Question

Arrange the elements in order of increasing ionization energy: \(\mathrm{Te}, \mathrm{Pb}, \mathrm{Cl}, \mathrm{S}, \mathrm{Sn}\).

Short answer

Increasing ionization energy order: Pb < Sn < Te < S < Cl.

Step-by-step solution

Understand Ionization Energy

Ionization energy is the energy required to remove the most loosely bound electron from a neutral atom in its gaseous state. It is influenced by the atomic radius, nuclear charge, and electron shielding. Generally, ionization energy increases across a period from left to right and decreases down a group on the periodic table.

Locate the Elements on the Periodic Table

Find the positions of the given elements on the periodic table. This will help you determine the trend in their ionization energies based on their group and period.

Apply the Periodic Trends of Ionization Energy

Considering the periodic trends, we can predict the ionization energy order. Ionization energy tends to be lower for elements situated towards the left and bottom of the periodic table and higher for those towards the right and top.

Arrange the Elements by their Groups and Periods

Te (Tellurium) and S (Sulfur) are in the same group with Sulfur being above Tellurium, hence S has a greater ionization energy. Cl (Chlorine) is to the right of Sulfur, suggesting even higher ionization energy. Sn (Tin) and Pb (Lead) belong to the same group, with Sn being above Pb, thus Sn has a higher ionization energy. However, both are below and to the left of the rest, indicating they have the lowest in the set.

Complete the Order

Starting with the lowest ionization energy: Lead (Pb), then Tin (Sn), next Tellurium (Te), followed by Sulfur (S), and finally Chlorine (Cl) with the highest ionization energy among the given elements.

Key concepts

Periodic Trends

Understanding periodic trends is fundamental when studying the properties of elements on the periodic table. Essentially, these trends can be viewed as the predictable changes that occur within the elements' properties as you move across a period (horizontal rows) or down a group (vertical columns). Ionization energy, for instance, typically increases as one moves from left to right across a period. This is because elements on the right side of the table tend to have a more significant hold on their electrons.

As you progress down a group, ionization energy usually decreases. This is due to the added electron shells that increase atomic size, thus reducing the pull that the nucleus exerts on the valence electrons, making them easier to remove. It is these trends that help in predicting elements' behavior and reactivity, and in arranging elements in order of their ionization energy, like with Te, Pb, Cl, S, and Sn.

Atomic Radius

Atomic radius refers to the size of an atom, commonly measured from the center of the nucleus to the outer boundary of the electron cloud. As we proceed down a group on the periodic table, the atomic radius increases because each subsequent element has an additional electron shell. In contrast, moving across a period from left to right, the radius typically decreases. This happens as the nuclear charge becomes more robust with the presence of more protons, pulling electrons closer to the nucleus.

The atomic radius is a crucial factor in determining ionization energy. Smaller atoms have higher ionization energy because the electrons are closer to the nucleus and thus more tightly bound. Consequently, atoms with larger radii have lower ionization energy, as their valence electrons are further from the nucleus and more easily removed.

Nuclear Charge

Nuclear charge is the total charge of the nucleus, which is the product of the number of protons within the nucleus and the fundamental charge. A higher nuclear charge means a stronger attraction between the nucleus and the electron cloud. As the nuclear charge increases with the addition of protons, electrons are pulled closer to the nucleus, and this enhances the ionization energy of the atoms.

Higher nuclear charge equates to a greater pull on electrons and, hence, an increased energy requirement to remove an electron. In the context of ionization energy trends on the periodic table, elements with higher nuclear charges (like those towards the right and top of the table) will tend to have higher ionization energies. When comparing the ionization energies of Pb, Sn, Te, S, and Cl, understanding nuclear charge helps to explain why elements such as Cl, with a high nuclear charge, also possess high ionization energy.

Electron Shielding

Electron shielding occurs when inner layers of electrons act as a 'shield', reducing the effective nuclear charge felt by the outermost electrons. The electrons in the inner orbitals repel those in the outer orbitals, making it easier for the outer electrons to be removed. In essence, electron shielding lessens the impact of the nucleus' positive charge on the electron cloud's outer edges.

This effect increases with the number of electron shells, meaning that elements lower in a group experience more substantial shielding and, as a result, have lower ionization energy. When applied to the elements Te, Pb, Cl, S, and Sn, electron shielding explains why Te has a lower ionization energy than S, even though they are in the same group. Te has more inner electron shells, leading to greater shielding and a reduced attraction between the nucleus and the outermost electron.