Chapter 8

1
Chapter 8Covalent Bonding
Ball-and-stick model
2
Section 8.1Molecular Compounds

• OBJECTIVES
• Distinguish between the melting points and
  boiling points of molecular compounds and ionic
  compounds.

3
Section 8.1Molecular Compounds

• OBJECTIVES
• Describe the information provided by a molecular
  formula.

4
Bonds are

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

1. Ionic bonds transfer of electrons (gained or
   lost makes formula unit)
2. Covalent bonds sharing of electrons. The
   resulting particle is called a molecule

5
Covalent Bonds

• The word covalent is a combination of the prefix
  co- (from Latin com, meaning with or
  together), and the verb valere, meaning to be
  strong.
• Two electrons shared together have the strength
  to hold two atoms together in a bond.

6
Molecules

• Many elements found in nature are in the form of
  molecules
• a neutral group of atoms joined together by
  covalent bonds.
• For example, air contains oxygen molecules,
  consisting of two oxygen atoms joined covalently
• Called a diatomic molecule (O2)

7
How does H2 form?

• The nuclei repel each other, since they both have
  a positive charge (like charges repel).

(diatomic hydrogen molecule)
8
How does H2 form?

• But, the nuclei are attracted to the electrons
• They share the electrons, and this is called a
  covalent bond, and involves only NONMETALS!

9
Covalent bonds

• Nonmetals hold on to their valence electrons.
• They cant give away electrons to bond.
• But still want noble gas configuration.
• Get it by sharing valence electrons with each
  other covalent bonding
• By sharing, both atoms get to count the electrons
  toward a noble gas configuration.

10
Covalent bonding

• Fluorine has seven valence electrons (but would
  like to have 8)

11
Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven

12
Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons

13
Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons

14
Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons

15
Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons

16
Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons

17
Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons
• both end with full orbitals

18
Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons
• both end with full orbitals

F
F
8 Valence electrons
19
Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons
• both end with full orbitals

F
F
8 Valence electrons
20
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21
Molecular Compounds

• Compounds that are bonded covalently (like in
  water, or carbon dioxide) are called molecular
  compounds
• Molecular compounds tend to have relatively lower
  melting and boiling points than ionic compounds
  this is not as strong a bond as ionic

22
Molecular Compounds

• Thus, molecular compounds tend to be gases or
  liquids at room temperature
• Ionic compounds were solids
• A molecular compound has a molecular formula
• Shows how many atoms of each element a molecule
  contains

23
Molecular Compounds

• The formula for water is written as H2O
• The subscript 2 behind hydrogen means there are
  2 atoms of hydrogen if there is only one atom,
  the subscript 1 is omitted
• Molecular formulas do not tell any information
  about the structure (the arrangement of the
  various atoms).

24
- Page 215
3. The ball and stick model is the BEST, because
it shows a 3-dimensional arrangement.
These are some of the different ways to represent
ammonia
1. The molecular formula shows how many atoms of
each element are present
2. The structural formula ALSO shows the
arrangement of these atoms!
25
Section 8.2The Nature of Covalent Bonding

• OBJECTIVES
• Describe how electrons are shared to form
  covalent bonds, and identify exceptions to the
  octet rule.

26
Section 8.2The Nature of Covalent Bonding

• OBJECTIVES
• Demonstrate how electron dot structures represent
  shared electrons.

27
Section 8.2The Nature of Covalent Bonding

• OBJECTIVES
• Describe how atoms form double or triple covalent
  bonds.

28
Section 8.2The Nature of Covalent Bonding

• OBJECTIVES
• Distinguish between a covalent bond and a
  coordinate covalent bond, and describe how the
  strength of a covalent bond is related to its
  bond dissociation energy.

29
Section 8.2The Nature of Covalent Bonding

• OBJECTIVES
• Describe how oxygen atoms are bonded in ozone.

