Chemical Bonding

1
Chemical Bonding
2
What is a Chemical Bond?

• a bond is a force of attraction between two atoms
• the source of the force are attractions and
  repulsions between electrons and the atomic
  nucleii
• before we get into detail, do the Ball Stick
  activity, to get an idea of what molecules look
  like.

3
Review of Grade 11

• Quantum theory
• electron configuration
• orbital diagrams
• Lewis diagrams

4
Quantum Mechanics and Atomic Orbitals

• Quantum mechanics is a mathematical treatment
  into which both the wave and particle nature of
  matter could be incorporated.

5
Quantum Mechanics

• since the electron is both a wave and a particle
  it is impossible to give its location or speed
  with certainty.
• gives a probability density map of where an
  electron has a certain statistical likelihood of
  being at any given instant in time.

6
Quantum Numbers

• The probability map reveals the atomic orbitals,
  and their corresponding energies.
• An orbital is described by a set of three quantum
  numbers.

7
Principal Quantum Number, n

• This relates to the energy of the electron
• As n becomes larger, the atom becomes larger and
  the electron is further from the nucleus.
• This is directly related to the period of the
  atom on the Periodic Table

8
Angular momentum quantum number, l

• This quantum number defines the shape of the
  orbital.
• There are 4 shapes
• s - begins at n 1
• p - begins at n 2
• d - begins at n 3
• f - begins at n 4
• Theoretical g, h, i, etc. orbitals exist, but no
  atoms have been created to use them.

9
Magnetic Quantum Number, ml

• Magnetic quantum numbers give the
  three-dimensional orientation of each orbital.
• s - has 1 orientation
• p - has 3 orientations
• d - has 5 orientations
• f - has 7 orientations

10
s Orbitals

• Spherical in shape.

11
p Orbitals

• propeller shaped
• Have two lobes with a node between them.

12
d Orbitals

• Four of the five orbitals have 4 lobes the other
  resembles a p orbital with a doughnut around the
  center.

13
f orbitals (flowers)
14
(No Transcript)
15
Spin Quantum Number, ms

• electrons have spin, which creates a magnetic
  field
• there are two spin states possible, 1/2 and -1/2
• a single orbital can hold a maximum of two
  electrons, which must have opposite spin (Pauli
  Exclusion Principle).

16
Electron Address

• thus every electron location is defined in terms
  of 4 things
• Principal Quantum Number - 1 to 7
• Angular Quantum Number s, p, d or f
• Magnetic Quantum Number implied by number of
  electrons in each shape s has 2, p has 6, d has
  10 and f has 14
• Spin Quantum Number why each orbital can
  contain 2 electrons

17
Electron Configurations

• Electrons tend to occupy the lowest energy
  orbitals.
• Thus the electron configuration of an atom is the
  arrangement of the electrons from the lowest
  energy level to the highest.

18
(No Transcript)
19
Electron Configurations

• Consist of
• Number denoting the energy level.
• Letter denoting the type of orbital.
• Superscript denoting the number of electrons in
  those orbitals.
• For instance
• Iron (Fe) contains 26 electrons
• 1s22s22p63s23p64s23d6
• watch the order of filling

20
Electron Configuration

• Potassium - 19 electrons
• 1s22s22p63s23p64s1
• Silver - 47 electrons
• 1s22s22p63s23p64s23d104p65s24d9
• Tungsten - 74 electrons
• 1s22s22p63s23p64s23d104p65s24d105p66s24f145d4
• Plutonium - 94 electrons
• 1s22s22p63s23p64s23d104p65s24d105p66s24f145d10
  6p67s25f6

21

• Write the correct electron configuration for the
  following
• Si, S, P, Ca, As, Fe, Br, Kr, At, U, Na1, F1-,
  Ne

