Transcript Unit 2: Chemical Bonding - Corner Brook Regional High

Unit 2: Chemical Bonding
Chemistry2202
1
Outline
Bohr diagrams
 Lewis Diagrams
 Types of Bonding
 Ionic bonding
 Covalent bonding (Molecular)
 Metallic bonding
 Network covalent bonding

2
Types of Bonding (cont’d)
 London Dispersion forces
 Dipole-Dipole forces
 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
Cl
Be
4
Lewis Diagrams
LD provide a method for keeping
track of electrons in atoms, ions, or
molecules
 Also called Electron Dot diagrams
 the nucleus (p+& n0) and filled energy
levels are represented by the element
symbol

5
Lewis Diagrams

dots are placed around the element
symbol to represent valence electrons
6
Lewis Diagrams
eg. Lewis Diagram for F
lone pair
bonding
electron
••
•
•F •
••
lone pair
lone pair
7
Lewis Diagrams
lone pair – a pair of electrons not
available for bonding
bonding electron – a single electron
that may be shared with another atom
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Lewis Diagrams
eg. Lewis Diagram for C
•
•
•C
•
9
Lewis Diagrams
eg. Lewis Diagram for P
••
•
•P
•
10
Lewis Diagrams
eg. Lewis Diagram for Na
•
Na
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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
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Lewis Diagrams for Compounds
draw the LD for each atom in the
compound
 The atom with the most bonding
electrons is the central atom
 Connect the other atoms using single
bonds (1 pair of shared electrons)
 In some cases there may be double
bonds or triple bonds

13
Lewis Diagrams for Compounds
eg. Draw the LD for:
PH3
CF4
Cl2O
C2H6
C2H4
C2H2
14
Lewis Diagrams for Compounds
eg. Draw the LD for:
NH3
SI2
POI
N2
SiCl4 N2H4
CO2
N2H2
CH3OH
H2
O2
HCN
CH2O
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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
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Why is propane (C3H8) a gas at STP while
kerosene (C10H22) a liquid?
17
Why is graphite soft enough to write with while
diamond is the hardest substance known even
though both substances are made of pure
carbon?
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Why can you tell if it is ‘real gold’ or just ‘fool’s
gold’ (pyrite) by hitting it with a rock?
19
‘As Slow As Cold Molasses’
‘All Because of Bonding’
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‘liquids’ @ -30 ºC
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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
22
Bonding
Intramolecular forces are forces of
attraction between atoms or ions.
Intramolecular forces include:
1. ionic bonding
2. covalent bonding
3. metallic bonding
4. network covalent bonding
23
Bonding
Intermolecular forces are forces of
attraction between molecules.
Intermolecular forces include:
5. London Dispersion Forces
6. Dipole-Dipole forces
7. Hydrogen Bonding
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Ionic and Covalent Bonding
ThoughtLab p. 161
Identify #’s 1 - 6
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Ionic Bonding



Occurs between cations and anions –
usually metals and non-metals.
An ionic bond is the 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

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Ionic Bonding
In an ionic crystal
the ions pack tightly
together.
 The repeating 3-D
distribution of
cations and anions
is called an ionic
crystal lattice.

27
Ionic Bonding
Each anion can be
attracted to six or
more cations at
once.
 The same is true
for the individual
cations.

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Ionic Bonding
29
Covalent Bonding
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.

