Question

The electron affinities of the elements from aluminum to chlorine are \(-44,-120,-74,-200.4,\) and \(-384.7 \mathrm{kJ} / \mathrm{mol},\) respectively. Rationalize the trend in these values.

Short answer

The trend in electron affinities from aluminum to chlorine generally increases (becomes more negative), which can be attributed to the increase in effective nuclear charge, making it more favorable for atoms to gain electrons. The exception between silicon and phosphorus occurs due to electron-electron repulsion caused by the introduction of a new electron into the same orbital, which results in a lower electron affinity for phosphorus.

Step-by-step solution

Understand Electron Affinity

Electron affinity is the energy change that occurs when an electron is added to a neutral atom to form a negative ion. An element with a more negative electron affinity value indicates that it is more favorable for the atom to gain an electron.

Identify the Elements in the Periodic Table

To better understand the trend, let's identify the elements from aluminum to chlorine on the periodic table: - Aluminum (Al, atomic number 13) - Silicon (Si, atomic number 14) - Phosphorus (P, atomic number 15) - Sulfur (S, atomic number 16) - Chlorine (Cl, atomic number 17) These elements are in period 3 of the periodic table and belong to groups 13, 14, 15, 16, and 17, respectively.

Analyze the Trend in Electron Affinities

Now we'll look at the electron affinities in the given order and compare them: - Aluminum (Al): -44 kJ/mol - Silicon (Si): -120 kJ/mol - Phosphorus (P): -74 kJ/mol - Sulfur (S): -200.4 kJ/mol - Chlorine (Cl): -384.7 kJ/mol We can observe the following trend: - The electron affinity generally increases (becomes more negative) from aluminum to chlorine. - There is an exception between silicon and phosphorus where the electron affinity decreases (becomes less negative).

Explain the Trend Based on Atomic Structure

To rationalize the trend, we need to consider the electron configuration of the elements and their positions in the periodic table: 1. Aluminum and silicon are both located in the p-block but have half-filled (Al) and fully filled (Si) s orbitals. Silicon has a more stable electron configuration, which makes it more favorable for the atom to gain an electron, resulting in a higher electron affinity. 2. The electron affinity decreases from silicon to phosphorus due to the introduction of a new electron into the 3p orbital, which causes electron-electron repulsion. This repulsion makes phosphorus less willing to gain an additional electron, resulting in a lower electron affinity. 3. From phosphorus to chlorine, the increase in the effective nuclear charge (number of protons) is responsible for the increase in electron affinity. The more protons in the nucleus, the stronger the attraction for negatively charged electrons, which makes gaining an electron more favorable and electron affinity more negative. In summary, the electron affinities of the elements from aluminum to chlorine show a general trend of increasing (becoming more negative) due to the increase in effective nuclear charge. The exception between silicon and phosphorus can be explained by the electron-electron repulsion caused by adding a new electron into the same orbital.

Key concepts

Periodic Table

The periodic table is an essential tool for understanding the properties and behaviors of elements. It arranges all known elements in order of increasing atomic number and groups them into columns based on similar properties. This organization helps predict how elements will interact. For example, the elements from aluminum to chlorine, which are in period 3 of the periodic table, display a variety of chemical behaviors as they move across the table.

As you move from left to right across a period, atomic number increases, meaning additional protons and electrons are added to the atoms. This change influences key properties like electron affinity, which is our main focus. By understanding the periodic table layout, we can better predict the trends in element behavior such as changes in electron affinity.

Atomic Structure

Atomic structure refers to how atoms are composed, primarily concerning their protons, neutrons, and electrons. Understanding an atom's structure, particularly its electron configuration, illuminates why certain trends occur in element properties like electron affinity.

Each element has a specific arrangement of electrons around its nucleus, filling electron shells and subshells in an order that minimizes energy. Elements such as aluminum and chlorine have different configurations, affecting their ability to gain or lose electrons. For instance, chlorine has a nearly complete outer shell, making it highly favorable for one additional electron, hence its high electron affinity.

When analyzing trends in the periodic table, it's crucial to understand how these electron configurations influence atomic interactions. This includes examining half-filled or fully filled subshells, which can offer stability and affect the atom's energy levels and reactivity.

Effective Nuclear Charge

The effective nuclear charge is a measure of the net positive charge experienced by outer electrons in an atom. It is a concept that explains why electrons are attracted to the nucleus. As you move across a period in the periodic table, the effective nuclear charge increases because more protons are added to the nucleus, yet electron shielding does not increase at the same rate.

This concept helps explain why electron affinity becomes more negative as you move from aluminum to chlorine. Chlorine, with its higher atomic number, experiences a greater effective nuclear charge, strongly attracting additional electrons. Therefore, it releases more energy when gaining an electron, leading to a higher (more negative) electron affinity value compared to aluminum and others earlier in the period.

Electron-Electron Repulsion

Electron-electron repulsion occurs when negatively charged electrons come into close proximity, resulting in a force that pushes them apart. This repulsion can significantly affect an atom's ability to gain additional electrons and helps explain exceptions in electron affinity trends.

For example, between silicon and phosphorus, the electron affinity decreases rather than increases. This anomaly can be attributed to electron-electron repulsion as phosphorus has an additional electron entering an already partially filled subshell. The added electron increases repulsion among electrons, leading to decreased stability and a less favorable gain of an electron.

It is important to consider electron-electron repulsion when evaluating electron affinities because it often explains why some elements deviate from the expected trends based on effective nuclear charge alone.