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Chapter_8___9

VIEWS: 14 PAGES: 12

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									Chapter 8 & 9: Bonding                                                                              1

Chapter 8 and 9

Chemical Bond:
Bond Energy:
Valence Electrons:



Electronegativity: ability of an atom in a bond to __________________


Ionic Bond: Attraction of oppositely charged ions
Ionic Compound:

Coulomb’s Law: E = (2.31 x 10-19 J nm) (Q1Q2/r)




Monatomic ions:
Polyatomic ions:




                 Active metals tend to lose electrons to obtain previous noble gas configuration.
                 Nonmetals tend to gain electrons to obtain next noble gas configuration.

Isoelectronic:




Electronegativity Difference (EN): determines how likely two elements are to form an ionic
       compound.
                  : large EN, ionic compound is likely (approximately 1.7)
                  : Cs 0.8 F 4.0 EN =3.2 ionic


Lattice Energy: the change in energy associated with the formation of an ionic solid from gaseous
ions.
Chapter 8 & 9: Bonding                                                                                   2


Example: Use the following data to estimate Hfo for magnesium fluoride. (-2088 kJ/mol)

        Lattice Energy:                                    -3916 kJ/mol
        First Ionization Energy of Mg                      735 kJ/mol
        Second Ionization Energy of Mg                     1445 kJ/mol
        Electron affinity of F                             -328 kJ/mol
        Bond energy of F2                                  154 kJ/mol
        Enthalpy of sublimation of Mg                      150. kJ/mol




Covalent Bonding
Covalent Bond: two atoms sharing one or more pairs of electrons (small EN values)



Lewis Formula: shows all valence electrons in an atom, ion or molecule
   Atoms:


Atoms in molecules tend to have noble gas electron configurations (share in 8 electrons)-octet rule or
rule of eight.

    Molecules:
              H2
              O2
              N2

Single covalent bond: sharing of
Double covalent bond: sharing of
Triple covalent bond: sharing of


Bond Energies (B.E.)
The bond energy is the energy needed to break one mole of bonds in a gas substance to form 1 mole of
products in the gas state at constant T and P.

        H2(g)  2H(g)        Ho = 432 kJ/mol
        O2(g)  2O(g)        Ho = 495 kJ/mol

    When bonds are broken, H is always
    When bonds are formed, H is always

Calculation of Ho using bond energies: Ho = broken - formed
Chapter 8 & 9: Bonding                                                                               3


Example: Calculate Ho for the reaction, 2H2(g) + O2(g)  2H2O(g) using B.E. values from Table 8.4
(page 351).
       Ho = 2BE(H-H) + BE(O=O) – 4BE(HO)
       Ho =

Example: C2H4(g) + F2(g)  C2H4F2(g)




Diatomic Elements: BrINClHOF


Unshared electrons:



Writing Lewis Structures (Electron Dot Structures)

        1. Count number of valence electrons. (NH3)

        2. Draw a skeleton structure joining adjacent atoms with a single bond.

        3. Arrange remaining electrons to satisfy octet for each atom (except H).

        *Central atom is usually written first in formula.
        *Terminal atoms are usually the most electronegative.


        NH4+

        ClO3-

        Cl2CO

        HCN


Resonance: Sharing of a bonding pair of electrons between more than 2 atoms.

        NO3-
Chapter 8 & 9: Bonding                                                                               4

        SO2




Exceptions to the Octet Rule
   1) Be compounds (BeCl2)
   2) B compounds (BBr3)
   3) Odd # of valence electrons (NO2)
   4) Expanded Octets (SF6, atoms with available d orbitals)




Formal Charge: difference between the number of valence electrons in the atom and the number of
electrons assigned to the atom
            -used to determine most accurate Lewis Structure in cases where more than one possible
            structure exists.

        Formal Charge = Group Number – (number bonding pairs + nonbonding electrons)



Polar and Nonpolar Covalent Bonds

    Nonpolar covalent bonds:




    Polar covalent bonds:



         H-F                H-Cl         H-Br            H-I
    EN: 2.1-4.0             2.1-3.0      2.1-2.8         2.1-2.5
    EN: 1.9                0.9          0.7             0.4
         Most Polar                                      Least Polar


    Electric dipole: separation of charge
       Unequal sharing of electrons leads to + and -
Chapter 8 & 9: Bonding                                                                         5




                          Properties of Ionic versus Covalent Compounds

Ionic Compounds                                          Covalent Compounds
1. Solids (High melting points >400oC)             1. Gases, Liquids, and Solids

2. Generally soluble in polar solvents (water)     2. Generally insoluble in polar solvents
                                                            (water).

3. Most are insoluble in nonpolar solvents.        3. Most are soluble in nonpolar solvents.

4. Molten compounds conduct electricity.           4. Liquids and molten compounds do not
                                                             conduct electricity.

5. Aqueous solutions conduct electricity.          5. Aqueous solutions are usually poor
                                                            conductors of electricity.



Naming Compounds

Naming Ionic Compounds

    Binary-

    Simple Binary Ionic Compounds

    1.
    2.
    3.


    Binary Ionic Compounds with Variable Charge Cations

    The charge on the cation must be specified in the name of the compound
    The charge on the cation is determined from the charge on the anion
Chapter 8 & 9: Bonding                                                     6


   Ionic Compounds with Polyatomic Ions

    1. Name cation
    2. Name anion



            potassium fluoride

            lithium phosphate

            Cd3(PO4)2

            Al2(SO4)3


Naming Binary Covalent (Molecular) Compounds
  CO
  CO2
  *Name first element, name 2nd element with “ide” ending. Use prefixes.