22
Electron Configuration

• Si - 14 e1- 1s22s22p63s13p3
• S - 16 e1- 1s22s22p63s23p4
• P - 15 e1- 1s22s22p63s23p3
• Ca - 20 e1- 1s22s22p63s23p64s14p1
• As - 33 e1- 1s22s22p63s23p64s23d104p3
• Fe - 26 e1- 1s22s22p63s23p64s23d6
• Br - 35 e1- 1s22s22p63s23p64s23d104p5
• Kr - 36 e1- 1s22s22p63s23p64s23d104p6
• At - 85 e1- 1s22s22p63s23p64s23d104p65s24d105p
  66s24f145d106p5
• U - 92 e1- 1s22s22p63s23p64s23d104p65s24d105p66s
  24f145d106p67s25f4
• Na1 - 10 e1- 1s22s22p6
• F1- - 10 e1- 1s22s22p6
• Ne - 10 e1- 1s22s22p6

23
Electron Promotion

• promotion of an outer s electron to the
  adjacent p orbital.
• turns non-bonding electrons into bonding
  electrons
• allows atoms to make more chemical bonds and
  achieve a lower energy
• applies to elements from groups 2, 13 and 14 only
• for these elements promotion is the rule

24
Electron Promotion

• Element Unhybridized Hybridized
• beryllium 1s22s2 1s22s12p1
• boron 1s22s22p1 1s22s12p2
• carbon 1s22s22p2 1s22s12p3

25
Orbital Diagrams

• are another way to illustrate the position of
  electrons.
• They are best learned by comparison with electron
  configuration
• Na (11 protons, 11 electrons)
• electron configuration 1s22s22p63s1
• orbital diagram
• 1s 2s 2p 3s
• ?? ?? ?? ?? ?? ?

26
Orbital Diagrams

• Representative
• Group Element Electron configuration
  Orbital Diagram
• 1s 2s 2p
• 1 lithium 1s22s1 ?? ?
• 2 beryllium 1s22s12p1 ?? ? ?
  
• 13 boron 1s22s12p2 ?? ? ? ?
• 14 carbon 1s22s12p3 ?? ? ? ? ?
• 15 nitrogen 1s22s22p3 ?? ?? ? ? ?
• 16 oxygen 1s22s22p4 ?? ?? ?? ? ?
• 17 fluorine 1s22s22p5 ?? ?? ?? ?? ?
• 18 neon 1s22s22p6 ?? ?? ?? ?? ??

27
(No Transcript)
28
(No Transcript)
29
(No Transcript)
30

• Na1, F1-, Ne are all the same

31
Electron dot (Lewis) diagrams

• gives information only concerning the valence
  electrons.
• the electrons on the outside of an atom
• the electrons responsible for bonding
• the electrons gained or lost when an atom
  ionizes.
• the electrons in the s and p orbitals of the
  highest energy level reached by the electrons of
  an atom.
• In this class when valence electrons are
  mentioned, the only elements concerned are those
  in groups 1, 2, and 13 through 18.
• all atoms in the same group have the same Lewis
  diagram !!

32
Lewis diagrams

• group 1
• group 2
• group 13
• group 14
• group 15
• group 16
• group 17
• group 18

33
Lewis diagrams

• Repeat last assignment, making lewis diagrams.

34
Lewis Diagram Answers

• Fe and U have no lewis diagrams

35
What do Orbital and Lewis diagrams tell us?

• both give information about valence electrons.
• if valence electrons are paired, they cannot be
  used for bonding with other atoms. They are
  lone-pair electrons.
• unpaired valence electrons are bonding electrons.

36

• 5 valence electrons
• 1 lone pair
• 3 bonding electrons this atom makes 3 chemical
  bonds.

37

• 4 valence electrons
• 0 lone pair
• 4 bonding electrons this atom makes 4 chemical
  bonds

38

• 7 valence electrons
• 3 lone pair
• 1 bonding electrons this atom makes 1 chemical
  bond

39

• 8 valence electrons
• 4 lone pair
• 0 bonding electrons this atom makes 0 chemical
  bonds

40
Valence level expansion

• Some compounds occur which cannot be easily
  explained
• PF5, SF6, ClF7, ArF8
• in each case the number of chemical bonds is
  equal to the number of valence electrons.
• this can only happen if electrons are promoted to
  a higher energy level. In this case it is the
  adjacent d orbital.

