5.1: Covalent Bonds

Learning Objectives

• To describe how a covalent bond forms.
• To apply the octet rule or duet rule to covalent compounds.

Covalent and Ionic Compounds

Previously, we discussed ionic bonding where electrons can be transferred from one atom to another so that both atoms have an energy-stable outer electron shell. Because most filled electron shells have eight electrons in them, chemists call this tendency the octet rule. As a general rule of thumb, compounds that involve a metal atom and at least a non-metal or a semi-metal atom will display ionic bonding. Thus, the compound formed from sodium and chlorine will be ionic (a metal and a non-metal). Nitrogen monoxide (NO) will be a covalently bound molecule (two non-metals), silicon dioxide (SiO₂) will be a covalently bound molecule (a semi-metal and a non-metal) and MgCl₂ will be ionic (a metal and a non-metal).

However, there is another way an atom can achieve a full valence shell: atoms can share electrons. Water is not an ionic compound; it is not composed of a metal and a nonmetal. Consequently, its properties are different from those of ionic compounds.

When two or more nonmetals combine, they form covalent bonds. For example, both hydrogen and oxygen are nonmetals, and when they combine to make water, they do so by forming covalent bonds. Compounds that are composed of only non-metals or semi-metals with non-metals will display covalent bonding and will be classified as molecular compounds.

Electron Sharing

This concept can be illustrated by using two hydrogen atoms, each of which has a single electron in its valence shell. (For small atoms such as hydrogen atoms, the valence shell will be the first shell, which holds only two electrons.) We can represent the two individual hydrogen atoms as follows:

Figure \(\PageIndex{1}\): Individual Hydrogen Atoms

In contrast, when two hydrogen atoms get close enough together to share their electrons, they can be represented as follows:

Figure \(\PageIndex{2}\): H₂ Molecule

By sharing their valence electrons, both hydrogen atoms now have two electrons in their respective valence shells. Because each valence shell is now filled, this arrangement is more stable than when the two atoms are separate. The sharing of electrons between atoms is called a covalent bond, and the two electrons that join atoms in a covalent bond are called a bonding pair of electrons. A discrete group of atoms connected by covalent bonds is called a molecule.

Chemists frequently use electron dot diagrams to represent covalent bonding in molecular substances. For example, the electron dot diagram of two separate hydrogen atoms are as follows:

The electron dot diagram of two hydrogen atoms sharing electrons looks like this:

This depiction of molecules is simplified further by using a dash to represent a covalent bond. The hydrogen molecule is then represented as follows:

Remember that the dash, also referred to as a single bond, represents a pair of electrons.

The bond in a hydrogen molecule, measured as the distance between the two nuclei, is about 7.4 × 10⁻¹¹ m, or 74 picometers (pm; 1 pm = 1 × 10⁻¹² m). This particular bond length represents a balance between several forces: the attractions between oppositely charged electrons and nuclei, the repulsion between two negatively charged electrons, and the repulsion between two positively charged nuclei. If the nuclei were closer together, they would repel each other more strongly; if the nuclei were farther apart, there would be less attraction between the positive and negative particles.

Fluorine is another element whose atoms bond together in pairs to form diatomic (two-atom) molecules. Two separate fluorine atoms have the following electron dot diagrams:

Each fluorine atom contributes one valence electron, making a single bond and giving each atom a complete valence shell, which fulfills the octet rule:

Two F's with two dots in between them, and two dots on the top, left/right side, and bottom of each F.

The circles show that each fluorine atom has eight electrons around it. As with hydrogen, we can represent the fluorine molecule with a dash in place of the bonding electrons:

Two F's with a long dash in between them, and two dots on the top, left/right side, and bottom of each F.

Each fluorine atom has six electrons, or three pairs of electrons, that are not participating in the covalent bond. Rather than being shared, they are considered to belong to a single atom. These are called nonbonding pairs (or lone pairs) of electrons.

Covalent Bonds between Different Atoms

Now that we have looked at electron sharing between atoms of the same element, let us look at covalent bond formation between atoms of different elements. Consider a molecule composed of one hydrogen atom and one fluorine atom:

Each atom needs one additional electron to complete its valence shell. By each contributing one electron, they make the following molecule:

An H and F connected by two dots, and the F has 2 dots on the top, right, and bottom. It also shows the same combination but with a long dash in between the H and F rather than two dots.

In this molecule, the hydrogen atom does not have nonbonding electrons, while the fluorine atom has six nonbonding electrons (three lone electron pairs). The circles show how the valence electron shells are filled for both atoms.

Example \(\PageIndex{1}\)

Use Lewis diagrams to indicate the formation of the following:

1. Cl₂
2. HBr

Solution

a. When two chlorine atoms form a chlorine molecule, they share one pair of electrons. In Cl₂ molecule, each chlorine atom is surrounded by an octet number of electrons.

The Lewis diagram for a Cl₂ molecule is similar to the one for F₂ (shown above).

