Problem 7 - Chemistry The Molecular Science

Question

In each of these reactions assign an oxidation number to each atom in reactants and products. Identify which substance is oxidized and which is reduced. Identify the oxidizing agent and the reducing agent. (a) \(\mathrm{Fe}(\mathrm{s})+\mathrm{Br}_{2}(\ell) \longrightarrow \mathrm{FeBr}_{2}(\mathrm{~s})\) (b) \(2 \mathrm{Al}(\mathrm{s})+3 \mathrm{Cl}_{2}(\mathrm{~g}) \longrightarrow 2 \mathrm{AlCl}_{3}(\mathrm{~s})\) (c) \(8 \mathrm{HI}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq}) \longrightarrow\) $$ \mathrm{H}_{2} \mathrm{~S}(\mathrm{aq})+4 \mathrm{I}_{2}(\mathrm{~s})+4 \mathrm{H}_{2} \mathrm{O}(\ell) $$

Step-by-Step Solution

Answer

(a) Fe oxidized, Br2 reduced; Fe reducing agent, Br2 oxidizing agent.

Step 1: Assign Oxidation Numbers (a)

For the reaction \( \mathrm{Fe}(\mathrm{s})+\mathrm{Br}_2(\ell) \longrightarrow \mathrm{FeBr}_2(\mathrm{~s}) \):* \( \mathrm{Fe} \) in reactants is 0 (elemental form).* \( \mathrm{Br}_2 \) in reactants is 0 (elemental form).* In \( \mathrm{FeBr}_2 \), \( \mathrm{Fe} \) is +2, \( \mathrm{Br} \) is -1.

Key Concepts

Oxidation Numbers

Understanding oxidation numbers is an important skill when studying chemistry, especially in the context of redox reactions. An oxidation number is a figure used to keep track of electrons in atoms during chemical reactions. It helps us identify which atoms are losing electrons and which ones are gaining.

• In their elemental state, such as in \( \mathrm{Fe}(\mathrm{s}) \) or \( \mathrm{Br}_2(\ell) \), the oxidation number is always 0.
• When elements form compounds, their oxidation numbers can vary. For example, in \( \mathrm{FeBr}_2 \), iron has an oxidation number of +2, and each bromine has an oxidation number of -1.
• The sum of oxidation numbers in a compound must equal the total charge of the compound.

By keeping these rules in mind, you'll be able to determine how electron transfer occurs in chemical reactions.

Oxidizing Agent

An oxidizing agent is a substance that gains electrons in a redox reaction, causing another substance to be oxidized. While it is itself reduced, its role in the reaction is crucial.

• In any redox process, the oxidizing agent facilitates the oxidation of another substance by accepting electrons from it.
• In reaction (a), \( \mathrm{Br}_2 \) is the oxidizing agent because it accepts electrons from \( \mathrm{Fe} \) to form \( \mathrm{Br}^- \).
• Remember: the oxidizing agent is reduced during the reaction!

Recognizing the oxidizing agent in a reaction helps us understand the flow of electrons and the changes that occur to different species involved in the reaction.

Reducing Agent

A reducing agent is the opposite of an oxidizing agent. It donates electrons in a redox reaction, causing another substance to be reduced. In the process, it is oxidized itself.

• The reducing agent loses electrons and increases its oxidation number through this donation.
• In reaction (a), \( \mathrm{Fe} \) serves as the reducing agent because it donates electrons to \( \mathrm{Br}_2 \).
• Create a mental note: the reducing agent is always oxidized, as it loses electrons!

By identifying the reducing agent, we can trace back which element or compound is losing electrons and contributing to the reduction of another component in the reaction.

Redox Reactions

Redox reactions, short for oxidation-reduction reactions, are chemical processes where electron transfer occurs. One substance gives up electrons (oxidation), while another gains them (reduction).

• Redox reactions are common, including processes such as rusting of iron or the burning of fuels.
• In essence, wherever there are changes in oxidation numbers of the substances involved, you're likely dealing with a redox reaction.
• For a balanced redox reaction, the amount of electrons lost must equal the amount of electrons gained by atoms in the system.

Understanding redox reactions involves recognizing the interplay of oxidation and reduction, and determining how substances interact through the transfer of electrons.