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Unit 2: Chemical Bonding

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Title: Unit 2: Chemical Bonding


1
Unit 2 Chemical Bonding
  • Chemistry2202

2
Outline
  • Bohr diagrams Lewis Diagrams
  • Types of Bonding
  • Ionic
  • Covalent (molecular)
  • Metallic
  • Network covalent bonding
  1. London Dispersion forces
  2. Dipole-Dipole forces
  3. Hydrogen Bonding
  • VSEPR Theory (Shapes)
  • Physical Properties

3
Bohr Diagrams (Review)
  • How do we draw a Bohr Diagram for
  • The F atom?
  • The F ion?
  • Draw Bohr diagrams for the atom and the ion for
    the following
  • Al S C l Be

4
Lewis Diagrams
  • Lewis Diagrams provide a method for keeping track
    of electrons in atoms, ions, or molecules.
  • AKA Electron Dot diagrams
  • the nucleus (p no), and filled energy levels
    are represented by the symbol
  • dots are placed around the element symbol to
    represent valence electrons

5
Lewis Diagrams
  • eg. Lewis Diagram for F

lone pair



bonding electron
F
lone pair




lone pair
6
Lewis Diagrams
  • lone pair a pair of electrons not available
    for bonding
  • bonding electron a single electron that may be
    shared with another atom

7
Lewis Diagrams
  • eg. Draw Lewis Diagrams for
  • carbon

phosphorus





P
C




sodium

Na
8
Lewis Diagrams
  • For each atom draw the Lewis diagram and state
    the number of lone pairs and number of bonding
    electrons
  • Li Be Al Si
  • Mg N B O

9
Lewis Diagrams for Compounds
  • To draw the LD for a molecule
  • draw the LD for each atom in the molecule
  • the atom with the most bonding electrons is the
    central atom
  • connect the other atoms using single bonds (1
    pair of shared electrons)
  • some molecules may have double bonds or triple
    bonds

10
Lewis Diagrams for Compounds
  • eg. Draw the LD for
  • PH3
  • CF4
  • Cl2O
  • C2H6
  • C2H4
  • C2H2

11
Lewis Diagrams for Compounds
  • eg. Draw the LD for
  • NH3 SiCl4 N2H4 HCN
  • SI2 CO2 N2H2 CH2O
  • POI CH3OH
  • N2 H2 O2

12
Lewis Diagrams for Compounds
  • A structural formula shows how the atoms are
    connected in a molecule.
  • To draw a structural formula
  • replace the bonded pairs of electrons with short
    lines
  • omit the lone pairs of electrons

13
  • Why is propane (C3H8) a gas at STP while
    kerosene (C10H22) a liquid?

14
  • Why is graphite soft enough to write with while
    diamond is the hardest substance? (both are C)

Graphite
15
(No Transcript)
16
  • Other forms of carbon

17
  • How can you which is real gold and which is
    fools gold (pyrite) by hitting it with a rock?

18
As Slow As Cold Molasses
All Because of Bonding
19
Viscosity of liquids
20
liquids _at_ -30 ºC
21
Malleability, Ductility and Conductivity
A single gram of gold can be stretched into a
wire 3.2km long.
 A gram of gold can be flattened into a sheet
with an area of 6.7 sq ft.
 Silver has the highest electrical
conductivity of any element and the
highest thermal conductivity of any metal. 
22
Melting and Boiling Points
23
  • dihydrogen monoxide
  • pepper demo
  • dd HB
  • teacher tube - IF

24
Bonding
  • Bonding between atoms, ions and molecules
    determines the physical and chemical properties
    of substances.
  • Bonding can be divided into two categories
  • - Intramolecular forces
  • - Intermolecular forces

25
Bonding
  • Intramolecular forces are forces of attraction
    between atoms or ions.
  • Intramolecular forces include
  • ionic bonding
  • covalent bonding
  • metallic bonding
  • network covalent bonding

26
Bonding
  • Intermolecular forces occur BETWEEN different
    molecules.
  • Intermolecular forces include
  • London Dispersion Forces
  • Dipole-Dipole forces
  • Hydrogen Bonding

27
Ionic and Covalent Bonding
ThoughtLab p. 161 Identify s 1 - 6
28
Ionic Bonding
Mar. 11
  • Occurs between cations and anions usually
    metals and non-metals.
  • An ionic bond is a force of attraction between
    positive and negative ions.
  • Properties
  • conduct electricity as liquids and in solution
  • hard crystalline solids
  • high melting points and boiling points
  • brittle

29
Ionic Bonding
Mar. 11
  • In an ionic crystal the ions pack tightly
    together. (p. 167)
  • The repeating 3-D distribution of cations and
    anions is called an ionic crystal lattice.

