Covalent Bonds and Bond Formation

Covalent bonds and bond formation

Covalent bonding involves the sharing of electrons to form a chemical bond. The outermost orbitals of the atoms overlap so that unpaired electrons in each of the bonding atoms can be shared. By overlapping orbitals, the outer energy shells of all the bonding atoms are filled. The shared electrons move in the orbitals around both atoms. As they move, there is an attraction between these negatively charged electrons and the positively charged nuclei. This attractive force holds the atoms together in a covalent bond.

Definition: Covalent bond

A form of chemical bond where pairs of electrons are shared between atoms.

Tip:

Covalent bonds are examples of interatomic forces.

We will look at a few simple cases to deduce some rules about covalent bonds.

Tip:

Remember that it is only the valence electrons that are involved in bonding, and so when diagrams are drawn to show what is happening during bonding, it is only these electrons that are shown. Dots or crosses represent electrons in different atoms.

Case 1: Two atoms that each have an unpaired electron

For this case we will look at hydrogen chloride and methane.

Example: Lewis Diagrams: Simple Molecules

Question

Represent hydrogen chloride (\(\text{HCl}\)) using a Lewis diagram.

Step 1: For each atom, determine the number of valence electrons in the atom, and represent these using dots and crosses.

The electron configuration of hydrogen is \(1\text{s}^{1}\) and the electron configuration for chlorine is \([\text{He}]2\text{s}^{2}2\text{p}^{5}\).The hydrogen atom has \(\text{1}\) valence electron and the chlorine atom has \(\text{7}\) valence electrons.

The Lewis diagrams for hydrogen and chlorine are:

Notice the single unpaired electron (highlighted in blue) on each atom. This does not mean this electron is different, we use highlighting here to help you see the unpaired electron.

Step 2: Arrange the electrons so that the outermost energy level of each atom is full.

Hydrogen chloride is represented below.

Notice how the two unpaired electrons (one from each atom) form the covalent bond.

The dot and cross in between the two atoms, represent the pair of electrons that are shared in the covalent bond. We can also show this bond using a single line:

Note how we still show the other electron pairs around chlorine.

From this we can conclude that any electron on its own will try to pair up with another electron. So in practise atoms that have at least one unpaired electron can form bonds with any other atom that also has an unpaired electron. This is not restricted to just two atoms.

Example: Lewis Diagrams: Simple Molecules

Question

Represent methane (\(\text{CH}_{4}\)) using a Lewis diagram

Step 1: For each atom, determine the number of valence electrons in the atom, and represent these using dots and crosses.

The electron configuration of hydrogen is \(1\text{s}^{1}\) and the electron configuration for carbon is \([\text{He}]2\text{s}^{2}2\text{p}^{2}\).Each hydrogen atom has \(\text{1}\) valence electron and the carbon atom has \(\text{4}\) valence electrons.

Remember that we said we can place unpaired electrons at any position (top, bottom, left, right) around the elements symbol.

Step 2: Arrange the electrons so that the outermost energy level of each atom is full.

The methane molecule is represented below.

Or:

Case 2: Atoms with lone pairs

We will use water as an example. Water is made up of one oxygen and two hydrogen atoms. Hydrogen has one unpaired electron. Oxygen has two unpaired electrons and two electron pairs. From what we learnt in the first examples we see that the unpaired electrons can pair up. But what happens to the two pairs? Can these form bonds?

Example: Lewis Diagrams: Simple Molecules

Question

Represent water (\(\text{H}_{2}\text{O}\)) using a Lewis diagram

Step 1: For each atom, determine the number of valence electrons in the atom, and represent these using dots and crosses.

The electron configuration of hydrogen is \(1\text{s}^{1}\) and the electron configuration for oxygen is \([\text{He}]2\text{s}^{2}2\text{p}^{4}\).Each hydrogen atom has \(\text{1}\) valence electron and the oxygen atom has \(\text{6}\) valence electrons.

Step 2: Arrange the electrons so that the outermost energy level of each atom is full.

The water molecule is represented below.

Tip:

Notice how in this example we wrote a \(\text{2}\) in front of the hydrogen? Instead of writing the Lewis diagram for hydrogen twice, we simply write it once and use the \(\text{2}\) in front of it to indicate that two hydrogens are needed for each oxygen.