Cl with 7 dots surrounding it plus Cl with 7 dots surrounding it means they are chlorine atoms. To turn into a chlorine molecule, have two dots in between each Cl and 6 dots surrounding each Cl on the sides that aren't connected.

b. When a hydrogen atom and a bromine atom form HBr, they share one pair of electrons. In the HBr molecule, H achieves a full valence of two electrons (duet) while Br achieves an octet. The Lewis diagram for HBr is similar to that for HF shown above.

H with one dot plus Br with 7 dots turns into H connected to Br by two dots, with 6 other dots surrounding Br.

Exercise \(\PageIndex{1}\)

Draw the Lewis diagram for each compound.

1. a molecule composed of one chlorine atom and one fluorine atom
2. a molecule composed of one hydrogen atom and one iodine atom

Answer a:

a- b-

Covalent Bonds in Larger Molecules

The formation of a water molecule from two hydrogen atoms and an oxygen atom can be illustrated using Lewis dot symbols (shown below).

Figure \(\PageIndex{3}\): H with one dot + O with 6 dots + H with one dot turns into H connected to O with two dots connected to H with two dots. O also has two dots on top and top dots on bottom.

The structure shown below is the electron dot structure for \(\ce{H2O}\). With two bonding pairs and two lone pairs, the oxygen atom has now completed its octet. Moreover, by sharing a bonding pair with oxygen, each hydrogen atom now has a full valence shell of two electrons. Chemists usually indicate a bonding pair by a single line, as shown (below).

or

Other large molecules are constructed in a similar fashion, with some atoms participating in more than one covalent bond. For example, methane (\(\ce{CH4}\)), the central carbon atom bonded to four hydrogen atoms, can be represented using either of the Lewis structures below. Again, sharing electrons between C and H atoms results in C achieving an octet while H achieves a duet number of electrons.

Figure \(\PageIndex{4}\): shows the number of covalent bonds various atoms typically form.

How Many Covalent Bonds Are Formed?

The number of bonds formed by an element is the same as the number of electrons missing to complete the octet (eight electrons for most atoms, but for H, it is only two). This is summarized in Table \(\PageIndex{1}\). In each case, the sum of the number of bonds and the number of lone pairs is 4, which is equivalent to eight (octet) electrons.

Because hydrogen only needs two electrons to fill its valence shell, it follows the duet rule. It is an exception to the octet rule. Hydrogen only needs to form one bond. This is the reason why H is always a terminal atom and never a central atom.

Table \(\PageIndex{1}\): lists atoms and their corresponding group numbers, as well as the number of bonds and lone pairs each has.

Atom (Group number)	Valence Electrons	Number of Bonds	Number of Lone Pairs of e-
Hydrogen (Group 1A)	1	1	0
Carbon (Group 4A)	4	4	0
Nitrogen (Group 5A)	5	3	1 (2 electrons are not part of any bond = 1 pair)
Oxygen (Group 6A)	6	2	2
Fluorine (Group 7A)	7	1	3

For example, each atom of a group 4A element has four electrons in its outermost shell and therefore requires four more electrons to reach an octet. These four electrons can be gained by forming four covalent bonds, as illustrated here for carbon in CH₄ (methane Figure \(\PageIndex{4}\)).

Group 5A elements such as nitrogen have five valence electrons in the atomic Lewis symbol: one lone pair and three unpaired electrons. To obtain an octet, these atoms form three covalent bonds, as in NH₃ (ammonia).
Oxygen and other atoms in group 6A obtain an octet by forming two covalent bonds. Fluorine and the other halogens in group 7A have seven valence electrons and can obtain an octet by forming one covalent bond.

Transition elements and inner transition elements do not follow the octet rule because they have d and f electrons involved in their valence shells.

Example \(\PageIndex{2}\)

Examine the Lewis structure of OF₂ below. Count the number of bonds formed by each element. Based on the element's location in the periodic table, does it correspond to the expected number of bonds shown in Table 4.1? Does the Lewis structure below follow the octet rule?

F surrounded by 6 dots, connected to O with a long dash. O is surrounded by 4 dots and connects to another F with a long dash. F surrounded by 6 dots.

Solution

Yes. F (group 7A) forms one bond, and O (group 6A) forms 2 bonds. Each atom is surrounded by 8 electrons. This structure satisfies the octet rule.

Exercise \(\PageIndex{2}\)

Examine the Lewis structure of NCl₃ below. Count the number of bonds formed by each element. Based on the element's location in the periodic table, does it correspond to the expected number of bonds shown in Table 4.1? Does the Lewis structure below follow the octet rule?

Answer

Both Cl and N form the expected number of bonds. Cl (group 7A) has one bond and 3 lone pairs. The central atom N (group 5A) has 3 bonds and one lone pair. Yes, the Lewis structure of NCl₃ follows the octet rule.

