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Chapter 18: Problem 1

Which of the following statements is(are) true? Explain. a. Oxidation and reduction cannot occur independently of each other. b. Oxidation and reduction accompany all chemical reactions. c. A substance that reacts with oxygen gas will always be oxidized.

Short Answer

Expert verified
Statement a is true. Oxidation and reduction are complementary processes that occur simultaneously in redox reactions, where one substance loses electrons (oxidation) and another gains electrons (reduction). Statement b is false. Not all chemical reactions involve oxidation and reduction. There are other types of reactions, such as acid-base, precipitation, and complexation reactions, that do not involve electron transfer. Statement c is true. When a substance reacts with oxygen gas, it typically loses electrons to the oxygen molecule and gets oxidized, while oxygen itself gets reduced by gaining electrons. This is a characteristic of redox reactions involving oxygen.

Step by step solution

01

Statement a: Oxidation and reduction cannot occur independently of each other.

The statement is true. In a redox (reduction-oxidation) reaction, one substance is oxidized by losing electrons, while the other substance is reduced by gaining electrons. Therefore, they occur simultaneously, and an oxidation process can't take place without its corresponding reduction process happening and vice versa.
02

Statement b: Oxidation and reduction accompany all chemical reactions.

The statement is false. While oxidation and reduction participate in many chemical reactions, they are not involved in every chemical reaction. There are other types of reactions, such as acid-base reactions, precipitation reactions, and complexation reactions, that do not involve the transfer of electrons.
03

Statement c: A substance that reacts with oxygen gas will always be oxidized.

The statement is true. In a redox reaction, the presence of oxygen will typically oxidize other substances by taking their electrons. When a substance reacts with oxygen, it loses its electrons to the oxygen molecule, hence it gets oxidized. Oxygen itself gets reduced by gaining electrons. This is a typical characteristic of a redox reaction involving oxygen.

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Most popular questions from this chapter

Consider the galvanic cell based on the following halfreactions: $$\begin{array}{ll}{\mathrm{Zn}^{2+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{Zn}} & {\mathscr{E}^{\circ}=-0.76 \mathrm{V}} \\ {\mathrm{Fe}^{2+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{Fe}} & {\mathscr{E}^{\circ}=-0.44 \mathrm{V}}\end{array}$$ a. Determine the overall cell reaction and calculate $\mathscr{E}_{\text { cell }}$ b. Calculate \(\Delta G^{\circ}\) and \(K\) for the cell reaction at $25^{\circ} \mathrm{C}$ c. Calculate \(\mathscr{E}_{\text { cell }}\) at \(25^{\circ} \mathrm{C}\) when \(\left[\mathrm{Zn}^{2+}\right]=0.10 M\) and $\left[\mathrm{Fe}^{2+}\right]=1.0 \times 10^{-5} \mathrm{M} .$

Consider a galvanic cell based on the following theoretical half-reactions: $$\begin{array}{ll}{\text {}} & { \mathscr{E}^{\circ}(\mathbf{V}) } \\\ \hline{M^{4+}+4 e^{-} \longrightarrow M} & {0.66} \\ {N^{3+}+3 e^{-} \longrightarrow N} & {0.39}\end{array}$$ What is the value of \(\Delta G^{\circ}\) and \(K\) for this cell?

The Ostwald process for the commercial production of nitric acid involves the following three steps: $$4 \mathrm{NH}_{3}(g)+5 \mathrm{O}_{2}(g) \longrightarrow 4 \mathrm{NO}(g)+6 \mathrm{H}_{2} \mathrm{O}(g)$$ $$2 \mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{NO}_{2}(g)$$ $$3 \mathrm{NO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 2 \mathrm{HNO}_{3}(a q)+\mathrm{NO}(g)$$ a. Which reactions in the Ostwald process are oxidation–reduction reactions? b. Identify each oxidizing agent and reducing agent.

Sketch the galvanic cells based on the following half- reactions. Show the direction of electron flow, show the direction of ion migration through the salt bridge, and identify the cathode and anode. Give the overall balanced equation, and determine \(\mathscr{E}^{\circ}\) for the galvanic cells. Assume that all concentrations are 1.0 \(M\) and that all partial pressures are 1.0 atm. a. $\mathrm{H}_{2} \mathrm{O}_{2}+2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \rightarrow 2 \mathrm{H}_{2} \mathrm{O} \quad \mathscr{E}^{\circ}=1.78 \mathrm{V}$ $\mathrm{O}_{2}+2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \rightarrow \mathrm{H}_{2} \mathrm{O}_{2} \quad \quad \mathscr{E}^{\circ}=0.68 \mathrm{V}$ b. $\mathrm{Mn}^{2+}+2 \mathrm{e}^{-} \rightarrow \mathrm{Mn} \quad \quad \mathscr{E}^{\circ}=-1.18 \mathrm{V}$ $\mathrm{Fe}^{3+}+3 \mathrm{e}^{-} \rightarrow \mathrm{Fe} \quad \mathscr{E}^{\circ}=-0.036 \mathrm{V}$

A galvanic cell is based on the following half-reactions: $$\begin{array}{ll}{\mathrm{Fe}^{2+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{Fe}(s)} & {\mathscr{E}^{\circ}=-0.440 \mathrm{V}} \\ {2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{H}_{2}(g)} & {\mathscr{E}^{\circ}=0.000 \mathrm{V}}\end{array}$$ where the iron compartment contains an iron electrode and \(\left[\mathrm{Fe}^{2+}\right]=1.00 \times 10^{-3} M\) and the hydrogen compartment contains a platinum electrode, $P_{\mathrm{H}_{2}}=1.00 \mathrm{atm},$ and a weak acid, HA, at an initial concentration of 1.00 \(\mathrm{M}\) . If the observed cell potential is 0.333 \(\mathrm{V}\) at \(25^{\circ} \mathrm{C},\) calculate the \(K_{\mathrm{a}}\) value for the weak acid HA.
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