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Chemical Bonding I: The Covalent Bond. Lewis dot symbols The covalent bond Electronegativity Writing Lewis structures Formal charge and Lewis structure The concept of resonance Exceptions to the octet rule Strength of the covalent bond. Lewis dot symbols. Lewis dot symbol

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Chemical bonding i the covalent bond
Chemical Bonding I: The Covalent Bond

  • Lewis dot symbols

  • The covalent bond

  • Electronegativity

  • Writing Lewis structures

  • Formal charge and Lewis structure

  • The concept of resonance

  • Exceptions to the octet rule

  • Strength of the covalent bond


Lewis dot symbols
Lewis dot symbols

  • Lewis dot symbol

    • By American chemist Gilbert Lewis

    • Consists of the symbol of an element and one dot for each valence electron in an atom of the element.


Chemical bonding i the covalent bond

  • The dots (representing electrons) are placed on the four sides of the atomic symbol (top, bottom, left, right)

  • Each side can accommodate up to 2 electrons

  • The number of valence electrons in the table below is the same as the column number of the element in the periodic table (for representative elements only)


  • The covalent bond
    The Covalent Bond

    • The covalent bond

      • results from the sharing of electrons between two atoms.

      • typically involves one nonmetallic element with another


    Chemical bonding i the covalent bond

    • The diatomic hydrogen molecule (H2) is the simplest model of a covalent bond, and is represented in Lewis structures

    • The shared pair of electrons provides each hydrogen atom with two electrons in its valence shell (the 1s) orbital.


    Chemical bonding i the covalent bond


    Chemical bonding i the covalent bond

    • The structures we use to represent H covalent electrons.2and Cl2 is called Lewis structures.

    • Lewis Structures

      • Is a representations of covalent bonding using Lewis dot symbols in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots on individual atoms.


    Chemical bonding i the covalent bond

    • Octet Rule covalent electrons.

      • An atom other than hydrogen tends to form bonds until it is surrounded by eight valence electrons.


    Chemical bonding i the covalent bond

    • Single bond covalent electrons.

      • Two atoms held together by one electron pair are said to be joined by single bond.

    • Multiple bonds

      • In many molecules atoms attain complete octets by sharing more than one pair of electrons between them

      • Two electron pairs shared a double bond

      • Three electron pairs shared a triple bond


    Chemical bonding i the covalent bond

    • Bond length covalent electrons.

      • The distance between the nuclei of two bonded atoms in a molecule.

      • Multiple bonds are shorter than single covalent bonds.


    Bond polarity and electronegativity
    Bond Polarity and Electronegativity covalent electrons.

    • The electron pairs shared between two atoms are not necessarily shared equally

      • Two extreme examples

        • In Cl2 the shared electron pairs is shared equally

        • In NaCl the 3s electron is stripped from the Na atom and is incorporated into the electronic structure of the Cl atom - and the compound is most accurately described as consisting of individual Na+ and Cl- ions

      • For most covalent substances, their bond character falls between these two extremes


    Chemical bonding i the covalent bond

    • Bond polarity is a useful concept for describing the sharing of electrons between atoms

      • A nonpolar covalent bond is one in which the electrons are shared equally between two atoms

      • A polar covalent bond is one in which one atom has a greater attraction for the electrons than the other atom. If this relative attraction is great enough, then the bond is an ionic bond


    Chemical bonding i the covalent bond

    • Electronegativity of electrons between atoms

      • A quantity termed 'electronegativity' is used to determine whether a given bond will be nonpolar covalent, polar covalent, or ionic.

      • Electronegativity is defined as the ability of an atom in a particular molecule to attract electrons to itself


    Drawing lewis structures
    Drawing Lewis Structures of electrons between atoms

    • The general procedure

      • 1. Sum the valence electrons from all atoms

        • Use the periodic table for reference

        • Add an electron for each indicated negative charge, subtract an electron for each indicated positive charge

      • 2. Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond

        • You may need some additional evidence to decide bonding interactions

        • If a central atom has various groups bonded to it, it is usually listed first: CO32-, SF4

        • Often atoms are written in the order of their connections: HCN


    Chemical bonding i the covalent bond

    • 3.Complete the octets of the atoms bonded to the central atom (H only has two)

