Question

Arrange the following in order of increasing first ionization energy: \(\mathrm{Na}, \mathrm{Cl}, \mathrm{Al}, \mathrm{S},\) and \(\mathrm{Cs}\).

Short answer

Cs < Na < Al < S < Cl.

Step-by-step solution

Understanding Ionization Energy

Ionization energy is the energy required to remove the outermost electron from an isolated atom in its gaseous phase. Generally, the ionization energy increases across a period (from left to right in the periodic table) due to an increase in nuclear charge and decreases down a group (from top to bottom) as the electrons are further from the nucleus.

Locate Elements on the Periodic Table

The given elements are Na (Sodium), Cl (Chlorine), Al (Aluminum), S (Sulfur), and Cs (Cesium). Sodium, Chlorine, Aluminum, and Sulfur are in the third period while Cesium is in the sixth period. Cesium is also in Group 1, making it likely to have a low ionization energy due to its large atomic size.

Analyze Period Trends

In the third period, Sodium has the lowest ionization energy as it is the first element, followed by Aluminum, then Sulfur, and finally Chlorine, which should have the highest ionization energy due to increased nuclear charge without significant change in shielding effect compared to the others.

Compare Across Periods and Groups

Cesium, being in a lower period and in Group 1, will have the lowest ionization energy among all these elements due to its large atomic radius and higher electron shielding effect, which decreases the nuclear hold on the outer electron. The order of increasing ionization energy is, therefore: Cs < Na < Al < S < Cl.

Key concepts

Periodic Table Trends

Periodic table trends play a crucial role in determining the characteristics of elements, such as their ionization energy. Ionization energy tends to increase as you move from left to right across a period. This is because elements on the right have a greater nuclear charge, meaning they have more protons exerting force on the electrons, thus holding them more tightly. Consequently, it requires more energy to remove an electron from such atoms.

On the other hand, as you move down a group, ionization energy decreases. This trend occurs because although the nuclear charge increases, the effect is outweighed by increased electron shielding and a larger atomic radius. As atoms become bigger and have more filled electron shells, the outermost electrons are further from the nucleus and are also shielded by inner electrons, making them easier to remove.

Atomic Radius

The atomic radius is defined as the distance from the nucleus of an atom to the outermost electron shell. This size influences ionization energy because the further away the outer electrons are from the nucleus, the less energy is required to remove an electron. Hence, larger atomic radii typically correspond to lower ionization energies.

As we move across a period from left to right, the atomic radius decreases. This is due to the increase in nuclear charge pulling the electron cloud closer to the nucleus. Conversely, as we move down a group in the periodic table, the atomic radius increases as additional electron shells are added, making the atom larger and reducing the nuclear pull on the outermost electrons.

Nuclear Charge

Nuclear charge refers to the total charge of the nucleus, which is determined by the number of protons an atom contains. A higher nuclear charge means that more protons are exerting an attractive force on the electrons. This concept explains why elements on the right side of the periodic table generally have higher ionization energies; the increased positive charge holds the electrons more tightly.

This increased attraction not only pulls the electrons closer, reducing the atomic radius, but also means that more energy will be required to remove an electron from the atom. Therefore, the ionization energy will be higher for elements with greater nuclear charge, assuming similar shielding effects.

Electron Shielding

Electron shielding, or the shielding effect, occurs when inner electrons block the effective nuclear charge experienced by outer electrons. This effect plays a significant role in periodic trends, as it influences both atomic radius and ionization energy.

When we descend a group in the periodic table, new electron shells are added, which increases shielding. This results in a decreased effective nuclear charge felt by the outermost electrons, making them easier to remove despite an increase in the nuclear charge. As a consequence, the ionization energy decreases.

Across a period, additional electrons fill the same energy level or shell without significantly increasing shielding, enabling the increased nuclear charge to exert a stronger pull on the outer electrons, thereby raising the ionization energy.