BONDING
TOPIC 4
Terms
• Bonds
– Breaking them takes energy
– Making them gives off energy
Covalent Bonding
• Exothermic
– More energy is given off than put in
• Endothermic
– More energy is absorbed than given off
• Intramolecular Forces
– Forces within molecules (ionic, covalent and metallic)
• Intermolecular Forces (IMF)
– Forces between particles
Ionic
Bonding
+
Less e- = Less e- repulsion
-
More e- = e- more repulsion.
• If the electronegative difference between
the atoms involved is =>1.8
– There are always exceptions to this rule!
• Will conduct electricity in its molten or
aqueous state (This test proves ionic)
Metal: K
Non-Metal: Cl
Drawing Ionic Bonding
Lewis Dot
Diagram
X
Na
-
+
Cl
Intramolecular
Forces
Electrons are
in pairs
Special Note:
The ionic bond is the electrostatic attraction
between oppositely charged ions!
• Just use the valence shell
• Be sure to include square brackets
and charge after electron exchange.
Ionic
Bonding
Combine
Mg
Fe
Al
Be
C
O
Cl
F
Br
Cl
Lewis Dot diagrams us the atoms
valance shell electrons
Lewis
Diagrams
Decomposition
CATHODE (-)
+
-
-
-
+
+
-
+
+
-
+
• When in molten orConductivity
aqueous state,
is ionic
substances WILL conduct
electricity, by the
FINITE
movement of (+) and (-) ions.
• This is different from how METALS conduct
electricity!
++
++ ++
ANODE (+)
2Na+(aq) + 2Cl-(aq)  2Na(s) + Cl2(g)
Intramolecular
Forces
NaCl
Giant Ionic Lattices
Anion
Cation
Ionic
Compounds
Force
Like charges
repel
Physical
characteristics
••• No
bonds
are made!!!
When
a
force
is applied, ionic
• Hard and brittle
• Solid doesn’t conduct
Electricity
compounds
will make
a clean
• More attractions
soluble in waterholds
than other
solvents
• Static
them
break.
• High MP and
BP
together.
(opposites
attract)
Metal: K
Non-Metal: Cl
Giant Ionic Lattices
Cubic or
Isometric
Table Salt
NaCl
Tetragonal
Giant Ionic Lattices
Cassiterite
SnO2
Orthorhombic
Giant Ionic Lattices
• Also found in mollusk shells and
coral
Agagonite
CaCO3
Hexagonal
Giant Ionic Lattices
Beryl
Be3Al2(SiO 3)6
Trigonal
Giant Ionic Lattices
Quartz
SiO2
Giant Ionic Lattices
Ionic
Bonding
Beryl
Be3Al2(SiO 3)6
Triclinic
Giant Ionic Lattices
Copper(II)
Sulfate
CuSO4
Transition Metals
+
Fe2+
Cu
Iron(II) Oxide
Oxide
Copper(I)
Intramolecular
Forces
2+
Fe3+
Cu
Iron (III) Oxide
Oxide
Copper(II)
• Transition metals can have multiple
ions.
• Ones you should know.
Multiple Ions
Reminder
NO3OHHCO3-
SO4-2
CO3-
Ions
PO4-3
2
NH4+
• Be sure to review your polyatomic
ions!!!
Polyatomic
Ions
Covalent Bonding
Topic 4
Covalent Bonding
2.1
3.0
Intramolecular
Forces
X
H
Cl
Differences
|3-2.1|
=0.9
Special Note:
The covalent bond is the electrostatic
attraction between pairs of e- and positively charged nuclei!
• If the electronegative difference
between the atoms involved is <1.8
• Will NOT conduct electricity
• Electrons are shared
COVALENT
BONDING
Questions
Review
Na + Cl
Li + O
K +
Ca + CO3
Na
+ SO
3
NO3
• For
What
ionic
is the
compounds
chemical formula?
to form the
shells
of both
metal and
• valance
What is the
names
for each?
non-metal must be full!!
Covalent Bonding
Intramolecular
Forces
H
H
X
H
C
H
H
X
Cl
X
H
X
H
H
• Structural formula
• Lewis structure
COVALENT
BONDING
Lewis Structures
Intramolecular
Forces
H2O
H
1
H
1
O
6
8
- 4=4
Hydrogen can only
hold 2e- remaining
must be paired on
Oxygen
• 1) Sum all valence e• 2) Subtract 2e- for every bond
• 3) Place e- around periphery atoms to
form octets. The remaining around
central atom
• 4) All atoms MUST be paired!!!!!!
