Lesson Explainer: Covalent Bonding Chemistry • Second Year of Secondary School

In this explainer, we will learn how to describe covalent bonding in terms of the electrostatic attraction between atomic nuclei and pairs of shared electrons.

Covalent compounds are molecules that contain covalent bonds. The oceans are full of covalently bonded water molecules that are about 0.3 nm wide, and the atmosphere is mostly made up of covalent molecules that are no bigger than 0.5 nm. Small covalent compounds can be found all over the world, and there must be a rule that explains why nonmetal atoms tend to form covalent compounds so often. Scientists usually use the octet rule to explain why nonmetal atoms form covalently bonded compounds. They use the octet rule to explain why nonmetal atoms form small and stable covalent compounds so frequently.

The octet rule is an incredibly simple scientific hypothesis. It states that atoms tend to share or transfer electrons because this helps them get eight valence electrons and the same electron configuration as a noble gas atom. Many nonmetal atoms can effectively attain the same electron configuration as a noble gas atom if they share one or more electrons with other nonmetal atoms. This statement could be reworded to say that nonmetal atoms can gain the same electron configuration as a noble gas atom if they make one or more covalent bonds, because covalent bonds are nothing more than shared pairs of electrons. The following figure shows the electron configuration of the three noble gas atoms that are the basis for the octet rule. The noble gas atoms have a very stable electron configuration because they have eight valence electrons.

The following figure shows that a single oxygen atom can attain the same electron configuration as a noble gas atom if it makes single covalent bonds with two hydrogen atoms. The figure shows that oxygen atoms have six valence electrons when they are not bonded and eight valence electrons when they make covalent bonds with two hydrogen atoms. Oxygen atoms effectively end up with a more stable electron configuration if they make covalent bonds with two hydrogen atoms. Oxygen atoms effectively end up with the same electron configuration as neon if they make covalent bonds with two hydrogen atoms.

Covalent bonds are stable due to a combination of attractive and repulsive electrostatic forces between the protons and electrons of covalently bonded atoms. There is electrostatic repulsion between particles that have the same charge sign and electrostatic attraction between particles that have opposite charge signs. The attractive interactions pull the atoms close together and the repulsive interactions drive the atoms apart. The atoms end up being at a stable distance where the attractive and repulsive electrostatic forces balance each other out.

Lewis structure diagrams are simple schematic illustrations that show how valence shell electrons are shared between atoms in covalently or ionically bonded compounds. Each valence shell electron is represented as a single small dot or cross and the atomic nuclei with core electrons is represented with chemical symbols, such as or .

Hydrogen gas () has one of the simplest Lewis structure diagrams, because hydrogen molecules only contain two hydrogen atoms () and each hydrogen atom has a single valence shell electron. The following figure shows a hydrogen molecule with its associated Lewis structure diagram.

One of the valence electrons is represented as a red dot and the other valence electron is represented as a black dot. The contrasting red and black colors are used to help readers differentiate between the valence electron of one hydrogen atom and the valence electron of the other hydrogen atom. Contrasting red and black dots will be used throughout this explainer to help readers differentiate between the valence electrons of covalently bonded atoms.

Lewis structure diagrams can also be drawn so that covalent bonds are represented as single straight lines. The following figure compares Lewis structure diagrams for hydrogen () and ammonia () molecules. The hydrogen molecule is drawn with one straight line because it contains a single hydrogen–hydrogen () covalent bond. The ammonia molecule is drawn with three straight lines because it contains three nitrogen–hydrogen () covalent bonds. Each straight line is used to represent a single covalent bond. You should notice here that the unbonded valence electrons on the nitrogen atom are drawn as a single pair of small dots. An unbonded pair of valence electrons is called a lone pair of electrons.

The following figures show the Lewis structure diagrams for two other common simple molecular compounds. The first figure shows Lewis structure diagrams for water () molecules and the second figure shows the Lewis structure diagrams for methane () molecules. The Lewis structure diagrams show that oxygen and carbon atoms can get the same electron configuration as a noble gas atom if they make the right number of covalent bonds with hydrogen atoms. Oxygen atoms can get the same electron configuration as a noble gas atom if they make two oxygen–hydrogen () bonds. Carbon atoms can get the same electron configuration as a noble gas atom if they make four carbon–hydrogen () bonds.

Some atoms form so-called double covalent bonds when they share two valence electrons. The following figure shows a double covalent bond that is formed between two oxygen atoms. The double bond contains a total of four valence electrons. Two of the electrons come from one oxygen atom and the other two electrons come from the other oxygen atom. The formation of the double covalent bond can be explained by the octet rule. The oxygen atoms have six valence electrons when they are not bonded and they have eight valence electrons when they make a double bond with another oxygen atom. Oxygen atoms gain the same electron configuration as a neon atom when they make one double bond with another oxygen atom.

It should not be too surprising that some atoms can also share three valence electrons with other atoms to gain the same electron configuration as a noble gas atom. The formation of double and triple covalent bonds helps atoms to get eight valence electrons and achieve a more stable electron configuration. The following figure shows how one nitrogen atom can share three of its valence electrons with another nitrogen atom to make a triple-bonded nitrogen () molecule.