    • 4. Place any leftover electrons on the central atom (even if it results in more than an octet

    • 5. If there are not enough electrons to give the central atom an octet, try multiple bonds (use one or more of the unshared pairs of electrons on the atoms bonded to the central atom to form double or triple bonds


    Chemical bonding i the covalent bond


    Chemical bonding i the covalent bond

    • 3. atom (H only has two)Completing the octets of the Cl atoms bonded to the central P atom:

    • 4. This gives us a total of (18 electrons) plus the 6 in the three single bonds, or 24 electrons total. Thus we have 2 extra valence electrons which are not accounted for. We will place them on the central element:


    Chemical bonding i the covalent bond


    Formal charge and lewis structure
    Formal Charge and Lewis Structure to invoke any double or triple bonds to achieve an octet for the central atom. We are finished.

    • Formal charge

      • The difference between the valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure is called formal charge

      • Two examples

        • O3 ( on textbook, P 258)

        • CO2


    Chemical bonding i the covalent bond

    • Example: Carbon Dioxide (CO to invoke any double or triple bonds to achieve an octet for the central atom. We are finished.2)

      • Carbon has 4 valence electrons. Each oxygen has 6 valence electrons, therefore our Lewis structure of CO2 will have 16 electrons.

      • Both structures fulfill the octet rule.

      • Which one is true?

      • Use formal charge to determine


    Chemical bonding i the covalent bond


    Chemical bonding i the covalent bond

    • In the above case, the second structure is the one with the smallest formal charges (i.e. 0 on all the atoms).

    • Furthermore, in the first possible Lewis structure the carbon has a formal charge of 0 and one of the oxygens it is bonded to has a formal charge of +1.

    • Oxygen is more electronegative than Carbon, so this situation would seem unlikely


    The concept of resonance
    The Concept of Resonance smallest formal charges (i.e. 0 on all the atoms).

    • Resonance structure

      • Is one of two or more Lewis structures for a single molecule that cannot be described fully with only one Lewis structure.

    • Resonance

      • Means the use of two or more Lewis structures to represent a particular molecule.


    Chemical bonding i the covalent bond


    Chemical bonding i the covalent bond

    • The important points to remember about resonance forms are: smallest formal charges (i.e. 0 on all the atoms).

      • The molecule is not rapidly oscillating between different discrete forms

      • There is only one form of the ozone molecule, and the bond lengths between the oxygens are intermediate between characteristic single and double bond lengths between a pair of oxygens

      • We draw two Lewis structures (in this case) because a single structure is insufficient to describe the real structure


    Exceptions to the octet rule
    Exceptions to the Octet Rule smallest formal charges (i.e. 0 on all the atoms).

    • There are three general ways in which the octet rule breaks down:

      • 1. Molecules with an odd number of electrons(The Odd-Electron Molecules)

      • 2. Molecules in which an atom has less than an octet (The Incomplete Octet)

      • 3. Molecules in which an atom has more than an octet( The Expanded Octet)


    Chemical bonding i the covalent bond

    • Less than an octet smallest formal charges (i.e. 0 on all the atoms). (most often encountered with elements of Boron and Beryllium)

      • It has 6 electrons

      • the structure of BF3, with single bonds, and 6 valence electrons around the central boron is the most likely structure


    Chemical bonding i the covalent bond

    • Odd number of electrons smallest formal charges (i.e. 0 on all the atoms).

      • Total electrons: 6+5=11

      • There are currently 5 valence electrons around the nitrogen.

      • We appear unable to get an octet around each atom


    Chemical bonding i the covalent bond

    • More than an octet smallest formal charges (i.e. 0 on all the atoms). (most common example of exceptions to the octet rule)

      • The orbital diagram for the valence shell of phosphorous is:

      • Third period elements occasionally exceed the octet rule by using their empty d orbitals to accommodate additional electrons


    Strengths of covalent bonds
    Strengths of Covalent Bonds smallest formal charges (i.e. 0 on all the atoms).

    • Bond-dissociation energy (i.e. "bond energy")

      • Bond energy is the enthalpy change (DH, heat input) required to break a bond (in 1 mole of a gaseous substance)