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    Basic Concepts of Chemical Bonding (Con't) : (4 + 6) 10 Valence Electrons

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    The document summarizes concepts of chemical bonding including carbon monoxide, formal charge, Lewis structures, resonance structures, exceptions to the octet rule, strength of covalent bonds, enthalpies of reactions, bond enthalpy and bond length, and oxidation numbers. Carbon monoxide forms a triple bond between carbon and oxygen to allow both atoms to complete their octets. Formal charge is calculated based on assigning electrons in Lewis structures. Resonance structures show equivalent Lewis structures for molecules like ozone and benzene. Exceptions to the octet rule include radicals and compounds with boron or elements in period 3 or higher that allow for expanded octets. Bond strength is measured by bond enthalpy and increases with shorter bond lengths

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    Basic Concepts of Chemical Bonding (Con't) : (4 + 6) 10 Valence Electrons

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    41 views6 pages
    The document summarizes concepts of chemical bonding including carbon monoxide, formal charge, Lewis structures, resonance structures, exceptions to the octet rule, strength of covalent bonds, enthalpies of reactions, bond enthalpy and bond length, and oxidation numbers. Carbon monoxide forms a triple bond between carbon and oxygen to allow both atoms to complete their octets. Formal charge is calculated based on assigning electrons in Lewis structures. Resonance structures show equivalent Lewis structures for molecules like ozone and benzene. Exceptions to the octet rule include radicals and compounds with boron or elements in period 3 or higher that allow for expanded octets. Bond strength is measured by bond enthalpy and increases with shorter bond lengths

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    Ch 8-3(con't)

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    The document summarizes concepts of chemical bonding including carbon monoxide, formal charge, Lewis structures, resonance structures, exceptions to the octet rule, strength of covalent bonds, enthalpies of reactions, bond enthalpy and bond length, and oxidation numbers. Carbon monoxide forms a triple bond between carbon and oxygen to allow both atoms to complete their octets. Formal charge is calculated based on assigning electrons in Lewis structures. Resonance structures show equivalent Lewis structures for molecules like ozone and benzene. Exceptions to the octet rule include radicals and compounds with boron or elements in period 3 or higher that allow for expanded octets. Bond strength is measured by bond enthalpy and increases with shorter bond lengths

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    41 views6 pages

    Basic Concepts of Chemical Bonding (Con't) : (4 + 6) 10 Valence Electrons

    Uploaded by

    Pineraser
    The document summarizes concepts of chemical bonding including carbon monoxide, formal charge, Lewis structures, resonance structures, exceptions to the octet rule, strength of covalent bonds, enthalpies of reactions, bond enthalpy and bond length, and oxidation numbers. Carbon monoxide forms a triple bond between carbon and oxygen to allow both atoms to complete their octets. Formal charge is calculated based on assigning electrons in Lewis structures. Resonance structures show equivalent Lewis structures for molecules like ozone and benzene. Exceptions to the octet rule include radicals and compounds with boron or elements in period 3 or higher that allow for expanded octets. Bond strength is measured by bond enthalpy and increases with shorter bond lengths

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    of 6

    Miles Eckles March 24, 2011

    Chem 101 A Chapter 8 (con’t)

    Basic Concepts of Chemical Bonding (con’t)

    Carbon Monoxide

    CO (4 + 6) = 10 valence electrons

    C–O

    That leaves 8 valence electrons, but no way to obtain octets for both C and O without multiple
    bonding. Triple bond allows for the octets to complete on both and use 4 nonbonding valence
    electrons.

    :C ≡ O:

    Formal Charge

     Formal charge of an atom is the charge that an atom in a molecule would have if all
    atoms had the same electronegativity.
     How to assign electrons for formal charge calculation:
    1. All nonbonding (unshared) electrons are assigned to atom on which they are found.
    2. Half of the bonding electrons are assigned to each atom in the bond.
     Formal charge equals the number of valence electrons in the isolated atom minus the
    number of electrons assigned to the atom in the final Lewis structure.
     Lewis Structures
    o After drawing the skeleton, count how many electrons are left. Then, count the #
    of electrons you need to fill octets with nonbonding electrons.
    o Compare:

