What a covalent bond is

A covalent bond is not just "sharing electrons". Sharing is part of the picture, but the bond exists because charged particles attract each other.

This definition has three required ingredients:

• a shared pair of electrons
• attraction to the nuclei of the bonded atoms
• the attraction is strong electrostatic attraction

The electrons in the shared pair are negatively charged. The nuclei are positively charged because they contain protons. The shared pair is attracted to both nuclei, so the two atoms are held together.

Full outer shells can be a useful checking pattern for many simple dot-and-cross diagrams, but "atoms want full shells" is not a definition of covalent bonding. The bond is an electrostatic attraction, not a desire or a rule that every species must obey.

Worked example: explaining the H-Cl bond

In hydrogen chloride, HCl, hydrogen and chlorine share one pair of electrons. That shared pair is attracted to the hydrogen nucleus and to the chlorine nucleus. The strong electrostatic attraction between the shared pair and the two nuclei is the covalent bond.

What dot-and-cross diagrams show

A dot-and-cross diagram shows the outer-shell electrons involved in bonding. Dots and crosses record which atom the electrons came from before bonding.

They do not mean there are two different types of electron. Once electrons are in a molecule, a dot electron and a cross electron are still just electrons. The notation is a bookkeeping tool.

Use these rules when constructing a dot-and-cross diagram:

1. Count the outer-shell electrons supplied by each atom.
2. If the species is an ion, adjust the total for the charge.
3. Place shared pairs between bonded atoms.
4. Place remaining outer electrons as lone pairs.
5. For an ion, put brackets around the whole diagram and show the charge.
6. Check the electron total and the number of pairs around the central atom.

The OCR boundary for this lesson includes dot-and-cross diagrams of molecules and ions with up to six electron pairs, including lone pairs, surrounding a central atom. That limit tells you the scale of diagram you may need to construct. The shapes and bond angles of those species are a separate lesson.

Worked example: hydrogen chloride, HCl

Hydrogen has 1 outer electron. Chlorine has 7 outer electrons.

• One electron from hydrogen and one electron from chlorine form one shared pair.
• Chlorine has three lone pairs left.
• Hydrogen has a pair in its first shell.

The dot-and-cross diagram should therefore show one bonding pair between H and Cl and three lone pairs on Cl.

Single and multiple covalent bonds

A single covalent bond contains one shared pair of electrons. A double covalent bond contains two shared pairs. A triple covalent bond contains three shared pairs.

Worked example 1: oxygen, O2

Each oxygen atom has 6 outer electrons. Together, the two oxygen atoms have:

In O2, the oxygen atoms share two pairs of electrons. That gives an O=O double bond. Each oxygen also has two lone pairs.

Check the electron total:

• two shared pairs contain 4 electrons
• four lone pairs in total contain 8 electrons
• total = 12 electrons

Worked example 2: nitrogen, N2

Each nitrogen atom has 5 outer electrons. Together, the two nitrogen atoms have:

In N2, the nitrogen atoms share three pairs of electrons. That gives an N≡N triple bond. Each nitrogen also has one lone pair.

Check the electron total:

• three shared pairs contain 6 electrons
• two lone pairs in total contain 4 electrons
• total = 10 electrons

Up to six pairs around a central atom

Some dot-and-cross diagrams have a central atom with more than four electron pairs around it. The lesson boundary allows up to six electron pairs around the central atom, counting both bonding pairs and lone pairs. For example, a diagram for SF6 would show sulfur in the centre with six S-F shared pairs and no lone pairs on sulfur.

Do not use this as a shape rule. At this stage you are only accounting for outer-shell electrons and shared pairs.

Dot-and-cross diagrams for ions

For a molecular ion, the dot-and-cross diagram must show the charge clearly. Put square brackets around the whole ion and write the charge outside the brackets.

The charge also changes the electron count:

• add one electron for each negative charge
• subtract one electron for each positive charge

Worked example: hydroxide ion, OH-

Oxygen has 6 outer electrons and hydrogen has 1. The 1- charge means one extra electron is present:

O and H share one pair of electrons. The remaining 6 electrons form three lone pairs on oxygen. The whole diagram is placed in brackets with a 1- charge.

Worked example: ammonium ion, NH4+

Nitrogen has 5 outer electrons. Four hydrogen atoms supply 4 more. The 1+ charge means one electron has been removed from the total:

Those 8 electrons form four shared pairs between nitrogen and hydrogen atoms. The whole NH4+ diagram is placed in brackets with a + charge.

Do not add a separate lone pair to nitrogen in NH4+. The electron count gives four pairs total around nitrogen, all used as bonding pairs.

Dative covalent bonding

A dative covalent bond is also called a coordinate bond.

The key point is electron origin. In an ordinary covalent bond, each bonded atom often contributes one electron to the shared pair. In a dative covalent bond, one atom supplies both electrons to the shared pair. After the bond has formed, it is still a covalent bond: a shared pair attracted to nuclei.

Worked example: NH3 forming NH4+

In ammonia, NH3, nitrogen has one lone pair. A hydrogen ion, H+, has no electrons to contribute.

Nitrogen donates both electrons from its lone pair to H+. This forms an N-H coordinate bond. In a dot-and-cross diagram of NH4+, that donated shared pair is shown with both electrons from nitrogen.

The overall ion is NH4+, so the whole dot-and-cross diagram is bracketed and labelled with a + charge.

Common mark-loss trap

Do not say that a coordinate bond is weaker just because both electrons came from one atom. The specification treats it as dative covalent bonding: it is a covalent bond, with a shared pair of electrons attracted to nuclei. The dot-and-cross diagram records where the electrons came from.

Average bond enthalpy and bond strength

Average bond enthalpy is used here as a measurement of covalent bond strength. The larger the average bond enthalpy, the stronger the covalent bond.

That is the whole requirement for this lesson. The detailed definition of average bond enthalpy and calculations using bond enthalpies are taught later in the energetics topic.

Worked example: comparing bond strength

Suppose a data table says:

Bond B has the larger average bond enthalpy, so bond B is the stronger covalent bond.

Do not calculate an enthalpy change from these values in this lesson. Do not write a full definition involving breaking bonds in gaseous molecules here. For now, use average bond enthalpy only to compare covalent bond strength.