Question

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

Short answer

Based on their positions in the periodic table, the atom with the larger first ionization energy for each pair is: (a) Ar, (b) Be, (c) Co, (d) S, and (e) Te.

Step-by-step solution

(a) Cl and Ar

Both Chlorine (Cl) and Argon (Ar) are in the same period, with Cl in group 17 and Ar in group 18. Since ionization energy increases from left to right across a period, Argon (Ar) has a larger first ionization energy.

(b) Be and Ca

Beryllium (Be) and Calcium (Ca) are present in the same group (Group 2), but Be is in period 2 whereas Ca is in period 4. Since ionization energy decreases down a group in the periodic table, Beryllium (Be) has a larger first ionization energy.

(c) K and Co

Potassium (K) is in group 1 and period 4, while Cobalt (Co) is in group 9 and period 4. As they are both in the same period but K is to the left of Co, Potassium (K) will have a smaller first ionization energy compared to Cobalt (Co).

(d) S and Ge

Sulfur (S) is in group 16 and period 3, while Germanium (Ge) is in group 14 and period 4. Ge is to the right of S but in a lower period. The effect of being in a lower period is stronger than the effect of being to the right, so Sulfur (S) has a larger first ionization energy.

(e) Sn and Te

Tin (Sn) is in group 14 and period 5, while Tellurium (Te) is in group 16 and period 5. Both elements are in the same period, but Sn is to the left of Te. Consequently, Tin (Sn) has a smaller first ionization energy compared to Tellurium (Te).

Key concepts

Periodic Table Trends

The periodic table is not just a tabular arrangement of the chemical elements; it is a treasure trove of information about the elements' properties. One of the key insights it offers is the trend in first ionization energies, which is the energy needed to remove one electron from a neutral atom. This property is critical in understanding an atom's reactivity. Generally, ionization energy increases as you move from left to right across a period. This is due to the increasing nuclear charge, which draws electrons closer, making them harder to remove. Conversely, ionization energy decreases as you move down a group because the electrons are further from the nucleus, shielded by additional layers of electrons, and thus are easier to remove.

For students to grasp concepts fully, always consider the two primary movements in the periodic table: across periods (horizontal) and down groups (vertical). When comparing elements, it's necessary to analyze their position in relation to these trends to predict which one will have a larger first ionization energy.

Atomic Structure

Understanding atomic structure is fundamental to grasping why ionization energy patterns exist. Atoms consist of protons, neutrons, and electrons. The electrons occupy 'shells' or energy levels, and the outermost electrons, or valence electrons, dictate an atom's chemical behavior. The first ionization energy is associated with the removal of these valence electrons. Factors like the effective nuclear charge (net positive charge experienced by valence electrons) and electron shielding (inner electrons blocking the pull of the nucleus on valence electrons) play crucial roles. An atom with a strong effective nuclear charge or minimal electron shielding will have a higher ionization energy because its valence electrons are held more tightly by the nucleus.

An intuitive way for students to visualize these concepts is by imagining the nucleus as a magnet and electrons as metal balls. Some balls (electrons) are closer to the magnet (nucleus) and harder to remove; this translates to higher ionization energy.

Ionization Energy Comparison

Comparing ionization energies across different elements helps us infer their likely chemical reactions and stability. When elements are within the same period and have an increasing number of protons (like Cl and Ar), the ionization energy tends to increase because the electrons are being pulled more strongly towards the nucleus. On the other hand, when elements belong to the same group (like Be and Ca), the element with fewer electron shells (and therefore electrons closer to the nucleus) will generally have a higher ionization energy. The idea is that the closer the valence electron to the nucleus, the more energy is required to overcome the attraction and remove the electron.

It's important for students to consider all factors, including atomic radius, nuclear charge, and shielding effect, when comparing ionization energies of different elements. This holistic view can sometimes reveal exceptions to the main trends, often in situations where electron subshells are of different types (such as 's' and 'd' subshells).

Periodic Groups and Periods

The periodic table's structure underpins much of chemistry and is divided into periods (rows) and groups (columns). A period corresponds to the number of electron shells an element has. In contrast, a group contains elements with the same number of electrons in their outer shell and hence similar chemical properties. When discussing ionization energy in the context of periods and groups, remember that a higher period number indicates more electron shells, which usually equates to lower ionization energy within a group. Within a period, however, as the group number increases, the ionization energy also generally increases.

For clarity, think of the periodic table as a map, where periods represent the different floors of a building and groups represent rooms on each floor. The higher the floor, the farther you are from the ground floor—akin to how the outermost electrons are farther from the nucleus in atoms with more shells (higher periods). Similarly, rooms (groups) further along the hallway have different characteristics. These analogies can help students better visualize and thus remember the trends in ionization energy throughout the periodic table.