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Chapter 2: Problem 16

\(14 \mathrm{H}^{+}(a q)+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+3 \mathrm{Ni}(s) \rightarrow\) \(2 \mathrm{Cr}^{3+}(a q)+3 \mathrm{Ni}^{2+}(a q)+7 \mathrm{H}_{2} \mathrm{O}(l)\) In the above reaction, a piece of solid nickel is added to a solution of potassium dichromate. Which species is being oxidized and which is being reduced? \(\quad\) Oxidized \(\quad\) Reduced (A) \(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q) \quad \mathrm{Ni}(s)\) (B) \(\mathrm{Cr}^{3+}(a q) \quad \mathrm{Ni}^{2+}(a q)\) (C) \(\mathrm{Ni}(s) \quad \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)\) (D) \(\mathrm{Ni}^{2+}(a q) \quad \mathrm{Cr}^{3+}(a q)\)

Short Answer

Expert verified
The species getting oxidized is Ni and the species getting reduced is Cr2O72-. Hence, the correct option is (C) Ni, Cr2O72-.

Step by step solution

01

Identifying Oxidation

First, identify the oxidation reaction. The oxidation is a process in which an atom loses its electrons. Compare the Ni atoms before and after the reaction. Before the reaction, the Ni s is neutral, but after the reaction, it forms Ni2+ aq, indicating that it has lost two electrons. Thus, Ni is the species getting oxidized.
02

Identifying Reduction

Next, identify the reduction reaction. Reduction is a process in which an atom gains electrons. Compare the Cr2O72- molecules before and after the reaction. Before the reaction, each Cr atom had an average oxidation state of +6 but after the reaction, each Cr forms Cr3+ aq, meaning its oxidation state decreases to +3. Therefore, it has gained electrons and Cr2O72- is the species getting reduced.

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Key Concepts

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

Oxidation
Oxidation is often thought of as the addition of oxygen to a substance. However, in chemistry, oxidation refers to when an atom loses electrons during a reaction. This is an essential part of redox reactions, where one substance loses electrons, known as oxidation, and another gains them, known as reduction.

Let's consider the reaction involving nickel and potassium dichromate. Nickel, represented as Ni(s), starts in its metallic form and is neutral with no charge. After the reaction, Ni ends up as Ni虏鈦�(aq), a charged ion with a +2 oxidation state. This change occurs because nickel loses two electrons in the process, thus undergoing oxidation. Hence, in this context, nickel is the substance being oxidized.
Reduction
Reduction can be thought of as the opposite of oxidation. It is the process where a substance gains electrons. During redox reactions, reduction occurs simultaneously with oxidation as electrons transfer from one species to another.

In the provided example, the dichromate ion, \( ext{Cr}_2 ext{O}_7^{2-}(aq)\), is reduced. Initially, the chromium within this ion has a high oxidation state of +6. After the reaction, the chromium reduces to form \( ext{Cr}^{3+}(aq)\), indicating a decrease in its oxidation state to +3. This change is due to the gain of electrons, showing that \( ext{Cr}_2 ext{O}_7^{2-}(aq)\) is indeed undergoing reduction.
Oxidation State
The oxidation state, or oxidation number, is a theoretical charge on an atom if all bonds to the atom were completely ionic. It serves as a helpful tool to keep track of electron transfers in redox reactions.

For example, in the compound \( ext{Cr}_2 ext{O}_7^{2-}\), each chromium atom initially holds an oxidation state of +6. After the reaction, it transforms into \( ext{Cr}^{3+}\) with an oxidation state of +3, highlighting a significant reduction. These changes are crucial for identifying which substances are oxidized or reduced. By observing the shift in oxidation states, one can determine the electron movement among atoms during the reaction.
Electron Transfer
Electron transfer is at the core of redox reactions and involves the movement of electrons from one substance to another. This transfer results in changes in the oxidation states, thus defining the nature of the redox process.

In our example with nickel and potassium dichromate, electrons move from nickel atoms to the chromium ions within the dichromate species. Nickel loses electrons, getting oxidized, while chromium gains those electrons, getting reduced. This electron movement is the driving force behind the chemical transformation, making it critical in understanding redox processes. Identifying the direction of electron flow helps discern which species act as oxidizing agents and which as reducing agents.

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Most popular questions from this chapter

Hydrogen fluoride, HF, is a liquid at \(15^{\circ} \mathrm{C}\) . All other hydrogen halides (represented by HX, where \(\mathrm{X}\) is any other halogen) are gases at the same temperature. Why? (A) Fluorine has a very high electronegativity; therefore, the H鈥揊 bond is stronger than any other H鈥揦 bond. (B) HF is smaller than any other H鈥揦 molecule; therefore, it exhibits stronger London dispersion forces. (C) The dipoles in a HF molecule exhibit a particularly strong attraction force to the dipoles in other HF molecules. (D) The H鈥揊 bond is the most ionic in character compared to all other hydrogen halides.

Which of the following best explains why the ionization of atoms can occur during photoelectron spectroscopy, even though ionization is not a thermodynamically favored process? (A) It is an exothermic process due to the release of energy as an electron is liberated from the Coulombic attraction holding it to the nucleus. (B) The entropy of the system increases due to the separation of the electron from its atom. (C) Energy contained in the light can be used to overcome the Coulombic attraction between electrons and the nucleus. (D) The products of the ionization are at a lower energy state than the reactants.

\(2 \mathrm{Ag}^{+}(a q)+\mathrm{Fe}(s) \rightarrow 2 \mathrm{Ag}(s)+\mathrm{Fe}^{2+}(a q)\) Which of the following would cause an increase in potential in the voltaic cell described by the above reaction? (A) Increasing \(\left[\mathrm{Fe}^{2+}\right]\) (B) Adding more \(\mathrm{Fe}(s)\) (C) Decreasing \(\left[\mathrm{Fe}^{2+}\right]\) (D) Removing some \(\mathrm{Fe}(s)\)

Which of the following is true for all bases? (A) All bases donate \(\mathrm{OH}^{-}\) ions into solution. (B) Only strong bases create solutions in which \(\mathrm{OH}^{-}\) ions are present. (C) Only strong bases are good conductors when dissolved in solution. (D) For weak bases, the concentration of the \(\mathrm{OH}^{-}\) ions exceeds the concentration of the base in the solution.

Neutral atoms of chlorine are bombarded by high-energy photons, causing the ejection of electrons from the various filled subshells. Electrons originally from which subshell would have the highest velocity after being ejected? (A) 1s (B) 2p (C) 3p (D) 3d
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