Skip to main content
Chemistry LibreTexts

4.2: Covalent Bonds and the Periodic Table

  • Page ID
    258854
  • Learning Objectives

    • Predict the number of covalent bonds formed based on the elements involved.

    How Many Covalent Bonds Are Formed?

    The number of bonds that an atom can form can often be predicted from the number of electrons needed to reach an octet (eight valence electrons). In the Lewis structure, the number of bonds formed by an element in a neutral compound is the same as the number of unpaired electrons it must share with other atoms to complete its octet of electrons. For example, each atom of a group 4A (14) 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). Group 5A (15) 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 (16) obtain an octet by forming two covalent bonds. Fluorine and the other halogens in group 7A (17) have seven valence electrons and can obtain an octet by forming one covalent bond.

    clipboard_edc1e143f3b3082bc589a665a4830d1bf.png

    Typically, the atoms of group 4A form 4 covalent bonds; group 5A form 3 bonds; group 6A form 2 bonds; and group 7A form one bond. The number of electrons required to obtain an octet determines the number of covalent bonds an atom can form. This is summarized in the table below. 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.

    Atom (Group number) Number of Bonds Number of Lone Pairs
    Carbon (Group 14 or 4A) 4 0
    Nitrogen (Group 15 or 5A) 3 1
    Oxygen (Group 16 or 6A) 2 2
    Fluorine (Group 17 or 7A) 1 3

    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. Figure \(\PageIndex{1}\) shows the number of covalent bonds various atoms typically form.

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

    Number.jpg

    Figure \(\PageIndex{1}\): How Many Covalent Bonds Are Formed? In molecules, there is a pattern to the number of covalent bonds that different atoms can form. Each block with a number indicates the number of covalent bonds formed by that atom in neutral compounds.

    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?

    OF2.jpg

    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?

    NCl3.jpg

    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.

    Concept Review Exercises

    1. How is a covalent bond formed between two atoms?

    2. How does covalent bonding allow atoms in group 6A to satisfy the octet rule?

    Answers

    1. Covalent bonds are formed by two atoms sharing electrons.

    2. The atoms in group 6A make two covalent bonds.

    • Was this article helpful?