Covalent Bonds

1
Covalent Bonds
Other Bonds

• James Treasury Sen. Kit
• Bond Bond Bond

2
Bonding

• A metal and a nonmetal form an ionic bond by the
  transfer of e-.
• Nonmetal atoms form covalent bonds by sharing e-.

3
Bonding

• A metal and a nonmetal form an ionic bond by the
  transfer of e-.
• Nonmetal atoms form covalent bonds by sharing e-.

thats not the whole story.
4
Consider electronegativity
5

• Call a bond with electronegativity difference of
• 0-.4 -- nonpolar covalent bond
• .5-1.8 -- polar covalent bond
• 1.9 and above -- ionic bond
• (We can also define ionic character)

6

• What kind of bond forms between chlorine and
  phosphorus atoms?

7

• What kind of bond forms between chlorine and
  phosphorus atoms?
• Cl P
• 3.0 - 2.1.9
• This is a polar covalent bond
• (use an absolute value for the difference)

8
What kind of bond forms between these atoms?

• H and Cl Cl and C
• Cl and F F and F
• Na and O Mg and N
• Mg and Mg N and O

9
The Lewis Diagram

• Bar represents a covalent bond,
• ?2 shared electrons
• Unshared pairs fill out the octets.
• Double and triple bars represent double and
  triple bonds.

10
Lewis diagrams

• Step 1 Count the total valence electrons
  available
• --use the columns of the periodic chart
• --negative ions have extra electrons,
• --positive ions are missing electrons
• Step 2 Count the total valence electrons needed
• --duet rule for hydrogen, or the
• --octet rule for everything else
• Step 3 Number of bonds (electrons
  needed-electrons available) / 2 electrons per
  bond

11
Lewis diagrams

• Step 4 Choose the central atom (almost always
  the unique one), surround it with the others.
• Step 5 Connect with one bond to each outer atom.
  (PS Recheck your formula!)
• Step 6 Fill in enough multiple bonds to satisfy
  step 3
• Step 7 Draw in unshared pairs to fill valence
  levels.

12
Dont

• Dont try to figure out whose electrons are
  whose. Electrons are identical.
• Dont string the atoms along. Put one atom in
  the center, unless you have 6 or more atoms.
• Dont EVER put two bonds or an unshared pair on
  H.

13
Lewis diagrams

• Draw a Lewis diagram of hydrogen cyanide, HCN

14
Lewis diagrams

• Draw a Lewis diagram of hydrogen cyanide, HCN
• H C N

15
Try an ion.

• Draw a Lewis diagram for the nitrite, NO2- , ion

16
Lewis diagrams

• Draw a Lewis diagram for the nitrite, NO2- , ion
• O N O -

17
Resonance

• Which one is preferable?
• O N O -
• O N O -

18
Resonance

• Each is valid. The multiple bond exists in both
  locations. This is called resonance.
• O N O -
• O N O -
• (the double-headed arrow signifies resonance)

19
Resonance

• Draw three resonance structures for carbon
  dioxide.

20
Formal Charge

• (valence e-) - (bonds/2 unshared e-)
• Try to minimize each. The more electronegative
  atom gets a more negative formal charge.

21
Formal Charge

• Draw two valid arrangements each for the atoms
  in
• H2CO
• H2O2
• N2O
• (count bonds, put them in, fill in lone pairs)

22
Formal Charge

• Which is preferable?

H
H
H
C
O
H
C
O
23
Formal Charge

• Which is preferable?

H
H
H
C
O
H
C
O
0 1 -1 0 or 0 0 0 0
24
Formal Charge

• Which is preferable?

H
H
H
C
O
H
C
O
0 1 -1 0 or 0 0 0 0
Choose this one!
25
Formal Charge

• Which is preferable?

H
H
H
O
O
H
O
O
26
Formal Charge

• Which is preferable?

