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Chapter 6: Problem 219

The reaction, \(3 \mathrm{ClO}^{-}(\mathrm{aq}) \rightarrow \mathrm{ClO}_{3}(\mathrm{aq})+2 \mathrm{Cl}^{-}(\mathrm{aq})\) is an example of a. oxidation reaction b. reduction reaction c. disproportionation reaction d. decomposition reaction

Short Answer

Expert verified
The reaction is a disproportionation reaction (c).

Step by step solution

01

Identify the Change in Oxidation States

In the reaction \(3 \text{ClO}^{-} \rightarrow \text{ClO}_3 + 2 \text{Cl}^{-}\), we need to determine the oxidation states of chlorine in the reactants and products. Initially, in \(\text{ClO}^{-}\), chlorine has an oxidation state of +1. In the product \(\text{ClO}_3\), chlorine has an oxidation state of +5, and in \(\text{Cl}^{-}\), chlorine has an oxidation state of -1.
02

Analyze the Changes in Oxidation States

We observe that one chlorine atom in \(\text{ClO}^{-}\) is oxidized to +5 in \(\text{ClO}_3\) (an increase in oxidation state), while another chlorine atom is reduced to -1 in \(\text{Cl}^{-}\) (a decrease in oxidation state).
03

Define Disproportionation Reaction

A disproportionation reaction is a type of redox reaction in which a single element undergoes both oxidation and reduction simultaneously. Since chlorine is both oxidized and reduced in this reaction, it fits the definition of a disproportionation reaction.

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Key Concepts

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

Oxidation States
In chemistry, oxidation states (also known as oxidation numbers) are crucial for understanding how electrons are distributed in molecules. They help us track the transfer of electrons in redox reactions. Oxidation state of an element within a compound is typically calculated based on a set of rules:
  • The oxidation state of a free element is always zero.
  • The oxidation state of a monoatomic ion is equal to the charge of the ion.
  • Oxygen has an oxidation state of -2 in most compounds (except in peroxides and some other cases).
  • Hydrogen is usually +1, except when bonded to metals as in hydrides where it's -1.
  • The sum of oxidation states for all atoms in a molecule or polyatomic ion equals the total charge of that molecule or ion.
Examining compounds in the provided reaction, chlorine in \(\text{ClO}^{-}\) has an oxidation state of +1, aligning with these rules. Balancing the changes in oxidation states helps us confirm the presence of redox processes in reactions.
Oxidation Reaction
An oxidation reaction is part of a redox process in which an element loses electrons, leading to an increase in its oxidation state. This is often accompanied by the element bonding to more electronegative atoms (like oxygen) or losing bonds to less electronegative atoms (such as hydrogen). In the classroom exercise with the reaction \(3 \text{ClO}^- \rightarrow \text{ClO}_3 + 2\text{Cl}^-\), one of the chlorine atoms in \(\text{ClO}^-\) experiences an oxidation as its oxidation state changes from +1 to +5 in \(\text{ClO}_3\).Oxidation is crucial in various chemical processes:
  • Production of energy in biological systems (e.g., during cellular respiration).
  • Burning of fuels, where carbon is oxidized to carbon dioxide.
  • Corrosion of metals, such as the rusting of iron.
Understanding these reactions helps us appreciate both natural processes and industrial applications.
Reduction Reaction
Reduction involves the gain of electrons by an atom, which results in a decrease in its oxidation state. Importantly, in any redox reaction, oxidation and reduction occur simultaneously. They are paired in a way that electrons lost by one element (in the oxidation) are gained by another (in the reduction). In the disproportionation reaction provided, a chlorine atom in \(\text{ClO}^-\) is reduced as its oxidation state moves from +1 to -1 in \(\text{Cl}^-\). Reduction is a fundamental concept in:
  • Metallurgy, such as converting metal ores to pure metals.
  • The function of electrochemical cells and batteries, where they allow for energy storage and release.
  • Biochemical pathways like photosynthesis, where carbon dioxide is reduced to glucose.
By understanding reduction reactions, students learn the balancing act of electrons in chemical reactions and the pivotal roles these play in both technology and nature.

