Covalent bonding

1
Chapter 16

• Covalent bonding

2
How does H2 form?

• The nuclei repel

3
How does H2 form?

• The nuclei repel
• But they are attracted to electrons
• They share (or fight over) the electrons

4
Covalent bonds

• Bonding between two nonmetals.
• Nonmetals hold onto their valence electrons, so
  they cant give away electrons to bond.
• They still want noble gas configurations.
• Share (or fight over) valence electrons with each
  other.
• By sharing both atoms get to count the electrons
  toward noble gas configuration.

5
Covalent bonding

• Fluorine has seven valence electrons

6
Covalent bonding

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

7
Covalent bonding

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

8
Covalent bonding

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

9
Covalent bonding

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

10
Covalent bonding

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

11
Covalent bonding

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

12
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
13
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
14
Single Covalent Bond

• A sharing of two valence electrons.
• Only nonmetals and hydrogen.
• Different from an ionic bond because they
  actually form groups called molecules.
• Two specific atoms are joined in covalent
  bonding.
• In an ionic solid you cant tell which atom the
  electrons moved from or to.

15
How to show how they formed

• Its like a jigsaw puzzle.
• I have to tell you what the final formula is.
• You put the pieces together to end up with the
  right formula.
• For example- show how water is formed with
  covalent bonds.

16
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 happy

17
Water

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

H
18
Water

• The second hydrogen attaches
• Every atom has full energy levels

H
H
19
(No Transcript)
20
Diatomic Molecules

• Two atoms of same element joined together
• There are elements that come in pairs they NEVER
  go out alone.
• These are hydrogen, nitrogen, oxygen, fluorine,
  chlorine, bromine, iodine.
• 1 7 on periodic table
• Double your fun with HOFBrINCl

21
Multiple Bonds

• Sometimes atoms share more than one pair of
  valence electrons.
• A double bond is when atoms share two pair (4) of
  electrons.
• A triple bond is when atoms share three pair (6)
  of electrons.

22
Carbon dioxide

• CO2 - Carbon is central atom ( I have to tell
  you)
• Carbon has 4 valence electrons
• Wants 4 more
• Oxygen has 6 valence electrons
• Wants 2 more

C
23
Carbon dioxide

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

C
24
Carbon dioxide

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

C
25
Carbon dioxide

• The only solution is to share more

C
26
Carbon dioxide

• The only solution is to share more

C
27
Carbon dioxide

• The only solution is to share more

C
O
28
Carbon dioxide

• The only solution is to share more

C
O
29
Carbon dioxide

• The only solution is to share more

C
O
30
Carbon dioxide

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

C
O
O
31
Carbon dioxide

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

8 valence electrons
C
O
O
32
Carbon dioxide

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

8 valence electrons
C
O
O
33
Carbon dioxide

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

8 valence electrons
C
O
O
34
How to draw them

• Add up all the valence electrons.
• Write the symbols of all elements in their
  places.
• Place two electrons (dots) between all atoms
  (they MUST be bonded).
• Subtract the electrons you have used from the
  total.
• Place remaining electrons (dots) on individual
  atoms to make an octet.
• If you cant do this with the remaining
  electrons, add another bond and try again.

35
Examples

• NH3
• N - has 5 valence electrons
• H each has 1 valence electron
• NH3 has 53(1) 8

N
H
36
Examples

• Draw in the bonds
• 8e- - 6e- 2e-
• Place remaining 2 electons on N, since it is not
  full.

H
H
H
N
37
Examples

• All 8 electrons are accounted for
• Everything is full

H
N
H
H
38
Examples

• HCN C is central atom
• N - has 5 valence electrons
• C - has 4 valence electrons
• H - has 1 valence electron
• HCN has 541 10

39
HCN

• Put in single bonds
• Remaining 6 electrons will not complete both
  octets.
• Add a bond - must go between C and N, because H
  is full.

N
H
C
40
HCN

• Put in double bond
• Leaves 4 more electrons still wont fill C and
  N
• Add another bond - must go between C and N

N
H
C
41
HCN

• Put in triple bonds
• 2 electrons left
• Must go on N, C is full now
• Molecule is complete (all atoms full, 10
  electrons used.)

N
H
C
42
Another way of indicating bonds

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

H
H
O
H
H
O
is the same as
43
Structural Examples

• C has 8 electrons because each line is 2
  electrons
• Ditto for N
• Ditto for C here
• Ditto for O

H C N
H
C O
H
44
Polar Molecules

• Molecules with ends

45
Polar Molecules

• Molecules with a positive and a negative end
• Requires two things to be true
• The molecule must contain polar bonds
• This can be determined from differences in
  electronegativity.
• Symmetry can not cancel out the effects of the
  polar bonds.
• Must determine geometry first.

46
Polar Bonds

• One atom in the bond is pulling harder on the
  electrons that are shared than the other. As a
  result, the electrons are shared unequally.
• In a nonpolar bond, both atoms have equal or
  nearly equal strength, so electrons are shared
  equally.

47
Polarity

• Measured in terms of differences in
  electronegativity
• Remember that electronegativity is the amount of
  pull on the electrons!
• Find electronegativity of elements on periodic
  table.
• Subtract to find difference
• 0.0 to 0.4 nonpolar covalent bond
• 0.5 and up polar covalent bond
• Ionic bonds will have greater differences yet,
  usually above 1.7

48
Showing Polarity

• There are two ways to show that a bond is polar
• Use an arrow with a positive end
• You can use a partial positive symbol and a
  partial negative symbol
• ? ?-
• Both mean the same thing!!!

49
Is it polar?

• HF
• H2O
• NH3
• CH4
• CO2
• N2

50
Intermolecular Forces

• What holds molecules to each other

51
Intermolecular Forces

• These are the forces BETWEEN molecules that hold
  them together.
• INTRAmolecular bonding is what holds the atoms
  within the molecule together
• They are what make solid and liquid molecular
  compounds possible.

52
van der Waal Forces

• The weakest are called van der Waals forces
• EVERY molecule has these forces no exceptions
• There are two kinds
• Dispersion forces
• Dipole Interactions
• You just need to know the main group

53
Disperson forces

• Temporary dipole is formed by moving electrons
• Bigger molecules more electrons
• More electrons means stronger forces, since
  electrons are more likely to shift and create
  temporary dipole
• Fluorine is a gas
• Bromine is a liquid
• Iodine is a solid

54
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.

55
Dipole Interactions
d d-
56
Hydrogen bonding

• Are the attractive force caused by hydrogen
  bonded to F, O, or N.
• F, O, and N are very electronegative so it is a
  very strong dipole.
• The hydrogen partially share with the lone pair
  in the molecule next to it.
• The strongest of the intermolecular forces.

57
Hydrogen Bonding
58
Hydrogen bonding
59
Circle those elements with Hydrogen Bonding

• H2O
• CH4
• NH3

HCl HF H2S
60
Circle those elements with Hydrogen Bonding

• H2O
• CH4
• NH3

HCl HF H2S