Question

Arrange these elements in order of increasing first ionization energy: \(\mathrm{Si}, \mathrm{F}, \mathrm{In}, \mathrm{N} .\)

Short answer

In order of increasing first ionization energy: \text{In}, \text{Si}, \text{N}, \text{F}.

Step-by-step solution

Understand Ionization Energy

First ionization energy is the energy required to remove the most loosely bound electron from an atom in its gaseous state. It generally increases across a period (from left to right on the periodic table) and decreases down a group (from top to bottom on the periodic table).

Locate the Elements on the Periodic Table

Find the position of Si (silicon), F (fluorine), In (indium), and N (nitrogen) on the periodic table to understand their group and period trends. Silicon is in group 14 and period 3, Fluorine is in group 17 and period 2, Indium is in group 13 and period 5, and Nitrogen is in group 15 and period 2.

Apply Period and Group Trends

Considering the periodic trends, we know that elements in the same period increase in ionization energy as we move from left to right, and elements in the same group decrease in ionization energy as we move from top to bottom. Since Si and N are in the same period but N is to the left of Si, N has lower ionization energy than Si. In is below Si in the same group, so it has lower ionization energy than Si. F, being to the right of N and in the same period, has the highest ionization energy of all.

Arrange the Elements

According to the trends described, the elements in order of increasing first ionization energy are In, Si, N, F. Indium being the lowest since it's farthest to the left and bottom, followed by Silicon, then Nitrogen, and Fluorine having the highest ionization energy.

Key concepts

Ionization Energy Definition

Ionization energy is a fundamental concept in chemistry that refers to the amount of energy required to remove the most loosely bound electron from an isolated atom in the gaseous state. Essentially, it measures how strongly an atom holds onto its electrons. The first ionization energy is particularly significant because it deals with removing the first electron, which is typically easier to remove than the subsequent ones due to less electrostatic repulsion.
First ionization energy is often measured in electron volts (eV) or kilojoules per mole (kJ/mol) and varies widely across different elements due to their atomic structure. Understanding ionization energy is crucial because it plays a key role in predicting the reactivity of elements, the formation of ions, and the chemical bonding behavior of atoms.

Periodic Table Trends

The periodic table is an organized display of chemical elements, structured by their atomic number, electron configuration, and recurring chemical properties. Trends in the periodic table are patterns that emerge when we look at the changes in certain elemental properties, such as ionization energy, as we move across periods or down groups. These trends can be attributed to the effective nuclear charge and the distance of valence electrons from the nucleus.
For example, as we move from left to right across a period, the ionization energy generally increases. This is because the number of protons (and thus the nuclear charge) increases, which pulls the electrons closer, making them harder to remove. Conversely, as we move down a group, the ionization energy decreases due to the increasing distance of the outermost electrons from the nucleus, even though the nuclear charge increases.

Atomic Structure

The atomic structure of an element is central to understanding its chemical behavior. An atom consists of a nucleus containing protons and neutrons, surrounded by electrons in orbital shells. The arrangement of these electrons, particularly those furthest from the nucleus (valence electrons), plays a significant role in an element's ionization energy.
The energy required to remove an electron depends partly on how tightly it is held by the nucleus. This is influenced by the number of protons in the nucleus (atomic number) and the size of the atom itself. Electrons in closer shells experience a greater effective nuclear charge, making them harder to remove compared to those in larger atoms where electrons are farther from the nucleus and experience more electron shielding.

Electron Removal

Electron removal, or ionization, is the process of detaching an electron from an atom, creating a positively charged ion. The ease with which an electron can be removed from an atom is attributed to the element’s atomic structure and position on the periodic table. Various factors come into play when an electron is removed:

• The effective nuclear charge, which is the net positive charge experienced by an electron in a multi-electron atom.
• The distance between the nucleus and the electron, with electrons in higher energy levels being easier to remove.
• The electron shielding effect, where inner shell electrons shield outer electrons from the full charge of the nucleus.
• The exchange energy which stabilizes half-filled and fully-filled orbitals, making their electrons more difficult to remove.

By considering these factors, one can deduce the relative ease of electron removal and thus the ionization energy of an element. The knowledge of how readily an atom can lose an electron is essential for understanding chemical reactivity and the formation of compounds.