Question

In Table 7.8 , the bonding atomic radius of neon is listed as \(58 \mathrm{pm},\) whereas that for xenon is listed as \(140 \mathrm{pm}\). A classmate of yours states that the value for Xe is more realistic than the one for Ne. Is she correct? If so, what is the basis for her statement?

Short answer

The classmate's statement is correct. The basis for her statement is the understanding of periodic trends and the relationship between the atomic radius and the position of elements in the periodic table. Xenon (Xe) has a larger atomic radius (140 pm) compared to Neon (Ne) with an atomic radius of 58 pm, due to its lower position in the periodic table, having more electron shells.

Step-by-step solution

The atomic radius is the distance from the nucleus to the outermost electron shell in an atom. As we move across a period in the periodic table, the atomic radius tends to decrease because the number of protons in the nucleus increases, leading to an increased attraction between the electrons and the nucleus. On the other hand, as we move down a group in the periodic table, the atomic radius increases due to additional electron shells. #Step 2: Compare Neon and Xenon in the periodic table#

Neon (Ne) and Xenon (Xe) are both noble gases. Neon is in the second period of the periodic table, while Xenon is in the fifth period. Since Xenon is farther down the group than Neon, it has more electron shells, which results in a larger atomic radius. #Step 3: Evaluate the classmate's statement#

The classmate's statement suggests that the bonding atomic radius value for Xenon (Xe) is more realistic than that of Neon (Ne). Considering the periodic trends, Xenon's atomic radius of 140 pm is expected to be larger than Neon's atomic radius of 58 pm, since Xenon is lower in the group and has more electron shells. #Step 4: Determine the correctness of the classmate's statement#

Since the periodic trends justify the larger bonding atomic radius for Xenon compared to Neon, we can conclude that the classmate's statement is correct. The basis for her statement is likely the understanding of periodic trends and the relationship between the atomic radius and the position of elements in the periodic table.

Key concepts

Periodic Trends

Periodic trends refer to patterns or regularities that can be observed across different periods and groups in the periodic table. One of the most prominent trends is the atomic radius, which changes as you move across periods and down groups.
The atomic radius typically decreases as you move from left to right across a period. This occurs because the number of protons in the nucleus increases. As a result, the attraction between the nucleus and the electrons becomes stronger, pulling the electrons closer and reducing the size of the atom.
Conversely, as you move down a group, the atomic radius generally increases. This increase is due to the addition of electron shells, which places the outermost electrons farther from the nucleus despite the increase in proton number. Thus, the shielding effect reduces the effective nuclear charge felt by the outermost electrons, allowing the atomic radius to expand. Understanding these trends is crucial to predicting and explaining the properties of elements, such as bonding and reactivity.

Noble Gases

Noble gases are a group of elements in Group 18 of the periodic table. They are known for their remarkable stability and lack of reactivity under standard conditions. This is because they have completely filled electron shells.
Neon (Ne) and Xenon (Xe) are both noble gases, with Neon being lighter and positioned in the second period, while Xenon resides in the fifth period. The filled outer shells make these elements chemically inert because they don't readily accept or donate electrons. This stability is why noble gases are often used in situations that require non-reactive environments, such as in lighting and in certain industrial processes.
The atomic radii of noble gases increase as you move down the group. This is due to the addition of electron shells rather than changes in nuclear charge, which geometrically configures from smaller in Ne to larger in Xe. Despite their inert nature, elements like Xenon can form compounds under extreme conditions due to their larger atomic size, which allows for more accessible electron perturbation.

Electron Shells

Electron shells are the orbitals around an atom's nucleus where electrons are most likely to be found. Each shell can hold a limited number of electrons and is associated with a particular energy level. As you move outwards from the nucleus, these shells increase in energy.
The distribution of electrons among these shells determines the chemical properties of an element, including its reactivity and atomic size. When looking at Neon (Ne) versus Xenon (Xe), a key difference is the number of electron shells. Neon has two shells, while Xenon has five.

• Neon's first shell can contain up to 2 electrons, and the second can hold up to 8.
• Xenon's five shells have a much larger capacity, accommodating more electrons overall.

This difference in electron shells is a major factor contributing to the difference in atomic radius between these two noble gases. The more shells there are, the larger the atomic radius, since electrons are positioned further from the nucleus.

Periodic Table Groups

Periodic table groups are columns that categorize elements with similar chemical properties. Elements in the same group typically have the same number of valence electrons, which govern their chemical behaviors and trends.
Each group in the periodic table is a powerhouse of common characteristics shared among the elements. For instance:

• Group 1 elements (alkali metals) are known for being highly reactive.
• Group 17 elements (halogens) are very reactive non-metals.
• Group 18 contains the noble gases, known for their stability and lack of reactivity.

Moving down a group results in increasing atomic radii due to additional electron shells, enhancing the variation among elements in terms of size and energy levels.
In the case of noble gases within Group 18, the increase in having more electron shells in lower positioned elements like Xenon is manifest in their larger atomic radius relative to lighter ones like Neon. These properties underline the importance of understanding group trends for predicting the general behavior of elements.