H
H
H
O
O
H
O
O
0 0 0 0 or -1 1 0 0
27
Formal Charge

• Which is preferable?

H
H
H
O
O
H
O
O
0 0 0 0 or -1 1 0 0
Choose this one!
28
Formal Charge

• Which is preferable?

N
N
O
N
N
O
29
Formal Charge

• Which is preferable?

N
N
O
N
N
O
-1 1 0 or 1 -2 1
30
Formal Charge

• Which is preferable?

N
N
O
N
N
O
-1 1 0 or 1 -2 1
Choose this one!
31
Formal Charge

• Which is preferable?

N
N
O
N
N
O
32
Formal Charge

• Which is preferable?

N
N
O
N
N
O
-1 1 0 or 0 1 -1
33
Formal Charge

• Which is preferable?

N
N
O
N
N
O
-1 1 0 or 0 1 -1
Choose this one!
34
Coordinate covalent bonds

• How did this
• become this?

O
C
O
C
35
Coordinate covalent bonds

• How did this
• become this?
• Carbon monoxide really does have the third bond.
  The oxygen donates both electrons to share. This
  is a coordinate covalent bond

O
C
O
C
O
C
36
Coordinate covalent bonds

• Draw a Lewis diagram of the ozone (O3) molecule.
  Count the formal charge for each atom and mark a
  coordinate covalent bond

37
Coordinate covalent bonds

• Draw a Lewis diagram of the ozone (O3) molecule.
  Count the formal charge for each atom and mark a
  coordinate covalent bond

O
O
O
-1 1 0
38
Exceptions to the octet rule

• Draw a Lewis diagram for the triiodide ion, I3-

39
Exceptions to the octet rule

• Draw a Lewis diagram for the triiodide ion, I3-
• Gadzooks!
• When you try to find the number of bonds,
  (24-22)/21 bond.
• Thats not enough to tie the ion together.

40
Exceptions to the octet rule

• When that happensgo old school. Circle your
  electrons
• I I I -

41
Exceptions to the octet rule

• When that happensgo old school. Circle your
  electrons
• I I I -

42
Exceptions to the octet rule

• When that happensgo old school. Circle your
  electrons
• I I I -
• Two single bonds will satisfy the outer two
  iodine atoms, the middle one breaks the octet
  rule (with 10 electrons).

43
Exceptions to the octet rule

• Draw a Lewis diagram for XeF4
• (The point here is to find out how many unshared
  pairs are on the central atom)

44
Exceptions to the octet rule

• Draw a Lewis diagram for XeF4
• (The point here is to find out how many unshared
  pairs are on the central atom)

F
Xe
F
F
F
45
Polar bonds

• We use a symbol to show a polar covalent
  bond.
• The arrow points toward the more electronegative
  atom, the () end is less electronegative

O
H
H
H
46
Polar bonds

• Or, mark the molecules () and (-) parts
• The d is the small Greek delta. It indicates a
  small change. In this case, a partial charge

d-
O
d
H
H
d
H
47
Three properties of polar bonds

• The less electronegative end of a polar bond
  d d-
• H Cl
• --is more positive
• --cannot attract the electrons as well
• --is farther from the shared pair of electrons

48
Molecular Shapes

• Most molecules have a central atom that follows
  the octet rule. This allows the following
  shapes.
• Tetrahedral
• Trigonal pyramid (trigonalhaving three
• Bent corners)
• Linear and
• Trigonal planar

49
Molecular Shapes

• Four bonds in four directions makes a tetrahedral
  shape

50
Molecular Shapes

• Three bonds and one lone pair in four directions
  makes a trigonal pyramid shape

51
Molecular Shapes

• Two bonds and two lone pairs in four directions
  makes a bent shape

52
Molecular Shapes

• A double bond holds two electron pairs in the
  same direction. With no lone pairs, this makes a
  trigonal planar molecule

53
Molecular Shapes

• One lone pair, with a single and a double bond
  gives a bent molecule.

