Covalent Bonding

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Covalent Bonding

• Chapter 16

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Single Covalent Bonds

• A single covalent bond is one in which two atoms
  share a pair of electrons.
• Structural formulas are chemical formulas that
  show the arrangement of atoms in molecules and
  polyatomic ions.
• Chemical Formula Structural Formula
• H2 H-H
• H2O
• NH3
• CH4

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Single Covalent Bonds

• Unshared pairs (lone pairs) are pairs of valence
  electrons that are not shared between atoms.
• When carbon forms bonds with other atoms, it
  usually forms four bonds.
• CH4
• C2H6

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Double Triple Covalent Bonds

• Double covalent bonds share two pairs of
  electrons.
• CO2 OCO
• Triple covalent bonds share three pairs of
  electrons.
• N2 NN
• See table 16.1 p. 443

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Coordinate Covalent Bonds

• Two nuclei are attracted to a lone pair of
  electrons (e-)
• A bond in which one atom shares an entire pair of
  electrons (lone pair) with another atom.
• Typically forms polyatomic ions
• In structural formulas, you can show coordinate
  covalent bonds as arrows that point from the atom
  donating the pair of electrons to the atom
  receiving them.

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Bond Dissociation Energies

• Bond dissociation energy is the total energy
  required to break the bond between two covalently
  bonded atoms.
• A typical carbon-carbon single bond has a bond
  dissociation energy of 347 kJ. The ability of
  carbon to form strong bonds helps to explain its
  stability.
• They are unreactive partly because the
  dissociation energy of these bonds is high.
• See table 16.3 p. 448

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Resonance

• Resonance structures are structures that occur
  when it is possible to write 2 or more valid
  electron dot formulas that have the same of
  electron pairs for a molecule or ion.
• Electrons do NOT change position
• Bond lengths are an average of possible bond
  lengths.
• Bonds become a hybrid of all the possible bonds.
• Increases stability within the atom

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Exceptions to the Octet

• Some molecules exist with an odd of e-
• NO2 has 7 valence e-
• Paramagnetic containing 1 or more unpaired e-
• Appear heavier in magnetic fields
• Diamagnetic contain only paired e-
• O2
• Recently discovered its paramagnetic
• Suggesting unpaired e- but double bonds make
  senseResonance is suggested

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VSEPR Theory

• Valence Shell Electron Pair Repulsion Theory
• Because electrons repel, molecules adjust shape
  so the electron pairs are as far apart as
  possible.
• Unshared pairs of e- are also important when
  trying to predict shape.
• Methane forms a tetrahedron with angles of 109.5O
  (tetrahedral angle).

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VSEPR Theory (cont)

• Some common shapes linear triatomic, trigonal
  planar, bent triatomic, pyramidal, tetrahedral,
  trigonal bipyramidal

• Linear triatomic

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• Trigonal planar
• Formaldehyde

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• Water
• 105o

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• Ammonia
• 107o

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• Tetrahedral
• Methane

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• Trigonal bipyramidal

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Polar Bonds

• When e- pairs are not shared equally between
  atoms
• Due to differences in electronegativity
• Polar covalent bond an uneven sharing of e-
  between 2 nonmetals
• Non-polar covalent bond an even sharing of e-
  between 2 nonmetals
• See table 16.4 p. 462

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Polar Bonds (cont)

• The lower case Greek letter delta shows that
  atoms involved in the covalent bond acquire only
  partial charges, much less than 1 or 1-
• The polarity of a bond may also be represented
  with an arrow pointing to the more
  electronegative atom

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Polar Molecules

• In a polar molecule, one end of the molecule is
  slightly negative and the other end is slightly
  positive.
• The electrically charged regions are called
  poles.
• A molecule that has two poles is called a dipolar
  molecule or dipole.
• The effect of polar bonds on the polarity of a
  molecule depends on its shape.
• CO2
• H2O

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Polar or Non-Polar Molecule?
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Despite having two polar covalent bonds, carbon
dioxide is non-polar. The reason for this is the
molecule's linear geometry The electrons are
pulled in equal but opposite directions, causing
them to cancel one another.
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Water has two polar covalent bonds, but the polar
covalent bonds pull electrons to the oxygen's
region of the molecule, to make the molecule
polar. The reason for this is the molecule's bent
geometry The electrons are not pulled in equal
and opposite directions, despite the bonds being
identical.
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Attractions Between Molecules

• There are several kinds of molecular attractions.
• The weakest attractions are collectively called
  van der Waals forces, which are named after the
  chemist who discovered them, Johannes van der
  Waals.
• Van der Waals forces consist of two types
  dispersion forces and dipole interactions.

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Van der Waals Forces

• Dispersion forces are the weakest of all
  molecular interactions and they are caused by the
  motion of electrons.
• The strength of dispersion forces increases as
  the of electrons increases
• Dipole interactions occur when polar molecules
  are attracted to one another.
• The slightly negative region of a polar molecule
  is attracted to the slightly positive region of
  another polar molecule.
• Similar but much weaker than ionic bonds

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Hydrogen Bonds

• Hydrogen bonds are attractive forces in which a
  hydrogen covalently bonded to a very
  electronegative atom is also weakly bonded to an
  unshared electron pair of another electronegative
  atom.

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Hydrogen Bonds

• Usually the electronegative atom is oxygen,
  nitrogen, or fluorine, which has a partial
  negative charge. The hydrogen then has the
  partial positive charge.
• Hydrogen bonding is usually stronger than normal
  dipole forces between molecules.
• Because oxygen has two lone pairs, two different
  hydrogen bonds can be made to each oxygen.

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• http//www.northland.cc.mn.us/biology/Biology1111/
  animations/hydrogenbonds.html

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Hydrogen Bonds in DNA