Chapter 11: Chemical Bonding

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Chapter 11 Chemical Bonding
Chemistry 1020 Interpretive chemistry Andy
Aspaas, Instructor
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Chemical bonds

• Chemical bonds forces that hold atoms together
• Ionic bonds forces of attraction between
  oppositely charged ions
• Electron transfer forms two oppositely charged
  ions
• Electrostatic forces opposite charges attract
• Covalent bonds forces of attraction between two
  atoms which are sharing electrons
• Molecule group of covalently bonded atoms

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Bond polarity

• Covalent bonding between unlike atoms
• Unequal sharing of electrons
• One end of bond has larger electron density than
  other
• Bond polarity result of uneven electron sharing
• End with larger electron density gets partial
  negative charge (?-)
• End that is electron deficient gets partial
  positive charge (?)

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Electronegativity

• Electronegativity ability of an atom to attract
  shared electrons
• Values from 0.7 - 4.0
• Large values atom attracts electrons more
  strongly
• Periodic table electronegativity increases left
  to right (across a period)
• Decreases top to bottom (down a group)
• Larger differences in electronegativity between
  covalently bonded atoms mean a more polar bond

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Dipole moment

• If a molecule has a center of positive charge and
  a center of negative charge in different points,
  it has a dipole moment
• If there are more than one partial negative or
  positive charges in a molecule, they may
  partially cancel each other out
• Combine to form a single dipole moment for the
  molecule
• Molecules with a large dipole moment are polar

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Charges on ions

• Stable ions form noble gas configurations
• Cations have lost electrons to get noble gas
  configuration
• Anions have added electrons
• Use charges to predict ionic formulas
• Ionic compounds have no net charge

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Lewis structures

• Chemical bonding involves only valence electrons
  of atoms
• Lewis structure shows valence electrons as dots
  around atoms
• Cations have no dots, anions have 8
• Duet rule Hydrogen forms stable molecules when
  it shars two electrons
• Octet rule Second-row nonmetals form stable
  molecules when valence orbitals are full, 8
  electrons

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Writing Lewis structures

• Find sum of all valence electrons in molecule
• Use one pair of electrons to connect each pair of
  bound atoms
• Arrange remaining elecctrons to satisfy the duet
  rule for hydrogen or the octet rule for 2nd-row
  elements

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Multiple-bonds

• Its possible for a pair of atoms to share 2 or 3
  pairs of electrons in order to satisfy the octet
  rule
• Double bond 2 pairs of electrons are shared
• Triple bond 3 pairs of electrons are shared

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A few exceptions to the octet rule

• Boron compounds are stable with 6 electrons in
  borons valence shell
• BF3
• Some third row elements can expand their octet
• S, P

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3-dimensional molecular structure

• VSEPR valence shell electron pair repulsion
  model
• 3-dimensional molecular structure is determined
  by minimizing repulsions between electron pairs
• Count electron pairs as well as bonds

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Electron pair arrangements

• Linear only 2 total lone pairs and/or bonds
• 180 bond angles
• BeCl2
• Trigonal planar 3 total lone pairs and/or bonds
• 120 bond angles
• BF3
• Tetrahedral 4 total lone pairs and/or bonds
• 109.5 bond angles
• CH4

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Molecule shape

• Arrangement of bonds indicate shape of molecule
• 3 bonds 1 lone pair
• Trigonal pyramid
• NH3
• 2 bonds 2 lone pairs
• Bent
• H2O