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Chapter 17: Problem 16

Balance these redox reactions, and identify the oxidizing agent and the reducing agent. (a) \(\mathrm{CO}(\mathrm{g})+\mathrm{O}_{3}(\mathrm{~g}) \longrightarrow \mathrm{CO}_{2}(\mathrm{~g})\) (b) \(\mathrm{H}_{2}(\mathrm{~g})+\mathrm{Cl}_{2}(\mathrm{~g}) \longrightarrow \mathrm{HCl}(\mathrm{g})\) (c) \(\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq})+\mathrm{Ti}^{2+}(\mathrm{aq}) \longrightarrow \mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{TiO}_{2}(\mathrm{~s})\) in acidic solution (d) \(\mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{MnO}_{4}^{-}(\mathrm{aq}) \longrightarrow \mathrm{Cl}_{2}(\mathrm{~g})+\mathrm{MnO}_{2}(\mathrm{~s})\) in acidic solution (e) \(\mathrm{FeS}_{2}(\mathrm{~s})+\mathrm{O}_{2}(\mathrm{~g}) \longrightarrow \mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{~s})+\mathrm{SO}_{2}(\mathrm{~g})\) (f) \(\mathrm{O}_{3}(\mathrm{~g})+\mathrm{NO}(\mathrm{g}) \longrightarrow \mathrm{O}_{2}(\mathrm{~g})+\mathrm{NO}_{2}(\mathrm{~g})\) (g) \(\mathrm{Zn}(\mathrm{s})+\mathrm{HgO}(\mathrm{s}) \longrightarrow \mathrm{Zn}(\mathrm{OH})_{2}(\mathrm{~s})+\operatorname{Hg}(\ell)\) in basic solution

Short Answer

Expert verified
Reactions balanced. Identified oxidizing and reducing agents for all reactions.

Step by step solution

01

Write the Unbalanced Reaction for (a)

The given reaction is \( \mathrm{CO}(\mathrm{g}) + \mathrm{O}_{3}(\mathrm{g}) \rightarrow \mathrm{CO}_{2}(\mathrm{g}) \).
02

Identify Oxidation and Reduction for (a)

- Carbon in \( \mathrm{CO} \) is oxidized to \( \mathrm{CO}_{2} \).- Oxygen in \( \mathrm{O}_{3} \) is reduced to \( \mathrm{O}_{2} \).- \( \mathrm{O}_{3} \) is the oxidizing agent and \( \mathrm{CO} \) is the reducing agent.
03

Balance Atoms and Charge for (a)

The balanced reaction is already presented in the simplest form:\[ \mathrm{CO}(\mathrm{g}) + \mathrm{O}_{3}(\mathrm{g}) \rightarrow \mathrm{CO}_{2}(\mathrm{g}) + \mathrm{O}_{2}(\mathrm{g}) \]
04

Write the Unbalanced Reaction for (b)

The given reaction is \( \mathrm{H}_{2}(\mathrm{g}) + \mathrm{Cl}_{2}(\mathrm{g}) \rightarrow \mathrm{HCl}(\mathrm{g}) \).
05

Identify Oxidation and Reduction for (b)

- Hydrogen is oxidized.- Chlorine is reduced.- \( \mathrm{Cl}_{2} \) is the oxidizing agent, and \( \mathrm{H}_{2} \) is the reducing agent.
06

Balance Atoms and Charge for (b)

The balanced reaction is:\[ \mathrm{H}_{2}(\mathrm{g}) + \mathrm{Cl}_{2}(\mathrm{g}) \rightarrow 2\mathrm{HCl}(\mathrm{g}) \]
07

Write the Unbalanced Reaction for (c)

The given reaction is \( \mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq}) + \mathrm{Ti}^{2+}(\mathrm{aq}) \rightarrow \mathrm{H}_{2} \mathrm{O}(\ell) + \mathrm{TiO}_{2}(\mathrm{s}) \).
08

Identify Oxidation and Reduction for (c)

