Bonding

1
Bonding

• Chapters 7-8

2
Octet Rule

• Atoms tend to lose or gain electrons to achieve a
  full valence shell (8)
• Exception First Energy Level is full with 2
  electrons

3
Electron Dot Structures

• Diagrams that show valence electrons, usually as
  dots
• AKA Lewis Electron Dot Diagrams
• Rules
• Start on any side
• First two get paired together
• Next three are separated
• Fill in as needed

O
4
Examples
H
He
F
Ne
N
Ar
Na
Cl
5
Ions

• Atoms that have gained or lost electrons, and now
  have a charge
• Must show charge

Na
F-
O-2
6
Practice

• Draw Lewis Electron Dot Structures for the
  following atoms and ions
• Li, B, Mg, Al, P, S, Cl, Br, Kr
• H, Li, Mg2, Al3, O2-, F-, Cl-, S2-, P3-

7
Practice
8
Practice
H
Li
Mg2
Al3
9
(No Transcript)
10
Compounds

• Two Main Types of Compounds
• Ionic
• Molecular (Covalent)
• Based on type of bonding involved

11
Bonding

• Bond
• Shared or exchanged electrons that hold two atoms
  together
• Three Main Types
• Covalent
• Ionic
• Metallic

12
Covalent Bonds

• Electrons are shared between two atoms to hold
  them together
• Each atom will try to achieve a full valence
  shell
• 2 nonmetals
• Two types of covalent bonds
• Non-Polar Covalent Shared equally
• Polar Covalent Shared unequally

13
Covalent Bonding

• H2

14
Covalent Bonding

• H2O

15
More Examples

• O2

16
More Examples

• N2

17
More Examples

• HCl
• NH3

18
More Examples

• CH4
• CO2

C
19
Determining Bond Type

• Whether electrons are shared or exchanged is
  based on electronegativity difference between two
  bonding atoms
• Nonpolar Covalent Bond
• 2 same Nonmetals (no difference in
  electronegativity)
• Polar Covalent Bond
• 2 different Nonmetals (small difference in
  electronegativity)

20
(No Transcript)
21
Determining Bond Polarity

• The larger the difference in electronegativity,
  the more polar the bond.

22
Determining Bond Polarity

• Which is more polar?

?EN
1.8
1.0
0.8
0.5
23
Bonding

• Ionic Bond
• Electrons are transferred from one atom to
  another (one gives, one takes)
• Metal and nonmetal, NaCl
• Large electronegativity difference
• Polyatomic ion, Mg(NO3)2
• More than 2 elements

24
Properties

• Ionic Compounds
• Most ionic compounds are hard, crystalline solids
  at room temperature
• High melting points
• Mostly soluble in water
• Can conduct an electric current when melted or
  dissolved in water(aq).

25
Ionic Compounds

• Formula Unit
• is the lowest whole-number ratio of ions in an
  ionic compound
• Ionic Compounds are repeating lattices of
  positive and negative ions

26
Ionic Compounds

• NaCl

27
Properties

• Covalent Compounds
• Most molecular compounds tend to have relatively
  lower melting and boiling points than ionic
  compounds.

28
Ionic Compounds

• Ionic compounds are electrically neutral, even
  though they are composed of charged ions
• Total positive charge equals total negative charge

29
Determining Formulas

• Must be electrically neutral
• Total positive charge must equal total negative
  charge
• Use oxidation numbers from Periodic Table
• Group 1 ? 1 Group 2 ? 2
• Group 13 ? 3 Group 15 ? -3
• Group 16 ? -2 Group 17 ? -1

30
Determining Formulas

• Determine number of each ion to balance out
  charge
• Use as subscript for element symbol
• Ex CaCl2, Na3PO4, Mg(NO3)2
• Write Positive Ion First
• Formula must be smallest whole-number ratio

31
Example

• Sodium and Chlorine

Na
Cl-
Na1Cl1
NaCl
32
Example

• Calcium and Fluorine

Ca2
F-
F-
Ca1F2
CaF2
33
Examples

• Potassium and Oxygen

K
O-2
K
K2O1
K2O
34
Polyatomic Ions

• Group of atoms that collectively have gained or
  lost electrons (Table E)
• Sodium and Nitrate

