Question

For each reaction, identify the substance being oxidized and the substance being reduced. (a) \(\mathrm{Mg}\) (s) \(+\mathrm{Br}_{2}(g) \longrightarrow \mathrm{MgBr}_{2}\) (s) (b) \(2 \mathrm{Cr}^{3+}(a q)+3 \mathrm{Mn}(s) \longrightarrow 2 \mathrm{Cr}(s)+3 \mathrm{Mn}^{2+}(a q)\) (c) \(2 \mathrm{H}^{+}(a q)+\mathrm{Ni}(s) \longrightarrow \mathrm{H}_{2}(g)+\mathrm{Ni}^{2+}(a q)\)

Short answer

a) Oxidized: Mg, Reduced: Br; b) Oxidized: Mn, Reduced: Cr^{3+}; c) Oxidized: Ni, Reduced: H^+.

Step-by-step solution

Understand Oxidation and Reduction

Oxidation is the loss of electrons, while reduction is the gain of electrons. In any chemical reaction, the substance that gets oxidized loses electrons, and the one that gets reduced gains electrons.

Analyze Reaction (a)

For the reaction \(\mathrm{Mg}\) (s) + \(\mathrm{Br}_{2}(g) \) to \(\mathrm{MgBr}_{2}\) (s), \(\mathrm{Mg}\) is oxidized, going from 0 to a +2 oxidation state. Conversely, \(\mathrm{Br}_{2}\) is reduced as each \(\mathrm{Br}\) atom goes from 0 to a -1 oxidation state.

Analyze Reaction (b)

For the reaction \(2 \mathrm{Cr}^{3+}(a q)+3 \mathrm{Mn}(s) \longrightarrow 2 \mathrm{Cr}(s)+3 \mathrm{Mn}^{2+}(a q)\), \(\mathrm{Mn}\) is oxidized as it goes from 0 to a +2 oxidation state, while \(\mathrm{Cr}^{3+}\) is reduced to \(\mathrm{Cr}\) (s), going from a +3 to a 0 oxidation state.

Analyze Reaction (c)

In the reaction \(2 \mathrm{H}^{+}(a q)+\mathrm{Ni}(s) \longrightarrow \mathrm{H}_{2}(g)+\mathrm{Ni}^{2+}(a q)\), \(\mathrm{Ni}\) is oxidized from 0 to a +2 oxidation state, and \(\mathrm{H}^{+}\) is reduced as two \(\mathrm{H}^{+}\) ions gain an electron each to form \(\mathrm{H}_{2}(g)\) with an overall 0 oxidation state.

Key concepts

Oxidation and Reduction

Understanding the dynamics between oxidation and reduction is crucial when it comes to analyzing redox reactions. Put simply, oxidation occurs when an atom, molecule, or ion loses one or more electrons, while reduction happens when an atom, molecule, or ion gains electrons. Noteworthy is that these two processes occur simultaneously; it's impossible for a substance to be oxidized without another being reduced—the this is the principle of conservation of charge.

Consider a simple analogy: If oxidation were someone giving away apples (electrons), then reduction would be another person receiving them. This reciprocal action showcases the integral balance within chemical reactions. Delving into the fundamentals, the substance that donates electrons—the 'apple giver'—is known as the reducing agent, while the recipient is termed the oxidizing agent. These agents facilitate the essential electron flow driving the reaction forward.

Oxidation States

Oxidation states, often referred to as oxidation numbers, serve as a bookkeeping tool to keep track of electrons during chemical reactions, especially redox processes. It is a hypothetical charge that an atom would have if all bonds to atoms of different elements were completely ionic.

By utilizing oxidation states, we can discern whether an atom is oxidized or reduced during a reaction. For instance, if an element's oxidation state increases, it indicates electron loss and therefore, oxidation. Conversely, if an oxidation state decreases, it infers electron gain which corresponds to reduction. These states provide a systematic way to describe the electron transfer that is fundamental to redox reactions, allowing chemists to quickly assess the transformation of reactants to products.

Chemical Reactions Analysis

Analyzing chemical reactions, particularly redox reactions, involves examining the changes in oxidation states to identify which substances are oxidized and which are reduced. This analysis is grounded in conservation laws, such as the conservation of charge and mass. When approaching a chemical reaction, it's vital to balance it, ensuring that the number of atoms of each element, and the overall charge, is the same on both sides of the equation.

For a more effective examination, a step-by-step approach is recommended.

Identify the Reactants and Products

Begin by determining what's reacting and what's being formed.

Determine the Oxidation States

Before and after the reaction to see which atoms have gained or lost electrons.

Match Oxidation and Reduction Pairs

Pinpoint which reactants are linked in this electron exchange. With these steps, it becomes more manageable to understand how electrons are shuffled during a reaction and can also shed light on the reaction's spontaneity and energetics.

Electron Transfer Processes

At the heart of redox reactions are electron transfer processes. These are not merely the shuffling of electrons from one species to another but the driving force behind energy production, chemical synthesis, and even biological metabolism. During these processes, the movement of electrons links to changes in energy states, often releasing or consuming energy.

For example, in the electron transfer process of a galvanic cell, the flow of electrons from the reducing agent to the oxidizing agent occurs spontaneously, generating electrical energy. Conversely, in electrolysis, an external power source propels electrons against their spontaneous flow. Mastering the concept of electron transfer is essential for designing batteries, understanding rust formation, and grasping how our bodies metabolize food into energy.