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Chapter 10: Problem 50

When is an oxidizing agent also the species that gets oxidized? Explain.

Short Answer

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An oxidizing agent can also be the species that gets oxidized when it reacts with a substance that has the same or greater ability to accept or gain electrons, such as in disproportionation reactions. In these cases, both substances will undergo oxidation, making the oxidizing agent the species that gets oxidized as well. An example of this is the decomposition of hydrogen peroxide (H2O2) in an acidic solution, where it acts as both the oxidizing and reducing agents.

Step by step solution

01

Understand Oxidation and Reduction

In a redox (reduction-oxidation) reaction, one species loses electrons, and another species gains electrons. The species that loses electrons is said to be oxidized, while the species that gains electrons is said to be reduced. For example, in the reaction between hydrogen gas (H2) and oxygen gas (O2) to form water (H2O): \(2H_2 + O_2 \rightarrow 2H_2O\) Hydrogen loses electrons and is oxidized: \(H_2 \rightarrow 2H^+ + 2e^-\) Oxygen gains electrons and is reduced: \(O_2 + 4e^- \rightarrow 2O^{2-}\)
02

Identifying Oxidizing and Reducing Agents

The species that causes oxidation of another species is called the oxidizing agent, while the species that causes reduction of another species is called the reducing agent. In the example above, oxygen (O2) is the oxidizing agent because it causes hydrogen (H2) to lose electrons. Hydrogen (H2) is the reducing agent because it causes oxygen (O2) to gain electrons.
03

Oxidizing Agent as the Species that Gets Oxidized

For an oxidizing agent to also be the species that gets oxidized, it must participate in a reaction with another species that has the same or greater ability to accept or gain electrons. In other words, the oxidizing agent needs to interact with another substance that is equally or more effective at oxidizing. In such scenarios, both substances will be undergoing oxidation, and therefore, the oxidizing agent is also the species that gets oxidized. An example of this situation can be seen in the reaction between two different metal ions where both are competing to take electrons from one another. Such reactions are also known as "disproportionation reactions." Disproportionation reaction example: The decomposition of hydrogen peroxide (H2O2) in an acidic solution can be shown as: \(2H_2O_2 \rightarrow 2H_2O + O_2\) In this case, hydrogen peroxide is both the oxidizing and reducing agents as it causes the reduction of one oxygen atom (O) while causing the oxidation of the other. In conclusion, an oxidizing agent can also be the species that gets oxidized when it reacts with a substance that has the same or greater ability to accept or gain electrons. In these cases, both the substances will undergo oxidation, making the oxidizing agent the species that gets oxidized as well.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

oxidizing agent
An oxidizing agent is a substance that facilitates oxidation in a redox reaction by accepting electrons from another substance. By doing so, it effectively causes the other material to lose electrons and become oxidized. The oxidizing agent, therefore, is reduced as it gains these electrons. For example, in the reaction between hydrogen gas and oxygen gas:
  • Oxygen acts as the oxidizing agent because it gains electrons from hydrogen.
Oxygen is thus reduced while causing hydrogen to oxidize. It's important to note that while oxidizing agents are typically reduced, they have the potential to themselves be oxidized in unique situations such as disproportionation reactions.
oxidation
Oxidation is a chemical process involving the loss of electrons by a molecule, atom, or ion. Whenever a species undergoes oxidation, it typically becomes more positive due to the lost negative electrons. This process can be illustrated through:
  • Chemical reactions where metals are converted to their ionic forms, like when iron (Fe) rusts to form iron oxide (Fe2O3).
  • Biological processes such as cellular respiration, where glucose is oxidized to release energy.
It's essential to understand that oxidation cannot occur alone; it always accompanies reduction, where another substance gains electrons.
reduction
Reduction is the gain of electrons by a molecule, atom, or ion, which results in a decrease in its oxidation state. This electron gain makes the species more negative. As a partner to oxidation in redox reactions, reduction ensures balance.
  • For instance, when oxygen gas gains electrons to form oxide ions, it is reduced.
  • In photosynthesis, carbon dioxide is reduced to glucose.
Reduction is crucial in many life processes and industrial applications, where it is often paired with oxidation reactions.
disproportionation reactions
Disproportionation reactions are special types of redox reactions where a single substance undergoes both oxidation and reduction simultaneously, being transformed into two different products. This can seem paradoxical, as one might assume a substance could only either gain or lose electrons, not both.
  • An example of such a reaction is the decomposition of hydrogen peroxide (\(H_2O_2\)) into water (\(H_2O\)) and oxygen (\(O_2\)).
  • In this reaction, part of the \(H_2O_2\) molecule is reduced to water, while another part is oxidized to oxygen gas.
Understanding disproportionation reactions helps in recognizing how complex and versatile chemical reactions can be, and how a substance might play dual roles like acting as both an oxidizing and reducing agent.

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Most popular questions from this chapter

Hydrogen peroxide \(\left(\mathrm{H}_{2} \mathrm{O}_{2}\right)\) can be made to decompose into water and oxygen as shown below. (The hydrogen peroxide in the beaker is decomposing to give a cloud of water vapor. Oxygen gas is also being produced.) Is this a redox reaction? If so, prove it.

\(\mathrm{Ca}+2 \mathrm{H}^{+} \rightarrow \mathrm{Ca}^{2+}+\mathrm{H}_{2}\)

There are two definitions of oxidation, one involving oxygen, the other not. (a) State the two definitions of oxidation. (b) Are they compatible? Explain. (c) When we speak of corrosion of metals, what chemical reaction are we usually talking about? Use Fe as an example.

What happens when you place an active metal in a solution of ions of a less active metal?

Assign an oxidation state to each atom in HCN. (Hint: Begin with a dot diagram.)
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