Unit 6: Chemical Bonding - PowerPoint PPT Presentation

1 / 48
About This Presentation
Title:

Unit 6: Chemical Bonding

Description:

Unit 6: Chemical Bonding Refer to Ch. 8 & 9 for supplemental reading. – PowerPoint PPT presentation

Number of Views:256
Avg rating:3.0/5.0
Slides: 49
Provided by: rach2162
Category:

less

Transcript and Presenter's Notes

Title: Unit 6: Chemical Bonding


1
Unit 6 Chemical Bonding
  • Refer to Ch. 8 9 for supplemental reading.

2
  • 1. Chemical Bond an attractive force that holds
    2 atoms together
  • 3 types ionic, covalent, metallic

3
Valence Electrons
  • The electrons in the outer energy level of an
    atom.
  • They are like the front lines of an army. They
    are the electrons involved in bonding.

4
Review
  • How do you find valence electrons?
  • Hint there are two ways!
  • Examples Look at the group for the
    representative elements.
  • Mg ___
  • O ___
  • Ar ___
  • Si ___
  • Examples Write the electron configuration.
  • Mg
  • O

2
6
8
4
1s2 2s2 2p6 3s2
1s2 2s2 2p4
5
Electron Dot Structures
  • Depicts element symbol w/ valence e- shown as
    dots.

Na
Mg
Al
Si
Cl
Ar
O
N
6
Ionic Bonds
  • Occurs when ions of opposite charge (,-) attract
    each other.
  • Metal ion Nonmetal ion
  • Simplest attraction
  • NaCl MgF2
  • Polyatomic ions Molecules that are charged.
  • Al(PO4) (NH4)2(SO4)
  • PO4 is phosphate with a -3 charge
  • SO4 is sulfate with a -2 charge

7
Formation of Ionic Bond
  • Cation- positive ion () (Cats have paws)
  • Forms when a metal atom loses e- to become
    stable.
  • Anion- negative ion (-) (A Negative Ion)
  • Forms when a nonmetal atom gains e- to become
    stable
  • An ionic bond is formed when e- are transferred
    between atoms and the resulting ions stick
    together.

8
Examples
  • Formation of NaCl
  • Na Cl ? Na Cl - NaCl
  • Formation of MgF2
  • Mg F F ? Mg2 F - F -
    MgF2
  • How would a compound form between two aluminum
    and three oxygen?

9
Electron Configuration of Ions
  • Cation example (metal)
  • Ca atom 1s2 2s2 2p6 3s2 3p6 4s2
  • Ca2 ion 1s2 2s2 2p6 3s2 3p6 Lost 2 electrons
    to obtain noble gas configuration (octet)

10
Electron Configuration of Ions
  • Anion example (nonmetal)
  • N atom 1s2 2s2 2p3
  • N3- ion 1s2 2s2 2p6
  • Gained 3 electrons to obtain noble gas
    configuration (octet)

11
Properties of Ionic Compounds
IONIC
The force that holds ionic compounds is really
strong. So you see all of these characteristics.
e- transferred from metal to nonmetal
Bond Formation
Type of Structure
Crystal lattice
Physical State
Solid (hard and rigid)
Melting Point
high
Boiling Point
high
yes (solution or liquid)
Electrical Conductivity
Other Properties
Hard, rigid and brittle
12
Electrolyte
  • A substance that conducts electricity
  • Because of ionic bonds ionic (charged) nature,
    ionic compounds conduct electricity in the molten
    or aqueous (dissolved in water) forms.

13
Covalent Bonds
  • Occurs when 2 nonmetals share pairs of electrons
    to become stable. These are called molecular
    compounds.
  • Examples
  • H2O CO2 C6H12O6 PCl5
  • Notice all of these elements in the molecule are
    nonmetals!

14
Covalent Bonds
  • Covalent bonds can be
  • single (1 shared pair)
  • double (2 shared pairs)
  • or triple (3 shared pairs)
  • Bond strength triple gt double gt single
  • Bond length single gt double gt triple

15
Creating Lewis Structures
  • Follow this system
  • Example H2O
  • 1) Draw a skeleton of the molecule. It
    generally works to place the different atom in
    the center.
  • H O H

16
Creating Lewis Structures
  • Find the needed electrons (N) for each atom and
    add them up. N will be 8 for most elements, with
    these exceptions
  • N 12
  • H 2
  • O 8
  • H 2
  • H gets 2 valence e-
  • Be gets 4 valence e-
  • B gets 6 valence e-

17
8
2
NEED
4
8
8
8
8
6
8
18
  • 3) Find the available (valence) electrons (A)
    for each atom and then add them up.
  • A
  • special note when completing a Lewis structure
    for a polyatomic ion, you will need to correct A
    by adding the absolute value of the charge if
    negative, and subtracting the charge if positive.
    For example, for the ion PO43-, you would add 3
    to A. For the ion NH4, you would subtract 1
    from A. (You do the opposite of the charge.)

