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Chapter 20: Problem 16

Indicate whether each of the following statements is true or false: (a) If something is reduced, it is formally losing electrons. (b) A reducing agent gets oxidized as it reacts. (c) Oxidizing agents can convert \(\mathrm{CO}\) into \(\mathrm{CO}_{2}\).

Short Answer

Expert verified
(a) False: If something is reduced, it is formally gaining electrons, not losing them. (b) True: A reducing agent gets oxidized as it reacts, giving away its own electrons to reduce another substance. (c) True: Oxidizing agents can convert CO into CO鈧�, as seen in the reaction CO + 陆O鈧� 鈫� CO鈧�, where oxygen acts as an oxidizing agent.

Step by step solution

01

Statement (a)

"If something is reduced, it is formally losing electrons." To determine the truth of this statement, let's recall the definition of reduction and oxidation: - Reduction: Gaining electrons - Oxidation: Losing electrons So, if something is reduced, it is formally gaining electrons, not losing them. Therefore, this statement is: \( \boxed{\textbf{False}} \)
02

Statement (b)

"A reducing agent gets oxidized as it reacts." A reducing agent is a substance that causes another substance to be reduced by giving away its own electrons. As the reducing agent gives away its electrons, it gets oxidized. So, the statement is: \( \boxed{\textbf{True}} \)
03

Statement (c)

"Oxidizing agents can convert \(\mathrm{CO}\) into \(\mathrm{CO}_{2}\)." Oxidizing agents are substances that cause another substance to be oxidized by taking electrons from it. In the case of converting \(\mathrm{CO}\) (carbon monoxide) into \(\mathrm{CO}_{2}\) (carbon dioxide), it means that the carbon in \(\mathrm{CO}\) gets oxidized by gaining more oxygen: \(\mathrm{CO} + \frac{1}{2} \,\mathrm{O}_{2} \,\rightarrow \mathrm{CO}_{2}\) In this reaction, \(\mathrm{O}_{2}\) (oxygen) acts as an oxidizing agent, taking electrons from carbon to become \(\mathrm{CO}_{2}\). Thus, the statement is: \( \boxed{\textbf{True}} \)

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Key Concepts

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

Reduction and Oxidation
Understanding the relationship between reduction and oxidation, commonly known as 'redox' reactions, is key to grasping many chemical processes, including those in biochemistry and environmental chemistry.

In simple terms, oxidation is the process where an atom, ion, or molecule loses one or more electrons. Conversely, reduction is the gain of one or more electrons. These two processes always occur together; when one substance is oxidized, another is reduced. This pair of simultaneous reactions is what we refer to as a redox reaction.

For example, in the reaction where iron rusts:
  • Iron (Fe) is oxidized as it loses electrons to form iron oxide (Fe_2O_3).
  • Oxygen, in this case, is reduced as it gains electrons from iron.
Considering the education exercise at hand, when a substance is reduced, it 'gains' electrons, which can be counterintuitive since the term 'reduction' might imply a decrease, but it's actually referring to a decrease in oxidation state - an increase in electrons.
Reducing Agent
A reducing agent, often referred to as a reductant, plays a crucial role in redox reactions. It is the substance that donates electrons to another substance, causing the latter to be reduced. As it gives up electrons, the reducing agent is itself oxidized in the process.

A classic example comes from the reaction between hydrogen and fluorine:
  • Hydrogen (H_2) acts as the reducing agent as it donates electrons to fluorine, forming hydrofluoric acid (HF).
In the context of our textbook exercise, when a reducing agent gets oxidized, it is essentially losing electrons; this is why the statement 'A reducing agent gets oxidized as it reacts' is true. This is an integral concept for students to understand, as it defines the fundamental characteristic of a reducing agent in redox chemistry.
Oxidizing Agent
An oxidizing agent, or oxidant, is the chemical player that accepts electrons during a redox reaction, causing another substance to lose electrons, thus being oxidized.

One of the most common oxidizing agents is oxygen, which is involved in combustion and corrosion reactions. For instance, in the reaction mentioned in the problem where CO is converted to CO_2, oxygen functions as the oxidizing agent:
  • The carbon monoxide (CO) is oxidized as it gains an oxygen atom and transforms into carbon dioxide (CO_2).
This illustrates that the statement 'Oxidizing agents can convert CO into CO_2' stands correct, and highlights the role of oxidizing agents in increasing the oxidation state of other substances by taking electrons away.