41
PF5

• the normal orbital diagram looks like this
• with valence level expansion it looks like this

42
Valence level expansion

• includes
• elements of groups 15 to 18, from period 3 down
• periods 1 and 2 do not have a d orbital to
  promote to.
• Please note that valence level expansion is the
  exception, not the rule.

43
Chemical Bonding

• is all about the electrons.
• in most of our discussions we will concentrate
  the valence electrons.
• since there are 4 of these orbitals in any
  quantum it requires 8 electrons (an octet) to
  fill them.
• the tendency of atoms to try to fill out the
  outer s and p shells is the octet rule.
• exception atoms from groups 2 and 13

44
Chemical Bonding

• Bonding can happen in one of two ways
• To share electrons -
• The outer orbitals of 2 atoms overlap so that
  each atom is in the vicinity of a full set of
  valence electrons.
• This type of bonding is called covalent bonding.

45
(No Transcript)
46
Chemical bonding

• Atoms gain or lose electrons to arrive at a full
  set of valence electrons.
• When atoms gain or lose electrons they become
  ions.
• Ions are attracted to ions of opposite charge and
  repelled by ions of the same charge.
• This type of bonding is called ionic bonding

47
(No Transcript)
48
How Many Bonds ???

• the number of bonds made by atoms in the s and
  p blocks of the Periodic Table is determined by
  how many electrons they are away from an octet

49
How Many Bonds ???
Group Valence Electrons Covalent Bonds Ionic Bonds
1 1 1 1 (1)
2 2 21 2 (2)
13 3 31 3 (3)
14 4 41 -
15 5 3 (52) 3 (3-)
16 6 2 (62) 2 (2-)
17 7 1 (72) 1 (1-)
18 8 0 (82) 0
1electron promotion 2valence level expansion 1electron promotion 2valence level expansion 1electron promotion 2valence level expansion 1electron promotion 2valence level expansion
50

• most of our discussion will be centred on
  covalent bonding.

51
Electronegativity

• is a measure of how strongly an atom is holding
  on to its valence electrons.
• If an atom loses an electron fairly easily it has
  a low electronegativity (and tends to be a
  cation).
• If an atom tends not to lose electrons, but tends
  to steal them from other atoms (and become an
  anion) it has a high electronegativity.

52

• To determine what type of bond exists between two
  atoms you subtract their respective
  electronegativities
• if the electronegativity difference is 0.2 or
  less, the bond is covalent
• if the electronegativity difference is 1.7 or
  greater the bond is ionic.

53
Electronegativity

• If the electronegativity difference between two
  atoms is between 0.3 and 1.6 the bond is
  polar-covalent.
• The greater the electronegativity difference the
  greater the ionic character of the bond

54
Polar-Covalent Bonding

• Covalent bonding implies equal sharing of
  electrons.
• If sharing is not equal, the electrons in a bond
  will spend more time with one atom than the
  other.

55
Polar-Covalent Bonding

• The atom where the electrons spend more time will
  have a net negative charge, while the atom at the
  other end of the bond will be positive.
• This type of bond is polar-covalent.