30
Ionic Bonding
Mar. 11
  • Each anion is attracted to six or more cations at
    once.
  • The same is true for the individual cations.

31
Ionic Bonding
Mar. 11
32
Covalent Bonding
Mar. 11
  • Occurs between non-metals in molecular compounds.
  • Atoms share bonding electrons to become more
    stable (noble gas structure).
  • A covalent bond is a simultaneous attraction by
    two atoms for a common pair of valence electrons.

33
Covalent Bonding
Mar. 11
  • Molecular compounds
  • have low melting and boiling points.
  • exist as distinct molecules.

34
Covalent Bonding
Mar. 11
  • do not conduct electric current in any form

35
Property Ionic Molecular
Type of elements Metals and nonmetals Non-Metals
Force of Attraction Positive ions attract negative ions Atoms attract a shared electron pair
Electron movement Electrons move from the metal to the nonmetal Electrons are shared between atoms
State at room temperature Always solids Solids, liquids, or gas
Mar. 11
36
Property Ionic Molecular
Solubility Soluble or low solubility Soluble or insoluble
Conductivity in solid state None None
Conductivity in liquid state Conducts None
Conductivity in solution Conducts None
Mar. 11
37
Metallic Bonding (p. 171)
38
Metallic Bonding (p. 171)
39
Metallic Bonding (p. 171)
  • metals tend to lose valence electrons.
  • valence electrons are loosely held and frequently
    lost from metal atoms.
  • this produces in positive metal ions surrounded
    by freely moving valence electrons.
  • metallic bonding is the force of attraction
    between the positive metal ions and the mobile or
    delocalised valence electrons

40
Metallic Bonding
41
Metallic Bonding
  • This theory of metallic bonding is called the
    Sea of Electrons Model or Free Electron Model

42
Metallic Bonding
  • Metallic bonding theory accounts for properties
    of metals
  • electrical conductivity
  • - electric current is the flow of electrons
  • - metals are the only solids in which electrons
    are free to move
  • solids
  • attractive forces between positive cations and
    negative electrons are very strong

43
Metallic Bonding
  • malleability and ductility
  • metals can be hammered into thin
    sheets(malleable) or drawn into thin
    wires(ductile).
  • metallic bonding is non-directional such that
    layers of metal atoms slide past each other under
    pressure.

44
Network Covalent Bonding (p. 199)
  • occurs in 3 compounds (memorize these)
  • diamond Cn
  • carborundum SiC
  • quartz SiO2
  • produces large molecules with covalent bonding in
    3d
  • each atom is held in place in 3d by a network of
    other atoms

45
Network Covalent bonding
  • Properties
  • the highest melting and boiling points
  • the hardest substances
  • brittle
  • do not conduct electric current in any form

46
  • Strongest
  • 1. Network Covalent (Cn ,SiO2 , SiC)
  • 2. Ionic bonding(metal nonmetal)
  • 3. Metallic bonding (metals)
  • 4. Molecular (nonmetals)
  • Weakest

47
Valence Shell Electron Pair Repulsion or VSEPR
theory
  • The shape of molecules is determined by the
    arrangement of valence electron pairs around the
    atoms in a compound.
  • The shapes are the result of REPULSION between
    pairs of valence electrons.
  • Valence electron pairs move as far away from each
    other as possible.

48
Valence Shell Electron Pair Repulsion or VSEPR
theory
  • There are 5 shapes that can be determined by the
    of bonds and of lone pairs on the central
    atom.

