Question

Two elements, $A$ and $B$, have the electron configurations shown. $$\mathrm{A}=[\begin{array}{ll}\mathrm{Ar}& 4s^1\end{array}\mathrm{B}=\left[\begin{array}{l}\mathrm{Ar}\end{array}\right]3d^{10}4{\mathrm{s}}^{2}4p^3$$ (a) Which element is a metal? (b) Which element has the greater ionization energy? (c) Which element has the larger atomic radius? (d) Which element has the greater electron affinity?

Short answer

a) Element A (Potassium) is a metal, b) Element B (Phosphorus) has greater ionization energy, c) Element A (Potassium) has a larger atomic radius, d) Element B (Phosphorus) has a greater electron affinity.

Step-by-step solution

Identification of Element

Identify elements A and B based on the given electron configurations. Element A's configuration ends at 4s¹. It occupies a position of Group 1 (Period 4) in the Periodic Table. This shows that element A is Potassium (K). Element B's configuration ends at 4p³. It occupies a position of Group 15 (Period 4) in the Periodic Table, indicating that B is Phosphorus (P).

Determination of Character

Identify whether the elements are metals or non-metals. The Group 1 elements are known to be highly reactive metals, including potassium. Therefore, element A is metal.

Ionization Energy

Next, we analyze the ionization energy, which is the energy required to remove an electron from an atom. The higher the group number, the higher the ionization energy generally is. Hence, Phosphorus (B) has a greater ionization energy.

Atomic Radius

Review the atomic radius, which depends on the location of the element in the Periodic Table. Elements on the left and lower part of the table have larger atomic radii due to increase in energy levels. Hence, Potassium (A) has a larger atomic radius.

Electron Affinity

Finally, determine the electron affinity, which refers to the ability of the atom to accept an additional electron. Non-metals usually have a higher electron affinity since they readily accept electrons to achieve a full outer energy level. Therefore, Phosphorus (B) has the greater electron affinity.

Key concepts

Ionization Energy

Ionization energy is an important concept in chemistry that refers to the amount of energy needed to remove the outermost electron from an atom in its gaseous state. This property is crucial for understanding an element's reactivity. Generally, ionization energy increases as you move across a period from left to right on the Periodic Table. This is because the nuclear charge increases, pulling electrons closer to the nucleus and thereby requiring more energy to remove them.

For example, in the provided exercise, Phosphorus (Element B) has a higher ionization energy compared to Potassium (Element A). This is because Phosphorus is further to the right on the Periodic Table, in Group 15, whereas Potassium is in Group 1. The increased nuclear charge in Phosphorus leads to a greater hold on its outer electrons, making it harder to remove an electron, thus resulting in a greater ionization energy.

• Ionization energy increases across a period.
• Greater nuclear charge increases ionization energy.
• Phosphorus, with a higher group number, has a higher ionization energy than Potassium.

Atomic Radius

Atomic radius is the term used to describe the size of an atom. This property is essential for predicting and explaining various chemical behaviors. As a rule of thumb, the atomic radius increases as you move down a group on the Periodic Table due to addition of electron shells. Conversely, it decreases as you move from left to right across a period because increasing nuclear charge pulls the electron cloud closer to the nucleus.

In the case of our exercise, Potassium (Element A) has a larger atomic radius than Phosphorus (Element B). Located in Group 1 of the 4th period, Potassium has more electron shells compared to Phosphorus, which sits in Group 15. Consequently, even though Potassium and Phosphorus are in the same period, Potassium’s additional electron shell results in a larger atomic radius.

• Atomic radius increases down a group.
• Atomic radius decreases across a period.
• Potassium’s larger atomic radius results from its additional electron shell.

Electron Affinity

Electron affinity refers to the ability of an atom to accept an additional electron. This property is valuable for understanding the tendency of atoms to form negative ions. Generally, non-metals, which have incomplete electron shells, exhibit higher electron affinities because they tend to gain electrons to achieve stable octet configurations.

In the context of our exercise, Phosphorus (Element B) has a greater electron affinity than Potassium (Element A). Phosphorus, being a non-metal in Group 15, is more inclined to accept additional electrons, as it needs three more electrons to complete its outer shell. Meanwhile, metals like Potassium tend to lose electrons in reactions.

• Non-metals have higher electron affinities.
• Electron affinity is about gaining electrons to complete an outer shell.
• Phosphorus readily gains electrons, unlike Potassium.