56
(No Transcript)
57
Assignment

• Determine the electronegativity difference for
  each chemical bond. If the bond is polar
  covalent draw an arrow in the direction of the
  dipole, from positive to negative
• C - H
• END 2.5 - 2.1 0.4
• polar covalent bond
• N - H B - F S - O
• P - H Si - Cl Cu - Br N - I
• Br - Cl O - H C - Cl C - O

58

• N - H
• END 3.0 2.1 0.9
• polar covalent bond N - H
• B - F
• END 2.0 4.0 2.0
• ionic bond B3F1-
• S O
• END 2.5 3.5 1.0
• polar covalent bond S - O

59

• P - H
• END 2.1 2.1 0.0
• covalent bond P H
• Si - Cl
• END 1.8 3.0 1.2
• polar covalent bond Si - Cl
• Cu - Br
• END 1.9 2.8 0.9
• polar covalent bond Cu - Br

60

• N I
• END 3.0 2.5 0.5
• polar covalent bond N - I
• Br Cl
• END 2.8 3.0 0.2
• covalent bond Br Cl
• O H
• END 3.5 2.1 1.4
• polar covalent bond O H

61

• C Cl
• END 2.5 3.0 0.5
• polar covalent bond C Cl
• C O
• END 2.5 3.5 1.0
• polar covalent bond C O

62
Back to the Ball Stick

• determine the electronegativity difference and
  bond type for each bond in each of the molecules
  in the activity.

63
Covalently-Bonded Structures

• we now have to consider molecules made of several
  atoms.
• most of the following discussion will concern
  itself with molecules made with covalent or
  polar-covalent bonds.
• ionic bonds (and others) will return later in the
  unit.

64
Lewis Structures

• Lewis structures allow us to predict how atoms
  will come together to make molecules.
• Lewis structures are representations of molecules
  showing all electrons, bonding and nonbonding.

65
Writing Lewis Structures

• Find the sum of valence electrons of all atoms in
  the polyatomic ion or molecule.
• If it is an anion, add one electron for each
  negative charge.
• If it is a cation, subtract one electron for each
  positive charge.

• PCl3

5 3(7) 26
66
Writing Lewis Structures

1. The central atom is the least electronegative
   element that isnt hydrogen. Connect the outer
   atoms to it by single bonds.

Keep track of the electrons 26 ? 6 20
67
Writing Lewis Structures

1. Fill the octets of the outer atoms.

Keep track of the electrons 26 ? 6 20 ? 18 2
68
Writing Lewis Structures

1. Fill the octet of the central atom.

Keep track of the electrons 26 ? 6 20 ? 18
2 ? 2 0
69
Writing Lewis Structures

• If you run out of electrons before the central
  atom has an octet
• form multiple bonds until it does.

70
Exceptions to Octet Rule

• Electron Promotion
• group 2 central atom will have 4 electrons
• group 13 central atom will have 6 electrons
• Valence Level Expansion
• group 15 central atom will have 10 electrons
• group 16 central atom will have 12 electrons
• group 17 central atom will have 14 electrons
• group 18 central atom will have 16 electrons

71
Writing Lewis Structures

• Write Lewis Structures for the following
  molecules

F2 SiF4
BCl3 C3H8
SeH2 C3H6
MgH2 C3H4
72
Molecular Shapes

• The shape of a molecule plays an important role
  in its reactivity.
• By noting the number of bonding and nonbonding
  electron pairs we can easily predict the shape of
  the molecule.

73
What Determines the Shape of a Molecule?

• Simply put, electron pairs, whether they be
  bonding or nonbonding, repel each other.
• By assuming the electron pairs are placed as far
  as possible from each other, we can predict the
  shape of the molecule.

74
Valence Shell Electron Pair Repulsion Theory
(VSEPR)

• The best arrangement of a given number of
  electron domains is the one that minimizes the
  repulsions among them.

75
Molecular Geometries

• the shape of any molecule can be predicted based
  on two things
• Group number of the central atom
• number of atoms bonded to the central atom

76
Group 2 Geometries
Central Atom Bonding Electrons Lone Pair Electrons Bond Type Shape Example
Magnesium 2 0 all single linear MgI2
one double linear MgO

• In this domain, there is only one molecular
  geometry linear.
• NOTE If there are only two atoms in the
  molecule, the molecule will be linear no matter
  what the electron domain is.

77
Group 13 Geometries
Central Atom Bonding Electrons Lone Pair Electrons Bond Type Shape Example
Boron 3 0 all single trigonal planar BI3
one double linear BIO
one triple linear BN

• because there are no lone pair electrons the
  molecular is planar (flat).