     C O 10

     :C ≡ O: 2 / 8 left 12 needed for octet


    o If electrons needed is greater than the electrons left, there is multiple bonding.
     #MB = # of electrons needed -- # of electrons left / 2
    o 12 – 8 / 2 = 4 / 2 = 2
    o Used: 3 (2) = 6
    o Left: 10 – 6 = 4
     Based on the group #
     Formal Charge = Group # - NBe – ½ Bonding
     NBe = # of nonbonding electrons

    Example

    :C ≡ O:
    Valence electrons 4 6
    Electrons to atom 5 5
    Formal charge -1 +1

    When several Lewis structures are possible, the most stable one will be that in which 1) the
    atoms bear the smallest formal charges and 2) any negative charges reside on the more
    electronegative atoms.

    o Key: formal charges do not represent real charges on atoms.

    Resonance Structures

     Resonance structures are equivalent Lewis structures except for placement of electrons.

    Ozone, O3

    o Still the same molecule even written as a Lewis structure


    Resonance in Benzene (C6H6)

    H H H H

    C C C C

    H C C H H C C H

    C C C C

    H H H H

    or

    Exceptions to Octet Rule

    1. Molecules with an odd number of electrons.


    2. Molecules in which an atom has less than an octet.
    3. Molecules in which an atoms has more than an octet.

    Odd number of electrons

     Very RARE.
    o Examples include nitric oxide (NO), nitrogen dioxide (NO 2), and chlorine dioxide
    (ClO2).
    o An unpaired electron is called a radical species; radicals are implicated in a
    number of ill biological effects, such as cacrinogenesis.

    Less than an octet

     Most often encountered with compounds of beryllium (Be) and boron (B).
    o Beryllium is 1 metal in Group II that can form
    covalent bonds
    Stable as sextet around boron atom

    Stable as quartet around beryllium atom

    More than an Octet

     Largest class of exceptions.


    o This doesn’t satisfy the octet rule
    o If in the 3rd period or below this is possible

    Occurs for Period 3+ elements

     It is possible to have expanded octets for elements that might be in Period 3 or higher
    because some of the electrons can be accommodated into d orbitals (don’t have d
    orbitals for first or second period).

    Strength of Covalent Bonds

     Bond enthalpy is the enthalpy change, ΔH, for the breaking of a particular bond in a
    mole of gaseous substance.

    :C≡O: (g) (g) + (g)

    ΔH = 1072 kJ/mol
    Bond enthalpies are always positive.

    Enthalpies of Reactions

     Enthalpy of reaction, ΔHrxn

    ΔHrxn = Σ (bond enthalpies of bonds broken) – Σ (bond enthalpies of bonds made)

    What is ΔHrxn for

    CH4 (g) + Cl2 (g) CH3Cl (g) + HCl (g)

    Solution:

    Bonds broken 1 mol C–H, 1 mol Cl–Cl

    Bonds formed: 1 mol C—Cl, 1 mol H—Cl

    ΔHrxn = [D (C – H) + D (Cl – Cl)] – [D (C – Cl) + D (H – Cl)]

    ΔHrxn = [(413 kJ) + (242 kJ)] – [(328 kJ) + (431 kJ)] = -104 kJ

    = -104 kJ

    D = bond enthalpy refers to bond dissociation

    Coefficients matter in the thermochemical equation

    Determines the # of bonds broken or formed

    Bond Enthalpy and Bond Length

     Bond length – the average distance between the nuclei of the atoms involved in a bond.
    o In general, as the number of bonds between two atoms increases, the bond
    becomes shorter and stronger.

    C C 1.54 Å, 348 kJ/mol


    C C 1.34 Å, 614 kJ/mol
    C C 1.20 Å, 839 kJ/mol

    Oxidation Numbers

     Oxidation number of an atom is the charge that results when the electrons in a covalent
    bond are assigned to the more electronegative atoms; it is the charge atom would
    possess if it were ionic.

    Rules for Assigning Oxidation Numbers

    1. Oxidation number of an element in its elemental form is zero.


    a. E.g., H2, S8
    2. Oxidation number of a monatomic ion is the same as its.
    a. E.g., Na+ will be +1
    3. In binary compounds (two elements) the element with the greater electronegativity is
    assigned a negative oxidation number equal to its charge in simple ionic compounds of
    an element.
    4. Sum of oxidation numbers equals zero for an electrically neutral compound and equals
    overall charge of ionic species.

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