Question

Determine whether each reaction is a redox reaction. For each redox reaction, identify the oxidizing agent and the reducing agent. a. \(\mathrm{Al}(s)+3 \mathrm{Ag}^{+}(a q) \longrightarrow \mathrm{Al}^{3+}(a q)+3 \mathrm{Ag}(s)\) b. \(\mathrm{SO}_{3}(g)+\mathrm{H}_{2} \mathrm{O}(I) \longrightarrow \mathrm{H}_{2} \mathrm{SO}_{4}(a q)\) c. \(\mathrm{Ba}(s)+\mathrm{Cl}_{2}(g) \longrightarrow \mathrm{BaCl}_{2}(s)\) d. \(\mathrm{Mg}(s)+\mathrm{Br}_{2}(l) \longrightarrow \mathrm{MgBr}_{2}(s)\)

Short answer

a) Redox reaction, reducing agent: Al(s), oxidizing agent: Ag+(aq). b) Not a redox reaction. c) Redox reaction, reducing agent: Ba(s), oxidizing agent: Cl2(g). d) Redox reaction, reducing agent: Mg(s), oxidizing agent: Br2(l).

Step-by-step solution

Identify Redox Reaction in Part (a)

To determine if a reaction is a redox reaction, we check if the oxidation states of the elements change during the reaction. In reaction (a), Aluminum is going from 0 to +3 and Silver is going from +1 to 0 which indicates that Al is oxidized and Ag is reduced. Therefore, it is a redox reaction.

Identify Oxidizing and Reducing Agents in Part (a)

The reducing agent is the species that is oxidized, which is Al(s) as it loses electrons. The oxidizing agent is the species that is reduced, which is Ag+(aq) as it gains electrons.

Identify Redox Reaction in Part (b)

In reaction (b), there is no change in the oxidation states of sulfur or oxygen; the sulfur is in a +6 oxidation state both before and after the reaction and oxygen is in a -2 oxidation state in all compounds. Thus, reaction (b) is not a redox reaction.

Identify Redox Reaction in Part (c)

In reaction (c), Ba goes from an oxidation state of 0 to +2, and Cl goes from 0 to -1. Barium is oxidized, and chlorine is reduced, showing that reaction (c) is a redox reaction.

Identify Oxidizing and Reducing Agents in Part (c)

The reducing agent in part (c) is Ba(s), as it loses electrons, and the oxidizing agent is Cl2(g), as it gains electrons.

Identify Redox Reaction in Part (d)

In reaction (d), Mg goes from 0 to +2, and Br goes from 0 to -1. Magnesium is oxidized, and bromine is reduced, indicating that this reaction is a redox reaction.

Identify Oxidizing and Reducing Agents in Part (d)

The reducing agent is Mg(s), as it donates electrons, and the oxidizing agent is Br2(l), as it accepts electrons.

Key concepts

Oxidation States

Understanding the concept of oxidation states is crucial when analyzing redox reactions. Essentially, an oxidation state is a theoretical charge that an atom would have if all bonds to atoms of different elements were 100% ionic. It is a helpful tool for determining electron flow in a chemical reaction.

When identifying changes in oxidation states, we look for variations in the charge assigned to an element before and after the reaction. For example, in reaction (a) from the provided exercise, Al starts with an oxidation state of 0 and ends with +3, signaling it has lost electrons. Similarly, Ag starts with +1 and goes to 0. These changes signify a transfer of electrons, which is characteristic of a redox reaction.

It's important to recognize certain trends: in elemental form, atoms have an oxidation state of 0. During oxidation, the oxidation state increases, whereas during reduction, it decreases. Therefore, tracking these states during a chemical reaction helps determine the flow of electrons and the corresponding redox process.

Oxidizing Agent

An oxidizing agent, also known as an oxidant, gains electrons and is itself reduced in a redox reaction. This aspect of redox chemistry is about recognizing which reactant is receiving electrons.

In the example of reaction (a), the Ag+ ion is the oxidizing agent because it gains electrons from Al and its oxidation state decreases from +1 to 0. This gain of electrons is often linked to a corresponding decrease in oxidation state, and the species that undergoes this change is the one that is driving the oxidation of the other reactant.

Oxidizing agents are typically substances with elements in high oxidation states or with electronegative elements that can gain one or more electrons, thereby reducing their charge. Recognizing an oxidizing agent in a redox reaction is key to identifying the flow and transfer of electrons from one substance to another.

Reducing Agent

A reducing agent, or reductant, is the exact opposite of an oxidizing agent in a redox reaction. It is the chemical that donates electrons and is itself oxidized in the process.

Using reaction (a) as an illustration, Al is the reducing agent because it donates electrons to Ag+, increasing its oxidation state from 0 to +3. The role of a reducing agent is central to the process of reduction because it enables the other reactant to gain electrons.

Reducing agents often contain elements in low oxidation states or metals that readily lose electrons. Identifying a reducing agent involves detecting which reactant is losing electrons and undergoing an increase in its oxidation state during the redox reaction.

Electron Transfer

Electron transfer is the heart of every redox reaction. It involves the movement of electrons from one reactant to another, leading to a change in oxidation states of the elements involved. In redox chemistry, this is perhaps the most fundamental concept, as the entire definition of a redox reaction hinges on it.

In the given reactions, for instance, electron transfer is evident where the oxidation states change. Al loses three electrons to Ag+ in reaction (a), which is an explicit display of electron transfer, with Al serving as the donor and Ag+ as the recipient.

A deep understanding of electron transfer not only helps to categorize chemical reactions as redox but also enables students to track the flow of energy within the reaction, as electron transfer is often associated with energy release or absorption, which has practical applications in fields like electrochemistry and energy storage.