Lewis structure diagrams can be used to represent double- and triple-bonded covalent molecules. The double bonds are represented as four dots or two straight lines between chemical symbols. The triple bonds are represented as six dots or three straight lines between chemical symbols. The following figure shows the Lewis structure diagrams for some simple molecular compounds that contain double and triple covalent bonds. The first Lewis structure diagram is for oxygen () molecules and the next Lewis structure diagrams are for carbon dioxide () and nitrogen () molecules.

The Lewis structure diagram for the oxygen molecule has four dots or two straight lines between two oxygen chemical symbols (). The dots and line representation methods both show that the oxygen–oxygen () bond is made up of four valence electrons. The Lewis structure diagrams for the carbon dioxide and nitrogen molecules are slightly different. The Lewis structure diagram for carbon dioxide has four dots or two straight lines between carbon () and each oxygen () chemical symbol. The Lewis structure for the nitrogen molecule has three lines between two nitrogen () chemical symbols. The Lewis structure diagrams show that the nitrogen–nitrogen bond () in molecules contains six electrons and the carbon–oxygen () bonds in molecules contain four electrons each.

The previous paragraphs have shown how atoms can combine together to make different types of covalent bonds. It was shown that some atoms can make single covalent bonds when they share an electron and that other atoms can make double covalent bonds when they share a pair of electrons. It was also shown that some atoms can make triple covalent bonds when they share three electrons. The following table recapitulates this information. It shows how atoms form different types of covalent bonds by sharing different numbers of electrons. You should notice here that the single covalent bonds are represented as one straight line between chemical symbols. The double covalent bonds are represented as two straight lines between chemical symbols, and the triple covalent bond is represented as three straight lines between chemical symbols.

The octet rule is not a faultless scientific hypothesis. There are lots of atoms that do not seem to obey the octet rule. Lots of atoms do not gain the same electron configuration as a noble gas atom if they make covalent bonds and form covalently bonded compounds. Lots of atoms do not share valence electrons simply because it helps them to get eight electrons and the same electron configuration as a noble gas atom.

Boron trifluoride is a relatively small covalently bonded molecule that seems to contradict the octet rule. The molecule has the chemical formula, and its Lewis structure is shown in the following figure. We can see that boron in the molecule has six valence electrons rather than the eight that are predicted by the octet rule. The boron atom does not share valence electrons simply because it helps it to get the same electron configuration as a noble gas atom. The boron atom is sharing electrons for different reasons that are not covered in this explainer.

There are other covalently bonded molecules that seem to contradict the octet rule. Phosphorus pentachloride is a covalently bonded compound that contains one phosphorus atom that is covalently bonded to five chlorine atoms. The molecule has the chemical formula, and its Lewis structure is displayed in the following figure. We can see that the phosphorus in the molecule has ten electrons rather than the eight that are predicted by the octet rule. The phosphorus atom does not share electrons simply because it helps it to gain the same electron configuration as a noble gas atom.

The boron trifluoride and phosphorus pentachloride molecules show that nonmetal atoms do not always gain the same electron configuration as a noble gas atom when they form simple molecular compounds. Some nonmetal atoms end up with more than eight valence electrons when they form covalently bonded compounds, and some atoms end up with fewer than eight valence electrons when they form covalently bonded compounds. The octet rule clearly has its limitations, and it cannot be used to understand the bonding of all covalently bonded molecules. It is also limited in that it cannot be used to predict the geometry of a covalent compound. We usually use the valence shell electron pair repulsion (VSEPR) theory to predict molecular geometries.

There are lots of simple molecular compounds that are formed by covalent bonding and they all tend to have similar physical properties. Most simple covalent compounds have relatively low melting points and boiling points because there are not very strong intermolecular interactions between neighboring covalently bonded molecules. The covalent bond itself is very strong and difficult to break, but the intermolecular interactions between covalent molecules are very weak and they are easily overcome with a relatively small amount of energy. The following table shows the common physical properties of simple covalent molecular compounds.

The table indicates that covalent compounds are not effective conductors of heat or electricity. The low electrical conductivity of covalent compounds can be understood by thinking about mobile charge-carrying particles. It is known that materials will only conduct electricity if they contain some type of mobile charge-carrying particle. The charge-carrying particles conduct an electrical current and they make a material an effective conductor of electricity. Covalently bonded molecules do not generate the charge-carrying particles that can conduct an electric current, and simple covalent compounds end up having low electrical conductivity values.

Small covalently bonded molecules cannot conduct heat very well because there is a lot of space between neighboring covalent molecules and the molecules do not collide very often. This makes it difficult to transfer heat, because heat conduction happens when molecules collide with each other. The following figure shows how heat is transferred between two molecules when they collide with each other. One of the molecules is hot, while the other is cold (step 1). The molecules then collide, and this makes thermal energy pass from the hot molecule to the cold molecule (step 2). The cold molecule gains thermal energy, and the other molecule loses thermal energy. The two molecules end up with similar amounts of thermal energy after the collision has happened (step 3).

Let us summarize what we have learned in this explainer.