Problem 35 - Chemistry The Molecular Nature of Matter

Question

In each pair, choose the better oxidizing agent. (a) \(\mathrm{O}_{2}\) or \(\mathrm{F}_{2}\) (b) \(A s_{4}\) or \(P_{4}\) (c) \(\mathrm{Br}_{2}\) or \(\mathrm{I}_{2}\) (d) \(\mathrm{P}_{4}\) or \(\mathrm{S}_{8}\) (e) \(\mathrm{Se}_{8}\) or \(\mathrm{Cl}_{2}\) (f) \(\mathrm{As}_{4}\) or \(\mathrm{S}_{8}\)

Step-by-Step Solution

Answer

\(\mathrm{F}_{2}\), \(\mathrm{P}_{4}\), \(\mathrm{Br}_{2}\), \(\mathrm{S}_{8}\), \(\mathrm{Cl}_{2}\), and \(\mathrm{S}_{8}\) are the better oxidizing agents in their respective pairs.

Step 1: Understanding the concept of oxidizing agents

An oxidizing agent, or oxidant, gains electrons and is reduced in a chemical reaction. A better oxidizing agent is one that is more willing to gain electrons (and thus be reduced) than another. Generally, elements with greater electronegativity and a higher tendency to gain electrons act as stronger oxidizing agents.

Step 2: Comparing electronegativity and electron affinity

To determine the better oxidizing agent, compare the electronegativity and electron affinity of the elements involved. The element with a higher value is typically the better oxidizing agent. For nonmetals, being closer to fluorine in the periodic table usually indicates better oxidizing ability.

Step 3: Analyzing the choices

For each pair, choose the element or compound that is the better oxidizing agent based on its position on the periodic table, its electronegativity, and its ability to gain electrons.

Step 4: Making the selection for each pair

(a) Choose \(\mathrm{F}_{2}\) because it is the most electronegative element and has the strongest tendency to gain electrons. ewline(b) Choose \(\mathrm{P}_{4}\) as phosphorus is more electronegative than arsenic. ewline(c) Choose \(\mathrm{Br}_{2}\) because bromine is more electronegative than iodine. ewline(d) Choose \(\mathrm{S}_{8}\) because sulfur is more electronegative than phosphorus. ewline(e) Choose \(\mathrm{Cl}_{2}\) as chlorine is more electronegative than selenium. ewline(f) Choose \(\mathrm{S}_{8}\) over \(\mathrm{As}_{4}\) for the same reason as in comparison (d).

Key Concepts

Electronegativity

Electronegativity is a measure of an atom's ability to attract and hold onto electrons when it's part of a compound. Think of it like a magnet's pull—the higher the electronegativity, the stronger the pull on electrons.

Why does this matter for oxidizing agents? Because the better an element is at 'luring' electrons towards itself, the more effective it will be as an oxidizing agent. That's because oxidizing agents 'want' to gain electrons -- it's how they get reduced.

To help decide which element is the better oxidizing agent, we can use the Pauling scale, which assigns electronegativities from 0 to 4. Fluorine sits at the top of this scale with a value of 4, making it the most electronegative and, hence, a powerful oxidizing agent.

Electron Affinity

Electron Affinity and Oxidizing Agents

Electron affinity is the amount of energy released when an atom in the gas phase accepts an electron to form a negative ion. This trait is crucial when discussing oxidizing agents because an atom with a high electron affinity really 'desires' to gain an electron and release energy in the process.

In their quest to become more stable, elements with high electron affinity values will be more aggressive in seeking out electrons—just like a great oxidizing agent should. High electronegativity often goes hand-in-hand with high electron affinity, as seen in nonmetals, particularly the halogens.

Periodic Table Trends

Trending on the Table: Predicting Oxidizing Ability

The periodic table is a treasure map that shows how atoms behave, and trends across the table can predict properties like electronegativity and electron affinity.

As we move from left to right across a period, electronegativity and electron affinity typically increase due to the atoms desiring to complete their valence shells. Moving up a group, these values also increase. This is why elements like fluorine, found at the upper right of the table, are superstars in the oxidizing agent league.

The periodic trends serve as a guide. Therefore, when comparing elements as oxidizing agents, remember to look at their 'address' on the periodic table. If one element is to the right and/or above the other, it's likely to be the better electron-grabber and, thus, the stronger oxidizing agent.