Question

Arrange the following groups of atoms in order of increasing size. a. \(\mathrm{Te}, \mathrm{S}, \mathrm{Se}\) b. \(\mathrm{K}, \mathrm{Br}, \mathrm{Ni}\) c. \(\mathrm{Ba}, \mathrm{Si}, \mathrm{F}\)

Short answer

Arrange the groups of atoms in order of increasing size as follows: a. \(\mathrm{S}, \mathrm{Se}, \mathrm{Te}\) b. \(\mathrm{Br}, \mathrm{Ni}, \mathrm{K}\) c. \(\mathrm{F}, \mathrm{Si}, \mathrm{Ba}\)

Step-by-step solution

a. Arrange \(\mathrm{Te}, \mathrm{S}, \mathrm{Se}\) in order of increasing size

We can find these elements in the same group (Group 16) of the periodic table. Since atomic size increases from top to bottom in a group, we can arrange them in the order of their position in the periodic table: S < Se < Te So the order of increasing size for these atoms is \(\mathrm{S}, \mathrm{Se}, \mathrm{Te}\).

b. Arrange \(\mathrm{K}, \mathrm{Br}, \mathrm{Ni}\) in order of increasing size

These elements belong to different groups but they are in the same period (Period 4). Since atomic size decreases from left to right in a period, we can arrange them in the order of their position in the periodic table: K > Ni > Br So the order of increasing size for these atoms is \(\mathrm{Br}, \mathrm{Ni}, \mathrm{K}\).

c. Arrange \(\mathrm{Ba}, \mathrm{Si}, \mathrm{F}\) in order of increasing size

These elements are neither in the same group nor in the same period. Therefore, we have to use both the group and period trends to determine their order. Ba is in Group 2 and Period 6. Si is in Group 14 and Period 3. F is in Group 17 and Period 2. Comparing Ba and Si: Ba has more energy levels than Si since it is below Si in the periodic table. Thus, Ba has a larger atomic size than Si. Comparing Si and F: Si is left of F in the periodic table within Period 3. Thus, Si has a larger atomic size than F. From the above comparisons, the order of increasing size for these atoms is \(\mathrm{F}, \mathrm{Si}, \mathrm{Ba}\).

Key concepts

Periodic Table Group Trends

When studying the periodic table, it's important to understand how atoms behave within the same group or column. A group consists of elements that have similar chemical properties because they have the same number of valence electrons.
However, one key trend to remember is that atomic size tends to increase as you move down a group in the periodic table.

• Each row (or period) down a group indicates a higher principal energy level or shell. This means additional shells are added, making the atomic radius larger.
• This results in atoms being larger at the bottom of the group than those at the top.

For example, in Group 16 with elements like sulfur (S), selenium (Se), and tellurium (Te), Te will be the largest because it is at the bottom of the group, making it occupy more space. This is why atomic size increases from S to Se to Te.

Periodic Table Period Trends

In contrast to group trends, period trends involve comparing elements across a period or row on the periodic table. When moving from left to right across a period, the atomic radius usually decreases.
This can be explained by the fact that elements across a period share the same number of shells, but more protons are added to the nucleus without adding new energy levels.

• The additional protons increase the nuclear charge, pulling electrons closer to the nucleus.
• This tighter pull decreases the size of the electron cloud, thereby reducing the atomic radius.

Let's take the elements potassium (K), nickel (Ni), and bromine (Br) from Period 4. As we move from K to Ni to Br, the atomic size decreases because the atoms become more tightly bound by the increased nuclear charge without adding new outer orbitals.

Atomic Radius

The atomic radius is a key concept for understanding atomic size. It is essentially the distance from the center of the nucleus to the outermost shell of an electron in an atom. Atomic radius can vary based on several factors:

• Number of energy levels: More levels mean a larger radius.
• Effective nuclear charge: Greater charge results in a smaller radius as electrons are pulled closer.
• Shrinkage across a period: With an increase in protons but not energy levels, radius shrinks.

For instance, when comparing elements in the same group, the added energy level leads to larger atomic radii. Conversely, across a period, although the energy levels remain the same, the atomic radius decreases due to a stronger effective nuclear charge.

Elements Comparison

Comparing atomic sizes of elements is about understanding both their position in a group and across a period. Some elements fall neither in the same group nor the same period, and such cases require careful analysis. Let's look at barium (Ba), silicon (Si), and fluorine (F) as examples.

• Barium is further down in Group 2 and hence has a larger atomic radius than Si and F due to more energy levels.
• Silicon and fluorine are both in different periods; Si is left of F in Period 3, making it larger because the atomic size decreases across a period.

In such comparisons, it helps to consider the number of energy levels first (group trend) and then think about the effective nuclear charge (period trend) for an informed decision on atomic size. These principles help predict the order of atomic sizes across diverse elements.