30
A Single Covalent Bond is...

• A sharing of two valence electrons.
• Only nonmetals and hydrogen.
• Different from an ionic bond because they
  actually form molecules.
• Two specific atoms are joined

31
How to show the formation

• Its like a jigsaw puzzle.
• You put the pieces together to end up with the
  right formula.
• Carbon is a special example - can it really share
  4 electrons 1s22s22p2?
• Yes, due to electron promotion!
• Group 3A 4A

32
Some Definitions

• Structural Formula represents the covalent
  bonds by dashes and shows the arrangement of
  covalently bonded atoms. (similar to electron dot
  structure)
• Unshared Pair of Electrons a pair of valence
  electrons that is not shared between atoms.
• Another example lets show how water is formed
  with covalent bonds, by using an electron dot
  diagram

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Water

• Each hydrogen has 1 valence electron
• - Each hydrogen wants 1 more
• The oxygen has 6 valence electrons
• - The oxygen wants 2 more
• They share to make each other complete

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Water

• Put the pieces together
• The first hydrogen is happy
• The oxygen still needs one more

H
35
Water

• So, a second hydrogen attaches
• Every atom has full energy levels

Note the two unshared pairs of electrons
H
H
36
Examples

• Ammonia NH3
• Methane - CH4
• More examples can be found on page 220

37
Multiple Bonds

• Sometimes atoms share more than one pair of
  valence electrons.
• A double bond is when atoms share two pairs of
  electrons (4 total)
• A triple bond is when atoms share three pairs of
  electrons (6 total)
• Table 8.1, p.222 - Know these 7 elements as
  diatomic and be familiar with their properties
  and uses
• Br2 I2 N2 Cl2 H2 O2 F2

Whats the deal with the oxygen dot diagram?
38
Dot diagram for Carbon dioxide

• CO2 - Carbon is central atom ( more metallic )
• Carbon has 4 valence electrons
• Wants 4 more
• Oxygen has 6 valence electrons
• Wants 2 more

C
39
Carbon dioxide

• Attaching 1 oxygen leaves the oxygen 1 short, and
  the carbon 3 short

C
40
Carbon dioxide

• Attaching the second oxygen leaves both of the
  oxygen 1 short, and the carbon 2 short

C
41
Carbon dioxide

• The only solution is to share more

C
42
Carbon dioxide

• The only solution is to share more

C
43
Carbon dioxide

• The only solution is to share more

C
O
44
Carbon dioxide

• The only solution is to share more

C
O
45
Carbon dioxide

• The only solution is to share more

C
O
46
Carbon dioxide

• The only solution is to share more

C
O
O
47
Carbon dioxide

• The only solution is to share more
• Requires two double bonds
• Each atom can count all the electrons in the bond

C
O
O
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Carbon dioxide

• The only solution is to share more
• Requires two double bonds
• Each atom can count all the electrons in the bond

8 valence electrons
C
O
O
49
Carbon dioxide

• The only solution is to share more
• Requires two double bonds
• Each atom can count all the electrons in the bond

8 valence electrons
C
O
O
50
Carbon dioxide

• The only solution is to share more
• Requires two double bonds
• Each atom can count all the electrons in the bond

8 valence electrons
C
O
O
51
How to draw them?

• Use these guidelines
• Add up all the valence electrons.
• Count up the total number of electrons to make
  all atoms happy.
• Subtract then Divide by 2
• Tells you how many bonds to draw
• Fill in the rest of the valence electrons to fill
  atoms up.

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Example

• NH3, which is ammonia
• N central atom has 5 valence electrons, wants
  8
• H - has 1 (x3) valence electrons, wants 2 (x3)
• NH3 has 53 8
• NH3 wants 86 14
• (14-8)/2 3 bonds
• 4 atoms with 3 bonds

N
H
53
Examples

• Draw in the bonds start with singles
• All 8 electrons are accounted for
• Everything is full done with this one.

H
N
H
H
54
Example HCN

• HCN C is central atom
• N - has 5 valence electrons, wants 8
• C - has 4 valence electrons, wants 8
• H - has 1 valence electron, wants 2
• HCN has 541 10
• HCN wants 882 18
• (18-10)/2 4 bonds
• 3 atoms with 4 bonds this will require multiple
  bonds - not to H however

55
HCN

• Put single bond between each atom
• Need to add 2 more bonds
• Must go between C and N (Hydrogen is full)

N
H
C
56
HCN

• Put in single bonds
• Needs 2 more bonds
• Must go between C and N, not the H
• Uses 8 electrons need 2 more to equal the 10 it
  has

N
H
C
57
HCN

• Put in single bonds
• Need 2 more bonds
• Must go between C and N
• Uses 8 electrons - 2 more to add
• Must go on the N to fill its octet

N
H
C
58
Another way of indicating bonds

• Often use a line to indicate a bond
• Called a structural formula
• Each line is 2 valence electrons

H
O
H
O

H
H
59
Other Structural Examples
H C N
H
C O
H
60
A Coordinate Covalent Bond...

