Question

Choose the element with the higher first ionization energy in each pair. $$ \text {a. Br or Bi} \quad \text {b. Na or Rb} \quad \text {c. As or At} \quad \text {d. P or Sn} $$

Short answer

a. Br has higher first ionization energy than Bi. b. Na has higher first ionization energy than Rb. c. As has higher first ionization energy than At. d. P has higher first ionization energy than Sn.

Step-by-step solution

Comparing Ionization Energies for Br and Bi

The first ionization energy tends to increase across a period and decrease down a group in the periodic table. Since bromine (Br) is above bismuth (Bi) in the same group, Br has a higher first ionization energy.

Comparing Ionization Energies for Na and Rb

Sodium (Na) is above rubidium (Rb) in the same group in the periodic table. Therefore, Na has a higher first ionization energy due to being higher in its group.

Comparing Ionization Energies for As and At

Arsenic (As) is above astatine (At) in the same group in the periodic table. Thus, As has a higher first ionization energy because it is higher in its group.

Comparing Ionization Energies for P and Sn

Phosphorus (P) is above tin (Sn) in the same group in the periodic table. As a result, P has a higher first ionization energy since it is positioned higher in the group.

Key concepts

Periodic Table Trends in Ionization Energy

Understanding the trends of first ionization energy within the periodic table is essential for predicting the reactivity and bonding behavior of elements. Ionization energy is the energy required to remove an electron from a gaseous atom or ion. The first ionization energy refers specifically to the removal of the first electron.

Across periods, that is from left to right on the periodic table, the first ionization energy generally increases. This occurs because as we move to the right, the number of protons in the nucleus increases, leading to a stronger attraction between the positively charged nucleus and the negatively charged electrons. Consequently, more energy is required to remove an electron from an atom.

Conversely, as we move down a group in the periodic table, the first ionization energy decreases. This pattern exists because the outermost electrons are farther from the nucleus due to an increasing number of electron shells. This increased distance reduces the nuclear charge experienced by the outermost electrons, making them easier to remove.

By understanding these periodic trends, students can make predictions about ionization energies of various elements and understand the chemical properties that arise from these periodic variations.

Comparing Ionization Energies

When comparing the ionization energies of different elements, such as in the textbook exercise, several guidelines can be used to determine which element has a higher first ionization energy. As the exercise shows, elements that are higher up in the same group on the periodic table, like Br compared to Bi or Na compared to Rb, have a higher first ionization energy. This is a direct application of the trend which states that ionization energy decreases as you move down a group due to the increasing number of electron shells.

Another comparison factor is the effective nuclear charge, which increases across a period due to the addition of protons without significant increase in electron shielding. This trend is the reason why elements towards the right of a period, like As over At, have higher first ionization energies.

However, we must be cautious because this trend does not apply uniformly across all elements. Some exceptions occur due to electron configurations, particularly when comparing elements that are near the transition from one subshell to another, or when comparing elements with filled or half-filled subshells, which are usually more stable and have slightly higher ionization energies.

Group and Period Effects on Ionization

The position of an element within a group or period provides important clues about its first ionization energy. For groups, as the number of occupied electron shells increases, the outer electrons can be removed more easily, which means that elements lower down a group will typically have lower first ionization energies.

Furthermore, the electron configuration within the same period can cause variations in first ionization energy. This is largely due to the stability associated with certain electron configurations. For example, the p-block elements exhibit a small drop in first ionization energy at the Group 15 due to the extra stability of a half-filled p subshell.

Therefore, when comparing elements such as P and Sn, as was done in the textbook exercise, the position of phosphorus higher in the group signifies it would have a higher first ionization energy compared to tin, besides reflecting the stability of its electron configuration. These comparisons emphasize the nuanced understanding needed when discussing ionization energies and how group and period placement plays a vital role.