Question

Consider only the species (at standard conditions) $$\mathrm{Na}^{+}, \mathrm{Cl}^{-}, \mathrm{Ag}^{+}, \mathrm{Ag}, \mathrm{Zn}^{2+}, \mathrm{Zn}, \mathrm{Pb}$$ in answering the following questions. Give reasons for your answers. (Use data from Table 18.1.) a. Which is the strongest oxidizing agent? b. Which is the strongest reducing agent? c. Which species can be oxidized by \(\mathrm{SO}_{4}^{2-}(a q)\) in acid? d. Which species can be reduced by \(\mathrm{Al}(s)\) ?

Short answer

a. Cl2 is the strongest oxidizing agent with a standard reduction potential of +1.36 V. b. Na is the strongest reducing agent with a standard reduction potential of -2.71 V. c. Zn and Na can be oxidized by SO42- in an acidic solution. d. Ag+ and Pb2+ can be reduced by solid Al.

Step-by-step solution

Reviewing standard reduction potentials

Remember that the standard reduction potential is an intensive property that measures how easily a chemical species can be reduced to a lower oxidation state. A more positive standard reduction potential means that the species has a stronger tendency to be reduced (gain electrons) or, in other words, a stronger oxidizing agent. On the other hand, a more negative standard reduction potential means that the species has a stronger tendency to be oxidized (lose electrons) or is a stronger reducing agent.

Identifying the strongest oxidizing agent

To identify the strongest oxidizing agent, we will compare the given species' standard reduction potentials. The species with the most positive standard reduction potential will be the strongest oxidizing agent. After referring to Table 18.1, we have the following reduction potentials for the given species: Ag+ + 1e- -> Ag E° = +0.80 V Pb2+ + 2e- -> Pb E° = -0.13 V Zn2+ + 2e- -> Zn E° = -0.76 V Na+ + 1e- -> Na E° = -2.71 V Cl2 + 2e- -> 2Cl- E° = +1.36 V Comparing the values, Cl2 is the strongest oxidizing agent because it has the most positive standard reduction potential of +1.36 V.

Identifying the strongest reducing agent

To identify the strongest reducing agent, we will look for the species with the most negative standard reduction potential. In this case, the strongest reducing agent is Na with a standard reduction potential of -2.71 V.

Identifying species oxidized by SO42- in an acidic solution

We are given that SO42- is the oxidizing agent in the reaction. Since we are considering reactions in an acidic solution, we will look at the reduction of SO42- to H2S. The overall reaction is: SO42- + 10 H+ + 8 e- -> H2S + 4 H2O E° = +0.19 V To be oxidized by SO42-, the other species must have a more negative reduction potential. Comparing the values in Step 2, the species that can be oxidized by SO42- in an acidic solution are Zn and Na, as their reduction potentials are more negative than +0.19 V.

Identifying species that can be reduced by Al

Lastly, we need to identify species that can be reduced by solid Al. The reduction potential for Al is: Al3+ + 3e- -> Al E° = -1.66 V For a species to be reduced by Al, it should have a more positive reduction potential than Al. Comparing the values from Step 2, the species that can be reduced by Al are Ag+ and Pb2+ because both have more positive standard reduction potentials than Al. To sum up: a. The strongest oxidizing agent is Cl2. b. The strongest reducing agent is Na. c. Zn and Na can be oxidized by SO42- in an acidic solution. d. Ag+ and Pb2+ can be reduced by solid Al.

Key concepts

Oxidizing and Reducing Agents

Understanding how oxidizing and reducing agents work is crucial in the field of electrochemistry. These agents are substances that gain or lose electrons in chemical reactions, respectively.

An oxidizing agent, or oxidant, gains electrons and is thereby reduced in a redox reaction. It 'pulls' electrons away from another substance, causing the other substance to be oxidized. A strong oxidizing agent has a higher tendency to gain electrons and, therefore, a more positive standard reduction potential.

Conversely, a reducing agent, or reductant, loses electrons and is oxidized, allowing another compound to gain those electrons and be reduced. Reducing agents have more negative standard reduction potentials, indicating a preference to lose electrons and become oxidized. This characteristic allows them to donate electrons easily.

Identifying the Strength of Oxidizing and Reducing Agents

When comparing standard reduction potentials from tables such as Table 18.1, remember the general rule: the higher the potential, the stronger the oxidizing agent, and the lower the potential, the stronger the reducing agent. In our exercise, this comparison revealed that Cl₂ is the strongest oxidizing agent and Na is the strongest reducing agent among the given species.

Electrochemistry

Electrochemistry is a branch of chemistry that deals with the interrelation of electrical currents and chemical reactions. At the center of electrochemistry are redox reactions, which involve the transfer of electrons between substances.

Through these electron transfers, species undergo oxidation (losing electrons) and reduction (gaining electrons), often with associated energy changes. These changes are harnessed in a variety of applications, such as batteries, corrosion processes, and electroplating.

Standard Reduction Potentials in Electrochemistry

The standard reduction potential is a fundamental concept in electrochemistry, providing a measure of a substance’s tendency to be reduced. The more positive the reduction potential, the more likely it is to gain electrons and act as an oxidizing agent. Similarly, a negative reduction potential indicates a substance that readily loses electrons, functioning as a reducing agent. These potentials allow chemists to predict the direction of electron flow in redox reactions, crucial for designing cells and understanding reaction spontaneity.

Redox Reactions

Redox reactions, short for reduction-oxidation reactions, occur when there is a transfer of electrons from one reactant to another. These reactions are fundamental processes in chemistry and are characterized by changes in the oxidation states of the atoms involved.

In a redox reaction, one species is oxidized, meaning it loses electrons and its oxidation state increases, while another is reduced, meaning it gains electrons and its oxidation state decreases. The balance of electron transfer must always be maintained—no electrons are lost or gained overall in the reaction.

Redox Reactions in Real-Life Applications

Redox reactions are not just theoretical constructs but are omnipresent in many real-world applications, from the metabolism of food in biological systems to industrial processes for material synthesis. By understanding redox mechanisms, scientists can manipulate reactions for desirable outcomes, such as energy production in fuel cells or the prevention of metal corrosion in infrastructure.