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Chemistry LibreTexts

4.5: Lewis Dot and Bonding

  • Page ID
    96780
  • Skills to Develop

    • Realize that valence electrons can be represented by dots.
    • Know that there are two types of bonding: ionic and covalent.
    • Know that ionic bonding involves a metal and a nonmetal with a transfer of electrons.
    • Know that covalent bonding involves two nonmetals or a metalloid with a nonmetal sharing electrons.

    Bonding (Ionic and Covalent Basics)

    When two atoms approach each other, they have the potential to bond (or connect). If a metal and a nonmetal interact, then an ionic bond will result. These types of bonds involve the metal donating it(s) valence electron(s) to a nonmetal. As the electronic transfer occurs, both atoms will achieve more stabile confirmations. The end result will be a less reactive compound. These type of species are composed of both cations and anions. In addition, they are crystalline and solid in nature. A few examples of real-world ionic compounds would be NaCl (table salt) and NaF (active ingredient in toothpaste).

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    Figure \(\PageIndex{1}\): NaCl crystals. Image used with permission (Public Domain; NASA).

    If two nonmetals interact, then a covalent bond will result. The connection that forms is due to each atom sharing it(s) valence electron(s). In this type of bonding, each atom (except for H, B, and Be) obtains an octet and becomes stable. Metalloids and nonmetals can bond to form covalent compounds as well. Solid covalent compound are not crystalline in texture. Unlike ionic compounds. covalent compounds can be liquids or even gases.

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    Figure \(\PageIndex{2}\): Covalent compounds NO2 and N2O4 contained in glass bottles. Image taken from: https://commons.wikimedia.org/wiki/File:NO2-N2O4.jpg

    Table \(\PageIndex{1}\): Comparing ionic and covalent compounds
    Ionic Bonding Covalent Bonding
    Metal/nonmetal combination nonmetal/nonmetal or metalloid/nonmetal combination
    Metal loses valence electrons and nonmetal gains up to eight. Both elements share valence electrons to achieve stability
    Individual species have cation/anion charges No charges are present in these compounds
    All ionic compounds are solids Covalent compounds can be solid, liquid, or gas.
    Examples could include: Li2CO3, Fe2O3, and MgSO4 Examples include: CO2, H2O, and CH3CH2OH

    Lewis Symbols

    At the beginning of the 20th century, an American physical chemist G. N. Lewis (1875–1946) devised a system of symbols—now called Lewis electron dot symbols (often shortened to Lewis dot symbols) that can be used for predicting the number of bonds formed by most elements in their compounds. Each Lewis dot symbol consists of the chemical symbol for an element surrounded by dots that represent its valence electrons. Lewis knew that incorporating models into his teaching would enable students to visualize chemical bonding easier. Gilbert Lewis was known for his interactive teaching methods and his interest in student success. Unfortunately, his nervous personality limited his ability to lecture in front of large groups of people.

    Regarding his research, Lewis constructed his own theory to explain the nature of acids and bases. He was the first to synthesize D2O (heavy water) and wrote numerous textbooks for his courses. In World War I, he served as a Major in the Gas Service unit of the American Army. He trained over 200 soldiers a week on how to adequately protect themselves from gas warfare. From these efforts, he helped reduce fatalities due to chemical gas exposures and was awarded the Distinguish Service Medal in 1922.

    Lewis was nominated over thirty times for the Nobel Prize. He never received this award and does not have an element named in honor of him. For more on Gilbert Lewis, click on this link.

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    Figure \(\PageIndex{3}\): G. N. Lewis and the Octet Rule. (a) Lewis is working in the laboratory. (b) In Lewis’s original sketch for the octet rule, he initially placed the electrons at the corners of a cube rather than placing them as we do now.

    Lewis Dot symbols

    • convenient representation of valence electrons
    • allows scientists to keep track of valence electrons during bond formation
    • consists of the chemical symbol for the element plus a dot for each valence electron

    To write an element’s Lewis dot symbol, we place dots representing its valence electrons, one at a time, around the element’s chemical symbol. Up to four dots are placed above, below, to the left, and to the right of the symbol (in any order, as long as elements with four or fewer valence electrons have no more than one dot in each position). The next dots, for elements with more than four valence electrons, are again distributed one at a time, each paired with one of the first four. For example, the element sulfur has six valence electrons (note roman numeral above group on the periodic table) and its Lewis symbol would be:

    Fluorine, for example, has seven valence electrons, so its Lewis dot symbol is constructed as follows:

    imageedit_5_6530088793.jpg

    The number of dots in the Lewis dot symbol is the same as the number of valence electrons, which is the same as the last digit of the element’s group number in the periodic table. Lewis dot symbols and electron configurations (not a topic for this course) for the elements in period 2 are given in Table \(\PageIndex{2}\).

    Table \(\PageIndex{2}\): Lewis Dot Symbols for the Elements in Period 2

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    Ionic compounds are produced when a metal bonds with a nonmetal. Stability is achieved for both atoms once the transfer of electrons has occurred. The image below shows how sodium and chlorine bond to form the compound sodium chloride. Unlike a sodium atom, the resulting compound is not explosive and less corrosive than chlorine. Ionic bonding can be viewed by noting the donation of valence electrons from a metal atom to a nonmetal atom by using the Bohr model. From this theory, we will move on the Lewis structure with an understanding that metals will always lose valence electrons and nonmetals will gain up to eight in order to form stable compounds. 

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    Figure \(\PageIndex{4}\): Formation of an ionic bond between sodium and chlorine. Image courtesy of By BruceBlaus - Own work, CC BY-SA 4.0, https://commons.wikimedia.org/w/inde...curid=44968540

    Rules for Drawing IOnic Lewis structures

    When drawing Lewis structures for compounds formed from the combination of a metal/nonmetal, use the list of rules shown below:

    1. Locate the elements of interest and recall how many valence electrons each species has. 
    2. Draw dots around each element individually.
    3. To obtain successful octets, the metal must transfer all of its valence electrons to the nonmetal.
    4. At this point, the metal should have no electrons around it. In addition, the nonmetal will have an octet. If this does not occur, then refer to step e.
    5. Additional atoms of the same two elements may be brought in to accomplish rules a-c.
    6. Cation and anion charges must be on Lewis structure to receive full credit (this is an ionic compound).
    7. To check structure, quickly write a formula for the ionic compound. This involves crossing down the charges diagonally and losing signs. You must reduce subscripts if they are divisible by a factor.

    Example \(\PageIndex{1}\): Chloride Salts

     ionic1 (2).jpg

    In this example, the sodium atom is donating its 1 valence electron to the chlorine atom. This creates a sodium cation and a chlorine anion. Notice that the net charge of the resulting compound is 0. Only one atom of each element is needed to obtain a stable compound.

     ionic2 (1).jpg

    In this example, the magnesium atom is donating both of its valence electrons to chlorine atoms. Each chlorine atom can only accept 1 electron before it can achieve an octet; therefore, 2 atoms of chlorine are required to accept the 2 electrons donated by the magnesium. Notice that the net charge of the compound is 0. This Lewis structure requires two anions per one cation to produce a stable compound.

    For more rigorous examples, access the CHM101 Furman moodle page (chapter 4 documents), to view a lightboard video of your instructor.