Chemical Bonding: Types, Properties, and Structures

Posted on Dec 8, 2024 in Chemistry

Unit 2: Chemical Bonding

Ionic Bond – Lattice Energy

Where K is Coulomb’s constant, Z represents the charges of the ions, e is the absolute value of the electron charge, N is Avogadro’s number, d is the internuclear distance, m is the Madelung constant, and ε is the Born exponent.

Born-Haber Cycle and Lattice Energy

The Born-Haber cycle is a thermodynamic cycle that analyzes all the processes involved in the formation of one mole of an ionic compound from its constituent elements in their most stable thermodynamic state.

Formula

Covalent Bond

a) Lewis Structure

Lewis proposed that atoms strive to achieve a noble gas electron configuration. To do this, they share, donate, or accept electrons. To study these compounds according to this theory, we need to know the number of valence electrons of each atom. We, therefore, determine its electron configuration and indicate the number of valence electrons.

Then, we represent the Lewis structure by placing each atom with its valence electrons. The number of bonds formed or transferred depends on how many electrons each element needs to obtain the noble gas configuration (8 electrons, typically).

b) Molecular Geometry According to VSEPR

To study covalent compounds using the Valence Shell Electron Pair Repulsion (VSEPR) theory, we analyze the number of bonding and non-bonding electron pairs around the central atom of a molecule, which determines the molecule’s geometry.

Number of electrons in the outermost shell (E)
Number of electrons that could be present (T)
Number of bonding electrons (R) = T – E
Number of non-bonding electrons = E – R

Having x pairs of bonding electrons and y pairs of non-bonding electrons results in a possible structure of type AX_xE_y. The possible arrangement will be XYZ with a bond angle of θ.

Hydrogen Bonding

A hydrogen bond is an attraction between a hydrogen atom (with a partial positive charge) and an atom of fluorine (F), oxygen (O), nitrogen (N), or a halogen (X) that has a pair of free electrons (with a partial negative charge). Water is an example of a substance that exhibits this type of intermolecular bonding. A water molecule is formed between an oxygen atom with six valence electrons (two are shared, leaving two lone pairs) and two hydrogen atoms, each with one valence electron (both will donate their single electron to oxygen to complete its octet).

Van der Waals Forces

Van der Waals forces are electrostatic forces that bind polar and nonpolar molecules.

Polar molecules: Existing dipoles can interact among themselves, producing weak bonds and modifying melting and boiling points.
Nonpolar molecules: The mobility of electron clouds can cause temporary electrical asymmetry, leading to the formation of instantaneous dipoles. These, in turn, can induce dipoles in nearby molecules, resulting in attractions called London dispersion forces.

Metallic Bond

Electron Sea Model

In the early twentieth century, a model was proposed based on the idea that metals are formed by an accumulation of positive ions (metal atoms that have lost their valence electrons) located at a distance that minimizes electronic repulsion. These positive ions are immersed in a “sea” of electrons formed by the valence electrons of the participating atoms.

The electrons do not belong to individual atoms but are shared among all the atoms that form the lattice; that is, they are delocalized. This explains their relatively easy movement through the metal structure, resulting in high conductivity and electron emission effects.

Furthermore, the displacement of ionic layers in the structure does not produce a significant variation, so they do not break easily.

Band Theory

This model assumes that in a compact metal lattice, the constituent atoms are very close to each other, so their valence orbitals overlap, giving rise to molecular orbitals with very similar energies. These energy levels are described as a band of energy levels due to their large number. The number of molecular orbitals formed equals the number of atoms aligned multiplied by the number of orbitals in each atom.

For example, in a crystal formed by N lithium atoms:

There are N completely filled, delocalized molecular orbitals resulting from the interaction of the 1s orbitals.
There are N half-filled molecular orbitals from the 2s orbitals.
There are 3N empty molecular orbitals from the 2p orbitals.

The molecular orbitals from the 2s and 2p orbitals overlap in energy. When a band is half-filled, electrons can move throughout the metal by applying an external electric field since electrons can easily transition to unoccupied molecular orbitals with small energy gaps within the band. The half-filled band is called the valence band. Electrons are placed at the bottom of the band, where the energy is lowest. Providing a minimum amount of energy to the band leads to quick and easy electron mobility, explaining the high conductivity of metals.

It often happens that the outermost molecular orbital band is completely full. The conductive properties in these cases depend on the energy difference (ΔE) between this band and the next completely empty band:

If the bands overlap, the crystal is a conductor.
If the energy difference between them is large, the crystal is an insulator.
If this difference is small, electrons with sufficiently high energy can transition from the full band to the empty band. The crystal is a semiconductor.