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Most popular questions from this chapter

Elemental calcium is produced by the electrolysis of molten \(\mathrm{CaCl}_{2}\). (a) What mass of calcium can be produced by this process if a current of \(7.5 \times 10^{3} \mathrm{~A}\) is applied for \(48 \mathrm{~h}\) ? Assume that the electrolytic cell is \(68 \%\) efficient. (b) What is the minimum voltage needed to cause the electrolysis?

(a) What is electrolysis? (b) Are electrolysis reactions thermodynamically spontaneous? Explain. (c) What process occurs at the anode in the electrolysis of molten \(\mathrm{NaCl} ?\) (d) Why is sodium metal not obtained when an aqueous solution of \(\mathrm{NaCl}\) undergoes electrolysis?

Given the following reduction half-reactions: $$ \mathrm{Fe}^{3+}(a q)+\mathrm{e}^{-} \longrightarrow \mathrm{Fe}^{2+}(a q) $$ \(E_{\mathrm{red}}^{\circ}=+0.77 \mathrm{~V}\) $$ \mathrm{S}_{2} \mathrm{O}_{6}{ }^{2-}(a q)+4 \mathrm{H}^{+}(a q)+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{H}_{2} \mathrm{SO}_{3}(a q) $$ \(E_{\mathrm{red}}^{\circ}=+0.60 \mathrm{~V}\) $$ \begin{array}{r} \mathrm{N}_{2} \mathrm{O}(g)+2 \mathrm{H}^{+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{N}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) \\ E_{\mathrm{red}}^{\circ}=-1.77 \mathrm{~V} \\ \mathrm{VO}_{2}^{+}(a q)+2 \mathrm{H}^{+}(a q)+\mathrm{e}^{-} \longrightarrow \mathrm{VO}^{2+}+\mathrm{H}_{2} \mathrm{O}(l) \\ E_{\mathrm{red}}^{\circ}=+1.00 \mathrm{~V} \end{array} $$ (a) Write balanced chemical equations for the oxidation of \(\mathrm{Fe}^{2+}(a q)\) by \(\mathrm{S}_{2} \mathrm{O}_{6}{ }^{2-}(a q),\) by \(\mathrm{N}_{2} \mathrm{O}(a q),\) and by \(\mathrm{VO}_{2}{ }^{+}(a q) .(\mathbf{b})\) Calculate \(\Delta G^{\circ}\) for each reaction at \(298 \mathrm{~K}\). (c) Calculate the equilibrium constant \(K\) for each reaction at \(298 \mathrm{~K}\).

A voltaic cell is constructed with two \(\mathrm{Zn}^{2+}-\mathrm{Zn}\) electrodes. The two half-cells have \(\left[\mathrm{Zn}^{2+}\right]=1.8 \mathrm{M}\) and \(\left[\mathrm{Zn}^{2+}\right]=\) \(1.00 \times 10^{-2} M\), respectively. (a) Which electrode is the anode of the cell? (b) What is the standard emf of the cell? (c) What is the cell emf for the concentrations given? (d) For each electrode, predict whether \(\left[\mathrm{Zn}^{2+}\right]\) will increase, decrease, or stay the same as the cell operates.

A voltaic cell is constructed with two silver-silver chloride electrodes, each of which is based on the following halfreaction: $$\mathrm{AgCl}(s)+\mathrm{e}^{-} \longrightarrow \mathrm{Ag}(s)+\mathrm{Cl}^{-}(a q)$$ The two half-cells have \(\left[\mathrm{Cl}^{-}\right]=0.0150 \mathrm{M}\) and \(\left[\mathrm{Cl}^{-}\right]=\) \(2.55 \mathrm{M},\) respectively. (a) Which electrode is the cathode of the cell? (b) What is the standard emf of the cell? (c) What is the cell emf for the concentrations given? (d) For each electrode, predict whether \(\left[\mathrm{Cl}^{-}\right]\) will increase, decrease, or stay the same as the cell operates.
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