54
Molecular Shapes

• Two double bonds, or a single and a triple makes
  a linear molecule

55
Molecular Shapes

• Two atoms are always in a straight line, a linear
  molecule.

56
Look for double bonds and unshared pairs
If Acentral atom, Batoms bonded to it, Ee-
pairs

• AB4tetrahedral
• (no double bonds)
• AB3E-trigonal pyramid
• (no double bonds)
• AB2E2 bent
• (no double bonds)
• ABE3 linear
• (no double bonds)

• AB3 trigonal planar
• (one double bond)
• AB2E bent
• (one double bond)
• ABE2linear
• (one double bond)
• AB2linear
• (2 doubles or 1 triple)
• ABElinear ()

57

• Determine the shape of each molecule and ion on
  the lab that has a single central atom.

58
Polarity of molecules

• When polar bonds are not cancelled by symmetry,
  you get a polar molecule. A polar molecule has
  () and (-) parts.
• POLARITY is the first property to look for when
  analyzing a molecule !

59
Polarity

• CH4 has no polar bonds. It is symmetric
• PH3 has no polar bonds It is not symmetric
• CO2 has polar bonds. It is symmetric
• H2O has polar bonds. It is not symmetric

60
Polarity

• CH4 has no polar bonds. It is symmetric
• Not polar!
• PH3 has no polar bonds It is not symmetric
• Not polar!
• CO2 has polar bonds. It is symmetric
• Not polar!
• H2O has polar bonds. It is not symmetric
• Polar!

61

• Mark each molecule on the lab that is polar.
• For those that are not polarwhy not?
• (PSdont even look at the ions. If it has a
  whole charge, ignore the partial charges)

62
Hybridization

• Atomic orbitals combine to form hybrid orbitals
  before bonding
• (Hydrogen is the only exception)

63
Before bonding

• The first step is a hybridization of the valence
  level
• C
• H H H H
• forms

p orbitals
s orbitals
64
Hybridization

• The first step is a hybridization of the valence
  level
• C
• H H H H
• The s and p orbitals hybridize to form sp3
  orbitals. The sp3 designation shows one s
  orbital and 3 p orbitals make the new ones

sp3 orbitals
65
Hybridization

• The first step is a hybridization of the valence
  level
• C
• H H H H
• The number of orbitals is preserved
• (4 in ? 4 out)

sp3 orbitals
66
Hybridization

• C
• H H H H
• All four bonds are identical. Methane is a
  symmetrical molecule.

67
sp2 Hybridization

• When one p orbital is left out of the
  hybridization, it is used to make a double bond
• C
• H H O
• forms.

p orbitals
s orbitals
68
sp2 Hybridization

• When one p orbital is left out of the
  hybridization, it is used to make a double bond
• C
• H H O

sp2 orbitals
Unused p orbitalswill form the second bond
between C and O
69
sp2 Hybridization

• When one p orbital is left out of the
  hybridization, it is used to make a double bond
• C
• H H O

?Makes the double bond!
70
sp2 Hybridization
Carbon oxygen share electrons in unused p
orbitals

• H
• C O
• H
• Carbon shares electrons in sp2 orbitals

71
sp2 Hybridization
Carbon oxygen share electrons in unused p
orbitals

• H
• C O
• H
• Carbon shares electrons in sp2 orbitals

p bond
s bonds
72
sp Hybridization

• When two p orbitals are left out of the
  hybridization, it is used to make two double
  bonds, or a triple bond
• C
• O O
• forms.

