Question

What is a disproportionation reaction? Use the following reduction potentials $$\begin{array}{ll} \mathrm{ClO}_{3}^{-}+3 \mathrm{H}^{+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{HClO}_{2}+\mathrm{H}_{2} \mathrm{O} & \mathscr{C}^{\circ}=1.21 \mathrm{V} \\ \mathrm{HClO}_{2}+2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{HClO}+\mathrm{H}_{2} \mathrm{O} & \mathscr{C}^{\circ}=1.65 \mathrm{V} \end{array}$$ to predict whether \(\mathrm{HClO}_{2}\) will disproportionate.

Short answer

A disproportionation reaction is a redox reaction in which a single reactant acts as both the reducing and oxidizing agent, being simultaneously oxidized and reduced to form two different products. Using the given reduction potentials for reactions involving HClO₂, we calculated the potential difference for the disproportionation reaction, which came out to be negative (-2.09 V). Since the potential is negative, the disproportionation of HClO₂ is non-spontaneous and unfavorable under standard conditions. Therefore, HClO₂ will not disproportionate.

Step-by-step solution

Understand disproportionation reactions

Disproportionation reactions are redox reactions in which a single reactant takes part both as the reducing agent and the oxidizing agent. In these reactions, the reactant is simultaneously oxidized and reduced, leading to the formation of two different products. The species undergoing oxidation loses electrons, whereas the species undergoing reduction gains electrons.

Analyze the given reduction potentials

We are given the reduction potentials of two reactions involving the species HClO2: Reaction 1: ClO3- + 3H+ + 2e- → HClO2 + H2O E°₁ = 1.21 V Reaction 2: HClO2 + 2H+ + 2e- → HClO + H2O E°₂ = 1.65 V We can use these reduction potentials to determine whether the disproportionation of HClO2 into ClO3- and HClO is possible by calculating the overall spontaneous reaction direction.

Calculate the potential difference for disproportionation

First, we need to determine the half-reactions for the disproportionation of HClO2: Oxidation half-reaction (reverse of Reaction 2): HClO2 → HClO + H+ + e- E°₃ = -1.65 V Reduction half-reaction (Reaction 1): ClO3- + 6H+ + 2e- → 2HClO2 + 2H2O E°₄ = 1.21 V Now, we need to balance the two half-reactions so that the number of electrons is equal: Oxidation half-reaction (balanced): 2HClO2 → 2HClO + 2H+ + 2e- 2xE°₃ = -3.30 V Reduction half-reaction (unchanged): ClO3- + 6H+ + 2e- → 2HClO2 + 2H2O E°₄ = 1.21 V Next, we find the potential for the disproportionation reaction by adding the oxidation and reduction half-reactions: Overall reaction: 2HClO2 + ClO3- + 6H+ → 2HClO + 2H+ + 2H2O E° (overall) = E°₄ + 2xE°₃ = 1.21 - 3.30 = -2.09 V

Predict disproportionation of HClO2

As the potential difference for the disproportionation reaction (E° (overall)) is negative (-2.09 V), it indicates that the disproportionation of HClO2 is non-spontaneous and unfavorable. Therefore, HClO2 will not disproportionate under standard conditions.

Key concepts

Redox Reactions

Redox reactions, short for reduction-oxidation reactions, occur when one substance is reduced and another is oxidized simultaneously. A reduction reaction involves the gain of electrons by a molecule, atom, or ion, while oxidation refers to the loss of electrons. These redox reactions are essential in a variety of chemical and biological processes, ranging from photosynthesis to combustion.
In a redox reaction, the substance that loses electrons is known as the reducing agent, whereas the substance that gains electrons is called the oxidizing agent. Disproportionation is a unique type of redox reaction where a single species is both oxidized and reduced, forming two different products.
Understanding these reactions is crucial as they help in predicting the direction and feasibility of chemical reactions, often detailed through the analysis of reduction potentials.

Reduction Potentials

Reduction potentials are a measure of the tendency of a chemical species to acquire electrons and be reduced. They are measured in volts (V) under standard conditions. Each half-reaction in a redox process has a specific reduction potential. These potentials are used to determine whether a reaction will proceed spontaneously.
Reduction potentials can be used to arrange half-reactions in order of their electron affinity. For example, in the original problem, we are given two reduction potentials for reactions involving \[ \text{HClO}_{2} \]. A higher reduction potential indicates a greater tendency for the species to gain electrons and be reduced.
When calculating the potential for a full redox reaction, the overall potential difference combines the reduction potentials of the two half-reactions. Determining this overall potential helps in assessing the spontaneity of the reaction, which is crucial for predicting reaction behavior in electrochemical cells.

Spontaneous Reactions

A spontaneous reaction is one that occurs naturally without any external input of energy. In terms of electrochemistry, a reaction is considered spontaneous if the overall change in Gibbs Free Energy is negative. Most straightforwardly, if a redox reaction has a positive potential difference (E°), it tends to proceed spontaneously.
In the context of electrochemical cells, spontaneous reactions are fundamental for generating electrical energy. For example, in the problem, calculating the potential difference of the disproportionation decision for \( \text{HClO}_{2} \), the reaction showed a negative potential. This indicates that under standard conditions, the reaction would not proceed spontaneously, meaning additional energy or altered conditions would be needed to drive the reaction forward.

Oxidation and Reduction

Oxidation and reduction are core concepts in redox chemistry. Oxidation is the process of losing electrons, whereas reduction involves gaining electrons. In a redox reaction, these processes happen simultaneously to match electron transfer.
In the calculation of disproportionation reactions for \( \text{HClO}_{2} \), understanding which part is oxidized and which part is reduced is crucial. For the oxidation half-reaction, \( \text{HClO}_{2} \) loses electrons, thus forming \( \text{HClO} \). Meanwhile, for the reduction half-reaction, electrons are gained.
Applying these principles allows chemists to predict outcomes of complex reactions, balance chemical equations, and comprehend the underlying mechanisms in chemical transformations.

Electrochemistry

Electrochemistry is the branch of chemistry that studies the relationship between electricity and chemical reactions. It involves the movement of electrons and their interactions within chemical cells. Devices like batteries are essentially practical applications of electrochemical principles, turning chemical energy into electric energy.
The analysis of reduction potentials in reactions forms the cornerstone of electrochemical predictions. In the exercise example, the determined negative potential led to the conclusion that \( \text{HClO}_{2} \) would not disproportionate spontaneously under standard conditions. This knowledge is vital for designing processes in batteries, electroplating, and corrosion prevention.
Electrochemistry also provides insights into redox processes, facilitating advancements in energy storage and conversion technologies. Understanding these fundamentals enables further exploration into sustainable energy solutions and efficient electrical systems.