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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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Key Concepts

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

Oxidation
Oxidation is a fundamental concept in chemistry and plays a critical role in redox reactions. Originally, oxidation referred to the combination of a substance with oxygen. However, the definition has expanded over time. Currently, it means the loss of electrons by a molecule, atom, or ion.

When an element undergoes oxidation, it increases in oxidation state, implying that it loses electrons. For instance, when iron rusts, it reacts with oxygen and loses electrons, becoming iron oxide.

In any oxidation process, something must simultaneously be reduced. This is because electrons cannot exist freely; they must be transferred to another substance.
  • Loss of electrons
  • Increase in oxidation state
  • Occurs simultaneously with reduction
Reduction
Reduction is another key concept in redox reactions, complementing the process of oxidation. The term 'reduction' refers to the gain of electrons by a molecule, atom, or ion.

When a substance is reduced, it decreases in oxidation state, meaning it has gained electrons. A common example is the reduction of oxygen in water formation, whereby oxygen molecules gain electrons from hydrogen.

This process can鈥檛 occur alone; if one substance is reduced, another must be oxidized, perpetuating the cycle of electron transfer in redox reactions.
  • Gain of electrons
  • Decrease in oxidation state
  • Occurs alongside oxidation
Electron Transfer
Electron transfer is the mechanism that underlies redox reactions. In these reactions, electrons move from one substance to another.

This movement is what connects oxidation and reduction, since the electron lost by the oxidized substance is the one gained by the reduced substance. The exchange of electrons can involve simple ions or complex molecules.

Understanding electron transfer is crucial because it enables various applications, from generating energy in batteries to facilitating metabolic processes in living organisms.
  • Foundation of redox reactions
  • Involves movement of electrons between substances
  • Essential in energy production and biochemical processes

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Most popular questions from this chapter

A disproportionation reaction involves a substance that acts as both an oxidizing and a reducing agent, producing higher and lower oxidation states of the same element in the products. Which of the following disproportionation reactions are spontaneous under standard conditions? Calculate \(\Delta G^{\circ}\) and \(K\) at \(25^{\circ} \mathrm{C}\) for those reactions that are spontaneous under standard conditions. a. \(2 \mathrm{Cu}^{+}(a q) \longrightarrow \mathrm{Cu}^{2+}(a q)+\mathrm{Cu}(s)\) b. \(3 \mathrm{Fe}^{2+}(a q) \longrightarrow 2 \mathrm{Fe}^{3+}(a q)+\mathrm{Fe}(s)\) c. \(\mathrm{HClO}_{2}(a q) \longrightarrow \mathrm{ClO}_{3}^{-}(a q)+\mathrm{HClO}(a q) \quad\) (unbalanced) Use the half-reactions: \(\mathrm{ClO}_{3}^{-}+3 \mathrm{H}^{+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{HClO}_{2}+\mathrm{H}_{2} \mathrm{O} \quad \mathscr{E}^{\circ}=1.21 \mathrm{V}\) \(\mathrm{HClO}_{2}+2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{HClO}+\mathrm{H}_{2} \mathrm{O} \quad \mathscr{E}^{\circ}=1.65 \mathrm{V}\)

Consider only the species (at standard conditions) $$$\mathrm{Br}^{-}, \quad \mathrm{Br}_{2}, \quad \mathrm{H}^{+}, \quad \mathrm{H}_{2}, \quad \mathrm{La}^{3+}, \quad \mathrm{Ca}, \quad \mathrm{Cd}$$ in answering the following questions. Give reasons for your answers. a. Which is the strongest oxidizing agent? b. Which is the strongest reducing agent? c. Which species can be oxidized by \(\mathrm{MnO}_{4}^{-}\) in acid? d. Which species can be reduced by \(\mathrm{Zn}(s) ?\)

An electrochemical cell is set up using the following unbalanced reaction: $$\mathrm{M}^{a+}(a q)+\mathrm{N}(s) \longrightarrow \mathrm{N}^{2+}(a q)+\mathrm{M}(s)$$ The standard reduction potentials are: $$\mathrm{M}^{a+}+a \mathrm{e}^{-} \longrightarrow \mathrm{M} \quad \mathscr{E}^{\circ}=0.400 \mathrm{V}$$ $$\mathrm{N}^{2+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{N} \quad \mathscr{E}^{\circ}=0.240 \mathrm{V}$$ The cell contains 0.10\(M \mathrm{N}^{2+}\) and produces a voltage of 0.180 \(\mathrm{V}\) . If the concentration of \(\mathrm{M}^{a+}\) is such that the value of the reaction quotient \(Q\) is \(9.32 \times 10^{-3},\) calculate \(\left[\mathrm{M}^{a+}\right] .\) Calculate \(w_{\text { max }}\) for this electrochemical cell.

When copper reacts with nitric acid, a mixture of \(\mathrm{NO}(g)\) and \(\mathrm{NO}_{2}(g)\) is evolved. The volume ratio of the two product gases depends on the concentration of the nitric acid according to the equilibrium $$2 \mathrm{H}^{+}(a q)+2 \mathrm{NO}_{3}^{-}(a q)+\mathrm{NO}(g) \rightleftharpoons 3 \mathrm{NO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l)$$ Consider the following standard reduction potentials at \(25^{\circ} \mathrm{C} :\) $$3 \mathrm{e}^{-}+4 \mathrm{H}^{+}(a q)+\mathrm{NO}_{3}^{-}(a q) \longrightarrow \mathrm{NO}(g)+2 \mathrm{H}_{2} \mathrm{O}(l)$$ $$\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad \mathscr{E}^{\circ}=0.957 \mathrm{V}$$ $$\mathrm{e}^{-}+2 \mathrm{H}^{+}(a q)+\mathrm{NO}_{3}^{-}(a q) \longrightarrow \mathrm{NO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l)$$ $$\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad \mathscr{E}^{\circ}=0.775 \mathrm{V}$$ a. Calculate the equilibrium constant for the above reaction. b. What concentration of nitric acid will produce a NO and NO \(_{2}\) mixture with only 0.20\(\% \mathrm{NO}_{2}\) (by moles) at \(25^{\circ} \mathrm{C}\) and 1.00 atm? Assume that no other gases are present and that the change in acid concentration can be neglected.

The equation \(\Delta G^{\circ}=-\mathrm{nF} \mathscr{E}^{\circ}\) also can be applied to half-reactions. Use standard reduction potentials to estimate \(\Delta G_{\mathrm{f}}^{\circ}\) for \(\mathrm{Fe}^{2+}(a q)\) and \(\mathrm{Fe}^{3+}(a q) .\left(\Delta G_{\mathrm{f}}^{\circ} \text { for } \mathrm{e}^{-}=0 .\right)\)
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