HS Chemistry - Chemical Bonding

Covalent Bonding

Overview of The Page

This page will cover:

• How does covalent bonding work?
• What are coordinate covalent bonds?
• What are violations of the octet rule?

There are three types of bonding between atoms: ionic bonding, covalent bonding, and metallic bonding. This page looks at covalent bonding.

Covalent bonding occurs between two non-metal atoms, and involves the sharing of electrons so that the atoms can achieve a full valence shell (the octet rule). The shared pair is called the bonding pair.

In this diagram, the area covered in blue lines represents the area where the Hydrogen atom's nucleus will attract a proton. The electron isn't stuck in an orbit; it moves all around in the subshell, shown easily here because Hydrogen only has one subshell. Here, two Hydrogen atoms are sharing their electrons in order to complete their valence shell.

The pair in the middle of the overlapping area is the bonding pair. The bonding pair is in the area where both atoms' nuclei are pulling on them in order to try and complete their shell. The Hydrogen atoms aren't sharing them so much as they're trying to pull them away in order to complete their outer shell. Each nucleus is pulling on both of the electrons. Sometimes, one of the electrons will be attracted to one of the nuclei, while the other one is attracted to the other nuclei. Sometimes, they'll both be attracted to one nuclei, and sometimes they'll both be attracted to the other nuclei. And sometimes they'll be in the overlap area, between both atoms.

The electrons are free to move anywhere in the combined area of both atoms' orbitals, and the atoms, continually fighting over the electrons and attracting the electrons towards themselves, remain bound together. Not only do the nuclei attract the electrons, but they also repel each other (as they are both positively charged), resulting in a fixed molecule shape, as the atoms don't move any closer or any farther from each other than they currently are.

A combined area of attraction surrounding the nuclei of both atoms is formed. This area is called the molecular orbital between these atoms, and it applies to any two atoms bonding together.

The boundaries of the electron shells have disappeared. They no longer exist - both electrons in the bond are free to move in the combined area, called the molecular orbital, outlined in red here.

The final shape of the hydrogen molecule ends up being:

with the electrons being somewhere in the molecular orbital. To make the notation $note: this isn't a formal notation$ easier, we can just draw:

The red oval represents the molecular orbital. The atoms have bonded together in a covalent bond. Note that the nuclei are not free-floating - they remain in fixed positions relative to each other, repelling each other.

Coordinate Covalent Bonds

Most covalent bonds are like this, where both atoms fight over the shared electrons to complete their outer shell. However, coordinate covalent bonds are bonds where both electrons are donated by one atom. Functionally, they otherwise work the same; the atom with less than 8 electrons will attempt to steal some from the atom that does have them, while the atom that has 8 electrons will attempt to keep them, and thus they will remain bonded. Coordinate covalent bonds only form when the atoms/molecules bonding together have opposite dipoles $dipoles are covered in [Polarity](6-Polarity.md)$. An example of coordinate covalent bonding is NH₄⁺, shown below:

In the diagram, the red dots are electrons that have been donated from Hydrogen atoms towards the covalent bond between them and Nitrogen. The fourth bond, between NH₃ and the Hydrogen cation, shows that none of the electrons in the bond came from the Hydrogen cation; they both came from the Nitrogen atom.

Since Nitrogen has a higher electronegativity than Hydrogen, the Nitrogen atom in the NH₃ molecule had a slightly negative charge. The Hydrogen cation, which had a positive charge, was attracted towards it. It then bonded with the NH₃ molecule, which had a spare electron pair for the coordinate covalent bond, to form NH₄⁺.

Most of the time, non-metals bond covalently to complete their valence shells $octet rule$. However, there are circumstances where this rule is violated.

Violations of the Octet Rule

Sometimes, an atom will form a molecule where it does not have 8 electrons in its valence shell. There are three situations in which this happens:

• Odd-electron molecules: There are a few stable compounds in which there are an odd number of electrons in the valence shells. This means that at least one atom in the molecule doesn't have 8 electrons in their valence shell. For example, in NO $N=O$, two atoms are shared between Nitrogen and Oxygen, and while Oxygen gets 8 valence electrons, Nitrogen only has 7. Other stable odd-electron molecules are NO₂ and ClO₂.

• Electron-deficient molecules: There are a few stable compounds in which one atom in the molecule has less than 8 valence electrons. For example, Beryllium $Be$, with only 2 valence electrons, can form 2 covalent bonds to have up to four valence electrons. Yet BeCl₂ is a stable compound. Boron can also form electron-deficient molecules.

• Expanded Valence Shell Molecules: There are stable compounds in which one atom has more than 8 electrons in its valence shell. This happens in atoms which have 3 or more shells, as they have empty d orbitals that can also be used for bonding. For example, in the molecule XeF₄, Xenon has 12 valence electrons.