• When one atom donates both electrons in a
  covalent bond.
• Carbon monoxide (CO) is a good example

Both the carbon and oxygen give another single
electron to share
61
Coordinate Covalent Bond

• When one atom donates both electrons in a
  covalent bond.
• Carbon monoxide (CO) is a good example

Oxygen gives both of these electrons, since it
has no more singles to share.
This carbon electron moves to make a pair with
the other single.
O
C
62
Coordinate Covalent Bond

• When one atom donates both electrons in a
  covalent bond.
• Carbon monoxide (CO)

The coordinate covalent bond is shown with an
arrow as
O
C
C O
63
Coordinate covalent bond

• Most polyatomic cations and anions contain
  covalent and coordinate covalent bonds
• Table 8.2, p.224
• Sample Problem 8.2, p.225
• The ammonium ion (NH41) can be shown as another
  example

64
Bond Dissociation Energies...

• The total energy required to break the bond
  between 2 covalently bonded atoms
• High dissociation energy usually means the
  chemical is relatively unreactive, because it
  takes a lot of energy to break it down.

65
Magnetism

• Paramagnetism when molecules with unpaired
  electrons are placed in a magnetic field they
  tend to be drawn into the field
• Diamagnetic when molecules in which all
  electrons are paired are placed in a magnetic
  field they tend to be pushed away.

66
Resonance is...

• When more than one valid dot diagram is possible
  for the same number of electron pairs for a
  molecule or ion.
• Consider the two ways to draw ozone (O3)
• Which one is it? Does it go back and forth?
• It is a hybrid of both and shown by a
  double-headed arrow
• found in double-bond structures!

67
Resonance in Ozone
Note the different location of the double bond
Neither structure is correct, it is actually a
hybrid of the two. To show it, draw all
varieties possible, and join them with a
double-headed arrow.
68
Resonance

• Occurs when more than one valid Lewis structure
  can be written for a particular molecule (due to
  position of double bond)

• These are resonance structures of benzene.
• The actual structure is an average (or hybrid) of
  these structures.

69
Polyatomic ions note the different positions of
the double bond.
Resonance in a carbonate ion (CO32-)
Resonance in an acetate ion (C2H3O21-)
70
The 3 Exceptions to Octet rule

• For some molecules, it is impossible to satisfy
  the octet rule
• 1. usually when there is an odd number of
  valence electrons
• NO2 has 17 valence electrons, because the N has
  5, and each O contributes 6. Note N page 228
• It is impossible to satisfy octet rule, yet the
  stable molecule does exist

71
Exceptions to Octet rule

• Another exception Boron
• Page 228 shows boron trifluoride, and note that
  one of the fluorides might be able to make a
  coordinate covalent bond to fulfill the boron
• 2 -But fluorine has a high electronegativity (it
  is greedy), so this coordinate bond does not form
• 3 -Top page 229 examples exist because they are
  in period 3 or beyond

72
Section 8.3Bonding Theories

• OBJECTIVES
• Describe the relationship between atomic and
  molecular orbitals.

73
Section 8.3Bonding Theories

• OBJECTIVES
• Describe how VSEPR theory helps predict the
  shapes of molecules.

74
Section 8.3Bonding Theories

• OBJECTIVES
• Identify ways in which orbital hybridization is
  useful in describing molecules.

75
Molecular Orbitals are...

• The model for covalent bonding assumes the
  orbitals are those of the individual atoms
  atomic orbital
• Orbitals that apply to the overall molecule, due
  to atomic orbital overlap are the molecular
  orbitals
• A bonding orbital is a molecular orbital that can
  be occupied by two electrons of a covalent bond

76
Molecular Orbitals - definitions

• Sigma bond- when two atomic orbitals combine to
  form the molecular orbital that is symmetrical
  along the axis connecting the nuclei
• Pi bond- the bonding electrons are likely to be
  found above and below the bond axis (weaker than
  sigma)
• Note pictures on the next slide

77
- Page 231
Sigma bond is symmetrical along the axis between
the two nuclei.
Pi bond is above and below the bond axis, and is
weaker than sigma
78
VSEPR stands for...

