Question

Why does ionization energy increase regularly across the periodic table from group \(1 \mathrm{~A}\) to group \(8 \mathrm{~A}\), whereas electron affinity increases irregularly from group \(1 \mathrm{~A}\) to group \(7 \mathrm{~A}\) and then falls dramatically for group \(8 \mathrm{~A} ?\)

Short answer

Ionization energy increases due to stronger nuclear attraction across the period, while electron affinity varies because of subshell configurations and noble gas stability.

Step-by-step solution

Understanding Ionization Energy

Ionization energy is the energy needed to remove an electron from an atom. As you move from Group 1A to Group 8A across a period, the atomic number increases, leading to a greater positive charge in the nucleus. Therefore, the attraction between the nucleus and the electrons increases, requiring more energy to remove an electron.

Analyzing the Ionization Trend

The stronger pull of the nucleus on the electrons across a period results in more energy being required to remove an electron. This systematic increase as more protons are added to the nucleus results in a regular increase in ionization energy across the period.

Understanding Electron Affinity

Electron affinity refers to the energy change when an electron is added to a neutral atom. As you move from Group 1A to 7A, atoms tend to gain electrons to complete their valence shells, generally resulting in an increase in electron affinity.

Exploring Electron Affinity Irregularity

The irregular increase in electron affinity is due to several factors, such as subshell electron configurations and electron-electron repulsions. For instance, adding an electron to a half-filled or filled subshell is less favorable, causing inconsistencies.

Understanding Group 8A Electron Affinity

Group 8A elements, known as noble gases, have complete outer electron shells, making them stable and generally non-reactive. Therefore, these elements have a very low or even negative electron affinity because they do not gain energy by adding electrons.

Key concepts

Ionization Energy

Ionization energy is the energy required to remove an electron from an atom. As you move across a period in the periodic table from left to right, you will notice a regular increase in this energy.

This happens because the number of protons in the nucleus increases, which means there is a greater positive charge. Naturally, this stronger nuclear charge pulls the electrons in closer. In other words, the electrons are held more tightly, which requires more energy to remove one.

This systematic increase in ionization energy goes hand in hand with increasing atomic numbers within a period. Hence, as you progress from Group 1A to Group 8A, removing an outer electron becomes increasingly difficult due to the stronger pull from the nucleus.

Electron Affinity

Electron affinity is all about the energy change when an atom gains an electron. This property is a bit more complex than ionization energy because it does not follow a simple pattern across the period.

From Group 1A to 7A, there is an overall tendency for electron affinity to increase. This increase generally happens because these atoms are striving to achieve a complete valence shell. However, electron affinity does not increase consistently like ionization energy. There are irregularities due to several factors:

• Electron-electron repulsions: When an extra electron is added, it repels the existing electrons, affecting how easily it can be incorporated.
• Subshell configurations: Atoms with half-filled or fully-filled subshells are more stable, so adding another electron isn't always energy favorable.

Consequently, elements with a stable electron arrangement might display lower electron affinities.

Periodic Table Trends

The periodic table organizes elements in such a way that visible trends, known as periodic trends, can be seen as you move across and down the table. These include trends like ionization energy and electron affinity.

While ionization energy tends to increase regularly across a period, electron affinity displays irregular behavior. This is largely due to differences in electron configurations and reactions to added electrons. Group 8A, known as the noble gases, are a clear example of deviation. They have a fully filled outer shell, leading to a dramatic drop in electron affinity. This complete shell means they generally don't gain energy by adding more electrons, which explains their non-reactivity.

Understanding these trends is essential to predict the chemical behavior of elements and explain why they participate in certain reactions, contributing to a well-rounded understanding of how the periodic table is structured.