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					 Chapter 8

Covalent Bonding
             Let’s Review
• What do we already know?
  – What is a chemical bond?
  – What is an ionic bond?
  Section 1

The Covalent Bond
             Stability
• Lower energy is more stable
• Noble-Gas electron configuration
• Octet rule
           Covalent Bond
• Atoms in nonionic compounds share
  electrons
• Covalent bond is the bond that results
  from sharing valence electrons
• Molecule is formed when two or more
  atoms bond covalently
        Diatomic Molecules
• Two atom molecules are more stable than
  one atom
• H2, N2, O2, F2, Cl2, Br2, I2



          H
                        H
    Hydrogen
     They Pair!!




H                  H
Hydrogen
Oxygen
Fluorine
    Fluorine



F              F
      Single Covalent Bonds
• One pair of valence electrons is shared
  – Pair may be referred to as “bonding” pair
• Also called sigma bonds
  –σ
  – Occurs when the shared pair is centered
    between the two atoms
            Bonding Orbital
• Localized region
  where bonding
  electrons are most
  likely found
      Groups and Single Bonds
•   Group 17
•   Group 16
•   Group 15
•   Group 14
      Homework (due Tuesday)
• Draw the Lewis structures for the following
  molecules
  –   PH3
  –   H2S
  –   HCl
  –   CCl4
  –   SiH4
• Challenge
  – Draw a generic Lewis Structure for a molecule formed
    between atoms of group 1 and group 16
       Homework continued
• Draw LDS for
  – CH4
  – Br2
  – C6H14 also written as CH3(CH2)4CH3
     Multiple Covalent Bonds
• Bond Order
  – Refers to the type of bond
• Single Bond
  – Shares ONE pair of electrons
• Double Bonds
  – Two pairs of electrons are shared
• Triple Bonds
  – Three pairs of electrons are shared
              The Pi Bond
• Multiple covalent bonds
  – Consist of at least one sigma and one pi bond
   Strength of Covalent Bonds
• CB involve attractive and repulsive forces
• Balance of the force is upset the bond can
  break
• Several factors influence strength of cb
               Bond Length
• Length depends on distance between
  bonded nuclei
• Bond length is the distance two nuclei at
  the position of maximum attraction
  – Determined by:
     • Sizes of two bonding atoms
     • Number of electrons shared
          Bonds and Energy
• Energy changes occur
  – When bonds are broken
    • Energy is released
    • Need energy put in to break it
       – Bond-dissociation energy
          » is the energy required to break a specific bond
          » Indicates strength of the bond
  – When bonds are formed
         Length and Energy
• Shorter the length the greater the energy
Energies of Chemical Reactions
• Total energy is determined from energy of
  bonds broken and formed
• Two types
  – Endothermic
  – Exothermic
Energies of Chemical Reactions
• Endothermic Reaction occurs when a
  greater amount of energy is required to
  break existing bonds in the reactants than
  is released when the new bonds formed.
• Endothermic Reaction
  – More energy to break a bond than energy
    when bond is broken
Energies of Chemical Reactions
• Exothermic



     Energy     Energy
       in        out

               Bond
Energies of Chemical Reactions
• Exothermic reaction occurs when more
  energy is released during product bond
  formation than is required to break bonds
  in reactants.
• Exothermic reaction
  – More energy is released than required to
    break the bonds
Energies of Chemical Reactions
• Endothermic



   Energy       Energy
     in          out
                Bond
Section Two

Naming Molecules
  Binary Molecular Compounds
Example: N2O
1. First element in the formula is always
   named first, using the entire element
   name.
  •   What is the first element?
      •   Nitrogen
 Binary Molecular Compounds
2. The second element in the formula is
   named using its root and adding the
   suffix –ide.
  1. What is the second element?
    •   Oxygen
  2. What will the name be?
    •   Oxide
  Binary Molecular Compounds
3. Prefixes are used to indicate the number of
    atoms of each element are present in the
    compound.
   1. How many nitrogens do we have?
     •   Two
  2. What will the prefix be?
     •   Di-
  3. What is the prefix plus the element?
     – Dinitrogen
Binary Molecular Compounds
1. How many oxygens do we have?
  • One
2. What will the prefix be?
  • Mono
3. What is the prefix plus the element?
  • Monoxide
 Binary Molecular Compounds
• What is the final answer?
How do we know what we are
         naming?
                    Pop Quiz
Match the following correctly, also note if
the acid is binary or an oxyacid:
1.   HCl                    A.   Chlorous acid
2.   HClO3                  B.   Sulfuric acid
3.   H2S                    C.   Hydrosulfuric acid
4.   H2SO4                  D.   Chloric acid
5.   H2ClO2                 E.   Hydrochloric acid