78
Group 14 Geometries
Central Atom Bonding Electrons Lone Pair Electrons Bond Type Shape Example
Carbon 4 0 all single tetrahedral CH4
one double trigonal planar COH2
one triple, or two double linear HCN CO2
79
Summary for Groups 2, 13 14

• because there are no lone pair electrons
• central atom bonded to 1 atom linear
• central atom bonded to 2 atoms linear
• central atom bonded to 3 atoms trigonal
  planar
• central atom bonded to 4 atoms
  tetrahedral

80
Group 15 Geometries
Central Atom Bonding Electrons Lone Pair Electrons Bond Type Shape Example
Nitrogen 3 1 all single trigonal pyramidal NH3
one double angular NOH
one triple linear N2
5 0 all single trigonal bipyramidal NCl5
81
Group 16 Geometries
Central Atom Bonding Electrons Lone Pair Electrons Bond Type Shape Example
Oxygen 2 2 all single angular H2O
one double linear O2
6 0 all single octahedral OF6

82
Group 17 18 Geometries

• group 17
• normally makes 1 chemical bond (linear)
• with valence level expansion can make 7 bonds
  (ClF7).
• group 18
• normally makes no bonds
• with valence level expansion can make 8 bonds
  (ArF8).

83
Larger Molecules

• In larger molecules, it makes more sense to talk
  about the geometry about a particular atom rather
  than the geometry of the molecule as a whole.
• This molecule is tetrahedral about the first
  carbon, planar trigonal about the second and
  angular about the single-bonded oxygen.

84
Larger Molecules

• This approach makes sense, especially because
  larger molecules tend to react at a particular
  site in the molecule.

85
Using VSEPR Theory

• Determine the shape of each molecule, or atom
  within the molecule

F2 SiF4
PCl3 C3H8
SeH2 C3H6
MgH2 C3H4
86
Polarity

• previously we discussed bond dipoles.
• But just because a molecule possesses polar bonds
  does not mean the molecule as a whole will be
  polar.

87
Polarity

• By adding the individual bond dipoles, one can
  determine the overall dipole moment for the
  molecule.
• In other words we can see if the dipoles
  reinforce each other, or cancel out.

88
Polarity
89
Assignment

• Perform the following activities for the
  molecules in the Balls Sticks Activity
• Draw a lewis structure
• Draw a structural diagram and indicate shape.
• Determine the electronegativity difference for
  each bond type and determine whether the bond is
  covalent, polar covalent or ionic
• Determine if they will be polar, or non-polar.

90
Properties of Molecules

• Intermolecular Forces

91
Properties of Molecules

• are determined by how the molecules interact with
  each other.
• How they interact is determined by the forces of
  attraction between molecules.
• These are called intermolecular forces the
  forces which act between molecules, to draw them
  together, forming the various phases of matter.

92
How do we measure force ?

• the best way is by temperature.
• temperature is a measure of kinetic energy the
  lowest temperature represents zero kinetic
  energy.
• for a substance to melt or boil the kinetic
  energy must overcome the intermolecular force.
• the higher the melting temperature or boiling
  temperature, the greater the intermolecular force.

93
Van der Waals Forces

• occur between covalently bonded molecules.
• three kinds
• London disersion
• dipole-dipole attraction
• hydrogen bonding

94
London Dispersion Forces

• are the dominant forces between covalently
  bonded, non-polar molecules
• based on the formation of instantaneous dipoles.
• the more electrons in a molecule, the stronger
  the force.

