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Chemistry LibreTexts

12.05: Covalent Bonds

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  • Learning Objectives

    • Define covalent bond.
    • Illustrate covalent bond formation with Lewis electron dot diagrams.

    Ionic bonding typically occurs when it is easy for one atom to lose one or more electrons and another atom to gain one or more electrons. However, some atoms won't give up or gain electrons easily. Yet they still participate in compound formation. How? There is another mechanism for obtaining a complete valence shell: sharing electrons. When electrons are shared between two atoms, they make a bond called a covalent bond.

    Let us illustrate a covalent bond by using H atoms, with the understanding that H atoms need only two electrons to fill the 1s subshell. Each H atom starts with a single electron in its valence shell:

    \[\mathbf{H\, \cdot }\; \; \; \; \; \mathbf{\cdot \: H}\]

    The two H atoms can share their electrons:

    \[\mathbf{H}\: \mathbf{: H}\]

    We can use circles to show that each H atom has two electrons around the nucleus, completely filling each atom's valence shell:

    Because each H atom has a filled valence shell, this bond is stable, and we have made a diatomic hydrogen molecule. (This explains why hydrogen is one of the diatomic elements.) For simplicity's sake, it is not unusual to represent the covalent bond with a dash, instead of with two dots:


    Because two atoms are sharing one pair of electrons, this covalent bond is called a single bond. As another example, consider fluorine. F atoms have seven electrons in their valence shell:

    These two atoms can do the same thing that the H atoms did; they share their unpaired electrons to make a covalent bond.

    Note that each F atom has a complete octet around it now:

    We can also write this using a dash to represent the shared electron pair:

    There are two different types of electrons in the fluorine diatomic molecule. The bonding electron pair makes the covalent bond. Each F atom has three other pairs of electrons that do not participate in the bonding; they are called lone pair electrons. Each F atom has one bonding pair and three lone pairs of electrons.

    Covalent bonds can be made between different elements as well. One example is HF. Each atom starts out with an odd number of electrons in its valence shell:

    The two atoms can share their unpaired electrons to make a covalent bond:

    We note that the H atom has a full valence shell with two electrons, while the F atom has a complete octet of electrons.

    Example \(\PageIndex{1}\):

    Use Lewis electron dot diagrams to illustrate the covalent bond formation in HBr.


    HBr is very similar to HF, except that it has Br instead of F. The atoms are as follows:

    The two atoms can share their unpaired electron:

    Exercise \(\PageIndex{1}\)

    Use Lewis electron dot diagrams to illustrate the covalent bond formation in Cl2.


    More than two atoms can participate in covalent bonding, although any given covalent bond will be between two atoms only. Consider H and O atoms:

    The H and O atoms can share an electron to form a covalent bond:

    The H atom has a complete valence shell. However, the O atom has only seven electrons around it, which is not a complete octet. We fix this by including a second H atom, whose single electron will make a second covalent bond with the O atom:

    (It does not matter on what side the second H atom is positioned.) Now the O atom has a complete octet around it, and each H atom has two electrons, filling its valence shell. This is how a water molecule, H2O, is made.

    Example \(\PageIndex{2}\):

    Use a Lewis electron dot diagram to show the covalent bonding in NH3.


    The N atom has the following Lewis electron dot diagram:

    It has three unpaired electrons, each of which can make a covalent bond by sharing electrons with an H atom. The electron dot diagram of NH3 is as follows:

    Exercise \(\PageIndex{2}\)

    Use a Lewis electron dot diagram to show the covalent bonding in PCl3.



    Food and Drink App: Vitamins and Minerals

    Vitamins are nutrients that our bodies need in small amounts but cannot synthesize; therefore, they must be obtained from the diet. The word vitamin comes from "vital amine" because it was once thought that all these compounds had an amine group (NH2) in it. This is not actually true, but the name stuck anyway.

    All vitamins are covalently bonded molecules. Most of them are commonly named with a letter, although all of them also have formal chemical names. Thus vitamin A is also called retinol, vitamin C is called ascorbic acid, and vitamin E is called tocopherol. There is no single vitamin B; there is a group of substances called the B complex vitamins that are all water soluble and participate in cell metabolism. If a diet is lacking in a vitamin, diseases such as scurvy or rickets develop. Luckily, all vitamins are available as supplements, so any dietary deficiency in a vitamin can be easily corrected.

    A mineral is any chemical element other than carbon, hydrogen, oxygen, or nitrogen that is needed by the body. Minerals that the body needs in quantity include sodium, potassium, magnesium, calcium, phosphorus, sulfur, and chlorine. Essential minerals that the body needs in tiny quantities (so-called trace elements) include manganese, iron, cobalt, nickel, copper, zinc, molybdenum, selenium, and iodine. Minerals are also obtained from the diet. Interestingly, most minerals are consumed in ionic form, rather than as elements or from covalent molecules. Like vitamins, most minerals are available in pill form, so any deficiency can be compensated for by taking supplements.

    Figure \(\PageIndex{1}\) :Vitamins and Mineral supplements. Every entry down through pantothenic acid is a vitamin, and everything from calcium and below is a mineral.