COVALENT
BONDING
Lewis Structures
•
•
•
•
•
Draw the following Lewis structures
H2
Cl2
O2
N2
HL: PCl5, PCl4+, PCl6and XeF4
HCN
C2H6
C2H4
C2H2
Intramolecular
Forces
Covalent
Bonding
Special Lewis Structures
+
+
Intramolecular
Forces
Lone pair of
e-
H
H
N
H
Electrophile
H
• Coordinate or dative covalent bonds
• When both e- are shared from the same
atom. (Not one from each as before)
• Occurs when a non bonding e- pair
donates an e- to an e- deficient atom.
Covalent
Bonding
Special Lewis Structures
• Draw the following Lewis structures
• CO
• H3O+
Intramolecular
Forces
Covalent
Bonding
Length, Strength & Hybrid
Resonance
2-
2-
O
O
C
2-
O
O
Intramolecular
Forces
O
C
O
O
O
C
O
Don’t
forget to
show the
e- pairs!!
• More bonds = more strength & shorter bonds
• Resonance structures
– Bond length is longer than a double bond but
shorter than a single bond
CO32-
Length & Strength
Ethene
H
Carboxylic
Ethyne Acid
H
O
H
H CCC H
R
OH
C C
H
Intramolecular
Forces
R = Functional
Group
• Compare the two molecules
• Ethyne has stronger and shorter bonds
•
• C=O bond is stronger and shorter due to
Oxygen being more electronegative
CO32-
Bond Polarity
Intramolecular
Forces
δ+
δH
X
Cl
• Non-Metals are fighting for e• Atom with larger electronegativity will
hold the e- closer to itself.
• Atoms become slightly charged.
Dipole
Moment
Covalent
Bonding
Exceptions to the Octet Rule
Intramolecular
Forces
F
B
F
F
• BF3
• Actual structure: Boron is e- deficient
• This is known because of its reactivity towards
electron rich molecules such as NH3
• CNOF all obey the octet rule.
Covalent
Bonding
Formal Charge
Intramolecular
Forces
• SO42• Single bonds (8 e- around S)
• Double bonds (12 e- around S)
• Formal Charge = (# valence e- on free atom) – (# valence eassigned to the atom in the molecule)
• (Valence e-)assigned = (# lone pair e-) + ½ (# of shared e-)
• 1) Molecules attempt to achieve Formal Charge as close to
0 as possible.
• 2) Any negative Formal charge will reside on most
electronegative atom.
Covalent
Bonding
VSEPR (shape)
Intramolecular
Forces
3 Pairs of e120o
2 Pairs of e180o
O
C
O
F
2-
F
O
B
C
F
O
O
• VSEPR (Valence Shell Electron Pair Repulsion)
• Paired e- attempt to get as far away from
each other as possible.
• Multiple bonds still count only as 1 pair!!
Covalent
Bonding
VSEPR
4 Pairs of e109.5o
Lone pair
107o
Intramolecular
Forces
Lone pair
104.5o
H
C
H
N
H
H
H
O
H
H
H
H
• Tetrahedral
• Lone pair e- have increased charge density
and require more room
• More repulsion from lone pair will decrease
bond angle.
Covalent
Bonding
Home Work
•
•
•
•
Predict the shape AND bond angles
H2S
PbCl4
H2CO
SO2
NO3PH3
NO2NH2POCl3
CO2
Intramolecular
Forces
Covalent
Bonding
HL VSEPR
Molecule
Shape
Total valance
electrons
Bond Pairs
Non Bonding
Electron pair
Angle
BeF2
Linear
180
BeF3
Triangular
Planar
120
SO2
Bent
117
CH4
Tetrahedral
109.5
NH3
Trigonal
pyramidal
107
H2O
Bent
104
HL VSEPR
Molecule
Shape
Total
Valance
electrons
Bond Pairs
Non
Bonding
Electron
pair
Angle
PCl5
Triangular
Bipyramidal
90 & 120
SF4
Seesaw
90 & ≈117
T-Shape
90
CF6
Octahedral
180
IF5
Square
Pyramidal
90
XeF4
Square
Planar
≈88
Expanded Valance Shell (14.1)
• Molecules with more than 8 electrons
• Electron promotion:
Dipole
Moment
Molecule Polarity (4.2.6)
2δ-
H
Cl
δ+
O
H
δ+
H
δ+
δ-
Non Polar
H
H
δ-
C
C
H
Cl
Cl
H
δ+
H
• Polarity effects state change (physical change)
• Unequal sharing causes a dipole moment to
form
• Q: Why is BF3 non-polar whereas PF3 is polar?