            Prefixes
            mono
            di
            tri
            tetra
            penta
            hexa
            hepta
            octa
            nona
            deca

    NO2
    N2O4


Recognizing and Naming Acids (page 66)

Acid: substances that produce H+ when dissolved in water
Formulas:
Acids without oxygen:



Acids with oxygen:
Chapter 8 & 9: Bonding                                                                      7




Lewis Structures:


Major features of molecular geometry can be predicted based upon electron pair repulsion.
       VSEPR represents Valence Shell Electron Pair Repulsion


Bonding and lone pairs of electrons are regions of High Electron Density (HED)



Ideal Geometries
Two electron pairs around a central atom.
Example       HED            Geometry          Bond Angles
BeCl2


Three electron pairs around a central atom.
Example        HED            Geometry         Bond Angles
BCl3


Four electron pairs around a central atom.
Example        HED            Geometry         Bond Angles
CH4



Five electron pairs around a central atom.
Example        HED            Geometry                 Bond Angles
PF5


Six electron pairs around a central atom.
Example        HED            Geometry         Bond Angles
SF6
Chapter 8 & 9: Bonding                                                         8

Effect of Lone Pairs on Geometry
1.      Electron pair geometry is relatively unaffected by lone pairs
2.      Molecular geometry is different than EG when lone pairs are present.




Example         HED       EG             MG             Bond Angles
CH4


Example         HED       EG             MG                    Bond Angles
NH3


Example         HED       EG             MG                    Bond Angles
H2 O


Polarity of Molecules
Nonpolar covalent bond: equal sharing of electrons
Polar Covalent bond: unequal sharing of electrons

Polar molecules contain positive and negative poles (dipole)




Two Factors Determine Polarity of a Molecule
      1.     bond polarity
      2.     molecular geometry




Example:    BCl3
Example:    BCl2F
Example:    CH2Cl2
Example:    CH2O


VSEPR: electrons repel each other for maximum separation

Valence Bond (VB): ½ filled orbitals overlap to form covalent bonds
Chapter 8 & 9: Bonding                                                            9

Why does O typically form two bonds?
Why does N typically form three bonds?


Example         HED      EG        MG    Bond Angles   Polarity   Hybridization
BeH2

Lewis structure
Hybridization


Example         HED      EG        MG    Bond Angles   Polarity   Hybridization
BH3

Lewis structure
Hybridization



Example         HED      EG        MG    Bond Angles   Polarity   Hybridization
CH4

Lewis structure
Hybridization



Example         HED      EG        MG    Bond Angles   Polarity   Hybridization
PCl5

Lewis structure
Hybridization




Example         HED      EG        MG    Bond Angles   Polarity   Hybridization
SF6

Lewis structure
Hybridization



Example         HED      EG        MG    Bond Angles   Polarity   Hybridization
IF7
Chapter 8 & 9: Bonding                                                             10

Lewis structure
Hybridization



Example         HED      EG    MG      Bond Angles      Polarity   Hybridization
NH3

Lewis structure
Hybridization



Example         HED      EG   MG    Bond Angles      Polarity   Hybridization
AsF4-

Lewis structure
Hybridization



Example         HED      EG   MG    Bond Angles      Polarity   Hybridization
SeF5-

Lewis structure
Hybridization



Example         HED      EG   MG    Bond Angles      Polarity   Hybridization
BrF3

Lewis structure
Hybridization


Example         HED      EG   MG    Bond Angles      Polarity   Hybridization
KrF4

Lewis structure
Hybridization



Multiple Centers
Examples and Models
Chapter 8 & 9: Bonding                                                                                11




Multiple Bonds
Examples, Models, Overheads
C2H4
C2H2
CH2O




Valence Bond Theory (Localized Electron Model): Overlap of ½ filled hybridized or pure atomic
orbitals leads to bonding.
    (hybridization is the mathematical combination of atomic orbitals on a specific atom)

Molecular Orbitals (MO): Atomic orbitals on separate atoms combine to form MO.
  (molecular orbitals are mathematical combinations of orbitals on separate atoms, no atomic orbitals
  remain once molecule has formed)

VB theory is easier to visualize (VSEPR)

MO theory gives better description of electron cloud, magnetism and bond energies.

Electrons can be treated mathematically as waves
   -waves added in phase interact constructively                +              =
   -waves added out of phase (subtracted) interact destructively
                       +                    =



Bonding Orbital: Results when two orbitals interact in phase (electron density between atoms, lower
energy than atomic orbitals).



Antibonding Orbital: Results when two orbitals interact out-of-phase (electron density shifts to
unfavorable regions, higher in energy than atomic orbitals).




                   bonding e' s  antibonding e' s
    bond order                                    
                                  2
Chapter 8 & 9: Bonding                                                                        12


Increasing bond order correlates with increasing bond stability and decreasing bond length.

                         Bond Order        Bond Length (Å)      Bond Energy (kJ/mol)
        N2                   3                  1.09                   946
        O2                   2                  1.21                   498
        F2                   1                  1.43                   159



Magnetic Properties
Paramagnetic substances-contain unpaired electrons.
         -Weakly attracted by magnetic fields.

Diamagnetic substances-contain no unpaired electrons.
         -Weakly repelled by magnetic fields.

Ferromagnetism- (Fe, Co, Ni) permanent magnets
      -Electrons in adjacent atoms align and generate a magnetic field.

								
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