H 1 O 6 H 1 Total A 8
N 12 A 8
19
8
1
Available
2
4
5
6
7
3
Figure out through e config
20
  • 4) Find the shared (S) electrons for the entire
    molecule by this formula S N A
  • S

S 12 8 4
N 12 A 8 S 4
21
  • 5) The shared electrons are the bonding
    electrons. Place all of the shared electrons
    between the atoms.
  • H O H
  • 6) You must place all of the available (A)
    electrons in the picture. The shared electrons
    are part of the available. See how many of the
    available electrons still need to be placed, and
    put them in the picture as lone pairs (unshared
    pairs) so that every atom gets an octet (remember
    H only needs 2).
  • H O H

N 12 A 8 S 4
N 12 A 8 S 4
4
22
(No Transcript)
23
(No Transcript)
24
(No Transcript)
25
  • CF4

F F C F F
N A S
8(4x8) 40
4(4x7) 32
24
8
26
Cl Be Cl
  • BeCl2

N A S
4(2x8) 20
2(2x7) 16
12
4
27
  • CO2

O C O
N A S
8(2x8) 24
4(2x6) 16
8
8
28
Polyatomic Ions
  • To find total of valence e- (A)
  • Add 1e- for each negative charge.
  • Subtract 1e- for each positive charge.
  • Place brackets around the ion and label the
    charge.

29
Polyatomic Ions
  • ClO4-

N A S
8(4x8) 40
O O Cl O O
7(4x6) 31
1 32
24
8
30
Bond Polarity
View Bonding Atomic Bonding Chemistry
ZoneAnimations.
  • Nonpolar
  • Polar
  • Ionic

31
Bond Polarity
  • Most bonds are a blend of ionic and covalent
    characteristics.
  • Difference in electronegativity determines bond
    type. (subtract)

If DEN is Bond type is
lt 0.4 Nonpolar covalent
0.4 lt DEN lt 1.7 Polar covalent
gt 1.7 Ionic
32
Bond Polarity
  • Nonpolar Covalent Bond
  • e- are shared equally between the atoms
  • symmetrical e- density
  • usually occurs between identical atoms

33
Bond Polarity
  • Polar Covalent Bond
  • e- are shared unequally
  • One atom hogs the electrons
  • asymmetrical e- density
  • results in partial charges (dipole)

O is more electronegative than H and pulls the
electrons closer to itself.
34
Electronegativity Chart
If ?EN is Bond type is
lt 0.4 Nonpolar covalent
0.4 lt ? EN lt 1.7 Polar covalent
gt 1.7 Ionic
Use the Electronegativity Chart to determine if
the bond between atoms is nonpolar covalent,
polar covalent, or ionic. Larger EN minus the
smaller EN
35
If ?EN is Bond type is
lt 0.4 Nonpolar covalent
0.4 lt ? EN lt 1.7 Polar covalent
gt 1.7 Ionic
  • Examples
  • Cl2
  • HCl
  • NaCl

3.16-3.160.0 Nonpolar 3.16-2.20.96 Polar 3.16-..
932.23 Ionic
36
Metallic Bond
  • Bond between metal and metal.
  • An electron sea is created where electrons
    overlap into neighboring atoms.
  • The electrons move around.


37
Metallic Bond
  • Because the electrons are free to move around
    from atom to atom, we see the properties that we
    know of metals.

http//www.pbs.org/wgbh/nova/tech/structure-of-met
al.html Click on this for a cool interactive on
metal properties.
38
Metal Properties
  • Malleable and ductile the electrons can move
    past each other so the shape can change.
  • Free flowing electrons can conduct heat and
    electricity quickly to other atoms.

39
VSEPR Theory
  • Valence Shell Electron Pair Repulsion Theory
  • Electron pairs place themselves so that they are
    as far apart from each other as possible.

40
A. VSEPR Theory
  • Types of e- Pairs
  • Bonding pairs - form bonds between the atoms
  • Lone pairs - nonbonding e- (electrons that are
    not between atoms)

41
A. VSEPR Theory
  • Lone pairs reduce the bond angle between atoms.

VS
42
B. 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.

43
Common Molecular Shapes
2 bonding pairs 0 lone pairs
LINEAR
44
Common Molecular Shapes
2 bonding pairs 1 lone pair
BENT 120
45
Common Molecular Shapes
2 bonding pairs 2 lone pairs
BENT 109.5
46
Common Molecular Shapes
3 bonding pairs 0 lone pairs
TRIGONAL PLANAR
47
Common Molecular Shapes
3 bonding pairs 1 lone pair
PYRAMIDAL
48
Common Molecular Shapes
4 bonding pairs 0 lone pairs
TETRAHEDRAL
49
Examples
  • PF3

3 bond 1 lone
PYRAMIDAL
50
Examples
  • CO2

2 total 2 bond 0 lone
LINEAR 180
51
Examples
  • H2S CCl4
  • BF3 SiO2
Write a Comment
User Comments (0)
About PowerShow.com