Covalent
Bonding
Hybridization (14.2.2)
• Sigma bond: σ (single bond)
– Axial overlap of orbital’s
1s1
H
2px2 py2 pz2
Cl
Hybridization (14.2)
• Sigma bond: σ (single bond)
– Axial overlap of orbital’s
Cl
Cl
Hybridization (14.2)
• Pi bond: π(Double bond, one σ bond)
– Parallel overlap of orbital’s
O
N
NO
Hybridization (14.2.3)
• Hybridization electron promotion
– New Orbital sp3
2px2 py2 pz2
2s2
C
Ground
Excited
State
State
4 Equal orbital`s capable of holding
a maximum of 2 electrons each
Hybridization (14.2)
• How to determine Hybridized orbital`s
– Look at the shape
Shape
High Electron dense
regions
Hybridized Orbital
sp
2
Tetrahedral
Sp2
3
Trigonal planar
sp3
4
Linear
sp3d
5
Trigonal bi-pyramidal
sp3d2
6
Octahedral (Square bi-pyramidal)
Carbon
Allotropes
C
C
C
C
C
• 1) Diamond (Tetrahedron, localized e-)
– Very hard and does not conduct electricity
• 2) Fullerenes (C60) Hexagonal and
pentagonal rings
– Nanotubes
Giant
Covalent
Carbon
Allotropes
C
C
C
C
C
C
Weak Pi
Bonds
HL: sp hybrid
Delocalized electrons
able to move
• 3) Graphite (Planar, delocalized e-)
– Weak pi bonding between sheets cause it to
conduct electricity and be slippery.
– Bonds are shorter than a tetrahedral due to the
pi bonding
Giant
Covalent
Benzene (14.3)
C
C
C
C
C
C
Pi bonds overlap
allowing for electrons
to be delocalized over
the entire molecule.
• Planar, delocalized e– Regular bonding would predict an alternating
double bond (Resonance structure)
– Hybrid theory shows sp2 configuration
C6H6
Silicon
Intramolecular
Forces
SiSi
Si
Si
Si
Si
Si
Si
Si
Si
Tetrahedron Configuration
Similar to diamond
Silicon
Silicon & Silicon dioxide
O
SiO2 but based
on a network
of SiO4
Intramolecular
Forces
Si
O
O
O
• Single bonds formed between Oxygen to
satisfy the octet.
Quartz
• HL: Less overlap in the P-sub orbital due
to atomic size difference therefore Pi bonds
do not form.
Metallic Bonding
Topic 4
Metallic Bonding
+
+
+
Intramolecular
Forces
+
-
+
+
+
+
-
+
-
+
+
+
+
-
+
• In solid state
+
+
-
Sea of
electrons
+
Conductivity is
INFINITE
• Outer e- are delocalized and free to move about
• Bond is a result of electrostatic attraction between
Fixed positive metal ions and delocalized e-
Metallic
Bonding
Physical Properties
+
+
+
+
-
+
+
+
+
-
Malleability
+
-
+
+
+
+
-
+
+
+
-
+
• The ability for a material to be pounded into
thin sheets.
• Aluminum Foil
• Swords and Folding
Metallic
Bonding
Physical Properties
+
+
+
+
+
+
+
+
+
+
+
+
Ductility
Electrons have been
excluded
• The ability for a material to be pulled into
wire
• Or in this case extruded into a wire
Metallic
Bonding
Physical Properties
• Because e- can move easily it can conduct
energy. (Heat or electricity)
• MP related to attractive force (between
atoms)
• 1) Size of Cation(+)
• 2) # of valence e• 3) Atom packing
• Size increases MP decreases:
• Giant Covalent substances have very high
mp
Metallic
Bonding
Allotrops
• Same element but different structure
• Carbon
• Diamond
• Graphite
• Fulluron
Metallic
Bonding
INTERMOLECULAR FORCES
Topic 4
Intermolecular
van der Waals’
Forces
Forces
(4.3.1)
Charge Induction
d+
d-
d+
• Van der Waals Forces
Charge Induction
d-
d+
Intermolecular
Forces
d-
IMF
Dipole-Dipole (4.3.1)
Intermolecular
Forces
Cl
Cl
d-
dC
C
Cl
H
d+
H
Cl
H
d+
H
• Polar molecules (polar covalent) have
slightly charged ends
• Opposites attract.