• Valence Shell Electron Pair Repulsion
• Predicts the three dimensional shape of
  molecules.
• The name tells you the theory
• Valence shell outside electrons.
• Electron Pair repulsion electron pairs try to
  get as far away as possible from each other.
• Can determine the angles of bonds.

79
Determining the Number of Sigma and Pi Bonds

• Single Bonds 1 sigma bond
• Double Bonds 1 sigma and 1 pi bond
• Triple Bonds 1 sigma and 2 pi bonds

80
VSEPR

• Based on the number of pairs of valence
  electrons, both bonded and unbonded.
• Unbonded pair also called lone pair.
• CH4 - draw the structural formula
• Has 4 4(1) 8
• wants 8 4(2) 16
• (16-8)/2 4 bonds

81
VSEPR for methane (a gas)

• Single bonds fill all atoms.
• There are 4 pairs of electrons pushing away.
• The furthest they can get away is 109.5º

H
C
H
H
H
This 2-dimensional drawing does not show a true
representation of the chemical arrangement.
82
4 atoms bonded

• Basic shape is tetrahedral.
• A pyramid with a triangular base.
• Same shape for everything with 4 pairs.

H
109.5º
C
H
H
H
83
Other angles, pages 232 - 233

• Ammonia (NH3) 107o
• Water (H2O) 105o
• Carbon dioxide (CO2) 180o
• Note the shapes of these that are pictured on the
  next slide

84
- Page 232
Methane has an angle of 109.5o, called tetrahedral
Ammonia has an angle of 107o, called pyramidal
Note the unshared pair that is repulsion for
other electrons.
85
Steric Number

• Steric Number Number of lone pairs and ligands
  around the central atom.
• Helps determine the molecular shapes
• http//intro.chem.okstate.edu/1314F00/Lecture/Chap
  ter10/VSEPR.html

2 No lone pairs Linear 180
3 No lone pairs Trigonal Planar 120
3 One lone pair Bent Less than 120
4 No lone pairs Tetrahedral 109.5
4 One lone pair Pyramidal 107
4 Two lone pairs Bent 105
5 No lone pairs Trigonal Bipyramidal 90, 120, 180
5 Two lone pairs T-shaped 90, 180
6 No lone pairs Octahedral 90, 180
6 Two lone pairs Square Planar 90, 180
86
Hybridization

• Hybridization occurs when several atomic orbitals
  mix to form the same total number of equivalent
  hybrid orbitals.
• Can occur in single, double, or triple bonds
• We are talking about the s, p, d, f electron
  sublevels
• The amount matches the steric number

87
Section 8.4Polar Bonds and Molecules

• OBJECTIVES
• Describe how electronegativity values determine
  the distribution of charge in a polar molecule.

88
Section 8.4Polar Bonds and Molecules

• OBJECTIVES
• Describe what happens to polar molecules when
  they are placed between oppositely charged metal
  plates.

89
Section 8.4Polar Bonds and Molecules

• OBJECTIVES
• Evaluate the strength of intermolecular
  attractions compared with the strength of ionic
  and covalent bonds.

90
Section 8.4Polar Bonds and Molecules

• OBJECTIVES
• Identify the reason why network solids have high
  melting points.

91
Bond Polarity

• Covalent bonding means shared electrons
• but, do they share equally?
• Electrons are pulled, as in a tug-of-war, between
  the atoms nuclei
• In equal sharing (such as diatomic molecules),
  the bond that results is called a nonpolar
  covalent bond

92
Bond Polarity

• When two different atoms bond covalently, there
  is an unequal sharing
• the more electronegative atom will have a
  stronger attraction, and will acquire a slightly
  negative charge
• called a polar covalent bond, or simply polar
  bond.

93
Electronegativity?

• The ability of an atom in a molecule to attract
  shared electrons to itself.