95
(No Transcript)
96
Relationship of Boiling Point to Number of
Electrons and Molar Mass

• Melting Point Boiling Point
• e- C C
• CH4 10 - 182 -161
• C2H6 18 - 183 - 88
• C3H8 26 - 190 - 44
• C4H10 34 - 138 - 0.5
• C8H18 66 - 57 125

97
London dispersion forces

• are influenced by shape
• Normal Pentane (C5H12)
• m.p. -130C,
• b.p. 36C
• Neopentane (C5H12)
• m.p. -20C,
• b.p. 9C

98
(No Transcript)
99
(No Transcript)
100
(No Transcript)
101
Dipole-Dipole attraction

• is a force which acts between polar molecules
  (ex. H2S).
• results from the attraction of the opposite poles
  of the permanent molecular dipoles.
• These substances generally have higher melting
  and boiling points than non-polar molecules with
  similar molecular weights (or numbers of
  electrons).

102
(No Transcript)
103
Ion-Dipole Interactions

• A fourth type of force, ion-dipole interactions
  are an important force in solutions of ions.
• The strength of these forces are what make it
  possible for ionic substances to dissolve in
  polar solvents.

104
(No Transcript)
105
Dipole-Dipole Interactions

• The more polar the molecule, the higher is its
  boiling point.

106
How Do We Explain This?

• The nonpolar series (SnH4 to CH4) follow the
  expected trend.
• The polar series follows the trend from H2Te
  through H2S, but water is quite an anomaly.

107
Hydrogen Bonding

• a specialized form of dipole-dipole attraction.
• It occurs as when O, N, and F are bonded to H,
  owing to the large electronegativity difference
  
• O - H 3.5 - 2.1 1.4
• N - H 3.1 - 2.1 1.0
• F - H 4.1 - 2.1 2.0

108
Hydrogen Bonding

• This is a stronger force than standard
  dipole-dipole attraction.
• Molecules with hydrogen bonding will have boiling
  points and melting points quite a bit higher than
  molecules that have only dipole-dipole or London
  dispersion forces.
• Hydrogen bonding is responsible for many of the
  unusual properties of water.

109
(No Transcript)
110
(No Transcript)
111
Hydrogen bonding

• is responsible for folding, final structure and
  function of proteins
• holds the DNA strands together, but allows them
  to be unzipped for copying and gene expression.

112
Relationship Between Polarity, Force and Boiling
Point

• e- B.P. (C) Polarity Force
• C2H6 18 - 161 Non-polar London
• Dispersion
• H2S 18 - 60 Polar Dipole - Dipole
• H2O 10 100 Very polar Hydrogen Bonds

113
In Summary

• Non-Polar Molecule
• London Dispersion Force
• More electrons, higher B.P., M.P.
• If isoelectronic, more compact higher M.P., lower
  B.P.
• Polar Molecule
• Dipole-Dipole Attraction
• Higher B.P. M.P. than London
• If both polar, higher electronegativity
  difference has higher B.P. M.P
• If have O-H, N-H, or F-H bond, hydrogen bonding,
  much higher M.P. B.P

114
Ionic Bonding

• generally occurs in compounds of metals and
  non-metals (salts).
• It is the result of the attraction of oppositely
  charged ions.
• The structures formed are very orderly and are
  given the name crystal lattice.
• Ionic solids are called crystals.

115
Ionic Bonding

• No sharing of electrons occurs between the ions
  in the crystal lattice.
• As a result, ionic solids are brittle.
• Ionic solids conduct electricity only in the
  molten state, and not very well.
• Ionic solids are characterized by very high
  melting and boiling points.

116
(No Transcript)
117
Metallic Bonding

• is the bonding which occurs between metals in the
  Periodic Table.
• It is characterized by close packing of the
  atoms, with the electrons delocalized that is,
  they are free to jump from atom to atom, filling
  unoccupied orbitals.

118
(No Transcript)
119
Metallic Bonding

• This free sharing of electrons allows metals to
  conduct electricity freely (copper conducts
  electricity 100 000 times better than molten
  NaCl).
• The free electrons also act as a lubricant,
  allowing metal atoms to slide over one another
  without affecting the integrity of the material.
  Thus metals are malleable and ductile.
• This bond is strong, giving most metals high
  melting and boiling points.
• This bond is also variable.

