not multiply the individual ionic terms. Therefore, the total equivalent conductance is expressed directly as $\Lambda^{\circ}_{\mathrm{Al}^{3+}}+\Lambda^{\circ}_{\mathrm{SO}^{2-}_4}$.

The electromotive force of a galvanic cell is fundamentally determined by the difference between the reduction potentials of the cathode and the anode, or equivalently by the difference in their oxidation potentials. By converting the reduction potential of the anode into its corresponding oxidation potential using the sign-reversal property, the cell potential becomes the sum of the oxidation potential of the anode and the reduction potential of the cathode. Alternatively, expressing both electrodes in terms of their oxidation potentials yields the oxidation potential of the anode minus the oxidation potential of the cathode. These transformations validate the specific mathematical combinations where the anode's oxidation potential is either added to the cathode's reduction potential or paired with the cathode's oxidation potential.

The dissociation constant of a weak electrolyte is related to its degree of dissociation through the Ostwald dilution law expression $K = C\alpha^2$. The degree of dissociation $\alpha$ is first found by dividing the equivalent conductance at the given concentration by the equivalent conductance at infinite dilution, represented by $\alpha = \frac{\Lambda}{\Lambda_\infty }$. Using the given values, $\alpha$ evaluates to 8.0 divided by 400, which equals 0.02. Substituting this degree of dissociation and the concentration of 1/32 into the dissociation constant equation gives $K = \frac{1}{32} \times (0.02)^2$, resulting in $1.25 \times 10^{-5}$.