Chapter 5 Chemical Bonding

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Chapter 5Chemical Bonding
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Poison or Seasoning?

• How can two poisons (elemental sodium and
  elemental chlorine) combine to form a flavor
  enhancer (sodium chloride) that tastes great on
  steak?
• Answer By an exchange of electrons that
  stabilizes both atomsthe formation of a chemical
  bond.

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G. N. Lewis and Bonding

• Lewis theory fundamental ideas
• Valence electrons are the most important.
• Two bond types
• Ionic valance electrons are transferred
• Covalent valence electrons are shared
• Bond formation results in formation of full outer
  Bohr orbits. Because this stable configuration
  typically involves eight electrons, this is
  commonly known as the OCTET rule.

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Lewis Structure

• Element symbol surrounded by a number of dots
  equal to the number of valence electrons
• Ignore the inner or CORE electrons.
• The order in which electrons (dots) are drawn and
  their exact locations are not critical.
• Chemical bonding brings together elements in the
  correct ratios so that all of the atoms involved
  form an octet.

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Ionic Lewis Structures

• Since ionic bonding involves the transfer of
  electrons from a metal to a nonmetal, the Lewis
  structure for an ionic compound involves moving
  dots. The metal becomes a cation and the
  nonmetal becomes an anion.

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Ionic Lewis Structures Charges

• The metal and the nonmetal each acquire a charge
  in the formation of an ionic bond.
• We indicate the magnitude of the charge in the
  upper right corner of the symbol.
• We enclose the anion in brackets.
• Charges on anions and cations within an ionic
  formula sum to zero.

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Concept Check 5.1

• Draw the Lewis structure for CaO.

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Concept Check 5.1 Solution

• Calcium is a metal and when it forms ionic bonds
  with nonmetals, it loses both of its valence
  electrons to form a Ca2 ion with a stable noble
  gas configuration.
• Oxygen is a nonmetal and when it forms ionic
  bonds with metals, it gains electrons to form an
  O2- ion with a stable noble gas configuration.
• The Lewis structure for CaO

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Covalent Lewis Structures

• Covalent bonds involve the sharing of electrons.
• Covalent Lewis structures contain dots that count
  for the octet of more than one atom.

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Types of Electron Pairs

• The electrons between two atoms are called
  bonding pairs.
• Electrons on a single atom are called lone pairs.
• Only bonding electrons count toward the octet of
  both atoms.

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Water

• The Lewis structure shows a 21 ratio of hydrogen
  to oxygen.
• H2O
• Observed in nature to have two hydrogens for
  every oxygen

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Multiple Bonds

• Sometimes multiple bonding pairs are necessary to
  complete the octets for each atom in the Lewis
  structure.
• Two bonding pairs are called a double bond. Three
  bonding pairs are a triple bond.

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Writing Lewis Structures

• Write the skeletal structure of the molecule.
• Determine the total number of electrons for the
  molecule.
• Place the electrons as dots to give octets to as
  many atoms as possible.
• If the central atom has not obtained an octet,
  form multiple bonds as necessary to complete its
  octet.

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Concept Check 5.2

• Draw the Lewis structure for the covalent
  molecule methane, CH4.

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Concept Check 5.2 Solution

• Draw the skeletal structure of the molecule.
• Count the number of valence electrons in CH4.
• C 4 valence electrons 1C 4 electrons
• H 1 valence electrons 4H 4 electrons
• 8 electrons total in structure

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Concept Check 5.2 Solution

• 3. Put a single bond between each C and H atom,
  using all eight of the valence electrons. Place
  the electrons as dots to give octets to as many
  atoms as possible. In this case, there are no
  unpaired electrons.

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Chemical Bonding in Ozone

• Ozone is an atmospheric gas that protects life on
  Earth from excessive exposure to UV light.

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Ozone Lewis Structure

• There appear to be two equally valid Lewis
  structures.
• The ozone molecule is actually best represented
  with two identical bonds, each one shorter than a
  single bond, but longer than a double bond.

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Ozone Resonance Structures

• Resonance in ozone is the averaging of two
  identical Lewis structures.
• Resonance structures are usually represented by a
  double-centered arrow between them.

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Oxygen vs. Ozone

• Molecular oxygens double bond is stronger than
  ozones bond and a half.
• UV light is not energetic enough to break
  oxygens strong bond.
• Oxygen does not absorb UV light.
• UV light will break one of ozones two bonds.
• Ozone absorbs UV light, effectively preventing
  much of the UV light from reaching the Earths
  surface.

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The Shapes of Molecules

• Molecular shape is an important factor in
  determining the properties of substances.
• Valence shell electron pair repulsion theory
  (VSEPR theory) allows us to predict molecule
  shapes from their Lewis structures.