73
sp Hybridization

• When two p orbitals are left out of the
  hybridization, it is used to make two double
  bonds, or a triple bond
• C
• O O

sp orbitals
Unused p orbitals
sp2 orbitals
74
sp Hybridization

• When two p orbitals are left out of the
  hybridization, it is used to make two double
  bonds, or a triple bond
• C
• O O

75
sp Hybridization
p bonds

• When two p orbitals are left out of the
  hybridization, it is used to make two double
  bonds, or a triple bond
• C
• O O

76
sp Hybridization
Carbon oxygen share electrons in unused p
orbitals

• O C O
• Carbon shares electrons in sp orbitals

77
Look for multiple bonds!

• of Multiples Bonding patterns Hybridization
• None AB4, AB3E, sp3
• AB2E2, ABE3
• One AB3, AB2E sp2
• ABE2
• Two AB2, ABE sp

78
Look for multiple bonds!

• of Multiples Bonding patterns Hybridization
• None sp3
• One sp2
• Two sp

79
What is the hybridization of the carbon atoms in

• CCl4
• H2CO
• C2H6
• C2H4

• C2H2
• CO
• CH3OH
• HCOOH

80
Molecular Orbital Theory (MOT)

• Overlapping s orbitals, or hybridized orbitals
  makes a s (sigma) bond
• The electron density is on the SAME line as the
  nuclei

s
s
s
s
81
Molecular orbitals

• Overlapping p orbitals, makes a p (pi) bond
• The electron density is on a PARALLEL line to the
  line of the nuclei

p
p
p


p
82
Molecular orbitals

• A single bond is a s bond
• A double bond is a s bond, and a p bond above and
  below the s
• A triple bond is a s bond, with two p bonds
  above/below and front/back

83

• For every bonding molecular orbital (s or p) an
  antibonding orbital is formed (s or p)
• A bond is formed when there are more bonding than
  antibonding electrons

84
VSEPR

• Valence Shell Electron Pair Repulsion
• Theory

85
VSEPR

• Valence Shell Electron Pair Repulsion
• Theory
• --pronounced Vesper
• Electron pairs repel each other. Just as it says.

86

• VSEPR is used to predict bond angles. The pairs
  will space themselves out as far as possible.
• A lone pair will take as much room as a bond AND
  MORE!
• Consider sp3 hybridization

87

• AB4like methane. Tetrahedral 109.5o
• AB3Elike ammonia. Pyramidal 107o
• AB2E2like water. Bent 104.5o angles
• --the unshared pairs force the bonds closer
  togetherbond angles decrease

88
With sp2 hybridization

• AB3like carbonate. Trigonal planar 120o
• AB2Elike nitrite. Bent less than 120o
• ABE2like O2(2 atoms, has to be linear)

89
With sp hybridization

• AB2like carbon dioxide. Linear 180o
• ABElike carbon monoxide. Linear 180o

90
but thats just if you always follow the rules

• like the octet rule.

91
With dsp3 hybridization

• AB5trigonal bipyramid
• AB4Eseesaw
• AB3E2t-shaped
• AB2E3linear
• ABE4linear

92
With d2sp3 hybridization

• AB6 octahedral
• AB5Esquare pyramid
• AB4E2square planar
• AB3E3t-shaped
• AB2E4 , ABE5linear

93
What is the shape of

• All of the molecules and ions on the lab?
• I3-, SF6, XeF4, PCl5, IF5?

94

• Count the s and p bonds in the following
  molecule. Label each bond as s or p

H
H
C
C
C
C
C
H
H
H
O
H
95
11
3

• Count the s and p bonds in the following
  molecule. Label each bond as s or p

H
H
C
C
C
C
C
H
H
H
O
H
96

• Determine the hybridization of the carbons and
  the oxygen atom

H
H
C
C
C
C
C
H
H
H
O
H
97

• Determine the hybridization of the carbons and
  the oxygen atom
• sp sp sp3 sp2 sp3
• sp2

H
H
C
C
C
C
C
H
H
H
O
H
98
The molecular aufbau diagram
99
The molecular aufbau order

• s1s2s1s2s2s2s2s2s2px2p2py,z4p2py,z4s2px2.
• For example
• O2 has 16 electrons. Its electron configuration
  is
• O2 s1s2s1s2s2s2s2s2s2px2p2py,z4p2py,z2

