Co-ordinate Bonds,
Intermolecular Forces and
Metallic Bonding
Everything you EVER
wanted to know but
were afraid to ask!
Co-ordinate Bonding
Also referred to as DATIVE bonding.
Occurs when a PAIR of electrons is donated from one atom to
another.
Only happens when an empty orbital is present
(ie an electron deficient species)
Once the coordinate bond is formed, we treat it exactly the same
as a normal covalent bond.
This is how Group 13 atoms get full octets!
Amino-Borane
Huh!?!?!?
NH3 – BH3 (don’t worry about the name)
Draw the Lewis Structure for NH3, then for BH3
N can “donate” its lone
pair to B forming a
co-ordinate bond.
N has a full
octet an a lone
pair of
electrons.
NH3 is an
uncharged
molecule.
B only has 6
electrons, it is
deficient.
BH3 is an
uncharged
molecule.
Amino-Borane
Now we have a new molecule with a covalent bond between N and B.
While there are formal charges on N and B, the overall charge on the
molecule is neutral.
The bond formed between N and B is just as strong as any covalent
bond.
Co-ordinate Bonding
Draw the co-ordinate bonds between each of the following pairs:
GaCl3, ClH+, H2O
H+, NH3
BF3, NH3
NH3, GaCl3
PCl3, GaCl3
AlCl3, AlCl3 (to form Al2Cl6)
Intermolecular Forces
Covalent Bonds are NOT intermolecular forces. They are
INTRAMOLECULAR forces.
So what are INTERMOLECULEAR Forces?
The forces of attraction/repulsion that exist between molecules.
The ones we need to overcome to change the state of the
substance.
Relative Strengths of Bond Types
Ionic
Covalent
Intermolecular Forces
Van der Waals Forces
Discovered but the Dutch physicist, Van der Waals and named in his
honor.
They are weak interactions between molecules are divided into 2
basic types:
1. Dipole – Dipole
 When a molecule has a permanent dipole moment (molecule is
polar), the negative end of one molecule will sit closer to the
positive end of its neighbor.
2. Dispersion or London Forces (weakest of all intermolecular)
 Caused by temporary shifts in electron density within a
molecule. For instance the 2 e- shared in H2, located on 1 H
atom rather than the other.
Dipole – Dipole Forces
d+
d-
d+
d-
d+
d+
d-
d+
d-
d-
Ion – Dipole Forces
Exist when we dissolve an ionic compound in a polar solvent, like H2O.
The d- end of H2O is attracted to a cation (+).
The d+ end of H2O is attracted to an anion (-).
London Forces
d+
2e-
2e-
d-
d-
Caused by the natural vibrations of the
electrons in the bond, there is no permanent
dipole moment.
The dominant intermolecular force in nonpolar substances.
d+
Hydrogen Bonding
A specific case of Dipole – Dipole Forces.
Occurs when H is involved in a strongly polar bond with O, N or F.
The H nucleus (just a proton) is attracted to the lone pairs of these
highly electronegative atoms.
The result is a network of strong intermolecular interactions
between H atoms and available lone pairs.
Occurs extensively in H2O, NH3 and HF.
The primary reason H2O doesn’t behave as almost all other
substances in the known universe.
H2O expands when it freezes  everything else contracts
Hydrogen Bonding
Intermolecular Take-home Message
The stronger the intermolecular interaction, the
stronger molecules are held together.
When molecules are held together more tightly, we see the evidence
when we examine physical properties of the substance  melting and
boiling points increase
Practice Questions
For each of the following molecules, state which intermolecular
force will be dominant. Then state which (out of each pair) should
have the higher melting point.
Molecule 1 Dominant
Force
Melting
Point
Molecule 2 Dominant
Force
Melting
Point
H2O
H-Bonding
0 ºC
H2S
Dipole
-82 ºC
HCl
Dipole
-114.2 ºC
H2
London
-259 ºC
NH3
H-bonding
-78 ºC
H2O
H-bonding
O ºC
Br2
London
-7 ºC
I2
London
114 ºC
Metals
Metals, they’re an interesting bunch . . .
They play by their own set of rules, so to describe them we need to
talk about METALLIC bonding.
First of all what do we know about the properties of metals?
 they’re malleable and ductile
 they’re electron rich (they give up their electrons to make
cations)
 they conduct electricity
 the free flow of electrons yields many colorful solutions
 they’re shiny
 they conduct heat
Metallic Bonding
The “Sea of Electrons” Model
Due to low
electronegativities,
low effective nuclear
charges and large
diffuse orbitals,
electrons can flow
freely from one atom
to the next.
Metallic Bonding
The “Sea of Electrons” Model
Electrons carry
electrical current,
if electrons can
flow freely
throughout the
metal the metal will
conduct electricity.
Alloys
A mix of 2 or more metals.
2 types: Substitutional and Interstitial
Zn
Cu
Zn
Fe
Fe
C
Cu
Zn
Cu
Fe
Fe
C
Fe
Fe
Brass (Zn and Cu in various
proportions)
Steel (Fe and C in various
proportions)
Typically Zn is put in place of a Cu
atom.
C fills in the “holes” between the Fe
atoms.