Question

In the second row of the periodic table, \(\mathrm{Be}, \mathrm{N},\) and \(\mathrm{Ne}\) all have positive (unfavorable) electron affinities, whereas the other second-row elements have negative (favorable) electron affinities. Rationalize why Be, \(N,\) and Ne have unfavorable electron affinities.

Short answer

In the second row of the periodic table, Be, N, and Ne have unfavorable electron affinities due to their specific electronic configurations, which require energy input to accommodate an extra electron. Be and Ne have completely filled outer orbitals (\(2s^2\) for Be, and \(2s^2\) \(2p^6\) for Ne) while N has three unpaired electrons in its 2p orbitals (\(1s^2\) \(2s^2\) \(2p^3\)) which leads to increased electron-electron repulsion when adding an extra electron. These factors result in positive, unfavorable electron affinities for these elements compared to the rest of the second-row elements.

Step-by-step solution

Recall the definition of electron affinity

Electron affinity is the amount of energy released or absorbed when an electron is added to a neutral atom, resulting in a negatively charged ion. In general, favorable electron affinities are associated with a release of energy by the atom when adding electrons.

Understand electronic configurations

The electronic configuration of an element is a representation of the distribution of electrons across its atomic orbitals. Elements in the same group (vertical column) in the periodic table have similar electronic configurations because they have the same number of electrons in their outermost shells. With this in mind, we can analyze the electronic configurations of Be, N, and Ne to explain their electron affinities.

Determine the electronic configurations of Be, N, and Ne

Let's examine the electron configurations of Be, N, and Ne: 1. Be (Beryllium) - Atomic number = 4 Electronic configuration: \(1s^2\) \(2s^2\) 2. N (Nitrogen) - Atomic number = 7 Electronic configuration: \(1s^2\) \(2s^2\) \(2p^3\) 3. Ne (Neon) - Atomic number = 10 Electronic configuration: \(1s^2\) \(2s^2\) \(2p^6\)

Analyze the electronic configurations and explain the unfavorable electron affinities

Upon analyzing the electronic configurations of Be, N, and Ne, we can observe the following: 1. Be has a completely filled 2s orbital. Adding another electron would require promotion to the higher energy 2p orbital, which needs energy input, leading to an unfavorable electron affinity. 2. N has three unpaired electrons in its 2p orbitals. Adding another electron results in one of the 2p orbitals having two electrons with the same spin, which increases electrostatic repulsion. Therefore, energy must be provided to the system to allow this process to happen, making the electron affinity unfavorable. 3. Ne has a completely filled 2p orbital. Neon's electron configuration corresponds to a stable noble gas configuration, with completely filled s and p orbitals in the outer shell. To add an additional electron, it must be placed in a 3s orbital with higher energy, which requires energy input. Thus, Ne has an unfavorable electron affinity.

Conclusion

In summary, Be, N, and Ne have unfavorable electron affinities due to their specific electronic configurations, which require energy input to accommodate an extra electron. Be and Ne have completely filled outer orbitals, whereas N's electron affinity is affected by inducing an increase in electron-electron repulsion within its 2p orbitals.

Key concepts

Periodic Table

The periodic table is a structured arrangement of elements based on their atomic number, electron configurations, and chemical properties. One essential aspect of the periodic table is its organization into rows and columns. Each element in a row (also known as a period) has the same number of atomic orbitals. For example, elements in the second row (like Be, N, and Ne) each have two electron shells.

Columns, on the other hand, group elements with similar chemical properties, often because they share the same number of electrons in their outer shell or valence shell. These shared characteristics among groups are crucial in determining how elements interact in chemical reactions.

The periodic table helps predict the behavior of elements, including their electronegativity, ionization energy, and electron affinity. Knowing the position of an element on the table can give insights into its electron arrangement and how it might behave when gaining, losing, or sharing electrons.

Electronic Configurations

Electronic configurations describe the distribution of electrons of an atom in its atomic orbitals. This configuration follows a specific order of filling based on increasing energy levels, primarily determined by quantum mechanics. Each electron shell (or energy level) can hold a finite number of electrons: the first shell holds up to 2, the second up to 8, and so on, following the 2n^2 rule.

In the case of Beryllium (Be), the electronic configuration is

• 1s² 2s²

This indicates that its first shell is completely filled, and the 2s orbital in the second shell is also filled. For Nitrogen (N), the configuration is

• 1s² 2s² 2p³

meaning there are three unpaired electrons in the 2p orbital. Neon (Ne), with

• 1s² 2s² 2p⁶

has a fully filled valence shell, rendering it particularly stable.

The electronic configuration helps explain why these elements have specific electron affinities. Fully or half-filled orbitals often result in higher stability, making additional electron uptake energetically unfavorable for Be, N, and Ne.

Energy Levels

Energy levels, or electron shells, are essentially layers around an atom's nucleus where electrons are most likely to be found. As you move further from the nucleus, the energy level number increases, and the electrons in those levels have higher potential energy.

The energy required to add or remove an electron from an atom considerably depends on these levels. A fully filled or half-filled energy level imparts extra stability to an atom, as seen with noble gases like Neon or elements with half-filled sub-levels like Nitrogen. In such cases, adding electrons requires energy, as it disrupts this stability.

For Be and Ne, the attempt to add an electron would mean promoting an electron to a higher energy level (for instance, moving from a 2p to a 3s orbital in Ne). This counters their stable configuration, thereby making their electron affinities positive or unfavorable.

Understanding these energy levels and their impacts on different elements explains much about how different elements might or might not interact with incoming electrons, revealing much about their chemical behavior.