Problem 47 - Chemistry and Chemical Reactivity

Question

Answer the questions below about the elements \(A\) and \(B\), which have the electron configurations shown. $$A=[K r] 5 s^{1} \quad B=[A r] 3 d^{10} 4 s^{2} 4 p^{4}$$ (a) Is element A a metal, nonmetal, or metalloid? (b) Which element has the greater ionization energy? (c) Which element has the less negative electron affinity? (d) Which element has the larger atomic radius?

Step-by-Step Solution

Answer

(a) Metal; (b) Element B; (c) Element A; (d) Element A.

Step 1: Determine Element A's Type

Element A has the electron configuration \([Kr] 5s^1\). This involves its outermost electron being in the \(5s\) subshell, which is characteristic of metals. Specifically, it belongs to Group 1 (ALKALI METALS) in the periodic table. Therefore, element A is a metal.

Step 2: Determine Element B's Type

Element B's electron configuration is \([Ar] 3d^{10} 4s^2 4p^4\). The presence of electrons in the \(p\) subshell classifies it as a nonmetal. It is located in Group 16 (the Chalcogens group) and falls under the category of nonmetals.

Step 3: Compare Ionization Energies

Ionization energy generally increases across a period and decreases down a group. Element A is a Group 1 metal and typically has low ionization energy compared to nonmetals. Element B, a nonmetal in Group 16, generally has a higher ionization energy. Thus, element B has the greater ionization energy.

Step 4: Compare Electron Affinities

Both elements belong to different groups, but nonmetals (element B) often have more negative electron affinities than metals (element A). Therefore, element A, being a metal, has the less negative electron affinity.

Step 5: Compare Atomic Radii

Atomic radius tends to increase down a group and decrease across a period. Element A is in the 5th period, while element B is in the 4th period. Being a period below, element A will have a larger atomic radius than element B.

Key Concepts

Ionization Energy

Ionization energy is the energy required to remove an electron from an atom. It helps us understand how strongly an atom holds onto its electrons.
Generally, ionization energy:

• Increases across a period (from left to right).
• Decreases down a group (from top to bottom).

Element A, being in Group 1 as an alkali metal, has a low ionization energy. These metals do not hold their outermost electrons strongly because they prefer to lose an electron to achieve a stable noble gas configuration.
Element B, a nonmetal in Group 16, requires more energy to remove its outer electrons due to its desire to gain electrons and complete its valence shell. Therefore, B has greater ionization energy than A.

Metallic and Nonmetallic Properties

Metallic and nonmetallic properties are related to how elements behave in terms of conductivity, malleability, and reactivity with other substances. Element A, based on its electron configuration ([Kr] 5s¹), belongs to the alkali metals. Metals typically:

• Conduct electricity.
• Are malleable and ductile.
• Have a shiny appearance.

Conversely, element B, with the configuration ([Ar] 3d¹⁰ 4s² 4p⁴), is a nonmetal from Group 16. Nonmetals are usually:

• Poor conductors of electricity.
• Brittle rather than malleable.
• Not shiny.

Understanding these properties helps in predicting the element's behavior in chemical reactions.

Atomic Radius

The atomic radius is the distance between the nucleus of an atom and its outermost electron. This concept gives insight into the size of an atom.
The trends in atomic radius are:

• Increase down a group (atoms have more electron shells).
• Decrease across a period (greater nuclear charge pulls electrons closer).

Element A exists in the 5th period, while element B is in the 4th period. Since A is lower down the periodic table, it has an extra electron shell, resulting in a larger atomic radius than B. Larger atomic radii indicate elements with more widely spread outer electrons.

Electron Affinity

Electron affinity refers to how much energy is released when an atom gains an electron. It shows an element’s tendency to attract electrons. Typically, nonmetals such as element B have high electron affinities because they want to complete their outer electron shells.

• Metals usually have lower, less negative electron affinities.
• Nonmetals have more negative electron affinities.

For element A, a metal, the desire to gain an electron and release energy is lower, resulting in a less negative electron affinity compared to element B, a nonmetal. Understanding electron affinity aids in predicting reactivity and bonding behavior of elements.