Problem 69 - Chemistry and Chemical Reactivity

Question

Answer the following questions about first ionization energies. (a) Generally ionization energies increase on proceeding across a period, but this is not true for magnesium \((738 \mathrm{kJ} / \mathrm{mol})\) and aluminum \((578 \mathrm{kJ} / \mathrm{mol}) .\) Explain this observation. (b) Explain why the ionization energy of phosphorus \((1012 \mathrm{kJ} / \mathrm{mol})\) is greater than that of sulfur \((1000 \mathrm{kJ} / \mathrm{mol})\) when the general trend in ionization energies in a period would predict the opposite.

Step-by-Step Solution

Answer

(a) Magnesium has a filled 3s subshell, making it more stable than aluminum's 3p electron. (b) Phosphorus's half-filled 3p configuration is more stable, causing its higher ionization energy compared to sulfur's electron pairing repulsion.

Step 1: Understanding General Period Trends

Ionization energy generally increases across a period from left to right due to increased nuclear charge pulling electrons closer to the nucleus. As electrons are added to the same principal energy level, an effective nuclear charge also increases, pulling electrons closer and thereby increasing ionization energy.

Step 2: Analyze Magnesium and Aluminum

Magnesium has a higher ionization energy than aluminum despite being to the left in the periodic table. This occurs because magnesium's outer electron is in the 3s subshell, which is lower in energy compared to aluminum's 3p subshell electron. The 3s subshell is fully filled and more stable, requiring more energy to remove an electron from magnesium.

Step 3: Compare Phosphorus and Sulfur

Phosphorus has a higher ionization energy than sulfur, which contradicts the expected periodic trend in which sulfur, being further to the right, would typically have a higher ionization energy. This anomaly is due to electron-electron repulsion in sulfur. Sulfur has paired electrons in the 3p^4 configuration, while phosphorus has unpaired electrons in the 3p^3 configuration, resulting in greater repulsion among paired electrons in sulfur, making it easier to remove an electron compared to phosphorus.

Key Concepts

Periodic Trends

Periodic trends refer to patterns observed in the properties of elements as you move across periods (rows) or down groups (columns) in the periodic table. One such trend is the increase in ionization energy as you move from left to right across a period. This is because the nuclear charge (the positive charge of the nucleus) increases, pulling the electrons more strongly toward the nucleus.

However, there are exceptions to this trend, which is an essential part of understanding ionization energy behavior. As seen with magnesium and aluminum or phosphorus and sulfur, these exceptions occur due to specific atomic configurations. These exceptions remind us that while trends provide a general guide, the actual electron arrangements can lead to unexpected ionization energy values.

Understanding periodic trends alongside their exceptions gives insights into electronic arrangement and the peculiarities of elements that form the periodic table's rich and complex landscape.

Effective Nuclear Charge

The concept of effective nuclear charge is crucial for understanding ionization energy. Effective nuclear charge ( (Z_{eff}) ) refers to the net positive charge experienced by an electron in an atom. It's an important factor because it influences how tightly an electron is held by the nucleus.

• As you move across a period in the periodic table, the number of protons (positive charges) in the nucleus increases.
• Meanwhile, additional electrons are added to the same energy level (or shell) but do not completely shield each other from the increasing nuclear charge.

As a result, the effective nuclear charge experienced by outer electrons increases, pulling them closer to the nucleus and usually increasing ionization energy. This is the reason behind the general increase in ionization energies across periods. Yet, in cases like magnesium and aluminum, subshell configurations modify this trend, causing variations in expected values.

Subshell Electron Configuration

Subshell electron configurations are pivotal for explaining anomalies in ionization energies. Each energy level of an atom is divided into subshells, namely s, p, d, and f, each with specific electron capacities. The configurations within these subshells influence energy stability. For example:

• Magnesium’s outer electron is located in the filled 3s subshell whereas aluminum’s valence electron is in the 3p subshell.
• The 3s subshell is lower in energy and more stable than the 3p. This stability accounts for magnesium's higher ionization energy despite aluminum being further to the right on the periodic table.

These configurations show us how seemingly minor differences in electron location can have significant impacts on energy requirements, providing a nuanced understanding of ionization energies.

Electron-Electron Repulsion

Electron-electron repulsion is another factor affecting ionization energies, especially when considering deviations from periodic trends. This concept pertains to the forces exerted between electrons within an atom, influencing its overall energy structure. Take phosphorus and sulfur as examples:

• In phosphorus, electrons in the 3p subshell are unpaired, leading to relatively less repulsion among them.
• In sulfur, however, one of the electrons pairs up in the 3p subshell, leading to increased electron-electron repulsion.

This increased repulsion in sulfur's case makes it easier – surprisingly so – to remove an electron compared to phosphorus, as less energy is required to overcome the repulsion forces. This explains why sulfur has a lower ionization energy than phosphorus, despite trends suggesting otherwise. Electron-electron repulsion is therefore a key component in understanding these periodic inconsistencies.