CHAPTER 8 COVALENT BONDS

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CHAPTER 8 COVALENT BONDS 8.1 Molecular
Compounds 8.2 The Nature of Covalent Bonding 8.3
Bonding Theories 8.4 Polar Bonds and Molecules
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• 8.1 Molecular Compounds
• Key Concepts
• How are melting points and boiling points of
  molecular compounds different from ionic
  compounds?
• What information does a molecular formula
  provide?

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• Molecules and Molecular Compounds
• Covalent bond occurs when two or more atoms
  share valence electrons.
• Molecule is a neutral group of atoms joined
  together by covalent bonds.
• Diatomic molecule is a molecule consisting of
  two atoms.
• Compound a substance that contains two or more
  elements chemically combined in a fixed
  proportion
• Molecular compounds a compound composed of
  molecules
• a. Have low melting points
• b. Have low boiling points
• c. Most are gas or liquid at room
  temperature
• d. Composed of two or more non-metals
• 6. Using page 214 illustrate some differences
  between ionic and covalent compounds.

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• Molecular Formula
• Molecular Formula the chemical formula of a
  molecular compound
• a. Describes how many of each atom a
  molecule contains
• b. Subscripts are used after the elements
  symbol to indicate the number of atoms of each
  element in the molecule.
• c. Reflects the actual number of atoms in
  each molecule and are not necessarily the lowest
  whole-number ratios.
• d. Can describe molecules consisting of one
  element.
• e. Does not tell you about the molecules
  structure
• 2. Using page 215 state the molecular formula for
  Ammonia and describe the types of diagrams and
  models used to represent Ammonia.

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• 8.2 The Nature of Covalent Bonding
• Key Concepts
• How does electron sharing occur in forming
  covalent bonds?
• How do electron dot structures represent shared
  electrons?
• How do atoms form double or triple covalent
  bonds?
• How are coordinate covalent bonds different other
  covalent bonds?
• How is the strength of a covalent bond related to
  its bond dissociation energy?
• How are oxygen atoms bonded in ozone?
• What are some exceptions to the octet rule?

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• The Octet Rule in Covalent Bonding
• In forming covalent bonds, electron sharing
  usually occurs so that atoms attain the electron
  configurations of noble gases.
• That is to say the valence electrons arrange
  themselves so that each atom sees an octet.
• Hydrogen has a noble gas configuration with 2
  electrons
• Groups four to seven are likely to form covalent
  bonds

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Single Covalent Bonds 1. Single Covalent Bond
Two atoms held together by sharing a pair of
electrons. 2. Hydrogen is an example. 3. An
electron dot structure can be used to show the
shared pair of electrons of the covalent bond. 4.
Using page 218 use electron dots to combine two
Fluorine atoms then show the electron
configuration for each atom. 5. Structural
Formula represents the covalent bonds by using
dashes, each dash represents one electron
pair. 6. Unshared Pair are electrons not shared
between atoms also called lone pair, nonbonding
pair. 7. Draw the electron dot structure for
ammonia (NH3) show the unpaired bonds and the
shared pairs properly.
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8. Now draw the structure for methane 9. Draw
the electron configuration for Carbon then using
p220 of the text explain why Carbon usually forms
four bonds.
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• Double and Triple Covalent Bonds
• Atoms sometime bond by sharing more than one pair
  of electrons.
• Double Covalent Bond Shares two pair of
  electrons
• Triple Covalent Bond Shares three pairs of
  electrons
• Try showing bonding Carbon Dioxide

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• Coordinate Covalent Bonds
• Is a covalent bond in which one atom contributes
  both bonding electrons
• In a structural formula, you can show coordinate
  covalent bonds as arrows that point from the atom
  donating the pair of electrons.
• Once formed, a coordinate covalent bond is like
  any other covalent bond.
• Most polyatomic cations and anions contain both
  covalent and coordinate covalent bonds.
• Compounds containing polyatomic ions include both
  ionic and covalent bonding.
• Polyatomic ions have charge in order to satisfy
  the octet rule for each atom present in the
  group.
• Show the coordinate covalent bond of Carbon
  Monoxide.
• Show the formation of the Ammonium ion.
• Show the formation of Sulfate.

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• 10. Using page 224 of your text, show the
  chemical and structural formula for the following
  Molecular Compounds.
• Nitrous Oxide
• Sulfur Trioxide
• Hydrogen Fluoride
• Nitric Oxide
• Hydrogen Peroxide
• Nitrogen Dioxide
• Hydrogen Cyanide
• Hydrogen Chloride
• Sulfur Dioxide
• 11. The electron dot structure for a neutral
  molecule contains the same number of electrons as
  the total number of valence electrons in the
  combining atoms.
• 12. The negative charge of a polyatomic ion shows
  the number of electrons in addition to the
  valence electrons.
• 13. Because a negatively charged polyatomic ion
  is part of an ionic compound, the positive charge
  of the cation of the compound balances these
  additional electrons.

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• Bond Dissociation Energies
• The energy required to break the bond between two
  covalently bonded atoms.
• Usually expressed as the energy needed to break
  one mole of bonds.
• A large bond dissociation energy corresponds to a
  strong covalent bond.
• High dissociation energies tend to create very
  stable compounds that tend to be chemically
  unreactive.
• Units are measured in kJ/mo1
• A mol is a chemical quantity of an element or
  compound in which there are 6.02x1023 atoms or
  molecules present.