49
  • 1. Tetrahedral (4 bonds 0 lone pairs)

50
  • 2. Pyramidal (3 bonds 1 lone pair)

51
  • 3. V-shaped (2 bonds 2 lone pairs)

52
  • 4. Trigonal Planar (3 bonds 0 lone pairs)

53
  • 5. Linear (2 bonds 0 lone pairs)

54
  • For each molecule below draw the Lewis diagram
    and the shape diagram.
  • 1 central atom
  • HOCl H2Se H2SiO
  • NBr3 CHCl3 SiH4
  • PBr3 HCN I2
  • 2 central atoms
  • C2F4 C2H6 CH3OH
  • C2H2 H2O2

55
Electronegativity (EN - p. 174)
  • EN is a measure of the attraction that an atom
    has for shared electrons.
  • A higher EN means a stronger attraction or
    electrostatic pull on valence electrons
  • EN values increase as you move
  • from left to right in a period
  • up in a group or family

56
Increases
57
Electronegativity Covalent Bonds
  • polar covalent bond
  • a bond between atoms with different EN
  • the shared electron pair is attracted more
    strongly to the atom with the higher EN

d-
d
58
Electronegativity Covalent Bonds
  • nonpolar covalent bond
  • occurs between atoms with same EN
  • the shared electron pair is attracted more
    strongly to the atom with the higher EN
  • the separation of charge or bond dipole is shown
    using an arrow pointing toward the more
    electronegative atom.
  • the Greek letter delta (d) indicates partial
    charges
  • Complete s 7 9 on p.178

59
Electronegativity and Ionic Bonds
  • Because the EN of metals is so low, metals lose
    electrons to form cations
  • Nonmetals gain electrons to form anions because
    their EN is relatively high
  • When ions form, the resulting electrostatic force
    is an ionic bond

60
Electronegativity and Covalent Bonds
  • Atoms in covalent compounds can either have the
    same EN
  • eg. Cl2 , PH3, NCl3
  • OR different EN
  • eg. HCl

61
Electronegativity and Covalent Bonds
  • Atoms with the same EN have the same attraction
    for shared valence electrons.
  • Covalent bonds resulting from equal sharing of
    the bonding electron pairs are called Nonpolar
    Covalent Bonds
  • Atoms with different EN attract the shared
    valence electron pair at different strengths.
  • (higher EN has a stronger attraction
  • for the shared electron pair)

62
Electronegativity and Covalent Bonds
  • eg. HCl
  • Cl has a higher EN
  • the bonding electron pair is pulled closer to the
    chlorine atom
  • this produces slight positive and negative
    charges within the bond
  • these charges are referred to as partial
    charges and are denoted with the Greek letter
    delta (d).

63
Electronegativity and Covalent Bonds
  • The region around the chlorine atom will be
    slightly negative
  • The region around the hydrogen will be slightly
    positive.

64
Electronegativity and Covalent Bonds
  • Because the bond is polarized into a positive
    area and a negative area the bond has a bond
    dipole.
  • an arrow points to the atom with the higher EN.
  • Covalent bonds resulting from unequal sharing of
    bonded electron pairs are Polar Covalent Bonds.

65
Electronegativity and Covalent Bonds
  • eg. H2O

66
Electronegativity and Covalent Bonds
  • eg. HF

67
Electronegativity Homework
  • p. 178 s 7, 8, 9
  • p. 180 s 1, 2, 3

68
Bond Energy (pp. 179-180)
  • 1. Describe the forces of attraction and
    repulsion present in all bonds.
  • 2. What is bond length?
  • 3. Define bond energy.
  • 4. Which type of bond has the most energy?
  • 5. How can bond energy be used to predict
    whether a reaction is endothermic or exothermic?

69
Test Outline
  • Bohr Diagrams (atoms ions)
  • Lewis Diagrams (Electron Dot)
  • Ion Formation
  • Ionic Bonding, Structures Properties
  • Covalent Bonding, Structures Properties

70
Test Outline
  • Metallic Bonding Theory Properties
  • Network Covalent Bonding Properties
  • Electronegativity
  • Bond Dipoles Polar Molecules
  • VSEPR Theory
  • LD, DD, H-bonding
  • Predicting properties (bp, mp, etc.)

