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Chapter 17: Problem 38

Why is a systematic method for balancing oxidation-reduction reactions necessary? Why can't these equations be balanced readily by inspection?

Short Answer

Expert verified
A systematic method for balancing oxidation-reduction reactions is necessary because balancing these equations by inspection can lead to incorrect coefficients and unbalanced charges. Redox reactions involve complex changes in oxidation states and electron transfers, which can be challenging to account for without a systematic approach. The systematic method involves breaking down the reaction into half-reactions, balancing the atoms and charges, ensuring equal loss and gain of electrons, and finally adding the half-reactions together. This process ensures accurate and consistent balancing of redox reactions, unlike balancing by inspection.

Step by step solution

01

1. Break down the given redox reaction into half-reactions

To perform a systematic balancing of a redox reaction, start by breaking down the given reaction into two half-reactions - the oxidation half-reaction and the reduction half-reaction. This will allow us to focus on the individual processes without being overwhelmed by the full equation. Make note of the species that are being oxidized and reduced, and the number of electrons involved.
02

2. Balance the atoms

In each half-reaction, balance all the atoms apart from oxygen and hydrogen. This step ensures that the number of atoms of each element is the same on both sides of the half-reactions.
03

3. Balance the oxygen and hydrogen atoms

For each half-reaction, balance the oxygen atoms by adding H2O molecules to the side that lacks oxygen. Then, balance the hydrogen atoms by adding H+ ions to the side lacking hydrogen. This step helps in maintaining the balance of oxygen and hydrogen atoms in the equation.
04

4. Balance the charges

Since redox reactions involve transfer of electrons, it is crucial to balance the charges on both sides of each half-reaction. You can do this by adding electrons (e-) to the side of the half-reaction with a more positive charge. Make sure that the charges on both sides of each half-reaction are equal.
05

5. Ensure equal loss and gain of electrons

Since the electrons lost during oxidation must be equal to the electrons gained during reduction, multiply each half-reaction by an appropriate factor so that the number of electrons lost in the oxidation half-reaction matches the number gained in the reduction half-reaction. This step ensures conservation of charge in the overall redox reaction.
06

6. Add the half-reactions together

Once the number of electrons lost and gained are equal in both half-reactions, add the half-reactions together to obtain the balanced redox reaction. Cancel out any species that appear on both sides of the reaction. By following these systematic steps, we avoid confusion and incorrect balancing that may occur when trying to balance a redox reaction by inspection. Balancing by inspection may not account for the changes in oxidation states and the number of electrons involved in the reaction, leading to incorrect coefficients and an unbalanced equation. Therefore, a systematic method for balancing redox reactions is necessary for maintaining accuracy and consistency.

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Most popular questions from this chapter

Write the chemical equation for the overall cell reaction that occurs in a lead storage automobile battery. What species is oxidized in such a battery? What species is reduced? Why can such a battery be "recharged"?

In each of the following reactions, identify which element is oxidized and which is reduced. a. \(2 \mathrm{Al}(s)+6 \mathrm{HCl}(a q) \rightarrow 2 \mathrm{AlCl}_{3}(a q)+3 \mathrm{H}_{2}(g)\) b. \(2 \mathrm{HI}(g) \rightarrow \mathrm{H}_{2}(g)+\mathrm{I}_{2}(s)\) c. \(\mathrm{Cu}(s)+\mathrm{H}_{2} \mathrm{SO}_{4}(a q) \rightarrow \mathrm{CuSO}_{4}(a q)+\mathrm{H}_{2}(g)\)

Assign oxidation states to all of the atoms in each of the following: a. \(\mathrm{H}_{2} \mathrm{SO}_{4}\) b. \(\mathrm{MnO}_{4}^{-}\) c. \(\mathrm{NO}_{3}^{-}\) d. \(\mathrm{K}_{3} \mathrm{PO}_{4}\)

What process is represented by the corrosion of a metal? Why is corrosion undesirable?

"Jump-starting" a dead automobile battery can be dangerous if precautions are not taken, because of the production of an explosive mixture of _________ and _________ gases in the battery.
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