              ··Hint·· ClO3 is chlorate
Section Three

Molecular Structure
        Molecular Formula
• Shows the elements symbols and
  subscripts
• PH3
Lewis Structure


H   P    H

    H
Space-filling Molecular Model
Ball-and-stick Molecular Model
Structural Formula


   H   P    H

       H
        Molecular Formula
• CH4
Lewis Structure
    H

H   C    H

    H
Space-filling Molecular Model
Ball-and-stick Molecular Model
Structural Formula

      H

  H   C    H

      H
            Lewis Structures
•   BH3
•   Nitrogen trifluoride
•   C 2H 4
•   Carbon Disulfide
•   NH4+
•   ClO4-
            Announcement
• Print out chapter 8 review from teacher
  page.
• Complete by Friday (will have time in class
  tomorrow to work on it)
• Test Monday on sections 1,2,3
      Resonance Structures
• Resonance
  – A condition that occurs when more than one
    valid Lewis structure can be written for a
    molecule or ion
  – Molecules and ions that undergo resonance
    behave as if there is only one structure
                Classwork
• Page 260
  – #53
• Page 274
  – #84, 101, 102, 103, 104


• BONUS: 5 pts #137
  Exceptions to the Octet Rule
• Odd number of valence electrons
• Suboctets and coordinate covalent bonds
  – Stable configuration with fewer than eight
    electrons present
  – BH3
  – Coordinate Covalent bond
    • One atom donates both of the electrons to be
      shared with an atom or ion that needs two
      electrons to form a stable electron arrangement
      with lower potential energy.
  Exceptions to the Octet Rule
• Expanded Octets
  – Central atoms contain more than eight
    valence electrons
  – Considers the d orbital
  – Extra lone pairs are added to the central atom
    for more bonds
Section Four

Molecular Shape
        Importance of Shape
• The shape can determine
  – Physical properties
  – Chemical properties
• Electron densities created by overlap of
  orbitals of shared electrons determine
  molecular shape
                VSEPR Model
•   Valence
•   Shell
•   Electron
•   Pair
•   Repulsion
            VSEPR Model
• Arrangement that minimizes the repulsion
  of shared and unshared electron pairs
  around the central atom
• Bond Angle
  – Angle between bonds
              Hybridization
• Hybridization
  – A process in which atomic orbitals mix and
    form new, identical hybrid orbitals
             Hybridization
• With regards to molecules that have more
  than two atoms
• To determine the orbital hybrid
  – Determine the number of e- pairs shared, and
    lone pairs
                Hybridization
• Count like this. . . .
  –1=s
  – 2 = sp
  – 3 = sp2
  – 4 = sp3
  – 5 = sp3d
  – 6 = sp3d2
            Molecular Shapes
• Linear
    – Example BeCl2




Total       Shared    Lone    Hybrid     Bond
Pairs       Pairs     Pairs   Orbitals   Angle

2           2         0       sp         180
            Molecular Shapes
• Trigonal Planar
    – Example AlCl3




Total       Shared    Lone    Hybrid     Bond
Pairs       Pairs     Pairs   Orbitals   Angle

3           3         0       sp2        120
            Molecular Shapes
• Tetrahedral
    – Example CH4




Total      Shared   Lone    Hybrid     Bond
Pairs      Pairs    Pairs   Orbitals   Angle

4          4        0       sp3        109.5
            Molecular Shapes
• Trigonal Pyramidal
    – Example PH3




Total      Shared   Lone    Hybrid     Bond
Pairs      Pairs    Pairs   Orbitals   Angle

4          3        1       Sp3        107.3
            Molecular Shapes
• Bent
    – Example H2O




Total      Shared   Lone    Hybrid     Bond
Pairs      Pairs    Pairs   Orbitals   Angle

4          2        2       Sp3        104.5
            Molecular Shapes
• Trigonal
  Bipyramidal
    – Example NbCl5




