Chemical Bonding and Molecular Architecture

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Chemical Bonding and Molecular Architecture

• Structure and Shapes of Chemicals

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Bonds

• Forces that hold groups of atoms together and
  make them function as a unit.

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Bond Energy

• It is the energy required to break or released in
  making a bond.
• It gives us information about the strength of a
  bonding interaction.
• Ionic bondsstrong attractions between
  oppositely charged ions
• Covalent bondsattraction between non-metal
  atoms as both atoms share electrons

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Bond Length

• The distance where the system energy is a
  minimum.

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Ionic Bonds

• Formed from electrostatic attractions of closely
  packed, oppositely charged ions.
• Formed when an atom that easily loses electrons
  reacts with one that has a high electron affinity.

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Ionic Configuration and Size

• Ions are formed when electrons are gained or lost
  from an atom. The gain or loss follows the
  pattern called the octet rule, that an atom
  forms an ion in which it attains the same
  electron configuration as the nearest noble gas.
  Most metals therefore lose electrons, and as a
  result get smaller. The trend is greater ,
  smaller size.
• Likewise, nonmetals gain electrons to form ions,
  thus increasing in size by the opposite rule to
  metals.

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Isoelectronic Ions

• Ions containing the the same number of
  electrons, due to attaining the configuration of
  the same noble gas
• (O2?, F?, Na, Mg2, Al3)
• All attain to Ne
• O2??gt F? gt Na gt Mg2 gt Al3
• largest
  smallest

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Electronegativity

• The ability of an atom in a molecule to attract
  shared electrons to itself.
• Periodic trend increases across the table to the
  halogen column. Decreases down a group. Least
  at Cs (0.7), greatest at F (4.0).

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Polarity

• A molecule, such as HF, that has a center of
  positive charge and a center of negative charge
  is said to be polar, or to have a dipole moment.

Polar bonds shown as arrow with point toward
negative pole, toward the positive pole
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Electronegativity and Polarity of Bonds

• Subtract lower EN from higher
• EN Difference Ionic Character Type of
  Bond
• 0 0 Nonpolar Covalent
• 0.1-0.5 1-5 Slightly polar covalent
• 0.6-1.5 6-40 Polar Covalent
• gt 1.5 over 40 Ionic
• Compounds with over 50 ionic character are
  considered to be totally ionic solids. These
  compounds are often called salts.

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Achieving Noble Gas Electron Configurations (NGEC)

• Two nonmetals react They share electrons to
  achieve NGEC.
• A nonmetal and a representative group metal
  react (ionic compound) The valence orbitals of
  the metal are emptied to achieve NGEC. The
  valence electron configuration of the nonmetal
  achieves NGEC.

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Binary Ionic--Lattice Energy

• The change in energy when separated gaseous ions
  are packed together to form an ionic solid.
• M(g) X?(g) ? MX(s)
• Lattice energy is negative (exothermic) from the
  point of view of the system.

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Formation of an Ionic Solid

• 1. Sublimation of the solid metal
• M(s) ? M(g) endothermic
• 2. Ionization of the metal atoms
• M(g) ? M(g) e? endothermic
• 3. Dissociation of the nonmetal
• 1/2X2(g) ? X(g) endothermic
• 4. Formation of X? ions in the gas phase
• X(g) e? ? X?(g) exothermic
• 5. Formation of the solid MX
• M(g) X?(g) ? MX(s) quite exothermic

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Homework!!

• p. 395ff 11, 14, 15, 20

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Covalent Chemical Bonds

• Happen when collections of atoms are more stable
  than the separate atoms. They provide a method
  for dividing up energy when stable molecules are
  formed from atoms.
• Covalent bonds are due to shared electron pairs.
  One pair shared is a single bond, two makes a
  double bond, three make a triple bond.
• As bond order increases (single, double, triple),
  bond length shortens

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Bond Energies

• Bond breaking requires energy (endothermic).
• Bond formation releases energy (exothermic).
• ?H ?D(bonds broken) ? ?D(bonds formed)

energy required
energy released
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Example 8.5
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Localized Electron Model

• A molecule is composed of atoms that are bound
  together by sharing pairs of electrons using the
  atomic orbitals of the bound atoms.
• Two types of electron pairs bonding pairs and
  lone pairs. Bonding pairs are linkages between
  atoms, lone pairs are electrons solely owned by
  an atom.