30
Covalent Bonding
Molecular compounds
have low melting and
boiling points.
 Exist as distinct
molecules.

31
Covalent Bonding
Molecular
compounds do
not conduct
electric current
in any form
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Property
Type of
elements
Ionic
Molecular
Metals and
nonmetals
Non-Metals
Force of
Attraction
Positive ions attract Atoms attract a
negative ions shared electron
pair
Electrons move Electrons are
Electron
from the metal to
movement
shared
the nonmetal
between atoms
State at room
Solids, liquids,
Always solids
temperature
or gas
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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
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Metallic Bonding (p. 171)
metals tend to lose valence electrons.
 valence electrons are loosely held and
frequently lost from metal atoms.
 This results in 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

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Metallic Bonding
36
Metallic Bonding

This theory of metallic bonding is called
the ‘Sea of Electrons’ Model or ‘Free
Electron’ Model
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Metallic Bonding

This theory accounts for properties of metals
1. electrical conductivity
- electric current is the flow of electrons
- metals are the only solids in which electrons
are free to move
2. solids
- Attractive forces between positive cations and
negative electrons are very strong
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Metallic Bonding
3.
-
-
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.
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Network Covalent Bonding (p. 199)
occurs in 3 compounds (memorize these)
 diamond – Cn
 carborundum – SiC
 quartz – SiO2
 large molecules with covalent bonding in
3-d
 each atom is held in place in 3-d by a
network of other atoms

40
Network Covalent bonding

Properties:
 the highest melting and boiling points
 the hardest substances
 brittle
 do not conduct electric current in any
form
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2. Ionic bonding(metal & nonmetal)
3. Metallic bonding (metals)
MP & BP decreases
Strongest
1. Network Covalent (Cn ,SiO2 , SiC)
4. Molecular (nonmetals)
Weakest
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Valence Shell Electron Pair
Repulsion theory (VSEPR)
The shape of molecules is caused by
repulsion between valence electron pairs
around the atoms in a compound.
 Repulsion between valence electron pairs
force them to move as far away from each
other as possible.

Valence Shell Electron Pair
Repulsion theory (VSEPR)
To determine molecular shapes, count the
# of bonds and # of lone pairs on the
central atom(s).
 We will examine 5 molecular shapes

1. Tetrahedral (4 bonds; 0 lone pairs)
2. Pyramidal (3 bonds; 1 lone pair)
3. V-shaped (2 bonds; 2 lone pairs)
4. Trigonal Planar (3 bonds; 0 lone pairs)
5. Linear (2 bonds; 0 lone pairs)
For each molecule below draw the Lewis
diagram and the shape diagram.
HOCl
H2Se
H2O2
NBr3
C2F4
C2H6
CHCl3
CH3OH
PBr3
I2
SiH4
HCN
SiH2O
C2H2
Electronegativity & Covalent Bonds




-
Electronegativity - EN - p. 174
EN measures 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
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Increases
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Electronegativity & Covalent Bonds
 polar covalent bond
a bond between atoms with different EN
- the bonding pair is attracted more
strongly to the atom with the higher EN
-
δ+
bond dipole
δ−
H Cl
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Electronegativity & Covalent Bonds
 nonpolar covalent bond
a bond between atoms with the same
EN
- the bonding pair is shared equally
between the atoms
-
Complete:
#’s 7 – 9 on p.178
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Electronegativity & Covalent Bonds
polar molecule
- a molecule in which the bond dipoles
do not cancel each other
- a polar molecule has a molecular
dipole that points toward the more
electronegative end of the molecule.
eg. HCN
HC
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Electronegativity & Covalent Bonds
nonpolar molecule
- a molecule in which the bond dipoles
cancel each other
OR
- there are no bond dipoles
eg. CO2
PH3
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Electronegativity & Covalent Bonds
To determine whether a molecule is polar:
- draw the Lewis diagram & shape diagram
- draw the bond dipoles & determine
whether they cancel
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Intermolecular Forces
58
2. Ionic bonding(metal & nonmetal)
3. Metallic bonding (metals)
MP & BP decreases
Strongest
1. Network Covalent (Cn ,SiO2 , SiC)
4. Molecular (nonmetals)
Weakest
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To compare mp and bp in covalent
compounds you must use:
- London Dispersion forces (p. 204)
(in all molecules)
- Dipole-Dipole forces (pp. 202, 203)
(in polar molecules)
- Hydrogen Bonding (pp. 205, 206)
(when H is bonded to N, O, or F)
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Intermolecular Forces (p. 202)
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Intermolecular Forces
Covalent compounds have low mp and
bp because attractive forces between
molecules are very weak.
 Intermolecular forces were studied
extensively by the Dutch physicist
Johannes van der Waals
 In his honor, two types of intermolecular
force are called Van der Waals forces.

