Question

Make a simple sketch of the shape of the main part of the periodic table, as shown. (a) Ignoring \(\mathrm{H}\) and He, write a single straight arrow from the element with the smallest bonding atomic radius to the element with the largest. Ignoring \(\mathrm{H}\) and He, write a single straight arrow from the element with the smallest first ionization energy to the element with the largest. (c) What significant observation can you make from the arrows you drew in parts (a) and (b)? [Sections 7.3 and 7.4]

Short answer

Sketch the periodic table, marking the s, p, d, and f blocks. The atomic radius arrow goes from the top right to the bottom left, while the ionization energy arrow goes from the top left to the bottom right. The significant observation is that these arrows are diagonal with opposite directions, showing that atomic radius and ionization energy have opposite trends across the periodic table.

Step-by-step solution

Sketch the periodic table

Draw a simple sketch of the periodic table showing the s, p, d, and f blocks. The periodic table can be separated into these blocks based on their electron configuration.

Identify the elements with the smallest and largest atomic radius

Atomic radius generally decreases from left to right across a period and increases down a group. So, the element with the smallest atomic radius (ignoring H and He) would be on the top right of the periodic table, and the element with the largest atomic radius would be on the bottom left. Locate these elements and make a note of them.

Identify the elements with the smallest and largest ionization energy

Ionization energy generally increases from left to right across a period and decreases down a group. Therefore, the element with the smallest ionization energy would be on the top left of the periodic table (ignoring H and He) and the element with the largest would be on the bottom right. Locate these elements and make a note of them.

Draw arrows for atomic radius and ionization energy

On the sketch of the periodic table, draw a single straight arrow from the element with the smallest atomic radius to the element with the largest. Then, draw a single straight arrow from the element with the smallest ionization energy to the element with the largest.

Make a significant observation from the arrows

Observe the arrows drawn for atomic radius and ionization energy. The significant observation is that both arrows move diagonally across the periodic table. The atomic radius arrow moves from top right to bottom left, while the ionization energy arrow moves from top left to bottom right. This demonstrates that atomic radius and ionization energy exhibit opposite trends across the periodic table.

Key concepts

Atomic Radius

The atomic radius is a term that refers to the size of an atom. Specifically, it is the distance from the center of an atom's nucleus to the outer boundary of the surrounding electron cloud. Understanding the atomic radius is important because it helps us predict how atoms will interact in molecular structures and during chemical reactions.
Across the periodic table, the atomic radius trends in a predictable way. As you move from left to right across a period, the atomic radius decreases. This occurs due to the increased nuclear charge, meaning there are more protons attracting the electrons towards the nucleus, causing the atom to contract. Conversely, as you go down a group, the atomic radius increases. As new electron shells are added for each subsequent element, the outer electrons are found further from the nucleus, thus increasing the atom's size.
Key takeaways about atomic radius trends:

• Decreases across a period from left to right.
• Increases down a group as additional electron shells are added.
• Influenced by the balance between nuclear charge and electron shielding.

These concepts explain why the smallest atomic radii, excluding hydrogen and helium, are found in the top right of the periodic table, while the largest are found in the bottom left.

Ionization Energy

Ionization energy is the amount of energy required to remove an electron completely from an atom. This property is essential for understanding chemical reactivity and bonding since elements with high ionization energies are typically less likely to lose electrons and form positive ions.
When analyzing the periodic trends related to ionization energy, we find that as you progress from left to right across a period, ionization energy generally increases. Atoms hold their electrons more tightly due to the larger positive charge of the nucleus, which makes it more difficult to remove an electron.
Conversely, as you move down a group, the ionization energy decreases. Although the nuclear charge increases, the additional electron shells mean that the outer electrons feel less of the nucleus's pull due to shielding by inner-shell electrons.
Key takeaways about ionization energy trends:

• Increases across a period from left to right.
• Decreases down a group as the effective nuclear charge felt by outer electrons decreases due to shielding.
• Essential in predicting an element's ability to form positive ions.

This is why the elements with the largest ionization energies can typically be found on the top right of the periodic table (again, excluding hydrogen and helium), while those with the smallest are located on the bottom left.

Electron Configuration

Electron configuration describes the distribution of electrons in an atom's electron shells and subshells. It is fundamental in determining an element’s chemical behavior and properties. Each element has a unique electron configuration which can be depicted using various notations, such as the Aufbau principle, Hund's rule, and the Pauli exclusion principle.
Electronic configuration allows the categorization of elements into blocks on the periodic table:

• s-block: Groups 1 and 2, which include the alkali and alkaline earth metals.
• p-block: Groups 13 to 18, containing diverse elements including metals, metalloids, and nonmetals.
• d-block: Transition metals, characterized by partially filled d-orbitals.
• f-block: Lanthanides and actinides, often shown separated from the main body of the periodic table.

The configuration affects the atomic radius and ionization energy as well. For instance, elements with filled or half-filled subshells often have higher ionization energies due to their greater stability.
By understanding electron configurations, one can make informed predictions about an element’s place on the periodic table, its possible chemical bonds, and its behavior in different chemical environments.