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Localized Electron Model

• Elements of the Model
• 1. Description of valence electron arrangement
  (Lewis structure).
• 2. Prediction of geometry (VSEPR model).
• 3. Description of atomic orbital types used to
  share electrons or hold lone pairs.

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Lewis Structure

• Shows how valence electrons are arranged among
  atoms in a molecule.
• Reflects central idea that stability of a
  compound relates to noble gas electron
  configuration.

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Comments About the Octet Rule

• 2nd row elements C, N, O, F observe the octet
  rule.
• 2nd row elements B and Be often have fewer than 8
  electrons around themselves - they are very
  reactive.
• 3rd row and heavier elements CAN exceed the octet
  rule using empty valence d orbitals.
• When writing Lewis structures, satisfy octets
  first, then place electrons around elements
  having available d orbitals.

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Rules for Drawing Lewis Structures

• Add up all of the valence electrons for the atoms
  involved in the molecule. In polyatomic ions,
  subtract electrons for a charge, add for a -
  charge
• Select a most likely central atom and arrange
  other atoms around it. Place pairs of electrons
  between atoms.
• Arrange the remaining electrons around external
  atoms first. If the central atom is not
  satisfied, form double or triple bonds to make
  the molecule work.

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Example 8.6
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Resonance

• Occurs when more than one valid Lewis structure
  can be written for a particular molecule.
• These are resonance structures. The actual
  structure is an average of the resonance
  structures.

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Example 8.9
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Homework 8b

• p. 397ff 31, 36, 39, 42, 50, 57

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Molecular Architecture

• The structure of a molecule is important in how
  it reacts and to its physical properties
• Once the Lewis structure of a molecule is
  determined, the shape of the molecule then can be
  predicted according to the VSEPR model.

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VSEPR Model

• The structure around a given atom is determined
  principally by minimizing electron pair
  repulsions.

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Predicting a VSEPR Structure

• 1. Draw Lewis structure.
• 2. Count pairs, both bonding and lone pairs
  around the central atom.
• 3. Determine positions of atoms from the way
  electron pairs are shared.
• 4. Determine the name of molecular structure from
  the number of bonding and lone pairs and their
  necessary arrangements. Remember that lone pairs
  prefer to be at 120º or greater from each other.

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Sample 8.12
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Sample 8.13
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Sample 8.14
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Homework 8c

• p. 399ff 59, 62, 73, 78, 79, 91

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Hybridization

• The mixing of atomic orbitals to form special
  orbitals for bonding.
• The atoms are responding as needed to give the
  minimum energy for the molecule.
• To determine hybridization, count lone and
  bonding pairs, but count multiple bonds only
  once.

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• A sigma (?) bond centers along the internuclear
  axis.
• A pi (?) bond occupies the space above and below
  the internuclear axis.

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The Localized Electron Model

• Draw the Lewis structure(s)
• Determine the arrangement of electron pairs
  (VSEPR model).
• Specify the necessary hybrid orbitals.

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Sample 9.1,2
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Sample 9.3,4,5
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Molecular Orbitals (MO)

• Analagous to atomic orbitals for atoms, MOs are
  the quantum mechanical solutions to the
  organization of valence electrons in molecules.

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Types of MOs

• bonding lower in energy than the atomic
  orbitals from which it is composed.
• antibonding higher in energy than the atomic
  orbitals from which it is composed.

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Bond Order (BO)

• Difference between the number of bonding
  electrons and number of antibonding electrons
  divided by two.

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Find bond order and magnetic properties for
He22, F2-2, O2-
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Paramagnetism

• unpaired electrons
• attracted to induced magnetic field
• much stronger than diamagnetism

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Outcomes of MO Model

• 1. As bond order increases, bond energy increases
  and bond length decreases.
• 2. Bond order is not absolutely associated with a
  particular bond energy.
• 3. N2 has a triple bond, and a correspondingly
  high bond energy.
• 4. O2 is paramagnetic. This is predicted by the
  MO model, not by the LE model, which predicts
  diamagnetism.

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Combining LE and MO Models

• ? bonds can be described as being localized.
• ? bonding must be treated as being delocalized.

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Homework 9

• p. 432ff 5, 8, 11, 17, 18, 22, 25, 33, 37