71
Molecular Dipoles
  • The vector sum of all the bond dipoles in a
    molecule is a Molecular Dipole
  • A Polar Molecule has a molecular dipole that
    points toward the more electronegative end of the
    molecule.
  • eg. H2O

72
Molecular Dipoles
  • Nonpolar molecules DO NOT have molecular dipoles.
  • This occurs when
  • - bond dipoles cancel
  • - there are no bond dipoles
  • To determine whether a molecule is polar
  • - draw the LD and the shape diagram
  • - draw the bond dipoles and determine whether
    they cancel

eg. CO2
eg. PH3
73
Molecular Dipoles
  • See Handout 1

74
Intermolecular Forces
Mar. 31
75
Mar. 31
  • Strongest bonds Highest mp and bp
  • 1. Network Covalent (Cn SiO2 SiC)
  • 2. Ionic bonding(metal nonmetal)
  • 3. Metallic bonding (metals)
  • 4. Molecular (nonmetals)
  • Weakest bonds Lowest mp and bp
  • - Intermolecular forces present

76
Mar. 31
  • To compare mp and bp in covalent compounds you
    must use
  • - London Dispersion forces (p. 204)
  • (all molecules)
  • - Dipole-Dipole forces (pp. 202, 203)
  • (polar molecules)
  • - Hydrogen Bonding (pp. 205, 206)
  • (H bonded to N, O, or F)

77
Intermolecular Forces (p. 202)
Mar. 31
78
Intermolecular Forces
Mar. 31
  • Covalent compounds have low mp and bp because
    forces between molecules in covalent compounds
    are very weak.
  • Intermolecular forces were studied by the Dutch
    physicist Johannes van der Waals
  • In his honor, two types of intermolecular force
    are called Van der Waals forces.
  • Intermolecular forces can be used to explain
    physical properties of covalent compounds.

79
1. London Dispersion Forces
Apr. 1
  • LD forces exist in ALL molecular elements
    compounds.
  • The positive charges in one molecule attract the
    negative charges in a second molecule.
  • The temporary dipoles caused by electron
    movement in one molecule attract the temporary
    dipoles of another molecule.

80
1. London Dispersion Forces
Apr. 1
  • The strength of these forces depends on
  • the number of electrons
  • more electrons produce stronger LD forces that
    result in higher mp and bp
  • eg. CH4 is a gas at room temperature.
  • C8H18 is a liquid at room temperature.
  • C25H52 is a solid at room temperature.
  • Account for the difference.

81
1. London Dispersion Forces
Apr. 1
Two molecules that have the same number of
electrons are isoelectronic eg. C2H6 and
CH3F
82
1. London Dispersion Forces
Apr. 1
  • b) shape of the molecule
  • molecules that fit together better will
    experience stronger LD forces
  • eg. Cl2 vaporizes at -35 ºC while C4H10
    vaporizes at -1 ºC. Use bonding to account for
    the difference.

83
2. Dipole-dipole Forces
Apr. 5
  • occur between polar molecules
  • the d end of one polar molecule is attracted to
    the d- end of another polar molecule (
    vice-versa)
  • eg. Which has the higher boiling point CH3F
    or C2H6 ?

84
Apr. 5
  • In the liquid state, polar molecules are oriented
    such that oppositely charged ends of the
    molecules are close to each other.

p. 202
85
3. Hydrogen Bonds
Apr. 5
  • a special type of dipole-dipole force
  • (about 10 times stronger)
  • - only occurs BETWEEN MOLECULES that contain H
    directly bonded to F, O, or N
  • ie. the molecule contains at least one H-F,
    H-O, or H-N covalent bond.

86
3. Hydrogen Bonds
Apr. 5
  • the hydrogen bond occurs between the H atom of
    one molecule and the N, O, or F of a second
    molecule.
  • eg. Arrange these from highest to lowest
    boiling point
  • C3H8 C2H5OH C2H5F

87
Apr. 5
p. 206
88
  • NOTE To compare mp and bp in covalent
    compounds you must use
  • - London Dispersion forces
  • (all molecules)
  • - Dipole-Dipole forces
  • (polar molecules)
  • - Hydrogen Bonding
  • (H bonded to N, O, or F)

89
  • WorkSheet Bonding 4
  • p. 225 s 9 10
  • p. 226 s 12 14,

90
Intermolecular Forces
  • 1. Use intermolecular forces to explain the
    following
  • a) Ar boils at -186 C and F2 boils at -188 C .
  • b) Kr boils at -152 C and HBr boils at -67 C.
  • c) Cl2 boils at -35 C and C2H5Cl boils at 13 C
    .
  • 2. Examine the graph on p. 210
  • a) Account for the increase in boiling point for
    the hydrogen compounds of the Group IV elements.
  • b) Why is the trend different for the hydrogen
    compounds of the Group V, VI, and VII elements?
  • c) Why are the boiling points of the Group IVA
    compounds consistently lower than the others.