100
The molecular aufbau order

• What is the electron configuration of
• N2
• NO
• Ne2

Remember we couldnt do a Lewis diagram with an
odd number of electrons!
101
The molecular aufbau diagram
102
The molecular aufbau diagram
These two can switch placesno effect on bonding,
but it causes magnetic effects we can measure
103
The molecular aufbau diagram
These two can switch placesno effect on bonding,
but it causes magnetic effects we can measure
104
Bond order

• The order of a bond in a diatomic molecule is
  half the number of shared electrons not cancelled
  by antibonding electrons.
• Or
• (number of bonding e- in the atoms-antibonding
  e-)/2

105
Bond order

• What is the bond order of
• N2
• NO
• Ne2

Remember we couldnt do a Lewis diagram with an
odd number of electrons!
106
You will be responsible for

• Writing the molecular orbital electron
  configuration and
• Calculating the bond order
• of any pair of atoms from the second period as
  they attempt to form a diatomic molecule.

107
Bond Energies

• The energy it takes to break a bond is the amount
  of energy released as the bond is formed.
• --measured in kJ/mol
• --can be used to estimate DHrxn
• --can be absorbed or emitted as light.

108
What is the DHf of NH3?
109
What is the DHf of NH3?

• Write the reaction
• N2 3H2 ?2NH3

110
What is the DHf of NH3?

• Count the bonds made and broken
• N2 3H2 ?2NH3
• 1 NN triple bond, 3 HH single bonds broken
• 6 NH single bonds made

111
What is the DHf of NH3?

• Look up bond energies, and find a total
• N2 3H2 ?2NH3
• 1 molx941kJ/mol3 molx436kJ/mol 2249kJ used
• 6 molx393 kJ/mol2358 kJ released

112
What is the DHf of NH3?

• Find the difference, express as kJ/mol
• N2 3H2 ?2NH3
• 2358 kJ-2249kJ109 kJ more is released, as 2mol
  NH3 is produced, DHf-109kJ/2mol-55kJ/mol

113

• Its an estimate.
• My book claims -46 kJ/mole.

114
What is the heat of reaction for

• CH4 O2? H2O CO2

115
What is the heat of reaction for

• CH4 2O2? 2H2O CO2

116
What is the heat of reaction for

• CH4 2O2? 2H2O CO2
• Break 2 OO and 4 C-H
• Form 4 H-O and 2 CO

117
What is the heat of reaction for

• CH4 2O2? 2H2O CO2
• Break 2 OO and 4 C-H
• 2_at_500 kJ 4_at_393kJ2572kJ
• Form 4 H-O and 2 CO
• 4_at_464kJ 2_at_799kJ 3454 kJ

118
What is the heat of reaction for

• CH4 2O2? 2H2O CO2
• Break 2 OO and 4 C-H
• 2_at_500 kJ 4_at_393kJ2572kJ
• Form 4 H-O and 2 CO
• 4_at_464kJ 2_at_799kJ 3454 kJ
• -882kJ/mol

119
What is the heat of reaction for

• 2H2O2? 2H2O O2

120
Tasks
121
On your test, you will be asked to

• Describe how and why ionization of metals and
  non-metals occurs
• Write ECs for atoms and ions
• Show formation of ionic and covalent bonds by
  electron dot diagrams
• Describe metallic bonding
• Define and identify electrolytes

122
On your test, you will be asked to

• Identify particles and types of substances by
  bonding
• Draw Lewis diagrams
• Identify shapes of molecules and ions
• Identify types of bonds between atoms
• Describe polarity

123
On your test, you will be asked to

• Identify polar and nonpolar molecules
• Identify hybridizations
• Describe single and double bonds by MOT
• Estimate DHrxn by bond energies
• Use VSEPR to predict molecular shapes and bond
  angles.
• Calculate and justify bond orders for diatomics
  from the second period