62
Intermolecular Forces

Intermolecular forces can be used to
account for the physical properties of
covalent compounds.
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Intermolecular Forces
F
F
F
F
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1. London Dispersion Forces
• 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.
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1. London Dispersion Forces
The strength of these forces depends on:
a) 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.
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1. London Dispersion Forces
Two molecules that have the same number
of electrons are isoelectronic
eg. C2H6 and CH3F
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1. London Dispersion Forces
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.
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2. Dipole-dipole Forces
- occur between polar molecules
- the δ+ end of one polar molecule is
attracted to the δ- end of another polar
molecule (& vice-versa)
eg. Which has the higher boiling point
CH3F or C2H6 ?
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p. 202
70
3. Hydrogen Bonds
- a special type of dipole-dipole force
(about 10 times stronger)
- only occurs between molecules that
contain a H atom which is directly
bonded to F, O, or N
ie. the molecule contains at least one
H-F, H-O, or H-N covalent bond.
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3. Hydrogen Bonds
-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
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p. 206
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NOTE: To compare covalent compounds
you must use:
- London Dispersion forces
(all molecules)
- Dipole-Dipole forces
(polar molecules)
- Hydrogen Bonding
(H bonded to N, O, or F)
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
Alchem worksheet pp. G32, 33
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p. 226 #13
Omit parts g), j) – o), q), u), & v)
- Answers on p. 815 for #13
- Incorrect answers
c), d), & s)
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p. 210
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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.
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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
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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 other
compounds.
80
p. 226 #’s 13 & 14
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Electronegativity
Electronegativity is a result of the
space between the nucleus and
the electrons
 As the number of protons in the
nucleus increases, the attractive
force on the electrons increases,
pulling them closer to the nucleus

82
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 the EN of
nonmetals is relatively high
83
Electronegativity and Ionic Bonds

When ions form, the resulting
electrostatic force is an ionic bond
84
Electronegativity and Covalent Bonds

Atoms in covalent compounds can
either have:


the same EN
eg. Cl2 , PH3, NCl3
different EN
eg. HCl
85
Electronegativity and Covalent Bonds
Atoms that have different EN attract
the shared pair of valence electrons
at different strengths
 The atom with the higher EN exerts
a stronger attraction on the shared
electron pair
eg. H2O

86
Electronegativity and Covalent Bonds
Since the oxygen atom has a higher EN
the bonding electrons will be pulled
closer to the oxygen atom
 This results in slight positive and
negative charges within the bond.


These charges are referred to as
“partial charges” and are denoted
with the Greek letter delta (δ).
87
Electronegativity and Covalent Bonds

The region around the oxygen atom
will be slightly negative, and around
the hydrogens will be slightly positive
88
Electronegativity and Covalent Bonds
The symbol, δ+ represents a partial
positive charge (less than +1) and
δ− represents a partial negative
charge (less than −1).
 Since the bond is polarized into a
positive area and a negative area
the bond has a “bond dipole”.

89
Electronegativity and Covalent Bonds

The arrow points to
the atom with the
higher EN.
p.178
90
Electronegativity and Covalent Bonds
 Covalent
bonds resulting from
unequal (electronegativities)
sharing of bonding electron
pairs are called Polar
Covalent Bonds
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Electronegativity Homework
#’s 7, 8, & 9 - p. 178
#’s 1, 2, & 3 - p. 180
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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?
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Test Outline
Bohr Diagrams (atoms & ions)
 Lewis Diagrams (Electron Dot)
 Ion Formation
 Ionic Bonding, Structures & Properties
 Covalent Bonding, Structures & Properties

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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.)

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