Double and Triple Bonds

In covalent bonding, using double (four total electrons) or triple (six total electrons) bonds is an option to achieve a stable octet. If a single bond does not work, then you may incorporate a multiple bond. Please note the rules above on which atoms will never achieve octets (hydrogen and boron). Although it is mathematically feasible, halogens will not use multiple bonds to become stable either.

Let's look at a few molecules that utilize multiple bonds:

In case of carbon dioxide, more than one pair of electrons must be shared between two atoms for both atoms to have an octet.

In this arrangement, the carbon atom shares four electrons (two pairs) with the oxygen atom on the left and four electrons with the oxygen atom on the right. There are now eight electrons around each atom. Two pairs of electrons shared between two atoms make a double bond between the atoms, which is represented by a double dash:

Atoms that can form multiple bonds include oxygen, nitrogen, sulfur, phosphorus, and carbon. If an octet can be achieved with a single bond, do not incorporate multiple bonds into the structure.

Triple bonds are covalent bonds in which three pairs of electrons are shared by two atoms. A compound that has a triple bond is acetylene (C₂H₂), whose Lewis diagram is as follows:

Multiple bonds involve more than one type of connection. For this reason, it is more difficult to break or cleave a multiple bond than it would be a single bond or a bond of a lower order (when comparing the same connected atoms). For example, the triple bond found in carbon monoxide is harder to break than one of the double bonds found in carbon dioxide. When reading about smog production or ozone depletion, consider the bond strengths of the chemicals involved in these environmental problems.

Example \(\PageIndex{3}\)

Before drawing the structures below, access the Furman University CHM101 moodle documents, Here, you will see a lightboard video of your instructor drawing Lewis structures for a variety of covalent compounds. Then, try to construct structures of the other molecules in this example.

1. O₃ (ozone)
2. N₂ (nitrogen)
3. SO₃ (sulfur trioxide)
4. CCl₄ (carbon tetrachloride)

Solutions

Example \(\PageIndex{4}\)

The Lewis diagram for formaldehyde \(\ce{CH2O}\) is shown below. The carbon atom is the central atom. One application of formaldehyde, is the preservation of biological specimens. Aqueous solutions of CH₂O are called formalin and have a sharp, characteristic (pungent) odor.

How many double bonds, single bonds, and lone pairs of electrons are present in the molecule of formaldehyde? Are the octet/duet rules met for each atom?

Solution

There is one double bond between C and O atom, two single bonds between C and H atoms, and two lone pairs on the oxygen atom. The octet for the oxygen and carbon are complete. The two hydrogen atoms complete their duet.

Exercise \(\PageIndex{4}\)

Draw the Lewis diagram for each molecule.

1. \(\ce{O2}\)
2. \(\ce{C2H4}\)

Answer a:

or

Answer b:

or or

Example \(\PageIndex{5}\)

Which is the correct Lewis structure for N₂H₂?

A. B. C.

Solution

Lewis structure A is the correct answer. Both N atoms satisfy the octet requirement, and the H atoms follow the duet rule.

Structure B is electron-deficient. It has only 10 valence electrons, instead of the usual 12.

Structure C has 14 (2 extra) electrons. The N atoms do not satisfy the octet.

Example \(\PageIndex{6}\)

Which is the correct Lewis structure for NOCl?

A. B. C.

Answer

Structure A violates the octet rule; N is surrounded by only 6e^-.

Structure B violates the octet rule; Cl has 10e^- around it.

Structure C is the correct structure. It has a total of 6e^- + 5e^- + 7e^- = 18e^-. Each atom is surrounded by 8 electrons (octet rule).

Rules for Drawing Covalent Lewis Structures

These apply to either nonmetal/nonmetal or metalloid/nonmetal combinations.

1. Tally up all valence electrons that are to be used. Remember, valence electrons are the roman numerals above the element.
2. Determine which atom is to be central. Carbon is always central, and Hydrogen is a terminal atom. If these atoms are not present, then go with the atom you have least of in the formula.
3. Connect each atom to the central atom using single bonds first.
4. For simple molecules or ions (only have two atoms), try single bonds, then double, and finally triple to achieve octets around each atom (except for B, Be, and H)
5. For more complex molecules or ions (3 or more atoms), try one multiple bond at a time if single bonds achieve octets.
6. My series for molecules/ions that need more than single bonds can be seen below (keeping in mind H, B, and Be only have single bonds)

• Single/single (\(\ce{-}\), \(\ce{-}\)) ,
• then single/double (\(\ce{-}\), \(\ce{=}\)),
• then double/double (\(\ce{=}\), \(\ce{=}\)),
• and if that does not work try triple/single (\(\ce{#}\), \(\ce{-}\))

7. In our class, we do not use multiple bonds for halogens.

Key Takeaways

• A covalent bond is formed between two atoms by sharing electrons.
• The number of bonds an element forms in a covalent compound is determined by the number of electrons it needs to reach octet.
• Hydrogen is an exception to the octet rule. H forms only one bond because it needs only two electrons.