- \( \mathrm{Ti}^{2+} \) is oxidized to \( \mathrm{TiO}_{2} \).- \( \mathrm{H}_{2} \mathrm{O}_{2} \) is reduced to \( \mathrm{H}_{2} \mathrm{O} \).- \( \mathrm{H}_{2} \mathrm{O}_{2} \) is the oxidizing agent, and \( \mathrm{Ti}^{2+} \) is the reducing agent.
09

Balance Atoms and Charge for (c)

Assume acidic conditions:1. Balance oxygen using water: \[ \mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq}) + \mathrm{Ti}^{2+}(\mathrm{aq}) \rightarrow \mathrm{TiO}_{2}(\mathrm{s}) + 2\mathrm{H}_{2} \mathrm{O}(\ell) \]2. Balance hydrogen using \( \mathrm{H}^{+} \): \[ 2\mathrm{H}^{+} + \mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq}) + \mathrm{Ti}^{2+}(\mathrm{aq}) \rightarrow \mathrm{TiO}_{2}(\mathrm{s}) + 2\mathrm{H}_{2} \mathrm{O}(\ell) \]3. Balance charge with electrons: \[ \mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq}) + \mathrm{Ti}^{2+}(\mathrm{aq}) \rightarrow \mathrm{TiO}_{2}(\mathrm{s}) + 2\mathrm{H}_{2} \mathrm{O}(\ell) + 2\mathrm{e}^{-} \]
10

Write the Unbalanced Reaction for (d)

The reaction is \( \mathrm{Cl}^{-}(\mathrm{aq}) + \mathrm{MnO}_{4}^{-}(\mathrm{aq}) \rightarrow \mathrm{Cl}_{2}(\mathrm{g}) + \mathrm{MnO}_{2}(\mathrm{s}) \).
11

Identify Oxidation and Reduction for (d)

- \( \mathrm{Cl}^{-} \) is oxidized to \( \mathrm{Cl}_{2} \).- \( \mathrm{MnO}_{4}^{-} \) is reduced to \( \mathrm{MnO}_{2} \).- \( \mathrm{MnO}_{4}^{-} \) is the oxidizing agent, and \( \mathrm{Cl}^{-} \) is the reducing agent.
12

Balance Atoms and Charge for (d)

In acidic conditions:1. Balance \( \mathrm{Cl} \): \[ 2\mathrm{Cl}^{-}(\mathrm{aq}) + \mathrm{MnO}_{4}^{-}(\mathrm{aq}) \rightarrow \mathrm{Cl}_{2}(\mathrm{g}) + \mathrm{MnO}_{2}(\mathrm{s}) \]2. Add \( \mathrm{H}^{+} \) and \( \mathrm{H}_{2}\mathrm{O} \): \[ 4\mathrm{H}^{+} + 2\mathrm{Cl}^{-}(\mathrm{aq}) + \mathrm{MnO}_{4}^{-} \rightarrow \mathrm{Cl}_{2} + \mathrm{MnO}_{2} + 2\mathrm{H}_{2}\mathrm{O} \]3. Balance charge: \[ 2\mathrm{Cl}^{-} + \mathrm{MnO}_{4}^{-} + 4\mathrm{H}^{+} + 2\mathrm{e}^{-} \rightarrow \mathrm{Cl}_{2} + \mathrm{MnO}_{2} + 2\mathrm{H}_{2}\mathrm{O} \]
13

Write the Unbalanced Reaction for (e)

The reaction is \( \mathrm{FeS}_{2}(\mathrm{s}) + \mathrm{O}_{2}(\mathrm{g}) \rightarrow \mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{s}) + \mathrm{SO}_{2}(\mathrm{g}) \).
14

Identify Oxidation and Reduction for (e)

- Sulfur in \( \mathrm{FeS}_{2} \) is oxidized to \( \mathrm{SO}_{2} \).- Oxygen is reduced.- \( \mathrm{O}_{2} \) is the oxidizing agent and \( \mathrm{FeS}_{2} \) is the reducing agent.
15

Balance Atoms and Charge for (e)

The balanced reaction is:\[ 4\mathrm{FeS}_{2} + 11\mathrm{O}_{2} \rightarrow 2\mathrm{Fe}_{2} \mathrm{O}_{3} + 8\mathrm{SO}_{2} \]
16