And now we can answer the questions that we asked before the worked example. We see that oxygen forms two bonds, one with each hydrogen atom. Oxygen however keeps its electron pairs and does not share them. We can generalise this to any atom. If an atom has an electron pair it will normally not share that electron pair.

A lone pair is an unshared electron pair. A lone pair stays on the atom that it belongs to.

Tip:

A lone pair can be used to form a dative covalent bond.

In the example above the lone pairs on oxygen are highlighted in red. When we draw the bonding pairs using lines it is much easier to see the lone pairs on oxygen.

Case 3: Atoms with multiple bonds

We will use oxygen and hydrogen cyanide as examples.

Example: Lewis Diagrams: Molecules With Multiple Bonds

Question

Represent oxygen (\(\text{O}_{2}\)) using a Lewis diagram

Step 1: For each atom, determine the number of valence electrons that the atom has from its electron configuration.

The electron configuration of oxygen is \([\text{He}]2\text{s}^{2}2\text{p}^{4}\). Oxygen has \(\text{6}\) valence electrons.

Step 2: Arrange the electrons in the \(\text{O}_{2}\) molecule so that the outermost energy level in each atom is full.

The \(\text{O}_{2}\) molecule is represented below. Notice the two electron pairs between the two oxygen atoms (highlighted in blue). Because these two covalent bonds are between the same two atoms, this is a double bond.

Each oxygen atom uses its two unpaired electrons to form two bonds. This forms a double covalent bond (which is shown by a double line between the two oxygen atoms).

Example: Lewis Diagrams: Molecules With Multiple Bonds

Question

Represent hydrogen cyanide (\(\text{HCN}\)) using a Lewis diagram

Step 1: For each atom, determine the number of valence electrons that the atom has from its electron configuration.

The electron configuration of hydrogen is \(1\text{s}^{1}\), the electron configuration of nitrogen is \([\text{He}]2\text{s}^{2}2\text{p}^{3}\) and for carbon is \([\text{He}]2\text{s}^{2}2\text{p}^{2}\). Hydrogen has \(\text{1}\) valence electron, carbon has \(\text{4}\) valence electrons and nitrogen has \(\text{5}\) valence electrons.

Step 2: Arrange the electrons in the \(\text{HCN}\) molecule so that the outermost energy level in each atom is full.

The \(\text{HCN}\) molecule is represented below. Notice the three electron pairs (highlighted in red) between the nitrogen and carbon atom. Because these three covalent bonds are between the same two atoms, this is a triple bond.

As we have just seen carbon shares one electron with hydrogen and three with nitrogen. Nitrogen keeps its electron pair and shares its three unpaired electrons with carbon.

Case 4: Co-ordinate or dative covalent bonds

Definition: Dative covalent bond

This type of bond is a description of covalent bonding that occurs between two atoms in which both electrons shared in the bond come from the same atom.

A dative covalent bond is also known as a coordinate covalent bond. Earlier we said that atoms with a pair of electrons will normally not share that pair to form a bond. But now we will see how an electron pair can be used by atoms to form a covalent bond.

One example of a molecule that contains a dative covalent bond is the ammonium ion (\(\text{NH}_{4}^{+}\)) shown in the figure below. The hydrogen ion \(\text{H}^{+}\) does not contain any electrons, and therefore the electrons that are in the bond that forms between this ion and the nitrogen atom, come only from the nitrogen.

Notice that the hydrogen ion is charged and that this charge is shown on the ammonium ion using square brackets and a plus sign outside the square brackets.

We can also show this as:

Note that we do not use a line for the dative covalent bond.

Another example of this is the hydronium (hydroxonium) ion (\(\text{H}_{3}\text{O}^{+}\)).

Optional Video: PheT Simulation: Build a Molecule

PhET Simulation: Build a Molecule

To summarise what we have learnt:

Any electron on its own will try to pair up with another electron. So in theory atoms that have at least one unpaired electron can form bonds with any other atom that also has an unpaired electron. This is not restricted to just two atoms.
If an atom has an electron pair it will normally not share that pair to form a bond. This electron pair is known as a lone pair.
If an atom has more than one unpaired electron it can form multiple bonds to another atom. In this way double and triple bonds are formed.
A dative covalent bond can be formed between an atom with no electrons and an atom with a lone pair.

Covalent Bonds and Bond Formation