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Most popular questions from this chapter

The \(\mathrm{emf}, \mathrm{E}\), is related to the change in Gibbs free energy, \(\Delta \mathrm{G}: \Delta \mathrm{G}=-\mathrm{nFE}\), where is the number of electrons transferred during the redox process and \(F\) is a unit called the Faraday. The faraday is the amount of charge on \(1 \mathrm{~mol}\) of electrons: \(1 \mathrm{~F}=96,500 \mathrm{C} / \mathrm{mol}\). Because \(\mathrm{E}\) is related to \(\Delta \mathrm{G}\), the sign of \(\mathrm{E}\) indicates whether a redox process is spontaneous: \(\mathrm{E}>0\) indicates a spontaneous process, and \(\mathrm{E}<0\) indicates a non-spontaneous one. \(\mathrm{Ag}^{+}(\mathrm{aq})+\mathrm{e}^{-} \rightarrow \mathrm{Ag}(\mathrm{s}) \mathrm{E}^{\circ}=+0.800 \mathrm{~V}\) \(\mathrm{AgBr}(\mathrm{s})+\mathrm{e}^{-} \rightarrow \mathrm{Ag}(\mathrm{s})+\mathrm{Br}^{-}(\mathrm{aq})\) \(\mathrm{E}^{\mathrm{o}}=+0.071 \mathrm{~V}\) \(\mathrm{Br}_{2}(\mathrm{l})+2 \mathrm{e}^{-} \rightarrow 2 \mathrm{Br}^{-}(\mathrm{aq}) \mathrm{E}^{\circ}=+1.066 \mathrm{~V}\) Use some of the above data to calculate Ksp at \(25^{\circ} \mathrm{C}\) for \(\mathrm{AgBr}\). a. \(6.7 \times 10^{-12}\) b. \(2.8 \times 10^{-14}\) c. \(3.7 \times 10^{-19}\) d. \(4.9 \times 10^{-13}\)

\(4.5 \mathrm{~g}\) of aluminium (at. mass 27 amu) is deposited at cathode from \(\mathrm{Al}^{3+}\) solution by a certain quantity of electric charge. The volume of hydrogen produced at STP from \(\mathrm{H}^{+}\)ions is solution by the same quantity of electric charge will be a. \(44.8 \mathrm{~L}\) b. \(22.4 \mathrm{~L}\) c. \(11.2 \mathrm{~L}\) d. \(5.6 \mathrm{~L}\).

P, Q, R and S are four metals. P can displace R from its salt solution, but \(\mathrm{Q}\) and \(\mathrm{R}\) cannot displace \(\mathrm{S} . \mathrm{Q}\) can displace hydrogen (H) for a dilute solution of a mineral acid, but \(R\) cannot. The reduction potentials of \(\mathrm{P}, \mathrm{Q}, \mathrm{R}, \mathrm{S}\) and \(\mathrm{H}\) (hydrogen) are in the order a. \(\mathrm{H}>\mathrm{P}>\mathrm{R}>\mathrm{S}>\mathrm{Q}\) b. \(\mathrm{P}>\mathrm{S}>\mathrm{Q}>\mathrm{H}>\mathrm{R}\) c. \(\mathrm{P}>\mathrm{Q}>\mathrm{R}>\mathrm{S}>\mathrm{H}\) d. \(\mathrm{R}>\mathrm{H}>\mathrm{Q}>\mathrm{S}>\mathrm{P}\)

The standard reduction potential values of three metallic cations \(X, Y\), and \(Z\) are \(0.52,-3.03\) and \(-1.18\) respectively. The order of reducing power of the corresponding metal is a. \(\mathrm{Y}>\mathrm{Z}>\mathrm{X}\) b. \(\mathrm{X}>\mathrm{Y}>\mathrm{Z}\) c. \(Z>Y>X\) d. \(Z>X>Y\).

Match the following: Column I \(\quad\) Column II A. A gas in contact with an (p) electrode potential inert electrode B. the potential difference (q) \(\mathrm{O}_{2}\) at anode metal in volts between the two and the solution of electrodes \(\quad\) metal ion C. Li metal has the lowest (r) \(\mathrm{H}_{2}(\mathrm{~g}) / \mathrm{Pt}\) potential standard electrode D. electrolysis of aq. \(\mathrm{Na}_{2} \mathrm{SO}_{4}\) (s) strongest reducing using Pt electrodes agent (t) \(-3.06\) Volt
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