Linus Pauling 1901 - 1994
94
Table of Electronegativities
Higher electronegativity
95
Bond Polarity

• Refer to Table 6.2, p. 177
• Consider HCl
• H electronegativity of 2.1
• Cl electronegativity of 3.0
• the bond is polar
• the chlorine acquires a slight negative charge,
  and the hydrogen a slight positive charge

96
Bond Polarity

• Only partial charges, much less than a true 1 or
  1- as in ionic bond
• Written as
• H Cl
• the positive and minus signs (with the lower case
  delta ) denote partial charges.

d d-
d and d-
97
Bond Polarity

• Can also be shown
• the arrow points to the more electronegative
  atom.
• Table 8.3, p.238 shows how the electronegativity
  can also indicate the type of bond that tends to
  form

H Cl
98
Polar molecules

• Sample Problem 8.3, p.239
• A polar bond tends to make the entire molecule
  polar
• areas of difference
• HCl has polar bonds, thus is a polar molecule.
• A molecule that has two poles is called dipole,
  like HCl

99
Polar molecules

• The effect of polar bonds on the polarity of the
  entire molecule depends on the molecule shape
• carbon dioxide has two polar bonds, and is linear
  nonpolar molecule!

100
Polar molecules

• The effect of polar bonds on the polarity of the
  entire molecule depends on the molecule shape
• water has two polar bonds and a bent shape the
  highly electronegative oxygen pulls the e- away
  from H very polar!

101
Polar molecules

• When polar molecules are placed between
  oppositely charged plates, they tend to become
  oriented with respect to the positive and
  negative plates.
• Figure 8.24, page 239

102
Attractions between molecules

• They are what make solid and liquid molecular
  compounds possible.
• The weakest are called van der Waals forces -
  there are two kinds
• 1. Dispersion forces
• weakest of all, caused by motion of e-
• increases as e- increases
• halogens start as gases bromine is liquid
  iodine is solid all in Group 7A

103
2. Dipole interactions

• Occurs when polar molecules are attracted to each
  other.
• 2. Dipole interaction happens in water
• Figure 8.25, page 240
• positive region of one molecule attracts the
  negative region of another molecule.

104
2. Dipole interactions

• Occur when polar molecules are attracted to each
  other.
• Slightly stronger than dispersion forces.
• Opposites attract, but not completely hooked like
  in ionic solids.

105
2. Dipole Interactions
d d-
106
3. Hydrogen bonding

• is the attractive force caused by hydrogen
  bonded to N, O, F, or Cl
• N, O, F, and Cl are very electronegative, so this
  is a very strong dipole.
• And, the hydrogen shares with the lone pair in
  the molecule next to it.
• This is the strongest of the intermolecular
  forces.

107
Order of Intermolecular attraction strengths

1. Dispersion forces are the weakest
2. A little stronger are the dipole interactions
3. The strongest is the hydrogen bonding
4. All of these are weaker than ionic bonds

108
3. Hydrogen bonding defined

• When a hydrogen atom is a) covalently bonded to
  a highly electronegative atom, AND b) is also
  weakly bonded to an unshared electron pair of a
  nearby highly electronegative atom.
• The hydrogen is left very electron deficient (it
  only had 1 to start with!) thus it shares with
  something nearby
• Hydrogen is also the ONLY element with no
  shielding for its nucleus when involved in a
  covalent bond!

109
Hydrogen Bonding(Shown in water)
This hydrogen is bonded covalently to 1) the
highly negative oxygen, and 2) a nearby unshared
pair.
110
Hydrogen bonding allows H2O to be a liquid at
room conditions.
111
Attractions and properties

• Why are some chemicals gases, some liquids, some
  solids?
• Depends on the type of bonding!
• Table 8.4, page 244 (Know these!)
• Network solids solids in which all the atoms
  are covalently bonded to each other

112
Attractions and properties

• Figure 8.28, page 243
• Network solids melt at very high temperatures, or
  not at all (decomposes)
• Diamond does not really melt, but vaporizes to a
  gas at 3500 oC and beyond
• SiC, used in grinding, has a melting point of
  about 2700 oC

113
Covalent Network Compounds
Some covalently bonded substances DO NOT form
discrete molecules.
Graphite, a network of covalently bonded carbon
atoms
Diamond, a network of covalently bonded carbon
atoms
114
End of Chapter 8