Unit 2: Chemical Bonding
Chemistry2202
Outline


Bohr diagrams & Lewis Diagrams
Types of Bonding
5.
1.
2.
3.
4.


Ionic
Covalent (molecular)
6.
Metallic
7.
Network covalent bonding
VSEPR Theory (Shapes)
Physical Properties
London Dispersion
forces
Dipole-Dipole forces
Hydrogen Bonding
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
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
Lewis Diagrams
eg. Lewis Diagram for F
lone pair
bonding
electron
••
•
•F •
••
lone pair
lone pair
Lewis Diagrams
lone pair – a pair of electrons not available
for bonding
bonding electron – a single electron that
may be shared with another atom
Lewis Diagrams
eg. Draw Lewis Diagrams for:
carbon
phosphorus
••
•
•
•C
•
•
•P
•
sodium
•
Na
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
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
Lewis Diagrams for Compounds
eg. Draw the LD for:
PH3
CF4
Cl2O
C2H6
C2H4
C2H2
Lewis Diagrams for Compounds
eg. Draw the LD for:
NH3
SI2
POI
N2
SiCl4 N2H4
CO2
N2H2
CH3OH
H2
O2
HCN
CH2O
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
Why is propane (C3H8) a gas at STP while
kerosene (C10H22) a liquid?
Why is graphite soft enough to write with
while diamond is the hardest substance?
(both are C)
Graphite
Diamond
Graphite
Diamond
Other forms of carbon
Buckyball Carbon
Nanotube Carbon
How can you which is ‘real gold’ and which is
‘fool’s gold’ (pyrite) by hitting it with a rock?
‘As Slow As Cold Molasses’
‘All Because of Bonding’
Viscosity of liquids
‘liquids’ @ -30 ºC
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.
Melting and Boiling Points
F2
Cl2
Br2
I2
dihydrogen monoxide
pepper demo
dd & HB
teacher tube - IF
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
Bonding
1.
2.
3.
4.
Intramolecular forces are forces of attraction
between atoms or ions.
Intramolecular forces include:
ionic bonding
covalent bonding
metallic bonding
network covalent bonding
Bonding
5.
6.
7.
Intermolecular forces occur BETWEEN
different molecules.
Intermolecular forces include:
London Dispersion Forces
Dipole-Dipole forces
Hydrogen Bonding
Ionic and Covalent Bonding
ThoughtLab p. 161
Identify #’s 1 - 6
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
Ionic Bonding


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.
Mar. 11
Ionic Bonding


Each anion is
attracted to six or
more cations at once.
The same is true for
the individual cations.
Mar. 11
Ionic Bonding
Mar. 11
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.
Covalent Bonding


Molecular compounds
have low melting and
boiling points.
exist as distinct
molecules.
Mar. 11
Covalent Bonding
 do not conduct
electric current in
any form
Mar. 11
Property
Type of
elements
Force of
Attraction
Ionic
Molecular
Metals and
nonmetals
Non-Metals
Mar. 11
Positive ions attract Atoms attract a
negative ions
shared electron
pair
Electron
Electrons move
Electrons are
movement from the metal to
shared
the nonmetal
between atoms
State at room
Always solids
Solids, liquids,
temperature
or gas
Property
Solubility
Ionic
Molecular
Soluble or low Soluble or
solubility
insoluble
Conductivity in None
solid state
None
Mar. 11
Conductivity in Conducts
liquid state
None
Conductivity in Conducts
solution
None
Metallic Bonding (p. 171)
•
Na
Na•
•
Na
• Na
Na •
•
Na
• Na
Na
•
Metallic Bonding (p. 171)
•
•
+
•
Na
+
Na
•
+
Na
•
Na
•
+
+
Na
+
Na
+
Na
•
•
+
Na
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
Metallic Bonding
Metallic Bonding

This theory of metallic bonding is called the
‘Sea of Electrons’ Model or ‘Free Electron’
Model
Metallic Bonding
Metallic bonding 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
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.
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
Network Covalent bonding

Properties:
 the highest melting and boiling points
 the hardest substances
 brittle
 do not conduct electric current in any form
2. Ionic bonding(metal & nonmetal)
3. Metallic bonding (metals)
4. Molecular (nonmetals)
Weakest
MP & BP decreases
Strongest
1. Network Covalent (Cn ,SiO2 , SiC)
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.
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.
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.
1 central atom
HOCl
H2Se
H2SiO
NBr3
CHCl3
SiH4
PBr3
HCN
I2
2 central atoms
C2F4
C2H6
CH3OH
C2H2
H2O2
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
Increases
Electronegativity & Covalent Bonds
1. 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
δ+
δ−
H Cl
Electronegativity & Covalent Bonds
2.
-
-
-
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 (δ) indicates
‘partial’ charges
Complete:
#’s 7 – 9 on p.178
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
Electronegativity and Covalent Bonds

Atoms in covalent compounds can either
have the same EN
eg. Cl2 , PH3, NCl3
OR different EN
eg. HCl
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)
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 (δ).
Electronegativity and Covalent Bonds


The region around the chlorine atom will be
slightly negative
The region around the hydrogen will be
slightly positive.
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.
Electronegativity and Covalent Bonds
eg. H2O
Electronegativity and Covalent Bonds
eg. HF
Electronegativity Homework
p. 178 #’s 7, 8, & 9
p. 180 #’s 1, 2, & 3
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?
Test Outline





Bohr Diagrams (atoms & ions)
Lewis Diagrams (Electron Dot)
Ion Formation
Ionic Bonding, Structures & Properties
Covalent Bonding, Structures & Properties
Test Outline
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


Metallic Bonding Theory& Properties
Network Covalent Bonding & Properties
Electronegativity
Bond Dipoles & Polar Molecules
VSEPR Theory
LD, DD, & H-bonding
Predicting properties (bp, mp, etc.)
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
Molecular Dipoles


Nonpolar molecules DO NOT have molecular
dipoles.
This occurs when:
- bond dipoles cancel eg. CO2
- there are no bond dipoles eg. PH3
To determine whether a molecule is polar:
- draw the LD and the shape diagram
- draw the bond dipoles and determine whether
they cancel

Molecular Dipoles
See Handout #1
Mar. 31
Intermolecular Forces
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
Mar. 31
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)
Intermolecular Forces (p. 202)
Mar. 31
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.
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.
1. London Dispersion Forces
Apr. 1
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.
1. London Dispersion Forces
Apr. 1
Two molecules that have the same number of
electrons are isoelectronic
eg. C2H6 and CH3F
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.
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 ?
Apr. 5
Apr. 5

p. 202
In the liquid state,
polar molecules are
oriented such that
oppositely charged
ends of the
molecules are close
to each other.
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.
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;
C 3H 8
C2H5OH
C 2H 5F
Apr. 5
p. 206
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)
WorkSheet: Bonding #4
p. 225
p. 226
#’s 9 & 10
#’s 12 – 14,
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.
p. 210
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
Summary
Strongest Weakest -
Network Covalent
Ionic
Metallic
Covalent
↣ LD forces (all molecules)
↣ DD forces (polar molecules)
↣ H-Bonding (H bonded to N, O, or F)
Ion-Dipole Forces

An ion-dipole force is the force of
attraction between an ion and a polar
molecule (a dipole).
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.
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.
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.
Dispersion (London) Forces

1.
-
-
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.
Dispersion (London) Forces
2.
-
-
-
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.
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.
Hydrogen Bonding
δ+
δ−
δ+
δ+
δ−
…
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)
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.
δ+
δ−
δ+
δ+
δ−
…
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
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.
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.