Question

Refer to the periodic table and predict which element in each of the following pairs has the higher ionization energy: (a) \(\mathrm{Mg}\) or \(\mathrm{Ca}\) (b) S or Se (c) Sn or Pb (d) N or P

Short answer

(a) Mg, (b) S, (c) Sn, (d) N.

Step-by-step solution

Understanding Ionization Energy

Ionization energy is the energy required to remove an electron from an atom in the gaseous state. Elements with higher ionization energy require more energy to remove an electron. Generally, ionization energy increases across a period (left to right) and decreases down a group (top to bottom) on the periodic table.

Compare Mg and Ca

The elements magnesium (Mg) and calcium (Ca) are located in the same group (Group 2) of the periodic table, with Mg located above Ca. As you move down a group, ionization energy decreases. Therefore, since Mg is above Ca, it has a higher ionization energy.

Compare S and Se

Sulfur (S) and selenium (Se) are located in the same group (Group 16), with S located above Se. Following the trend that ionization energy decreases down a group, sulfur has a higher ionization energy than selenium.

Compare Sn and Pb

Tin (Sn) and lead (Pb) are in Group 14, with Sn above Pb. As you move down a group, ionization energy decreases. Therefore, tin has a higher ionization energy than lead.

Compare N and P

Nitrogen (N) and phosphorus (P) are in Group 15, with N located above P. Ionization energy decreases down a group, so nitrogen has a higher ionization energy than phosphorus.

Key concepts

Periodic Table

The periodic table is a systematic arrangement of all known chemical elements in order of increasing atomic number. It helps us understand various properties of elements, including their ionization energy, by showing trends across periods and groups.
The table is organized into rows called periods and columns called groups, which allow us to predict how an element will behave chemically.

• Elements within the same period have increasing ionization energy, as they move from left to right. This is because more protons hold the electrons more tightly.
• Within the same group, ionization energy usually decreases from top to bottom, since outer electrons are further from the nucleus and are not held as tightly.

Recognizing these patterns in the periodic table simplifies predicting how elements compare in terms of removing their outermost electron.

Atomic Structure

Understanding atomic structure is crucial to grasp why ionization energy behaves the way it does. Every atom consists of a nucleus, with protons and neutrons, surrounded by electrons in shells.
These electrons fill energy levels at increasing distances from the nucleus.

• Electrons in the outermost shell are known as valence electrons and have the most influence on ionization energy.
• As you move across a period, even though the electrons are added, they enter the same energy level. The increasing nuclear charge from added protons attracts these electrons more strongly, increasing ionization energy.
• Electrons in inner shells shield outer electrons from the nucleus, which plays a role when moving down a group.

The position of electrons and their interactions with the nucleus and each other direct the element's ionization energy.

Ionization Trends

Ionization trends explain how and why it's harder or easier to remove an electron from an atom based on its position in the periodic table.
These trends help us predict and rationalize differences in ionization energy between elements.

• Increases across a period: As more protons are added, the attraction to the outermost electron increases.
• Decreases down a group: Electrons are further from the nucleus and more shielded by inner electrons, making it easier to remove an outer electron.
• Exceptions: Some transition metals and heavy p-block elements might not strictly follow these trends due to additional complexities in electron configuration.

Understanding these trends helps in explaining differences in reactivity and bonding potential among elements, as ionization energy directly impacts an element's ability to form compounds.