120
(No Transcript)
121
Network Covalent Bonding

• regular covalent bond, expanded to 2 or 3
  dimensions in a network which is theoretically
  infinite
• Network solids include diamonds, graphite,
  quartz, and most rocks.
• very hard (diamonds are the hardest substance
  known).
• brittle,
• do not conduct electricity.
• crystals.
• very high melting and boiling points

122
(No Transcript)
123
(No Transcript)
124
Intermolecular Forces Affect Many Physical
Properties

• The strength of the attractions between
  particles can greatly affect the properties of a
  substance or solution.

125
Viscosity

• Resistance of a liquid to flow is called
  viscosity.
• It is related to the ease with which molecules
  can move past each other.
• Viscosity increases with stronger intermolecular
  forces and decreases with higher temperature.

126
Surface Tension

• Surface tension results from the net inward
  force experienced by the molecules on the surface
  of a liquid.

127
Vapor Pressure

• At any temperature, some molecules in a liquid
  have enough energy to escape.
• As the temperature rises, the fraction of
  molecules that have enough energy to escape
  increases.

128
Intermolecular Forces in Summary
129
Strength of Attraction
Molecular (London dispersion) 1 to 50
Molecular (dipole-dipole) 5
Molecular (hydrogen bonding) 15
Ionic 50
Metallic 20 to 80
Network Covalent 100
130
Melting and Boiling Point
Molecular (London dispersion) very low (nitrogen boils at - 196 ?C) VARIABLE
Molecular (dipole-dipole) low (H2S boils at -61 ?C)
Molecular (hydrogen bonding) medium (H2O boils at 100 ?C)
Ionic high (NaCl boils at 1413 ?C)
Metallic variable (Hg _at_ 357 ?C, W _at_ 5660 ?C)
Network Covalent very high (SiO2 boils at 2600 ?C)
131
Properties of Solids
Molecular (London dispersion) soft, waxy
Molecular (dipole-dipole) more rigid
Molecular (hydrogen bonding) crystalline, brittle
Ionic long-range crystalline, hard, brittle
Metallic short-range crystalline, ductile, malleable
Network Covalent long-range crystalline, hard, brittle
132
Conductance of Heat and Electricity
Molecular (London dispersion) non-conductor
Molecular (dipole-dipole) non-conductor
Molecular (hydrogen bonding) non-conductor
Ionic conductor in liquid or dissolved phase
Metallic conductor in solid or liquid phase
Network Covalent non-conductor in 3d form, some conductance in 2d
133
Solubility in H2O
Molecular (London dispersion) Not soluble
Molecular (dipole-dipole) Soluble
Molecular (hydrogen bonding) Soluble
Ionic Soluble
Metallic Not soluble
Network Covalent Not soluble
134
Other Stuff
135
Coordinate Covalent bonds

• is defined as a covalent bond where one atom
  provides both of the electrons

136
Polyatomic Ions

• we can use coordinate covalent bond theory to
  explain most ions

137
Resonance

• This is the Lewis structure we would draw for
  ozone, O3.

-
138
Resonance

• But this is at odds with the true, observed
  structure of ozone, in which
• both OO bonds are the same length.
• both outer oxygens have a charge of ?1/2.

139
Resonance

• One Lewis structure cannot accurately depict a
  molecule such as ozone.
• We use multiple structures, resonance structures,
  to describe the molecule.

140
Resonance

• Just as green is a synthesis of blue and yellow
• ozone is a synthesis of these two resonance
  structures.

141
Resonance

• In truth, the electrons that form the second CO
  bond in the double bonds below do not always sit
  between that C and that O, but rather can move
  among the two oxygens and the carbon.
• They are not localized, but rather are
  delocalized.

142
Resonance

• The organic compound benzene, C6H6, has two
  resonance structures.
• It is commonly depicted as a hexagon with a
  circle inside to signify the delocalized
  electrons in the ring.