Na
(NO3)-
Na1(NO3)1
NaNO3
35
More Examples

• Potassium and Sulfate
• Ammonium and Sulfur

K2SO4
(NH4)2S
36
Short-cut (criss-cross method)

• Magnesium and Phosphate

Mg2
PO4-3
Mg3(PO4)2
37
Short-cut (criss-cross method)

• Magnesium and Carbonate

Mg2
CO3-2
Mg2(CO3)2
Must Simplify
MgCO3
38
(No Transcript)
39
Do This Now

• Write out formulas for
• Sodium Sulfate
• Calcium Phosphate

Na2SO4
Ca3(PO4)2
40
Dot Structures

• Shows valence electrons
• Must show charge for Ions
• NaCl

41
Dot Structures

• MgO

42
Dot Structures

• CaF2

43
(No Transcript)
44
Polyatomics

• Compounds with polyatomic ions contain BOTH ionic
  and covalent bonds
• Example NaNO3

Na
45
Coordinate Covalent Bonds

• Shared pair of electrons comes from one atom in
  the bonding pair
• Usually found in Polyatomic Ions

H
H
46
Allotropes

• Two or more different molecular forms of the same
  element in the same physical state
• Different properties because they have different
  molecular structures
• O2 vs O3
• Diamond, Graphite, Fullerenes (pictured on next
  slide)

47
Allotropes
48
Metallic Bonding

• Bonding within metallic samples is due to highly
  mobile valence electrons
• Free flowing valence electrons
• Sea of Electrons

49
Network Solids

• All atoms in a network solid are covalently
  bonded together
• Network solids have very high melting and boiling
  points, since melting requires the breaking of
  many bonds throughout the compound.
• Some of the strongest materials known to man are
  network solids.

50
Network Solids

• Diamonds ( C )
• Graphite ( C )
• Silicon Dioxide (SiO2)
• Silicon Carbide (SiC)

51
Bond Energy

• When two atoms form a bond, energy is released
• Example Cl Cl ? Cl2 energy
• Energy needs to be added to break a bond
• Example Cl2 energy ? Cl Cl

52
(No Transcript)
53
Structural Formulas

• Shared electrons are written as a line, unshared
  electrons are not written
• Each line represents 2 electrons

54
Molecular Polarity

• Polar Molecule
• one end of a molecule is slightly negative(d-)
  and the other end is slightly positive(d).
• Asymmetrical charge distribution
• Nonpolar Molecule
• Can not be separated into different ends
• Symmetrical charge distribution

55
Polar Molecule

• H2O
• Polar Covalent Bond
• Electrons shared Unequally

d-
56
More Examples

• HCl
• NH3

d-
d
d-
57
Another Example

• CH4
• Nonpolar Molecule

d
d-
d
d
d
58
Polarity

• Ionic Compounds are Ionic
• Nonpolar Covalent Bonds always indicate Nonpolar
  Molecules
• Polar Covalent Bonds
• Determine Symmetry

59
Like Dissolves Like

• Polar and Ionic substances will dissolve in
  other Polar Substances
• Nonpolar substance will dissolve in other
  nonpolar substances

60
(No Transcript)
61
Intermolecular Forces

• Intermolecular Forces of Attraction
• attraction between two molecules or ions that
  hold them together (not a bond)
• Determines melting and boiling points of
  compounds
• Stronger intermolecular forces, higher melting
  and boiling points

62
Intermolecular Forces

• Van der Waals
• Dispersion
• Dipole-Dipole
• Molecule-Ion
• Hydrogen Bonding

63
Van der Waals

• Dispersion
• Electrons of one atom are attracted to the
  Protons of the next atom.
• Also called an induced dipole
• Attraction increases with increasing mass

e
e
p
p
p
p
e
e
64
Van der Waals

• Dipole-Dipole
• negatively charged end of polar molecule is
  attracted to positively charged end of another
  polar molecule

65
Molecule Ion

• Attraction between polar molecules and ions in
  solution

66
Hydrogen Bonding

• Hydrogen bonded to N, O, or F, is attracted to
  the N, O, or F of another molecule.
• Not actual bond, just attraction

67
Boiling Point of H compounds
68
Boiling Point of H compounds