Question

For the half-reaction \(\mathrm{Hg}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \longrightarrow \mathrm{Hg}(1)\) \(E^{\circ}=0.854 \mathrm{V} .\) This means that \((\mathrm{a}) \mathrm{Hg}(1)\) is more readily oxidized than \(\mathrm{H}_{2}(\mathrm{g}) ;\) (b) \(\mathrm{Hg}^{2+}(\) aq) is more readily reduced than \(\mathrm{H}^{+}(\mathrm{aq}) ;\) (c) \(\mathrm{Hg}(\) l) will dissolve in 1 M HCl; (d) Hg(l) will displace Zn(s) from an aqueous solution of \(\mathrm{Zn}^{2+}\) ion.

Short answer

(a) True, (b) True, (c) False, (d) True.

Step-by-step solution

Check if Hg(l) is more readily oxidized than H2(g)

The standard cell potential of hydrogen (E0_H2) is defined as 0V so it is less than the standard reduction potential of Hg^2+ which is 0.854V, meaning at standard conditions, Hg(l) is oxidized (loses electrons) more readily than Hg(g). So (a) is True.

Check if Hg^2+(aq) is more readily reduced than H+(aq)

The standard reduction potential of H+ (E0_H+) is 0V which is less than the standard reduction potential of Hg^2+ (E0_Hg) which is 0.854V, meaning at standard conditions, Hg^2+(aq) is reduced more readily than H+(aq). So (b) is True.

Check if Hg(l) will dissolve in 1 M HCl

Hg(l) won't dissolve in 1M HCl because Hg(l) is not easily oxidized by H+ ions whose (E0_H+) is 0V which less than the standard reduction potential of Hg^2+ (E0_Hg) which is 0.854V. So (c) is False.

Check if Hg(l) will displace Zn(s) from an aqueous solution of Zn^2+ ion

The standard reduction potential for Zn^2+/Zn (E0_Zn) is -0.76V which is less than the standard reduction potentials of Hg^2+/Hg (i.e., 0.854V). Thus, Hg(l) can displace Zn(s) from a solution of Zn^2+ ions since it has a higher reduction potential. So (d) is True.

Key concepts

Standard Electrode Potential

Standard electrode potential, denoted as \( E^{ ext{0}} \), is a critical value in electrochemistry.
It represents the voltage or electrical potential difference that can be generated by a half-cell under standard conditions when it is compared to the standard hydrogen electrode.
This potential is a measure of the tendency of a chemical species to be reduced, with higher values indicating a greater likelihood of reduction.

• The standard electrode potential is measured in volts (V).
• Standard conditions imply 1 M concentration for solutions, 1 atm pressure for gases, and pure solids or liquids.
• The standard hydrogen electrode (SHE) is the reference with an \( E^{ ext{0}} \) of 0 V.

In our exercise, the half-reaction for mercury has a standard electrode potential \( (E^{ ext{0}}) \) of 0.854 V. Comparing this to the hydrogen electrode (0 V), it can be observed that \( ext{Hg}^{2+} \) ions are more easily reduced than \( ext{H}^{+} \) ions. Thus, the greater the \( E^{ ext{0}} \) value, the higher the tendency of the reduction to occur.

Redox Reactions

Redox reactions are short for reduction-oxidation reactions.
These reactions involve the transfer of electrons between chemical species.
One species will gain electrons (reduction) and another will lose electrons (oxidation).
Redox reactions are essential in many natural and industrial processes, such as:

• Combustion
• Respiration
• Photosynthesis
• Batteries and electrochemical cells

In the exercise, the redox reaction occurs between mercury and hydrogen in aqueous solutions.
Hg(l) has a higher tendency to be oxidized than hydrogen gas based on the standard electrode potentials, leading to the conclusion that when comparing these potentials, we determine the direction of electron flow.

Oxidation-Reduction

Understanding oxidation-reduction is crucial to unravel the nature of redox reactions.
In these reactions:

• Oxidation refers to the loss of electrons by a molecule, atom, or ion.
• Reduction is the gain of electrons by a molecule, atom, or ion.

A useful mnemonic to remember this process is "OIL RIG": Oxidation Is Loss, Reduction Is Gain.
By examining the standard electrode potentials, it is easier to predict which species will undergo oxidation and which will undergo reduction.
In the textbook exercise, mercury \( ext{Hg}(l) \) is compared to hydrogen and other species. Its higher electrode potential enables it to be reduced more readily than \( ext{H}^{+} \), but less than \( ext{Zn(s)} \).
As electrons are transferred, energy changes occur that can be harnessed in various applications like electricity generation.