Covalent Bonding Fundamentals

In this section we consider the second of our three types of strong atomic bonds. Covalent bonding is what we call the strong attraction that holds together the atoms of nonmetallic elements. Covalent bonding is associated with a great variety of materials. It is found in elements and in compounds. It is found in networks and in molecules. It is also found within the polyatomic ions. Essentially it is found in any material in which nonmetallic atoms are bonded together.

In this section, we will look first at the basics of covalent bonding, specifically what happens to the electrons of the atoms involved in the bonds.

Sharing Electrons

However, what if we have two non-metal atoms placed together? For example, H and Cl. Despite being in group IA with sodium, hydrogen is a non-metal, and its preference is not to lose electrons but to gain them. If an electron is transferred from either of these atoms to the other, one atom will be "happy" but the other will be "unhappy." Anyone who has tried to care for small children knows that, if two children both want toys, taking a toy from one and giving it to the other is not a workable solution!

We can see above the symbolism we will use to represent this kind of sharing. As you can see, our starting point is the Lewis symbols for H and Cl. The electrons marked in red get shared between the atoms. Once they are shared, we stop representing them as separate dots and instead draw them as a single line. One of the most important things you must remember in this lesson is that a line like this represents two shared electrons.

Because the positive nuclei are both attracted to the negative electrons, they are pulled together, much more strongly and closely than they would normally come to each other. This creates a strong, durable link between them called a covalent bond. It is a bond in the same way an ionic bond is: an attractive force between atoms. And it is "covalent" because it is a sharing ("co-") of valence electrons ("-valent").

The electrons used to make the bond are called bonding electrons. In this case they are marked in red, though this will not usually be the case. The other electrons (the six others on Cl) are called non-bonding electrons. We will see that they almost always come in pairs; for this reason they are usually called "non-bonding pairs" or "lone pairs."

The Octet Rule

In the previous lesson, we applied the concept of the octet rule to ionic bonding, when we saw that for the representative elements (i.e. not the transition metals), atoms usually form ions in ways that give them eight valence electrons. This is why aluminum loses three electrons to get a +3 charge, and sulfur gains two electrons to get a -2 charge.

This rule can also be applied to covalent bonding. The octet rule says that atoms tend to gain, lose or share electrons so as to have eight electrons in their valence shell. It is a very useful rule but you should also know that there are many bonding situations where it does not apply. As you learn to use the octet rule, also learn to recognize situations where it does not apply and disregard it in those situations.

You can see the octet rule at work in the example above: chlorine has made a covalent bond, allowing it to increase its number of valence electrons from seven to eight (because both the electrons in the bond count toward its total). This satisfies the octet rule.

The Duet Rule

We'll see some exceptions to the octet rule later, but for now there is an important one that is worth mentioning early: hydrogen.

The octet rule for nonmetals is based on the fact that nonmetal atoms are driven to fill up their valence s and p subshells. However, hydrogen does not have p orbitals to fill. The first energy level contains only an s subshell. The maximum number of electrons it can fit in that energy level is 2.

So, hydrogen abides by the "duet rule" - it tends to gain or share electrons so as to have two electrons in its valence shell. We can see this rule in action above: while chlorine has eight valence electrons, hydrogen has two (both the electrons in the bond). This satisfies the octet and duet rules, respectively.

*Note: the duet rule technically also applies to the elements helium, lithium, and beryllium. However, helium is totally unreactive, and the duet rule isn't really needed to analyze the bonding of Li and Be. So the rule is functionally only about hydrogen.

Principles of Covalent Bonding

The example above shows an extremely simple case of covalent bonding: two atoms share a single pair of electrons, fulfilling the octet/duet rule for both; end of story. But covalent bonding is an incredibly diverse field that can produce an astounding variety of structures. Let's briefly develop some principles for what is and is not possible in covalently bonded systems.