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VSEPR Theory

• A simple model
• Based on the idea that centers of negative charge
  created by both bonding and lone pair electrons
  repel each other and determine the shape of the
  resulting molecule
• Minimization of repulsions (maximization of
  distance between like charges) determines
  three-dimensional molecular geometry.

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The Ideal Shape

• Three-dimensional geometry
• Greatest angle (minimum repulsions, maximum
  distance) between four bonding electron pairs on
  a central atom is tetrahedral.
• The combination of lone pairs and multiple bonds
  with bonded pairs makes possible other electron
  pair and molecular geometries.
• It is important to distinguish between electron
  geometry and molecular geometry.

In CH4, there are no lone pairs of electrons,
therefore, electron pair and molecular geometries
are the same.
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Water Electron vs. Molecular Geometry

• Electron geometry is tetrahedral.
• Molecular geometry is bent.

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Molecular Geometry
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Concept Check 5.3

• Use Lewis theory and VSEPR to predict the
  electron and molecular structure of BF3. Boron
  and other group 3A elements may have either six
  or eight electrons in their outer shells and
  still be relatively stable.

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Concept Check 5.3 Solution

• To draw the Lewis structure, first draw skeletal
  structure of the molecule.
• Count the number of valence electrons in BF3.
• B 3 valence electrons 1B 3 electrons
• F 7 valence electrons 3F 21 electrons
• 24 electrons total in structure

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Concept Check 5.3 Solution

• 3. Put a single bond between each B and F atom.
  Place the electrons as dots to give octets to as
  many atoms as possible. In this case, there are
  three lone pairs of electrons on each F.

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Concept Check 5.3 Solution

• 4. VSEPR can be applied to a correct Lewis
  structure to determine molecular geometry around
  the central atom. VSEPR is based on minimizing
  repulsions between electrons, either bonding or
  lone pairs, by maximizing the distance between
  them. BF3 has three bonding pairs of electrons
  and no lone pairs around B. A trigonal planar
  geometry best minimizes these repulsions.

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Bond Polarity

• The dots in covalently bonded Lewis structures
  appear to be equally shared but reality yields a
  less ideal picture.
• Unequal electron sharing between two elements
  (polar bonds) can result in polar molecules.
• Polar bonds arise whenever elements with
  different electron-attracting abilities form a
  bond. The symbols d- and d indicate partial
  negative and partial positive charges
  respectively.
• Electronegativity is a measure of an atoms
  ability to attract electrons in a bond.

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Electronegativity

• Electronegativity increases as you move right
  across a row on the periodic table and decreases
  as you move down a column.
• A polar bond is analogous to a bar magnet its
  uneven electron distribution results in a
  negative pole and a positive pole.
• Similarly, an entire molecule may be polar if
  uneven electron distribution within the molecule
  results in a negative pole and a positive pole.

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Examples

• Carbon monoxide, CO, is a polar molecule.
• Carbon dioxide, CO2, contains polar bonds but is
  a nonpolar molecule.
• VSEPR theory predicts a molecular shape for CO2
  in which the polar bonds cancel each other.

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Water

• Since the water molecule is polar, attractions
  between water molecules tend to hold them
  together (high boiling point).
• Molecules arrange themselves to maximize forces
  of attraction between them and minimize forces of
  repulsion (lower density when freezes).
• Water is immiscible with nonpolar molecules.

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Polar or Nonpolar?

• Look at the Lewis structure and ask two
  questions.
• Does the molecule contain polar bonds?
• Do the polar bonds together give overall polarity
  to the molecule?

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Polar or Nonpolar?

• Look at the Lewis structure and ask two
  questions.
• Does the molecule contain polar bonds?
• Do the polar bonds together give overall polarity
  to the molecule?

Yes, the CF bond is polar. No, all of the
individual CF bond dipoles cancel each other so
overall, the molecule is nonpolar.
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Concept Check 5.4

• Is BF3 a polar or nonpolar molecule

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Concept Check 5.4 Solution

• The first influence to consider is the polarity
  of the bonds. Are BF bonds polar? Referring to
  Figure 5-4, using the difference between the
  electronegativity values for B and F we get
  4.0(F) 2.0(B) 2.0.
• The geometry of the polar bonds determines
  whether the molecule is polar or not. Referring
  back to Concept Check 5.3, BF3 is trigonal
  planar.
• Table 5.2 indicates that a trigonal planar
  structure with identical polar bonds is nonpolar,
  therefore, BF3 is nonpolar.

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Chapter Summary

• Molecular Concept
• Ionic bonds
• Covalent bonds
• Lewis theory
• Valence shell electron pair repulsion (VSEPR)
• Polarity

• Societal Impact
• Life would be impossible without compounds.
• Chemical bonding has allowed us to make products
  such as plastic.
• The shapes of molecules determine their
  properties.