Question

Classify each of the following half-equations as oxidation or reduction. (a) \(\mathrm{Mn}^{2+}(a q) \longrightarrow \mathrm{MnO}_{4}^{-}(a q)\) (b) \(\mathrm{CrO}_{4}^{2-}(a q) \longrightarrow \mathrm{Cr}^{3+}(a q)\) (c) \(\mathrm{PbO}_{2}(s) \longrightarrow \mathrm{Pb}^{2+}(a q)\) (d) \(\mathrm{ClO}_{2}^{-}(a q) \longrightarrow \mathrm{ClO}^{-}(a q)\)

Short answer

Question: Classify the following half-equations as oxidation or reduction: (a) Mn²⁺(aq) → MnO₄⁻(aq) (b) CrO₄²⁻(aq) → Cr³⁺(aq) (c) PbO₂(s) → Pb²⁺(aq) (d) ClO₂⁻(aq) → ClO⁻(aq) Answer: (a) Oxidation (b) Reduction (c) Reduction (d) Reduction

Step-by-step solution

Identify the oxidation numbers for each element

First, we need to identify the oxidation number of the elements involved in the half-equations. The oxidation number indicates the number of electrons that an atom has either gained or lost. Here are the rules for determining the oxidation numbers: - The oxidation number of an element in its elemental state is 0 - The oxidation number of a monoatomic ion is equal to its charge - The oxidation number of oxygen in a compound is usually -2 (except in peroxides, where it is -1) - The oxidation number of hydrogen in a compound is usually +1 (except when it is bound to a metal, then it is -1) - The sum of the oxidation numbers in a compound is equal to the compound's overall charge

Determine if the oxidation numbers have increased or decreased

Next, we will compare the initial and final oxidation numbers for each reaction. If the oxidation number increases, it is an oxidation reaction, and if it decreases, it is a reduction reaction. (a) \(\mathrm{Mn}^{2+}(a q) \longrightarrow \mathrm{MnO}_{4}^{-}(a q)\) Initial oxidation number of Mn: +2 Final oxidation number of Mn: +7 (using the rules mentioned above) The oxidation number increased, so this is an oxidation reaction. (b) \(\mathrm{CrO}_{4}^{2-}(a q) \longrightarrow \mathrm{Cr}^{3+}(a q)\) Initial oxidation number of Cr: +6 (using the rules mentioned above) Final oxidation number of Cr: +3 The oxidation number decreased, so this is a reduction reaction. (c) \(\mathrm{PbO}_{2}(s) \longrightarrow \mathrm{Pb}^{2+}(a q)\) Initial oxidation number of Pb: +4 (using the rules mentioned above) Final oxidation number of Pb: +2 The oxidation number decreased, so this is a reduction reaction. (d) \(\mathrm{ClO}_{2}^{-}(a q) \longrightarrow \mathrm{ClO}^{-}(a q)\) Initial oxidation number of Cl: +3 (using the rules mentioned above) Final oxidation number of Cl: +1 The oxidation number decreased, so this is a reduction reaction.

Final classification

Based on the change in oxidation numbers in each half-reaction, we can classify them as follows: (a) Oxidation (b) Reduction (c) Reduction (d) Reduction

Key concepts

Oxidation Number

Understanding the oxidation number is crucial when examining chemical reactions, especially redox reactions. An oxidation number, also referred to as oxidation state, is a numerical representation that reflects the number of electrons an atom can gain or lose when participating in a chemical bond. It's important to apply certain rules to correctly assign oxidation numbers:

• The oxidation number of a pure element is always zero.
• For a monoatomic ion, the oxidation number equals its charge.
• Oxygen typically has an oxidation number of -2, while hydrogen is usually +1. Exceptions include peroxides for oxygen and hydrides for hydrogen.
• The sum of oxidation numbers in a neutral molecule is zero, and for an ion, it equals the ion's charge.

When comparing oxidation states before and after a reaction, you can identify if electrons were gained or lost and whether oxidation or reduction has occurred.

Redox Reactions

Redox reactions are the bread and butter of chemistry involving electron transfer processes. Redox is a combination of two processes: reduction and oxidation, which always occur simultaneously. An easy way to remember this is 'OIL RIG' - Oxidation Is Loss, Reduction Is Gain of electrons. In these reactions, one species (the reductant) donates electrons and is oxidized, while another species (the oxidant) accepts electrons and is reduced.

Importance of Redox Reactions

Redox reactions are fundamental in various biochemical processes like respiration and photosynthesis, as well as industrial applications such as batteries and metal extraction. Thus, grasping the concept of redox reactions allows us to understand not only academic exercises but a broad spectrum of real-world chemical processes.

Electron Transfer

Electron transfer is the core of redox reactions; it's the motion of electrons from one element or compound to another. This transfer of electrons leads to changes in oxidation numbers, signifying either oxidation or reduction. In academic exercises, tracking the flow of electrons can often be done through half-reactions that show only one of these processes.

Identifying which substance is getting oxidized and which one is getting reduced is essential: electrons 'move' from the oxidized species to the reduced one. In practice, this movement of electrons brings about significant energy changes, which is why redox reactions are commonly harnessed in energy production, such as in batteries where these reactions provide electrical energy.

Chemical Reactions

Chemical reactions are processes where substances (reactants) transform into new substances (products). This transformation involves breaking and forming chemical bonds, often accompanied by energy changes. Chemical reactions can be categorized into various types, but redox reactions are particularly significant due to their ubiquitous nature in chemical and biological systems.

The exercise of identifying oxidation and reduction in half-reactions is essential for understanding the broader scope of chemical reactions and their role in energy transfer, synthesis, and decomposition. These concepts form the foundation for more complex topics in chemistry such as stoichiometry, thermodynamics, and kinetics.