Chapter 9 Bonding and Molecular Structure: Fundamental Concepts

1
Chapter 9Bonding and Molecular
StructureFundamental Concepts

• Bond Formation
• Ionic Bonds
• Lewis Structures
• Molecular Shapes
• Polarity
• More Lewis Structures

2
I. Bond Formation

• Helpful picture Lewis Symbols
• show valence electrons in main group elements (IA
  VIIIA)
• General goal of a chemical bond
• Give an atom a noble gas electron configuration
  (a complete octet of electrons)
• Two ways to do this
• Transfer electrons from one atom to another
  (ionic bonding)
• Share electrons between atoms (covalent bonding)

3
II. Ionic Bonds

• Ion Pair Energy
• Strength of an ionic bond base on
• magnitude of ion charges (n and n-)
• distance between ion centers (d)

4

• Lattice Energy
• Now instead of an ion pair, make a complete
  crystal
• ?Elattice follows same trends as ?Eion pair
• ?Elattice cannot be measured, only calculated
• Use Born-Haber Cycle (application of Hesss
  Law) Considers all the steps that go to into
  making an ionic compound.

5
III. Lewis Structures

• Lewis structures show valence electrons and
  covalent bonds in a molecule
• Octet Rule Atoms in molecules and polyatomic
  ions tend to be surrounded by 8 electrons

6

• General Steps in Writing Lewis Structures
• Draw a skeletal structure of the molecule.
• Which atom is the central atom?
• central atom is often the element farthest to
  the left and/or lowest in the periodic table.
• central atom is often the element in the compound
  for which there is only one atom.
• hydrogen is never the central atom.
• Count the total number of valence electrons in
  all the atoms. Make sure to add/subtract
  electrons if the molecule is a polyatomic ion.
• Draw single bonds between the central atom and
  the surrounding atoms.

7

• Steps in Writing Lewis Structures (cont.)
• Fill the octets of the atoms bonded to the
  central atom.
• Remember that hydrogen only needs 2 electrons in
  its outer shell, not 8!
• Put leftover electrons on the central atom as
  lone pairs.
• Atoms in and beyond the third period may have
  an expanded octet (i.e., more than 8 electrons in
  their outer shell)!
• If central atom doesnt have an octet, make
  double or triple bonds.
• Remember that atoms in Groups IIA and IIIA
  dont need an octet.

8
IV. Molecular Shapes

• Valence Shell Electron Pair Repulsion (VSEPR)
  Model
• used to predict molecular shapes
• Two Keys
• Electron pairs repel each other
• Electron pairs want to stay as far away from each
  other as possible
• In VSEPR

Electron pair 1 bond pair (single, double, or
triple bond) OR 1 lone pair
9
(No Transcript)
10
(No Transcript)
11
Bent
Sourcehttp//www.molecules.org/Images/animations/
NO2_.html Example Nitrite (NO2) 2 bond pairs,
1 lone pair
12
(No Transcript)
13
(No Transcript)
14
(No Transcript)
15
(No Transcript)
16
(No Transcript)
17
(No Transcript)
18
(No Transcript)
19
(No Transcript)
20
(No Transcript)
21
(No Transcript)
22
V. Polarity

• Polar Covalent Bond
• unequal sharing of electrons results in atoms
  that are partially charged
• Electronegativity
• measure of an atoms ability to draw electrons
  toward itself in a chemical bond

23
(No Transcript)
24

• Polar Molecule
• unequal distribution of electrons results in a
  negatively charged end and a positively charged
  end
• lines up in a magnetic field

25
VI. More Lewis Structures

• Bond Order
• number of bonding electron pairs shared by two
  atoms in a molecule
• Formal Charge
• compares the number of valence electrons for an
  atom in a Lewis structure with the free atom
• sum of formal charges charge on molecule
• Always try to minimize formal charge in a Lewis
  structure
• Always place negative formal charge on the most
  electronegative element