Chemical%20Bonding - PowerPoint PPT Presentation

About This Presentation
Title:

Chemical%20Bonding

Description:

Chemical Bonding – PowerPoint PPT presentation

Number of Views:147
Avg rating:3.0/5.0
Slides: 48
Provided by: Mrs1172
Category:

less

Transcript and Presenter's Notes

Title: Chemical%20Bonding


1
  • Chemical Bonding

2
  • Chemical Bond
  • attractive force between atoms or ions that binds
    them together as a unit
  • bonds form in order to
  • decrease potential energy (PE)
  • increase stability

3
COMPOUND
more than 2 elements
2 elements
Ternary Compound
Binary Compound
NaNO3
NaCl
4
ION
1 atom
2 or more atoms
Polyatomic Ion
Monatomic Ion
NO3-
Na
5
Types of Bonds
COVALENT
IONIC
e- are transferred from metal to nonmetal
e- are shared between two nonmetals
Bond Formation
Type of Structure
true molecules
crystal lattice
Physical State
liquid or gas
solid
Melting Point
low
high
Solubility in Water
yes
usually not
yes (solution or liquid)
Electrical Conductivity
no
Other Properties
odorous
6
Types of Bonds
METALLIC
e- are delocalized among metal atoms
Bond Formation
Type of Structure
electron sea
Physical State
solid
Melting Point
very high
Solubility in Water
no
yes (any form)
Electrical Conductivity
malleable, ductile, lustrous
Other Properties
7
Ionic Bonds
8
Ionic Bonding - Crystal Lattice
9
Covalent Bonding - True Molecules
Diatomic Molecule
10
Metallic Bonding - Electron Sea
11
Bond Polarity
  • Most bonds are a blend of ionic and covalent
    characteristics.
  • Difference in electronegativity determines bond
    type.

12
Bond Polarity
  • Electronegativity
  • Attraction an atom has for a shared pair of
    electrons.
  • higher e-neg atom ? ?-
  • lower e-neg atom? ?

13
Bond Polarity
  • Electronegativity Trend (p. 151)
  • Increases up and to the right.

14
Bond Polarity
  • Nonpolar Covalent Bond
  • e- are shared equally
  • symmetrical e- density
  • usually identical atoms

15
  • Polar Covalent Bond
  • e- are shared unequally
  • asymmetrical e- density
  • results in partial charges (dipole)

16
  • Nonpolar
  • Polar
  • Ionic

17
Bond Polarity
  • Examples
  • Cl2
  • HCl
  • NaCl

3.0-3.00.0 Nonpolar 3.0-2.10.9 Polar 3.0-0.92.1
Ionic
18
  • Chemical Bond
  • attractive force between atoms or ions that binds
    them together as a unit
  • bonds form in order to
  • decrease potential energy (PE)
  • increase stability

19
Lewis Diagrams
  • Molecular Structure

20
Rule
  • Remember
  • Most atoms form bonds in order to have 8 valence
    electrons.

21
A. Octet Rule
  • Exceptions
  • Hydrogen ? 2 valence e-
  • Groups 1,2,3 get 2,4,6 valence e-
  • Expanded octet ? more than 8 valence e- (e.g. S,
    P, Xe)
  • Radicals ? odd of valence e-

22
B. Drawing Lewis Diagrams
  • Find total of valence e-.
  • Arrange atoms - singular atom is usually in the
    middle.
  • Form bonds between atoms (2 e-).
  • Distribute remaining e- to give each atom an
    octet (recall exceptions).
  • If there arent enough e- to go around, form
    double or triple bonds.

23
B. Drawing Lewis Diagrams
  • CF4

1 C 4e- 4e- 4 F 7e- 28e- 32e-
F F C F F
- 8e- 24e-
24
B. Drawing Lewis Diagrams
  • BeCl2

1 Be 2e- 2e- 2 Cl 7e- 14e- 16e-
  • Cl Be Cl

- 4e- 12e-
25
B. Drawing Lewis Diagrams
  • CO2

1 C 4e- 4e- 2 O 6e- 12e- 16e-
  • O C O

- 4e- 12e-
26
C. Polyatomic Ions
  • To find total of valence e-
  • Add 1e- for each negative charge.
  • Subtract 1e- for each positive charge.
  • Place brackets around the ion and label the
    charge.

27
C. Polyatomic Ions
  • ClO4-

1 Cl 7e- 7e- 4 O 6e- 24e- 31e-
O O Cl O O
1e- 32e-
- 8e- 24e-
28
C. Polyatomic Ions
  • NH4

1 N 5e- 5e- 4 H 1e- 4e- 9e-
H H N H H
- 1e- 8e-
- 8e- 0e-
29
D. Resonance Structures
  • Molecules that cant be correctly represented by
    a single Lewis diagram.
  • Actual structure is an average of all the
    possibilities.
  • Show possible structures separated by a
    double-headed arrow.

30
D. Resonance Structures
  • SO3

31
Molecular Geometry
32
VSEPR Theory
  • Valence Shell Electron Pair Repulsion Theory
  • Electron pairs orient themselves in order to
    minimize repulsive forces.

33
VSEPR Theory
  • Types of e- Pairs
  • Bonding pairs - form bonds
  • Lone pairs - nonbonding e-

34
VSEPR Theory
  • Lone pairs reduce the bond angle between atoms.

35
Determining Molecular Shape
  • Draw the Lewis Diagram.
  • Tally up e- pairs on central atom.
  • double/triple bonds ONE pair
  • Shape is determined by the of bonding pairs and
    lone pairs.

36
Common Molecular Shapes
  • 2 total
  • 2 bond
  • 0 lone

LINEAR 180
37
Common Molecular Shapes
  • 3 total
  • 3 bond
  • 0 lone

TRIGONAL PLANAR 120
38
Common Molecular Shapes
  • 3 total
  • 2 bond
  • 1 lone

BENT lt120
39
Common Molecular Shapes
  • 4 total
  • 4 bond
  • 0 lone

TETRAHEDRAL 109.5
40
Common Molecular Shapes
  • 4 total
  • 3 bond
  • 1 lone

TRIGONAL PYRAMIDAL 107
41
Common Molecular Shapes
  • 4 total
  • 2 bond
  • 2 lone

BENT 104.5
42
Common Molecular Shapes
  • 5 total
  • 5 bond
  • 0 lone

TRIGONAL BIPYRAMIDAL 120/90
43
Common Molecular Shapes
  • 6 total
  • 6 bond
  • 0 lone

OCTAHEDRAL 90
44
Examples
  • PF3

4 total 3 bond 1 lone
TRIGONAL PYRAMIDAL 107
45
Examples
  • CO2

2 total 2 bond 0 lone
LINEAR 180
46
(No Transcript)
47
(No Transcript)
Write a Comment
User Comments (0)
About PowerShow.com