Question

Compare the elements \(\mathrm{Na}, \mathrm{Mg}, \mathrm{O},\) and \(\mathrm{P}.\) (a) Place the elements in order of increasing ionization energy. (b) Place the elements in order of increasing electron affinity.

Short answer

The elements in order of increasing ionization energy are: Na, Mg, P, O. The elements in order of increasing electron affinity are: Na, Mg, P, O.

Step-by-step solution

Understanding Periodic Table Trends

The positioning of elements in the periodic table profoundly affects their ionization energy and electron affinity. In general, as you move across a period (left to right), the ionization energy and electron affinity increases. This occurs because the number of protons in the nucleus increases, attracting the electrons more strongly and making it harder to remove an electron (higher ionization energy) and easier to accept an electron (higher electron affinity). But as you move down a group (top to bottom), both ionization energy and electron affinity decrease due to increases in atomic radius and electron shielding.

Ordering Elements by Increasing Ionization Energy

Arrange the elements in order of increasing atomic number: Na, Mg, O, P. Using the rule that ionization energy increases as you move from left to right across a period and from bottom to top in a group, the following sequence from lowest to highest ionization energy would be Na < Mg < P < O.

Ordering Elements by Increasing Electron Affinity

Again using the atomic number, the elements are arranged as follows: Na, Mg, O, and P. As noted, electron affinity typically increases as you move from left to right across a period. However, it's important to note that in some cases there can be exceptions. Therefore, using general trends, the expected sequence from lowest to highest electron affinity would be Na < Mg < P < O.

Key concepts

Periodic Table Trends

Understanding periodic table trends is crucial for predicting and comparing the chemical behavior of elements.
These trends arise due to the atomic structure and the periodic arrangement of elements.

• Moving across a period (left to right), the ionization energy and electron affinity generally increase. This is because more protons are present in the nucleus, enhancing the nuclear charge. This stronger pull makes it harder to remove an electron (higher ionization energy) and easier to add an electron (higher electron affinity).
• Conversely, moving down a group (top to bottom), these properties generally decrease. A larger atomic size reduces the pull on outer electrons due to increased distance from the nucleus (increased atomic radius) and greater electron shielding.

These trends help us predict the order of ionization energy and electron affinity among elements, as seen in the comparison of elements like sodium ( \( \mathrm{Na} \)), magnesium ( \( \mathrm{Mg} \)), oxygen ( \( \mathrm{O} \)), and phosphorus ( \( \mathrm{P} \)).

Electron Affinity

Electron affinity measures how easily an atom accepts an additional electron. It reflects an atom's tendency to gain electrons and form negative ions.
As you move across a period, electron affinity generally increases because atoms tend to fill their outer electron shells.

• Atoms like oxygen have higher electron affinity as they are closer to completing their valence shell.
• However, there are exceptions to this trend due to subshell configurations, as seen with elements like phosphorus.

Therefore, in some cases, even if a trend exists, it is useful to consider specific electron configurations to account for any exceptions.

Atomic Radius

The atomic radius is the distance from the nucleus to the outermost electrons and serves as a good indicator of the atom’s size.

• As you move across a period, atomic radius decreases due to increased nuclear charge, which draws the electron cloud closer to the nucleus.
• Conversely, moving down a group, the atomic radius increases significantly due to addition of electron shells.

This goes hand-in-hand with other properties such as ionization energy and electron affinity, as larger atoms have loosely held outer electrons, influencing their chemical behavior. For example, sodium ( \( \mathrm{Na} \)) has a larger atomic radius compared to magnesium ( \( \mathrm{Mg} \)), leading to differences in their physical and chemical properties.

Electron Shielding

Electron shielding, also known as the screening effect, influences how tightly the outer electrons are held by the nucleus.

• Core electrons shield the outer electrons from the full effect of the nucleus's positive charge.
• This effect is more pronounced as you move down a group due to the increased number of inner electron shells.

Reduced shielding across a period maintains higher nuclear attraction on the valence electrons. This is why ionization energy and electron affinity trends are linked to the underlying changes in electron shielding, atomic radius, and nuclear charge.