Problem 23 - Chemistry Central Science

Question

How do the sizes of atoms change as we move (a) from left to right across a row in the periodic table. (b) from top to bottom in a group in the periodic table? (c) Arrange the following atoms in order of increasing atomic radius \(\mathrm{F}, \mathrm{P}, \mathrm{S}\), As.

Step-by-Step Solution

Answer

(a) The atomic radius decreases from left to right across a row due to increased effective nuclear charge. (b) The atomic radius increases from top to bottom in a group due to increased number of energy levels and electron shielding. (c) The order of increasing atomic radius is: \(\mathrm{F} < \mathrm{S} < \mathrm{P} < \mathrm{As}\).

Step 1: a) Change in atomic size from left to right across a row

As we move from left to right in a row in the periodic table, the number of protons in the nucleus increases. This results in a higher effective nuclear charge which means the nucleus has a stronger pull on the electrons. The electron shielding does not significantly change as electrons are added to the same energy level (shell) while moving across a row. Consequently, the atomic radius (or atomic size) decreases from left to right across a row.

Step 2: b) Change in atomic size from top to bottom in a group

In contrast, as we move from top to bottom in a group, the number of energy levels (shells) increases. Each additional energy level is further away from the nucleus, which results in a larger atomic radius. Even though the number of protons increases as we move in a group down the periodic table, the strong shielding effect of the additional energy levels causes the atomic radius to increase rather than decrease from top to bottom in a group.

Step 3: c) Arrange \(\mathrm{F}, \mathrm{P}, \mathrm{S}\), As in order of increasing atomic radius

To arrange these elements in order of increasing atomic radius, we must take into account both the row and group trends mentioned earlier. Let's start by locating these atoms in the periodic table. F is in Group 17 and Period 2. P is in Group 15 and Period 3. S is in Group 16 and Period 3. As is in Group 15 and Period 4. Now, considering that the atomic size decreases as we move left to right across a row and increases as we move top to bottom in a group, we can arrange the given elements in order of increasing atomic radius: F (smallest atomic radius) S P As (largest atomic radius) So, the order of increasing atomic radius is: \(\mathrm{F} < \mathrm{S} < \mathrm{P} < \mathrm{As}\).

Key Concepts

Periodic Table

The periodic table is a marvel of organization that arranges the chemical elements in a structured way. It’s organized by increasing atomic number from left to right and top to bottom. This layout allows us to see patterns in the properties of the elements, such as their atomic radius, reactivity, and electronegativity. Each row in the periodic table is called a "period," and each column is known as a "group."

Periodic trends emerge as you progress across a period or down a group:

• Across a period: Elements become less metallic and more non-metallic.
• Down a group: Elements become more metallic.

This systematic arrangement helps predict an element's properties based on its position, granting insight into its behavior in chemical reactions and bonding.

Atomic Size

Atomic size, also known as atomic radius, refers to the distance between the nucleus of an atom and its outermost electrons. The size of atoms varies systematically across the periodic table due to changes in the number of protons and electron shells.

When you move from left to right across a row in the periodic table, an increase in protons results in a tighter hold on the electrons, which shrinks the atomic radius. However, as you go down a group, the addition of energy levels outweighs the pull of additional protons, leading to an increase in atomic size.
Key points to remember include:

• The effective nuclear charge increases across a period, decreasing atomic size.
• Additional energy levels increase the atomic size down a group.

Overall, the atomic size is influenced by the balance between these trends as you navigate the table.

Nuclear Charge

The nuclear charge of an atom is the total charge of all the protons in its nucleus. It's a crucial factor determining atomic size and other properties. As you move across a period in the periodic table, nuclear charge increases because more protons are added to the nucleus.

The increasing nuclear charge has significant effects:

• It pulls electrons closer, reducing atomic radius across a period.
• It generally increases ionization energy, making it harder to remove electrons.

Effective nuclear charge is a concept that also takes shielding into account, and it represents the net positive charge experienced by valence electrons. Understanding nuclear charge helps explain why atoms across the periodic table behave as they do in chemical reactions.

Electron Shielding

Electron shielding, or screening, occurs when inner electrons block the attractive force of the nucleus from the outermost electrons. As more inner electron shells are added, they create a barrier that diminishes the full charge of the nucleus, shielding the valence electrons.

This phenomenon plays a vital role in the trends of the periodic table:

• Across a period, shielding remains relatively constant because electrons are added to the same shell.
• Down a group, additional shells increase shielding, lessening the effective nuclear charge.

Impact on atomic size: Due to increased electron shielding as you move down a group, the outer electrons are farther from the nucleus, expanding the atomic radius. Understanding electron shielding is key to predicting how element properties change throughout the periodic table.