Total       Shared    Lone    Hybrid     Bond
Pairs       Pairs     Pairs   Orbitals   Angle

5           5         0       sp3d       90; 120
            Molecular Shapes
• Octahedral
    – Example SF6




Total      Shared   Lone    Hybrid     Bond
Pairs      Pairs    Pairs   Orbitals   Angle

6          6        0       sp3d2      90; 90
         Practice Problems
• Page 264
  – #56 through60
     Section Five

Electronegativity & Polarity
 Electron Affinity, Electronegativity,
        and Bond Character
• Electron Affinity
  – The measure of the tendency of an atom to
    accept electrons
  – How attractive an atom is to electrons
  – Increases with atomic number within a period
  – Decreases with atomic number within a group
 Electron Affinity, Electronegativity,
        and Bond Character
• Electronegativity
  – Derived by comparing an atom’s attraction for
    shared electrons to that of a fluorine’s atom
    attraction for shared electrons
  – Ability of an atom to attract electrons to itself
    within a covalent bond
Electron Affinity, Electronegativity,
       and Bond Character
• Bond Character
  – Chemical bonds between atoms of different
    elements is never completely ionic or covalent
  – Four Types
    •   Mostly ionic
    •   Polar covalent
    •   Mostly covalent
    •   Nonpolar covalent
 Electron Affinity, Electronegativity,
        and Bond Character
• Bond Character
  – Can be predicted using the electronegativity
    difference of the elements that bond

Electronegativity           Bond Character
Difference
          > 1.7                    Mostly ionic
        0.4 – 1.7                Polar covalent
          < 0.4                  Mostly covalent
            0                  Nonpolar covalent
       Polar Covalent Bonds
• Polar Covalent Bonds
  – An unequal sharing of valence electrons
• Partial Charge
  – Represented by δ (Greek letter delta)
  – Due to unequal sharing, partial charges result
    • Partial positive—the atom with the lower electron
      affinity
    • Partial negative—the atom with higher electron
      affinity
           Molecular Polarity
• Covalently bonded molecules
  – Either polar or nonpolar
     • Depends on location and nature of bonds
• Nonpolar Molecules
  – Not attracted by electric field
• Polar Molecules
  – Dipoles, with charged ends
  – Uneven electron density = attracted by
    electric field
 Polarity and Molecular Shape
• Let’s look at H2O and CCl4
• What shape does water take?
  – Bent
• What shape does carbon tetrachloride
  take?
  – Tetrahedral
• Draw them
H2O & CCl4
  Polarity and Molecular Shape
• The symmetry in CCl4 allows for a
  nonpolar molecule.
• There is no symmetry in H2O, so it is
  polar.
• What about NH3?
  – It is polar.
Properties of Covalent Compounds
• Covalent compounds have strong bonds
  between atoms
• Attraction forces between molecules are
  relatively weak
• Intermolecular forces
  – Many types
Properties of Covalent Compounds
• Intermolecular Forces
  – Between nonpolar molecules
    • Force is weak
    • Called dispersion force or induced dipole
  – Between opposite charged ends of two polar
    molecules
    • dipole-dipole force
    • The more polar the molecule the stronger the force
Properties of Covalent Compounds
• Intermolecular Forces
  – Between hydrogen end of one dipole and a F,
    O, N atom on another dipole
     • Hydrogen bond
• Forces and Properties
  – Weak forces result in relatively low melting
    points
  – Molecular substances as gases at room
    temperature
     • O2, CO2, H2S
Properties of Covalent Compounds
• Forces and Properties
  – Hardness
    • Depends on strength of intermolecular forces
    • Many covalent compounds are soft
       – Example: Paraffin, found in candles
Properties of Covalent Compounds
• Forces and Properties
  – Solid Phase
    • Molecules align to form a crystal lattice
       –   Similar to ionic solid
       –   Less attraction between particles
       –   Shape affected by molecular shape
       –   Most information has been determined by molecular
           solids
     Covalent Network Solids
• Covalent Network Solids
  – Composed only of atoms interconnected by a
    network of covalent bonds
    • Example: Quartz and diamonds
  – Structure can explain properties
    • Diamond
       – Tetrahedral
       – Strong bonds
       – High melting point, extremely hard
The End

				
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