Link
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• Resonance
• A structure that occurs when it is possible to
  draw two or more valid electron dot structures
  that have the same number of electron pairs for a
  molecule or ion.
• Although no back-and-forth changes occur, double
  headed arrows are used to connect resonance
  structures.
• Show the structural formation of ozone.

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• Exceptions to the Octet Rule
• The octet rule cannot be satisfied in molecules
  whose total number of valence electrons is an odd
  number. There are also molecules in which an
  atom has fewer, or more, than a complete octet of
  valence electrons.
• Draw two resonance structures for Nitrogen
  Dioxide
• Other odd number electron molecules are Chlorine
  Dioxide and Nitric Oxide.
• Several molecules with an even number of
  electrons, such as some compounds of Boron, also
  fail to follow the octet rule.
• Draw the structure for Boron Trifluoride and show
  the significance of it reacting with ammonia.
• A few atoms, Phosphorus and Sulfur, can have ten
  or twelve electrons instead of eight
• Draw the structure for Phosphorus Pentachloride
  and Sulfur Hexafluoride

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Exceptions to the Octet Rule There are three
classes of exceptions to the octet rule 1)
Molecules with an odd number of electrons
2) Molecules in which one atom has less than
an octet 3) Molecules in which one
atom has more than an octet.   Odd Number of
Electrons Few examples. Generally
molecules such as ClO2, NO, and NO2 have an odd
number of electrons. Less than an Octet Less
Molecules with less than an octet are
typical for compounds of Groups 1A, 2A, and
3A. Most typical example is BF3. More electrons
than an Octet This is the largest
class of exceptions. Atoms from the 3rd period
onwards can accommodate more than
an octet. Beyond the third period, the d-
orbitals are low enough in energy to
participate in bonding and accept the extra
electron density.
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• 8.3 Bonding Theories
• How are atomic and molecular orbitals related?
• How does VESPR theory help predict the shapes of
  molecules?
• In what ways is orbital hybridization useful in
  describing molecules?

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Hybridization of atomic orbitals
Quantum mechanical approaches by combining the
wave functions to give new wavefunctions are
called hybridization of atomic orbitals.
Hybridization has a sound mathematical fundation,
but it is a little too complicated to show the
details here. We can say that an imaginary
mixing process converts a set of atomic orbitals
to a new set of hybrid orbitals that are a
combination of the two overlaping orbitals.
At this level, we consider the following hybrid
orbitals spsp2sp3
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• Molecular Orbitals
• Molecular orbitals are created when two atoms
  combine by the overlap of each atoms atomic
  orbital creating an orbital that applies to the
  entire molecule.
• Each atomic orbital is full when it contains two
  electrons.
• Bonding Orbitals in covalent bonds two
  electrons are also required to fill a molecular
  orbital.
• Sigma Bonds are created when two atomic
  orbitals combine to form a molecular orbital that
  is symmetrical around the axis connecting two
  atomic nuclei
• hybridization of single bonds
• Two examples of sigma bonds
• Are H2 and F2

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5. Pi Bonds Are created by the side by side
overlap of p orbitals the bonding electrons
are most likely to be found in sausage-shaped
regions above and below the bond axis.
Hybridization of double bonds Atomic
orbitals of pi bonding overlap less than in
sigma bonding therefore, pi bonds tend to
be weaker than sigma bonds. Example Ethene
Simply sigma are single bonds pi
are double bonds
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6. Hybridization of triple bonds
Example acetylene
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double bonds 1 pi bond 1 sigma.triple bond2
pi bonds 1 sigma single bonds 1 sigma.
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VSEPR Valence Shell Electron-Pair Repulsion
Theory The repulsion theory between electron
pairs causes molecular shapes to adjust so that
the valence electron pairs stay as far apart as
possible creating three dimensional structures
Therefore, VSEPR diagrams are characterized by
the number of lone pair electrons (unshared
electron pairs) and the angles between the shared
pairs of electrons
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The AXE system
American general chemistry textbooks adopt the
excellent AXmEn system, where A is the central
atom, m the number of ligands X, and n the number
of nonbonded lone-pairs of electrons, E, about
the central atom.
methane, CH4, is AX4ammonia, H3N, is
AX3E1water, H2O, is AX2E2
Note that different AXmEn designations can give
rise to the same overall geometry or shape
AX2E1 and AX2E2 both give rise to bent or angular
geometriesAX2 and AX2E3 both give rise to linear
geometries
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The AXE system gives rise to a pattern, from
which the various atomic geometric shapes can be
determined/assigned
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A Couple of More Advanced Examples
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Valence Electron Pair Geometry Number of Orbitals Hybrid Orbitals
Linear 2 sp
Trigonal Planar 3 sp2
Tetrahedral 4 sp3
Trigonal Bipyramidal 5 sp3d
Octahedral 6 sp3d2
Hybrid Orbitals -Provides information about both
molecular bonding and molecular shape unlike
VSEPR theory that just deals with molecular
shape.
Valence Electron Pair Geometry Number of Orbitals Hybrid Orbitals
Linear 2 sp
Trigonal Planar 3 sp2
Tetrahedral 4 sp3
Trigonal Bipyramidal 5 sp3d
Octahedral 6 sp3d2