HS Chemistry - Periodic Table Trends

Group 2 Elements' Properties

Overview of The Page

This page will cover:

• Common properties among Group 2 elements.

To be done.

The elements in Group $column$ 2 of the Periodic Table $aka the Beryllium Group, although you may also hear it referred to as the Magnesium Group, or even the Calcium Group$ exhibit specific properties and trends due to their similar electron configuration.

Atomic Radius

Each element further down a group in the Periodic Table has an additional electron shell. For example, the element in Period $row$ 3 of the Periodic Table has 3 electron shells, while the element in Period 5 has 5 electron shells. Since they have more electron shells, their atomic and ionic radii will be larger. Thus, Group 2 elements have larger atomic and ionic radii the further down the group they are.

Electrical Conductivity

Group 2 elements are all metals. Therefore, they can all conduct electricity.

Melting Points

Generally, metal atoms with a larger atomic radius have weaker metallic bonds, as the nuclei are farther away from each other. This causes metals with a larger atomic radius to have lower melting and boiling points than metals with a smaller atomic radius. Following this trend, Group 2 elements further down the Group have lower melting and boiling points than those higher up in the Group.

The only exception to this rule for Group 2 elements is Magnesium, which has a lower melting point than Calcium even though Calcium has a larger atomic radius.

Density

At first, going down the Group, the density of the Group 2 elements decreases. This is because as the atomic radius increases, so does the distance between each atom, and therefore less atoms can fit into a certain area, decreasing the density. This continues until Calcium, but then Strontium has a greater density than Calcium, Barium has a greater density than Strontium, and Radium has a greater density than Barium.

Why? After we reach Calcium, we start adding electrons to the d subshell as well instead of just the s and p subshells. That means that while the atomic radius doesn't increase by much between Calcium and Strontium $as only one more electron shell is added$, the number of protons and neutrons in the nucleus does. The difference between Strontium's atomic number $38$ and Calcium's atomic number $20$ is much greater than the difference between Calcium's atomic number $20$ and Magnesium's atomic number $12$. The larger amount of protons in the nucleus is more than enough to offset the decrease in density due to the larger atomic radius. This trend continues as you go further down the Group.

The density of the Group 2 elements decreases from Beryllium to Calcium, then starts increasing.

Electron Configuration

There is a pattern in the electron configurations of the Group 2 elements. Each of them have the same electron configuration as the previous noble gas element, plus two electrons in the next shell's s subshell.

To explain that through examples:

• Beryllium $Be$ has the same electron configuration as Helium, plus 2 2s electrons. It's electron configuration is [He]2s²

• Magnesium $Mg$ has the same electron configuration as Neon, plus 2 3s electrons. It's electron configuration is [Ne]3s²

• Calcium $Ca$ has the same electron configuration as Argon, plus 2 4s electrons. It's electron configuration is [Ar]4s²

And this pattern continues. Thus, each of the Group 2 elements has only 2 valence electrons.

1^st Ionization Energy

The 1^st ionization energy of an atom is the amount of energy required to remove the first valence electron from the atom $i.e. the amount of energy required to turn neutral atom X into ion X^\+^$. As the atomic radius and number of protons in the nucleus is different for each atom, and the first ionization energy depends on these two things, the first ionization energy is also different for each atom.

In general, as the atomic radius increases, the first ionization energy decreases, as it becomes easier to remove that first electron from the atom's valence shell. Thus, the first ionization energy decreases for Group 2 elements as you go down the group.

Electronegativity

The electronegativity of an atom is how tightly it can hold onto its atoms. This means that a larger first ionization energy corresponds to a larger electronegativity. Thus, the electronegativity for Group 2 elements also decreases for Group 2 elements as you go down the group.