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Chapter 8: Problem 82

Why is the oxidizing agent in a redox reaction reduced and the reducing agent oxidized?

Short Answer

Expert verified
Answer: In a redox reaction, the oxidizing agent gets reduced because it accepts electrons while causing the oxidation of another substance, and the reducing agent gets oxidized because it donates electrons while causing the reduction of another substance. This electron transfer between substances leads to a change in the oxidation state of the reactants involved in the redox process.

Step by step solution

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1. Defining Redox Reaction

A redox reaction is a chemical reaction in which the oxidation state of one or more of the reactants changes. This usually involves a transfer of electrons between reactants. Redox reactions consist of two half-reactions: one for the oxidation process and another for the reduction process.
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2. Defining Oxidation and Reduction

Oxidation is the process of losing electrons, whereas reduction is the process of gaining electrons. In a redox reaction, one reactant will undergo oxidation and another reactant will undergo reduction.
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3. The Roles of Oxidizing and Reducing Agents

An oxidizing agent is a substance that causes the oxidation of another substance by accepting electrons. A reducing agent is a substance that causes the reduction of another substance by donating electrons.
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4. Oxidizing Agent Gets Reduced

Since the oxidizing agent is responsible for oxidizing another substance, it must accept electrons in the process. By accepting electrons, the oxidizing agent itself gets reduced. This is because it gains electrons, which is the definition of reduction.
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5. Reducing Agent Gets Oxidized

On the other hand, the reducing agent is responsible for reducing another substance, and it does this by donating electrons. By donating electrons, the reducing agent itself gets oxidized. This is because it loses electrons, which is the definition of oxidation.
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6. Conclusion

In a redox reaction, the oxidizing agent gets reduced because it accepts electrons, and the reducing agent gets oxidized because it donates electrons. This is at the core of the redox process, involving the transfer of electrons between substances, ultimately leading to a change in the oxidation state of the reactants involved.

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Key Concepts

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

Oxidation
Oxidation is a fundamental concept in redox reactions, where a substance loses electrons. This is an essential part of how redox reactions work. When a substance undergoes oxidation, its oxidation state increases because there are fewer electrons around to balance out positively charged protons in its nucleus. Think of oxidation as losing parts: if you lose electrons, you become more positive because electrons carry a negative charge. It's like losing something you own; your overall value decreases. Every oxidation event is coupled with a reduction event, because the lost electrons have to be gained by another species.
Reduction
Reduction is the opposite of oxidation in redox reactions. It involves gaining electrons. When a substance is reduced, its oxidation state decreases. This is because adding electrons lessens its positive charge, effectively neutralizing some of the positive charges from the protons in the nucleus. To visualize reduction, imagine adding value to something you have. Gaining electrons means you become more negative, as electrons are negatively charged. Reduction always occurs alongside oxidation. Electrons lost from one species are gained by another, making redox a coupled process.
Oxidizing Agent
An oxidizing agent is crucial in redox reactions because it enables oxidation to occur. It does this by accepting electrons from another substance. In the process, the oxidizing agent becomes reduced, as taking in electrons means its oxidation state decreases. A simple way to think of an oxidizing agent is like a friend who takes responsibility off your shoulders. By taking electrons, it facilitates the oxidation of another substance. Some common examples of oxidizing agents include oxygen, hydrogen peroxide, and halogens like chlorine.
Reducing Agent
The reducing agent plays an indispensable role in redox reactions by donating electrons to another substance. By giving away electrons, the reducing agent helps the other substance to be reduced while itself undergoing oxidation. You can think of a reducing agent as someone giving a gift; by donating electrons, it helps reduce the oxidation state of another species. But in doing so, it itself becomes oxidized. Common reducing agents include hydrogen, carbon monoxide, and metals like zinc and copper.

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Most popular questions from this chapter

Chlorine was first prepared in 1774 by heating a mixture of \(\mathrm{NaCl}\) and \(\mathrm{MnO}_{2}\) in sulfuric acid: $$\begin{aligned} \mathrm{NaCl}(a q)+\mathrm{H}_{2} \mathrm{SO}_{4}(a q)+\mathrm{MnO}_{2}(s) & \rightarrow \\ & \mathrm{Na}_{2} \mathrm{SO}_{4}(a q)+\mathrm{MnCl}_{2}(a q)+\mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{Cl}_{2}(g) \end{aligned}$$a. Assign oxidation numbers to the elements in each compound, and balance the redox reaction in acid solution. b. Write a net ionic equation describing the reaction for the formation of chlorine. c. If chlorine gas is inhaled, it causes pulmonary edema (fluid in the lungs) because it reacts with water in the alveolar sacs of the lungs to produce the strong acid \(\mathrm{HCl}\) and the weaker acid HOC1. Balance the equation for the conversion of \(\mathrm{Cl}_{2}\) to \(\mathrm{HCl}\) and \(\mathrm{HOCl}\).

What is the difference between a strong acid and a weak acid?

Sodium fluoride is added to drinking water in many municipalities to protect teeth against cavities. The target of the fluoridation is hydroxyapatite, \(\mathrm{Ca}_{10}\left(\mathrm{PO}_{4}\right)_{6}(\mathrm{OH})_{2},\) a compound in tooth enamel. There is concern, however, that fluoride ions in water may contribute to skeletal fluorosis, an arthritis-like disease. a. Write a net ionic equation for the reaction between hydroxyapatite and sodium fluoride that produces fluorapatite, \(\mathrm{Ca}_{10}\left(\mathrm{PO}_{4}\right)_{6} \mathrm{F}_{2}\) b. The U.S. EPA currently restricts the concentration of \(\mathrm{F}^{-}\) in drinking water to \(4 \mathrm{mg} / \mathrm{L}\). Express this concentration of \(F^{-}\) in molarity. c. One study of skeletal fluorosis suggests that drinking water with a fluoride concentration of \(4 \mathrm{mg} / \mathrm{L}\) for 20 years raises the fluoride content in bone to \(6 \mathrm{mg} / \mathrm{g}\), a level at which a patient may experience stiff joints and other symptoms. How much fluoride (in milligrams) is present in a 100 mg sample of bone with this fluoride concentration?

Give the formulas of two strong bases and two weak bases.

Why is \(\mathrm{HSO}_{4}^{-}(a q)\) a weaker acid than \(\mathrm{H}_{2} \mathrm{SO}_{4} ?\)
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