Write the Unbalanced Reaction for (f)

The reaction is \( \mathrm{O}_{3}(\mathrm{g}) + \mathrm{NO}(\mathrm{g}) \rightarrow \mathrm{O}_{2}(\mathrm{g}) + \mathrm{NO}_{2}(\mathrm{g}) \).
17

Identify Oxidation and Reduction for (f)

- \( \mathrm{NO} \) is oxidized to \( \mathrm{NO}_{2} \).- \( \mathrm{O}_{3} \) is reduced to \( \mathrm{O}_{2} \).- \( \mathrm{O}_{3} \) is the oxidizing agent, and \( \mathrm{NO} \) is the reducing agent.
18

Balance Atoms and Charge for (f)

The balanced reaction:\[ \mathrm{O}_{3}(\mathrm{g}) + \mathrm{NO}(\mathrm{g}) \rightarrow \mathrm{O}_{2}(\mathrm{g}) + \mathrm{NO}_{2}(\mathrm{g}) \]
19

Write the Unbalanced Reaction for (g)

The reaction is \( \mathrm{Zn}(\mathrm{s}) + \mathrm{HgO}(\mathrm{s}) \rightarrow \mathrm{Zn(OH)}_{2}(\mathrm{s}) + \mathrm{Hg}(\ell) \).
20

Identify Oxidation and Reduction for (g)

- \( \mathrm{Zn} \) is oxidized.- \( \mathrm{HgO} \) is reduced to \( \mathrm{Hg} \).- \( \mathrm{HgO} \) is the oxidizing agent and \( \mathrm{Zn} \) is the reducing agent.
21

Balance Atoms and Charge for (g)

In basic conditions:1. Balance Zn and Hg: \[ \mathrm{Zn} + \mathrm{HgO} \rightarrow \mathrm{Zn(OH)}_{2} + \mathrm{Hg} \]2. Add \( \mathrm{OH}^{-} \) to balance O and H: \[ \mathrm{Zn} + \mathrm{HgO} + 2\mathrm{OH}^{-} \rightarrow \mathrm{Zn(OH)}_{2} + \mathrm{Hg} \]

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Key Concepts

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

Oxidizing Agent
In a redox reaction, an oxidizing agent is a substance that causes another substance to lose electrons, thereby becoming reduced itself. This means the oxidizing agent gains electrons. In reaction (a) from the exercise, \( \mathrm{O}_{3} \) acts as the oxidizing agent because it causes the oxidation of carbon monoxide (\( \mathrm{CO} \)) by accepting electrons and is reduced to oxygen (\( \mathrm{O}_{2} \)). Suspiciously, the name can imply that the oxidizing agent is itself oxidized, but it's important to remember that the oxidizing agent actually gets reduced.
To identify the oxidizing agent in a chemical reaction, you can:
  • Look for the species that gains electrons.
  • Identify which element's oxidation state decreases.
  • Check the reactant species that accepts electrons in the reaction.
Understanding the role of an oxidizing agent helps clarify the electron transfer process crucial in redox reactions.
Reducing Agent
The reducing agent in a redox reaction is the substance that donates electrons to another substance, leading to its own oxidation. In simpler terms, it loses electrons while causing the other substance to gain electrons and become reduced. For example, in reaction (b) we see that \( \mathrm{H}_{2} \) serves as the reducing agent.
The hydrogen (\( \mathrm{H}_{2} \)) donates electrons to chlorine (\( \mathrm{Cl}_2 \)), forming hydrochloric acid (\( \mathrm{HCl} \)), while oxidizing itself. This identification helps understand which species shifts its electrons.
  • Identifying the reducing agent involves finding the species that loses electrons.
  • Check for the element whose oxidation state increases.
  • Look for reactants that provide electrons in the reaction.
Remember, the ability to recognize a reducing agent is key in predicting electron flow and the direction of redox reactions.
Balancing Chemical Equations
Balancing chemical equations involves ensuring that the number of atoms for each element is equal on both sides of the reaction. This process obeys the law of conservation of mass. For example, in reaction (c), the reaction is comprehensively balanced using steps:1\. Balance oxygen first by adding water molecules.2\. Then balance hydrogen using hydrogen ions (\( \mathrm{H}^+ \)).3\. Finally, balance the charge by adding electrons.
During the process, it鈥檚 crucial to:
  • Understand the oxidation state changes to recognize the oxidized and reduced species.
  • Use \( \mathrm{H}^+ \) and water in acidic solutions or hydroxide (\( \mathrm{OH}^- \)) in basic solutions to balance hydrogen and oxygen.
  • Finish by ensuring all electrons are balanced, contributing to charges to either side of the equation.
These steps give a clear, linear approach to obtaining a balanced equation, essential for accurate stoichiometric calculations and analysis.

