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2.5: Lewis Dot and Bonding

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    Learning Objectives
    • Identify and describe covalent bonds and ionic bonds.
    • Draw the Lewis dot structure for an atom.

    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 its valence electron(s) to a nonmetal, forming an ionic compound. As the transfer of electrons occurs, both atoms will achieve a more stabile confirmation. The end result will be a less reactive compound. These ionic species are composed of both cations (positively charged ions) and anions (negatively charged ions). In addition, ionic compounds are crystalline and solid in nature. A few examples of real-world ionic compounds would be NaCl (sodium chloride, table salt) and NaF (sodium fluoride, the active ingredient in toothpaste).

    Photograph of sodium chloride crystals showing transparent, cubic crystal shapes with flat faces and sharp edges. The repeating cube structure reflects the crystalline lattice of NaCl, where sodium and chloride ions are arranged in a regular three-dimensional pattern
    Figure \(\PageIndex{1}\): NaCl crystals. (Public Domain; NASA).

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

    Photograph of two clear glass bottles, one containing a red liquid of dinitrogen tetroxide, the other containing a yellow liquid of nitrogen dioxide.
    Figure \(\PageIndex{2}\): Covalent compounds NO2 and N2O4 contained in glass bottles. Image taken from: commons.wikimedia.org/wiki/File:NO2-N2O4.jpg
    Table \(\PageIndex{1}\): Comparing ionic and covalent compounds
    Ionic Bonding Covalent Bonding
    Metal and nonmetal combination nonmetal and nonmetal or metalloid and nonmetal combination
    Metal loses valence electrons and nonmetal gains up to eight. Both elements share valence electrons to achieve stability
    Individual atoms have charges (anions or cations) 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

    ADAPT \(\PageIndex{1}\)

     

    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 more easily. 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 to read a bibliographic memoir published by the National Academy of Sciences.

    Two‑part figure illustrating the development of the octet rule. (a) Black‑and‑white photograph of a chemist working at a laboratory bench with glassware, tubing, and chemical apparatus. (b) Reproduction of an early hand‑drawn sketch related to the octet rule, showing electrons arranged at the corners of a cube around element symbols, reflecting an early model of valence electron placement before modern Lewis dot structures.
    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 (this can be determined by looking at the Roman numeral listed above the element's group on the periodic table) and its Lewis symbol would be:

    Lewis structure of sulfur with two dots above the S, two below the S, and one dot on each side. The dots represent electrons.

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

    Lewis dot structure of fluorine is constructed by placing the first dot above the atomic symbol and adding one each time moving counter clockwise.

    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

    Table comparing electron configurations and Lewis dot symbols for elements across Groups 1, 2, and 13 through 18 in Period 2. The top row lists group numbers. The middle row shows electron configurations using noble gas notation from lithium through neon, progressing from one to eight valence electrons. The bottom row displays corresponding Lewis dot symbols for Li, Be, B, C, N, O, F, and Ne, with dots around each element symbol indicating the number of valence electrons.

    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 sodium metal, the resulting compound is not explosive and is less corrosive than chlorine gas. 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 to the Lewis structure with an understanding that metals will always lose valence electrons, and nonmetals will gain up to eight to form stable compounds.

    Diagram illustrating the formation of an ionic compound. On the right side, a sodium atom labeled “Na” is shown above a chlorine atom labeled “Cl,” each surrounded by electron shells. An electron transfer is implied, resulting in sodium becoming a positively charged ion and chlorine becoming a negatively charged ion. The diagram emphasizes opposite charges attracting to form sodium chloride (NaCl).
    Figure \(\PageIndex{4}\): Formation of an ionic bond between sodium and chlorine. Image courtesy of By BruceBlaus - Own work, CC BY-SA 4.0, 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 and a nonmetal, use the list of rules shown below:

    1. Locate the elements of interest and recall how many valence electrons each species has (look at the Roman numeral listed above the element's group).
    2. Draw dots around each element individually, one dot for each valence electron.
    3. To obtain an octet, the metal atom must transfer all of its valence electrons to the nonmetal. 
    4. At this point, the metal should have no valence electrons around it (and have a positive charge). With the newly acquired electrons from the metal, the nonmetal atom will have an octet (and a negative charge). 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 indicate the number of electrons gained or lost (this is an ionic compound).
    7. To check the structure, write a formula for the ionic compound, the charges of all atoms in the compound must add up to zero. 

    ADAPT \(\PageIndex{2}\)

     

    ADAPT \(\PageIndex{3}\)

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

    Diagram showing Lewis dot representations for the formation of sodium chloride. On the left, a sodium atom with one valence electron is shown next to a chlorine atom with seven valence electrons. An arrow indicates the transfer of the electron from sodium to chlorine. On the right, sodium is shown as a positively charged ion (Na⁺) and chlorine as a negatively charged ion (Cl⁻), each with a full outer shell, illustrating ionic bond formation.

    In this example, the sodium atom donates its 1 valence electron to the chlorine atom. This creates a sodium cation (+ charge) and a chloride anion (- charge). Notice that the net charge of the resulting compound is 0 (+1 - 1 = 0). In this example, only one atom of each element is needed to obtain a stable compound.

    Diagram showing Lewis dot representations for the formation of magnesium chloride. In the middle, a magnesium atom with two valence electron is shown next to two chlorine atoms each with seven valence electrons. An arrow indicates the transfer of one electron from magnesium to each chlorine. On the right side of the image, magnesium is shown as a positively charged ion and each chlorine as a negatively charged ion (Cl⁻), each with a full outer shell, illustrating ionic bond formation.

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

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    2.5: Lewis Dot and Bonding is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by LibreTexts.