• Large electronegative difference =
stronger attraction.
IMF
van
der Waals’
Forces
Hydrogen
Bonding
(4.3.1)
d-
d+
Intermolecular
Forces
dd+
O
d+
H
H
d+
• Hydrogen Bonding (F, O or N bonded to H)
• Due to small size and high electronegativity
of non metals
• Creates a large charge difference
• Basically a super strong dipole-dipole bond
IMF
Boiling Point Trends (4.3.2)
Intermolecular
Forces
Get a picture of group 4,5,6,7 boiling points for hydrides
Key question is why does water have an abnormally high BP?
H bonding with O, F and N
• Phase change when IMF are overcome
• Be sure to explain using the words IMF and
how they affect the bonds BETWEEN particles.
• Van der Waals’ Forces are ALWAYS present!!!
IMF
Physical Properties
• Van der Waal’s: Lowest MP, Non polar
• Butane (C4H10)
• Dipole-dipole: Slightly miscible
• Propanone C3H6O
• H2O
• Ionic Bonding: Only conducts electricity when liquid
or aqueous. (Decomposition when it does)
• NaCl
• Metallic Bonding: Conducts electricity, not water
soluble, MP regulated by, valance, size and
packing.
• Fe
• Giant Covalent: Highest MP, Insoluble in both nonpolar and polar solvents. Does not conduct
electricity except for graphite.
• Diamond and Graphite (Allotropes)
Increasing Melting Point
• Hydrogen Bonding: Miscible with polar substances
Bonding Questions
• Compare the following for B.P
•
•
•
•
•
HF and HCl
H2O and H2S
NH3 and PH3
CH3OCH3 and CH3CH2OH
CH3CH2CH3, CH3CHO and CH3CH2OH
HL Material
Hybridization (14.2)
• Sigma bond: σ (single bond)
– Axial overlap of orbital’s
Hybridization (14.2)
• Sigma bond: σ (single bond)
– Axial overlap of orbital’s
Hybridization (14.2)
• Sigma bond: σ (single bond)
– Axial overlap of orbital’s
Hybridization (14.2)
• Sigma bond: σ (single bond)
– Axial overlap of orbital’s
Lattice Formation
• Where the heat comes from
• Route 1: A + B + C + E
• Route 2: F
• Hess’s law: A + B +C + E = F
+107 + 122 + 496 + (-349) + E = -411
E = -787 kJ mol-1
Lattice Enthalpy
Na(s) + ½ Cl2(g) 
Na+Cl-
1) Na(s)  Na(g)
2) Na(g)  Na+(g) + e-
or NaCl
Intramolecular
Forces
½ Cl2(g)  Cl(g)
Cl(g) + e-  Cl- (g)
3) Na+(g) + Cl-(g)  NaCl(s)
• 1) Production of Gaseous atoms
• 2) Formation of Gaseous ions
• 3) Production of solid ionic lattice
NaCl
Born-Haber Cycle
Na+
(g)
+
Cl-
ΔHθI.E.
1st Ionization of Na
+496 kJ mol-1
E
(g)
NaCl(s)
Lattice Enthalpy
ΔHθE.A.
1st electron affinity of Cl
-349 kJ mol-1
C
D
Na(g)
Endothermic
Cl(g)
Exothermic
ΔHθf
ΔHθat
ΔHθat
Atomization of Na
-107 kJ mol-1
A
Atomization of Cl
+122kJ mol-1
Na(s)
+ Cl2(g)
F
B
Formation of NaCl
-411 kJ mol-1
Spare Parts
N
C
O
-
O
C
C
O
H
H
δ+
δ-
H
H
Cl
H
C
C
+
H
C
N
H
X
H
C
C
C
C
H
H
H
Cl
O
H
H
H