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Most popular questions from this chapter

An electrolytic cell is set up with \(\mathrm{Cd}(\mathrm{s})\) in \(\mathrm{Cd}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})\) and \(\mathrm{Zn}(\mathrm{s})\) in \(\mathrm{Zn}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq}) .\) Initially both electrodes weigh \(5.00 \mathrm{~g}\). After running the cell for several hours the electrode in the left compartment weighs \(4.75 \mathrm{~g}\). (a) Which electrode is in the left compartment? (b) Does the mass of the electrode in the right compartment increase, decrease, or stay the same? If the mass changes, what is the new mass? (c) Does the volume of the electrode in the right compartment increase, decrease, or stay the same? If the volume changes, what is the new volume? (The density of \(\mathrm{Cd}\) is \(\left.8.65 \mathrm{~g} / \mathrm{cm}^{3} .\right)\)

In principle, a battery could be made from aluminum metal and chlorine gas. (a) Write a balanced equation for the reaction that would occur in a battery using \(\mathrm{Al}^{3+}(\mathrm{aq}) \mid \mathrm{Al}(\mathrm{s})\) and \(\mathrm{Cl}_{2}(\mathrm{~g}) \mid \mathrm{Cl}^{-}(\) aq \()\) half-cells. (b) Identify the half-reaction at the anode and at the cathode. Do electrons flow from the \(\mathrm{Al}\) electrode when the cell does work? Explain. (c) Calculate the standard potential, \(E_{\text {cell }}^{\circ}\), for the battery.

Identify each statement as true or false. Rewrite each false statement to make it true. (a) Oxidation always occurs at the anode of an electrochemical cell. (b) The anode of a discharging voltaic cell is the site of reduction and is negative. (c) Standard-state conditions for electrochemical cells are a concentration of \(1.0 \mathrm{M}\) for dissolved species and a pressure of 1 bar for gases. (d) The potential of a voltaic cell does not change with temperature. (e) All product-favored oxidation-reduction reactions have a standard cell potential \(E_{\text {cell }}^{\circ}\), with a negative sign.

Consider two different electrolytic cells; one cell contains aqueous \(\mathrm{Zn}^{2+}\) and the other contains \(\mathrm{Cr}^{3+}\). The initial metal ion concentration is the same in each cell and the metal ions are reduced to the metal during the electrolysis. Each cell operates at the same current. Without doing calculations, predict which cell has the greater mass of metal deposited after 5 min. Explain your prediction.

Fluorine, \(\mathrm{F}_{2}\), is made by the electrolysis of anhydrous \(\mathrm{HF}\). $$ 2 \mathrm{HF}(\ell) \longrightarrow \mathrm{H}_{2}(\mathrm{~g})+\mathrm{F}_{2}(\mathrm{~g}) $$ Typical electrolysis cells operate at 4000 to \(6000 \mathrm{~A}\) and 8 to \(12 \mathrm{~V}\). A large-scale plant can produce about \(9.0 \mathrm{met}-\) ric tons of \(\mathrm{F}_{2}\) gas per day. (a) Calculate the mass (g) of HF consumed. (b) Using the conversion factor of \(3.60 \times 10^{6} \mathrm{~J} / \mathrm{kWh}\), calculate how much energy in kilowatt-hours is transferred to a cell operating at \(6.0 \times 10^{3} \mathrm{~A}\) at \(12 \mathrm{~V}\) for \(24 \mathrm{~h}\).
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