Problem 101 - Chemistry The Molecular Science

Question

Answer these questions about the elements \(X\) and \(Z\), which have the electron configurations shown. $$ \mathrm{X}=[\mathrm{Kr}] 4 d^{10} 5 s^{1} \quad Z=[\mathrm{Ar}] 3 d^{10} 4 s^{2} 4 p^{4} $$ (a) Is element \(X\) a metal or a nonmetal? (b) Which element has the larger atomic radius? (c) Which element would have the greater first ionization energy?

Step-by-Step Solution

Answer

(a) Metal (b) Element X (c) Element Z

Step 1: Determine if Element X is a Metal or Nonmetal

To determine if element X with the configuration \([ ext{Kr}] 4d^{10} 5s^{1}\) is a metal or nonmetal, observe the outer electron configuration. Element X has a single electron in the 5s orbital following a filled 4d subshell, typical of transition metals. Since it fills a transition metal d-block, element X is likely a metal.

Step 2: Compare Atomic Radii of Elements X and Z

Element X has the configuration \([ ext{Kr}] 4d^{10} 5s^{1}\), while element Z has \([ ext{Ar}] 3d^{10} 4s^{2} 4p^{4}\). Element X is likely positioned in the 5th period, while Z is in the 4th period of the periodic table. Elements in the 5th period generally have larger atomic radii than those in the 4th period, yielding element X with the larger atomic radius.

Step 3: Determine First Ionization Energy Between X and Z

The first ionization energy typically decreases down a group and increases across a period. Since X is in a period above Z, it likely requires less energy to remove an electron from X, a characteristic of lower ionization energy as it is higher in the periodic table. Thus, element Z, being a period above, would have the greater first ionization energy.

Key Concepts

Transition Metals

Transition metals are a fascinating group of elements found in the d-block of the periodic table, known for their unique characteristics and roles in chemistry. They are typically defined by having partially filled d orbitals. This feature gives them the ability to form various oxidation states, which is key in many chemical reactions.
Moreover, transition metals, like element X discussed in the exercise, are usually metals. This means they exhibit properties such as good conductivity, malleability, and a shiny appearance. Specifically, element X has a configuration of \([ ext{Kr}] 4d^{10} 5s^{1}\), indicating it's probably a transition metal, thanks to its electron arrangement in the d-block.
Some important characteristics of transition metals include:

• Ability to form complexes with other molecules through dative bonding.
• Exhibit magnetic properties due to unpaired d electrons.
• Able to participate in catalytic activities in industrial reactions.

Atomic Radius

Understanding the atomic radius is crucial when comparing the size of atoms across the periodic table. It refers to the distance from the nucleus to the outer boundary of the electron cloud surrounding it. Typically, the atomic radius increases as you move from top to bottom within a group, because additional electron shells are added. Conversely, it decreases as you move from left to right across a period due to the increasing positive charge that pulls electrons closer.
In the exercise, element X, with its configuration of \([ ext{Kr}] 4d^{10} 5s^{1}\), is likely in the 5th period, whereas element Z, \([ ext{Ar}] 3d^{10} 4s^{2} 4p^{4}\), is in the 4th period. Period 5 elements generally have a greater atomic radius than those in period 4 because they have one more electron shell.
Key points about atomic radius are:

• Increases down a group as new electron shells are added.
• Decreases across a period due to stronger pull from a more positively charged nucleus.
• A larger atomic radius typically means the atom is more reactive, particularly for metals.

First Ionization Energy

First ionization energy is the energy required to remove the outermost electron from a neutral atom in its gaseous state. This fundamental property is a measure of the attraction between the nucleus and the outer electron, which is influenced by both the atomic structure and electron configuration.
Generally, ionization energy decreases as you move down a group in the periodic table because the outer electrons are further away from the nucleus and are shielded by the inner shells. This makes it easier to remove an electron. Conversely, it increases across a period due to a stronger attraction exerted by a higher positive nuclear charge.
Drawing from the exercise, element X is above element Z in the periodic table. X, requiring less energy to remove an electron because it’s in a higher period, implies lower first ionization energy than element Z. Therefore, element Z, being in the lower 4th period, has a higher first ionization energy.
Essential points to note about first ionization energy include:

• Decreases down a group as electrons are farther from the nucleus.
• Increases across a period due to greater nuclear attraction.
• A higher ionization energy indicates a stronger hold on the electrons, meaning lesser reactivity for metals.