Question

The standard \(E^{\circ}\) values of few redox couples are \(\mathrm{Zn}^{2+} / \mathrm{Zn}=-0.76 \mathrm{~V}, \mathrm{Ag}^{+} / \mathrm{Ag}=+0.80 \mathrm{~V}\) \(\mathrm{Cu}^{2+} / \mathrm{Cu}=0.34 \mathrm{~V}\). Choose the correct option. (a) Ag can oxidise \(\mathrm{Zn}\) and \(\mathrm{Cu}\). (b) \(\mathrm{Ag}\) can reduce \(\mathrm{Zn}^{2+}\) and \(\mathrm{Cu}^{2+}\). (c) Zn can reduce \(\mathrm{Ag}^{+}\)and \(\mathrm{Cu}^{2+}\). (d) Cu can reduce \(\mathrm{Zn}^{2+}\) and \(\mathrm{Ag}^{+}\).

Short answer

The correct option is (b) Ag can reduce \textrm{Zn}^{2+}\textrm{ and }\textrm{Cu}^{2+}\textrm{.}

Step-by-step solution

Understanding Standard Electrode Potentials

Standard Electrode Potentials (E°) indicate the tendency of a chemical species to be reduced; the more positive the E°, the greater the species' affinity for electrons, and thus the greater its tendency to be reduced. In redox reactions, the species with the higher E° will undergo reduction, while the one with the lower E° will undergo oxidation.

Comparing E° values

To predict which species can oxidize or reduce another, compare their E° values. A species with a higher E° value can be reduced by a species with a lower E° value, which will be oxidized. For example, comparing the E° values of Ag⁺/Ag (+0.80 V) and Zn²⁺/Zn (-0.76 V), it can be deduced that Ag⁺ can reduce Zn (since Ag⁺ has a higher E° value), meaning that Zn can oxidize to Zn²⁺.

Determining the Correct Option

(a) Ag can oxidise Zn and Cu. Incorrect, because Ag has a higher E° value than both Zn²⁺ and Cu²⁺, so it will be reduced, not act as an oxidizing agent.(b) Ag can reduce Zn²⁺ and Cu²⁺. Correct, since Ag⁺ has a higher E° value, it can be reduced by Zn and Cu, which will be oxidized.(c) Zn can reduce Ag⁺ and Cu²⁺. Incorrect, because Zn has a lower E° value, it will be oxidized by both Ag⁺ and Cu²⁺.(d) Cu can reduce Zn²⁺ and Ag⁺. Incorrect, because although Cu²⁺ can reduce Ag⁺, it cannot reduce Zn²⁺ since its E° is higher than that of Cu.

Key concepts

Redox Reactions

Redox reactions are a type of chemical reaction in which oxidation and reduction processes occur simultaneously. In these reactions, one substance transfers electrons to another substance. The principle behind a redox reaction is quite straightforward: for one species to gain electrons (reduction), another must lose them (oxidation).

To identify what is happening in a redox reaction, consider the oxidation states of the elements involved. An increase in oxidation state corresponds to oxidation, while a decrease correlates with reduction. Often, these reactions involve metals and nonmetals, where metals commonly lose electrons to become positive ions, and nonmetals gain these electrons.

An example is evident from the exercise provided, where different metals have the potential to either be oxidized or reduced. Understanding which metals can act as an oxidizing or reducing agent is key to solving such problems, and this depends on their standard electrode potentials.

Oxidation-Reduction

Oxidation-reduction, commonly referred to as 'redox', describes all chemical reactions in which atoms have their oxidation state changed. This can be simply explained: oxidation involves the loss of electrons or an increase in oxidation state by a molecule, atom, or ion, whereas reduction involves a gain of electrons or a decrease in oxidation state.

Take the concept further by bringing in the notion of 'oxidizing agents' and 'reducing agents'. An oxidizing agent, or oxidant, gains electrons and is reduced in a chemical reaction, while a reducing agent, or reductant, loses electrons and is oxidized. Hence, in the exercise, Ag⁺, with a high propensity for electron gain (a more positive standard electrode potential), is a good oxidizing agent and will tend to attract electrons from metals like Zn and Cu, which are reducing agents.

Electrochemical Series

The electrochemical series, also known as the activity series, is a list of elements organized according to their standard electrode potentials. It serves as a handy tool for predicting the outcome of redox reactions. Higher up in this series are elements with greater tendencies to gain electrons and be reduced, while those lower down more readily lose electrons and are oxidized.

In practical terms, by looking at the series or comparing the standard electrode potentials of different elements, one can predict whether a metal can displace another in a solution, as illustrated in the given exercise. Elements like Ag⁺ with higher positive standard electrode potentials are at the top of the series and are less likely to act as reducing agents in a reaction. Conversely, elements like Zn, which have lower or negative standard electrode potentials, are situated lower in the series and typically act as reducing agents.