Question

Based on their positions in the periodic table, predict which atom of the following pairs will have the smaller first ionization energy: (a) \(\mathrm{Cl}, \mathrm{Ar}\); (b) Be, Ca; (c) K, Co; (d) S, Ge; (e) Sn, Te.

Short answer

Based on the periodic table trends for ionization energy, we predict the following: (a) Cl has a smaller first ionization energy than Ar. (b) Ca has a smaller first ionization energy than Be. (c) K has a smaller first ionization energy than Co. (d) Ge has a smaller first ionization energy than S. (e) Sn has a smaller first ionization energy than Te.

Step-by-step solution

Identify the positions of elements in the periodic table

First, let's identify the positions of the given elements in the periodic table and their respective group and period numbers. The elements and their locations are: (a) Cl (Group 17, Period 3), Ar (Group 18, Period 3) (b) Be (Group 2, Period 2), Ca (Group 2, Period 4) (c) K (Group 1, Period 4), Co (Group 9, Period 4) (d) S (Group 16, Period 3), Ge (Group 14, Period 4) (e) Sn (Group 14, Period 5), Te (Group 16, Period 5)

Comparing ionization energies

Now, let's compare the ionization energies for each pair: (a) Cl and Ar: Both elements belong to the same period (Period 3), so we need to look at the trend for ionization energy within the period. Ionization energy increases across a period, therefore Cl will have smaller first ionization energy than Ar. (b) Be and Ca: Both elements belong to the same group (Group 2), so the trend is that ionization energy decreases down a group. Thus, Be will have a larger ionization energy, and Ca will have a smaller first ionization energy. (c) K and Co: Both elements belong to the same period (Period 4), so we need to look at the trend within the period. Ionization energy increases across the period, therefore K will have a smaller first ionization energy than Co. (d) S and Ge: These elements belong to different periods and groups. Sulfur is in Period 3, Group 16 and germanium is in Period 4, Group 14. Generally, as we move from left to right across a period and from top to bottom in a group, ionization energy decreases. This means that S will have a larger ionization energy and Ge will have a smaller first ionization energy. (e) Sn and Te: Both elements belong to the same period (Period 5), so we need to look at the trend for ionization energy within the period. Ionization energy increases across a period, therefore Sn will have a smaller first ionization energy than Te.

Summary of predicted ionization energies

Based on the periodic table trends for ionization energy we predict the following: (a) Cl has a smaller first ionization energy than Ar. (b) Ca has a smaller first ionization energy than Be. (c) K has a smaller first ionization energy than Co. (d) Ge has a smaller first ionization energy than S. (e) Sn has a smaller first ionization energy than Te.

Key concepts

Periodic Table Trends

Understanding the periodic table trends is essential when predicting properties such as ionization energy, which measures the energy required to remove an electron from an atom in its gaseous state. Ionization energy is influenced by an element's position in the periodic table, reflecting the underlying atomic structure.
One of the key trends is that ionization energy generally increases as we move from left to right across a period. This increase is because as protons are added to the nucleus, the positive charge increases, pulling the electrons closer to the nucleus and increasing the energy required to remove an electron.
Conversely, ionization energy tends to decrease as we move down a group because the added electron shells place electrons farther from the nucleus, reducing the positive nuclear charge's effect on the outermost electrons and therefore reducing the energy required to remove them.

Atomic Structure

The atomic structure of an element is a central factor in determining its ionization energy. Each atom consists of a nucleus containing protons and neutrons, surrounded by electrons in orbitals. The configuration of these electrons, particularly those in the outermost shell, or valence electrons, significantly affects an atom's ionization energy.
In elements with a highly charged nucleus (higher atomic numbers within a period), electrons are more strongly attracted to the nucleus, making their removal more energy-intensive. Thanks to the shielding effect, additional inner electron shells can shield the outer electrons from the nuclear charge in heavier elements (those further down a group), making it easier to remove an electron and resulting in lower ionization energies.

Comparing Elements

When comparing elements, it's important to analyze their position within the periodic table to understand how their ionization energies might differ. For instance, chlorine (Cl) and argon (Ar) are in the same period but different groups. Since Cl is to the left of Ar and ionization energy increases from left to right, Cl has a smaller first ionization energy.
For elements in the same group, such as beryllium (Be) and calcium (Ca), the factor to consider is the additional electron shell present in calcium, which causes Be to have a higher first ionization energy due to less shielding. When elements are in different periods and groups, like sulfur (S) and germanium (Ge), the comparison requires balancing both the period and group trends. Sulfur, being above and to the right of germanium, has a higher ionization energy because it is in a higher period (fewer shielding electron shells) and further to the right (greater nuclear charge affecting valence electrons).