91
p. 210
92
  • 3. Which substance in each pair has the higher
    boiling point. Justify your answers.
  • (a) SiC or KCl
  • (b) RbBr or C6H12O6
  • (c) C3H8 or C2H5OH
  • (d) C4H10 or C2H5Cl

93
Summary
  • Strongest - Network Covalent
  • - Ionic
  • - Metallic
  • Weakest - Covalent
  • ? LD forces (all molecules)
  • ? DD forces (polar molecules)
  • ? H-Bonding (H bonded to N, O, or F)

94
Ion-Dipole Forces
  • An ion-dipole force is the force of attraction
    between an ion and a polar molecule (a dipole).

95
Ion-Dipole Forces
  • NaCl dissolves in water because the attractions
    between the Na and Cl- ions and the partial
    charges on the H2O molecules are strong enough to
    overcome the forces that bind the ions together.

96
Dispersion (London) Forces
  • Bond vibrations, which are part of the normal
    condition of a non-polar molecule, cause
    momentary, uneven distribution of charge
  • a non-polar becomes slightly polar for an
    instant, and continues to do so in a random but
    constant basis.

97
Dispersion (London) Forces
  • At the instant that one non-polar molecule is
    in a slightly polar condition, it is capable of
    inducing a dipole in a nearby molecule
  • This force of attraction is called a dispersion
    force.

98
Dispersion (London) Forces
  • Two factors affect the magnitude of dispersion
    forces
  • of electrons in the molecule
  • Vibrations within larger molecules with more
    electrons than smaller molecules can easily cause
    an uneven distribution of charge.
  • dispersion forces between these larger molecules
    are thus stronger, which raises the boiling point
    for larger molecules.

99
Dispersion (London) Forces
  • The shape of the molecule
  • molecules with spherical shape have a smaller
    surface area than a straight chain molecule with
    the same number of electrons
  • substances with spherical moleculular shape will
    have weaker dispersion forces and a lower boiling
    point.
  • London dispersion forces are responsible for the
    formation and stabilization of the biological
    membranes surrounding every living cell.

100
Hydrogen Bonding
  • In order to form a hydrogen bond, a hydrogen atom
    must be bonded to a highly electronegative atom
    such as oxygen, nitrogen, or fluorine.
  • These bonds are very polar, and since hydrogen
    has no other electrons, the positive proton, H,
    is exposed and can become strongly attracted to
    the negative end of a nearby dipole.
  • A hydrogen bond is an electrostatic attraction
    between the nucleus of a hydrogen atom, bonded to
    fluorine, oxygen, or nitrogen and the negative
    end of a dipole nearby.

101
Hydrogen Bonding
H
H
d
d
H
H
O
d
d
O
H
d-
d-

102
Hydrogen Bonding
  • In biological systems, these polar bonds are
    often parts of much larger molecules (ie. N H
    bonds and C O bonds found in biological
    molecules)

103
Hydrogen Bonding in Water
  • Hydrogen bonds between the hydrogen atoms in one
    water molecule and the oxygen atom in another
    account for many unique properties of water.

H
H
d
d

H
H
O
d
O
d
H
d-
d-
104
Hydrogen Bonding in Water
  • In liquid water, each water molecule is hydrogen
    bonded to at least four other water molecules.
  • The large number of bonds between water
    molecules makes the net attractive force quite
    strong

105
Hydrogen Bonding in Water
  • the strong attractive forces are responsible for
    the relatively high boiling point of water.
  • The water molecules are farther apart in ice
    then they are in liquid water making ice less
    dense than liquid water.

106
Hydrogen Bonding in Water
  • Hydrogen bonds force water molecules into the
    special hexagonal, crystalline structure of ice
    when the temperature is below 4 degrees celcius.
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