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Chemical Bonding

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Chemical Bonds. Defn force holding two atoms together. How are they formed? Atoms gain, lose, or share valence electrons. Why does bonding occur? – PowerPoint PPT presentation

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Title: Chemical Bonding


1
Chemical Bonding
2
Chemical Bonds
  • Compound are formed from chemically bound atoms
    or ions
  • Bonding only involves the valence electrons

3
Chemical Bonds
  • Defn force holding two atoms together
  • How are they formed?
  • Atoms gain, lose, or share valence electrons
  • Why does bonding occur?
  • Stability achieve octet rule

4
Electron Dot Structure
  • Shows valence electrons around atomic symbol

hydrogen
(group 1)
H



N


nitrogen
(group 5)




(group 7)
Cl

chlorine



5
Types of Chemical Bonds
  • 3 Types
  • covalent bond
  • ionic bond
  • metallic bond

6
Covalent Bond
  • Defn two atoms share one pair of electrons

Electrons shared

A

B

A
B

7
Covalent Bonds
  • Where are these bonds found?
  • - molecules (molecular compounds)
  • - polyatomic ions

8
Ionic Bond
  • Defn force holding cations and anions together


A

B

A
B-

Ionic bond
9
Ionic Bond
  • Where are these bonds found?
  • Ionic Compounds

10
Metallic Bonding
  • Defn attraction of metallic cations

Occurs only in metals
11
Covalent Bonding
  • Whats going on?
  • Molecule formed when 2 or more atoms bond
    covalently

Sharing of electrons
12
Two Types of Covalent Bonds
  • i) nonpolar covalent equal sharing of e-
  • ii) polar covalent UNequal sharing of e-

13
Nonpolar vs. Polar
NONPOLAR
POLAR
14
Nonpolar vs. Polar
15
Nonpolar vs. Polar
16
Single Bond
  • Defn one pair (2) of e- shared
  • Lewis Structures represents how atoms in
    molecules are arranged
  • atoms MUST obey octet rule (except hydrogen)

17
Lewis Structures
  • bonded electrons occur between bonded atoms


A
B
A
B
or

single bond
18
Lewis Structures
  • Unshared or Lone Pairs electron pairs NOT
    involved in bonding






A
B
A
B





lone pairs
19
Lewis Structures Examples
  • H2O

(8 valence e- or 4 pairs)










O
H
H
H
O






H



O
H

H
20
Lewis Structures Examples
  • NHF2

(20 v.e. or 10 pairs)














N
F
F
F


















H
N
H





F











N
F
F






H
21
Multiple Covalent Bonds
  • Double Bond two pairs (4) e- shared



A
B
A
B


O2
(12 v.e. 6 pairs)














O









O
O
O
O
O



















O
O


22
Multiple Covalent Bond
  • Triple Bond three pairs (6) e- shared




A
B
A
B



N2
(10 v.e. 5 pairs)



















N


N


N
N
N
N














N
N



23
Comparing single, double, and triple bonds
  • Bond Strength
  • Bond Length

Triple gt Double gt Single
Single gt Double gt Triple
The shorter the bond, the stronger it is
24
Polyatomic Ions
  • Defn CHARGED group of atoms covalently bonded
  • - ex SO42-, NH41, NO31-

25
Polyatomic Ions
SO42-
(32 v.e. 16 pairs)
2-


2-




O




O


















O
O
S
O
O
S


















O




O






26
Polyatomic Ions
NH41
(8 v.e. 4 pairs)
1
1
H
H




H
H
N
H
H
N




H
H
27
Ionic Bonding
giving/taking of valence electrons
  • Whats going on?
  • If I gave you a compound, how can you tell if it
    is ionic or not?
  • combo of metal nonmetal

28
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29
Formation of Ionic Bonds
  • NaCl

1-






Na
Cl
Na1





Cl







2s22p63s1
3s23p5
2s22p6
3s23p6
8 v.e.
30
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31
Formation of Ionic Bonds
  • CaBr2

1-







Ca
Br




Br







Ca2





Br
1-







Br




32
Using electronegativity to determine bond type
  • Recall electronegativity how much an atom wants
    electrons
  • Each atom is assigned a number between 0-4.0 to
    determine electronegativity strength

33
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34
Using electronegativity to determine bond type
  • We know 3 types of bonds
  • - nonpolar covalent
  • - polar covalent
  • - ionic
  • To determine bond type, subtract
    electronegativity values and see scale

35
Scale
Using electronegativity to determine bond type
polar covalent
nonpolar covalent
ionic
1.7
0.3
0
4.0
36
Using electronegativity to determine bond type
H and Cl
3.0 2.1
0.9
polar covalent
C and S
2.5 2.5
0
nonpolar covalent
Na and F
4.0 0.9
3.1
ionic
37
Metallic Bonding
  • Defn bond formed from attraction between
    positive nuclei and delocalized electrons
  • holds metals together
  • Delocalized Electrons electrons detached from
    parent atom
  • lost electron away from home

38
Electron Sea Model
  • Defn electrons move freely within other
    molecular orbitals

39
Properties of Metals
  • Electron sea model gives metals certain physical
    properties
  • Shiny due to photoelectric effect
  • Conduct electricity and heat electrons move
    easily from one place to another
  • Malleable (pound into sheets)
  • Ductile (put into wires)

40
Why malleable and ductile?
atoms can also move from one place to another and
still remain in contact with and bonded to the
other atoms and electrons around them
Shape 1
Shape 2
shifted atoms
41
Dipole Moment
  • defn imbalance of electron density in a
    covalent bond
  • Due to electronegativity of atoms

?- (partial negative) signifies more EN atom
? (partial positive) signifies less EN atom
shows direction of dipole moment
42
Examples
?
?-
?-
?
H
O
Cl
C
H 2.2 C 2.6 N 3.0 Cl 3.2 O 3.4 F 4.0
?
?-
?-
?
C
F
N
H
43
Intermolecular Forces
  • Defn attractive forces between 2 molecules

44
Intermolecular Forces
  • Dipole-Dipole attraction between oppositely
    charged polar molecules

?-
?-
?
?
?-
?
45
Intermolecular Forces
  • London Dispersion Forces very weak, very brief
    dipole moment created in nonpolar molecules

46
Electrons evenly distributed
London force
Temporary dipole
47
Intermolecular Forces
  • Hydrogen Bonding strong bond between H and
    N,O, or F of another molecule
  • - Water is prime example

?-
O
O
H
H
?
?
48
?-
O
O
O
O
H
H
H
?
H
?
?-
O
hydrogen bond
O
H
H
?
?
O
O
O
O
H
H
H
H
49
Strength Ranking
Hydrogen gt dipole-dipole gt London
50
VSEPR
  • Valence Shell Electron Pair Repulsion
  • Defn determines the shape of molecule
  • Electron pairs try to stay far away as possible

51
lone pairs
atoms bonded to central atom
shape
4
0
tetrahedral
52
Tetrahedral
53
lone pairs
atoms bonded to central atom
shape
4
0
tetrahedral
trigonal pyramidal
1
3
54
Trigonal Pyramidal
55
lone pairs
atoms bonded to central atom
shape
4
0
tetrahedral
trigonal pyramidal
1
3
2
2
bent
56
Bent
57
lone pairs
atoms bonded to central atom
shape
4
0
tetrahedral
trigonal pyramidal
1
3
2
2
bent
trigonal planar
3
0
58
Trigonal Planar
59
lone pairs
atoms bonded to central atom
shape
4
0
tetrahedral
trigonal pyramidal
1
3
2
2
bent
trigonal planar
3
0
2
0
linear
60
Linear
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