Problem 63 - General Chemistry

Question

In the following reactions, label the oxidizing agent and the reducing agent. a. \(2 \mathrm{Al}(s)+3 \mathrm{~F}_{2}(g) \longrightarrow 2 \mathrm{AlF}_{3}(s)\) b. \(\mathrm{Hg}^{2+}(a q)+\mathrm{NO}_{2}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}(I) \longrightarrow\) $$ \mathrm{Hg}(s)+2 \mathrm{H}^{+}(a q)+\mathrm{NO}_{3}^{-}(a q) $$

Step-by-Step Solution

Answer

a. F_2 is the oxidizing agent, Al is the reducing agent. b. Hg^{2+} is the oxidizing agent, NO_2^{-} is the reducing agent.

Step 1: Write the Oxidation States

Identify the oxidation states for each element in the reactants and products for each reaction. For reaction (a): - Aluminum (Al) in its elemental form is 0. - In fluoride ion ( F_{2} ), each fluorine atom is 0. - In AlF_3 , aluminum is +3 and each fluorine is -1. For reaction (b): - Mercury (Hg^{2+} ) is +2 initially. - NO_2^{-} (nitrite ion): nitrogen is +3. - In Hg (mercury solid), the oxidation state is 0. - In NO_3^{-} (nitrate ion), nitrogen is +5.

Step 2: Determine Electron Transfer

Determine which atoms undergo a change in oxidation state; this change indicates electron transfer. For reaction (a): - Aluminum goes from 0 to +3, losing 3 electrons per atom. - Fluorine goes from 0 to -1, each gaining 1 electron. For reaction (b): - Mercury goes from +2 to 0, gaining 2 electrons. - Nitrogen in NO_2^{-} goes from +3 to +5, losing 2 electrons.

Step 3: Identify the Oxidizing and Reducing Agents

Label the substances that undergo reduction and oxidation. The species that is reduced is the oxidizing agent; the species that loses electrons (is oxidized) is the reducing agent. For reaction (a): - F_2 is reduced (gains electrons) and is the oxidizing agent. - Al is oxidized (loses electrons) and is the reducing agent. For reaction (b): - Hg^{2+} is reduced (gains electrons) and is the oxidizing agent. - NO_2^{-} is oxidized (loses electrons) and is the reducing agent.

Key Concepts

Oxidizing Agent

In a chemical reaction, the oxidizing agent is the substance that helps another component to oxidize by accepting electrons from it.
As the oxidizing agent accepts electrons, it becomes reduced itself, which means it gains electrons.
In the reactions we analyzed, the oxidizing agent is seen in the following ways:

• For reaction (a), the compound \(F_2\) acts as the oxidizing agent because it gains electrons from Aluminum (Al).
• In reaction (b), the species \(Hg^{2+}\) is the oxidizing agent because it accepts electrons from \(NO_2^{-}\).

Understanding the role of oxidizing agents is key, as they are crucial for enabling electron transfer in these reactions.

Reducing Agent

The reducing agent in a redox reaction is responsible for losing electrons, thereby getting oxidized itself.
As the reducing agent loses electrons, it donates them to the substance that is reduced, contributing to the overall exchange of electrons.
Let's look at the reactions:

• In reaction (a), Aluminum (Al) acts as the reducing agent because it loses electrons to the \(F_2\) and becomes oxidized in the process.
• For reaction (b), \(NO_2^{-}\) acts as the reducing agent because it loses electrons to \(Hg^{2+}\), causing it to be oxidized.

Recognizing the reducing agent is important because it signifies the substance that supports the reduction of an alternate reactant.

Electron Transfer

Electron transfer is the exchange of electrons between substances.
This is the essence of oxidation-reduction (redox) reactions where one substance loses electrons and another gains.
Here's how this occurs in our reactions:

• For reaction (a), Aluminum loses electrons and transforms from an oxidation state of 0 to +3, while fluorine gains electrons changing from 0 to -1. This electron exchange facilitates the formation of \(AlF_3\).
• In reaction (b), Mercury gains electrons going from a +2 to a 0 oxidation state, and nitrogen within \(NO_2^{-}\) loses electrons shifting its oxidation state from +3 to +5.

This dynamic of electron moving helps us see the fundamental processes at play, emphasizing changes in oxidation states that denote oxidation or reduction.

Oxidation States

The oxidation state is a practical concept used to keep track of electron transfer in chemical reactions.
It refers to the hypothetical charge an atom would have if all bonds to elements of different atoms were entirely ionic.
The rules for assigning these states can include:

• Atoms in their elemental form are assigned an oxidation state of 0.
• For reaction (a), Aluminum (Al) begins at 0 and moves to +3 in the compound \(AlF_3\). Each fluorine starts at 0 and shifts to -1 in the same compound.
• In reaction (b), Mercury (from \(Hg^{2+}\) starts at +2 and becomes 0, while nitrogen in \(NO_2^{-}\) goes from +3 to +5.

By analyzing these oxidation states, we can identify which species were oxidized or reduced and understand the flow of electrons in a given chemical reaction.