TOPIC: Intermolecular Forces
How do particle diagrams of
liquids & solids compare to those
of gases?
Describe relative positions and motions
of particles in each of 3 phases
SOLID
LIQUID
GAS
Why do some substances exist
as gases, some as liquids, and
some as solids at room temp?
 Part
of answer has to do with
forces between separate
molecules (called intermolecular
forces)
Intermolecular forces between molecules.
They are weaker. Intramolecular forces are
between individual atoms (we will learn this later)
Intermolecular
forces
Intramolecular
forces
Intermolecular Forces-IMF
means “between” or “among”
 Intermolecular forces = forces
between neighbouring compounds
 Inter
All molecules have Dispersion
forces (the regents calls
these Van der Waals)
2 other types of forces (IMF):
1. Dipole-Dipole forces
2. Hydrogen bonds
-if one of these are present,
they are more important.
Most atoms don’t have a charge,
unless they are ions, so we often
refer to them as having partial
charges and write them like this…
This separation of
Charge is responsible
For the forces
Between the
molecules
1. Dispersion Forces (van der waals):
●
●
weakest IMF
occur between nonpolar molecules
•
•
•
Nonpolar means no poles
(+/-)
Can’t tell one end of
molecule from other end
electrons are evenly
distributed
Click here for animation (slide 4 of 13)
●
•instantaneous and momentary
•fluctuate
•results from motion of electrons
if charge cloud not symmetrical will induce asymmetry in neighbor’s
charge cloud!
4 categories of Nonpolar Molecules
(you need to memorize)
 Noble
 He,
Gas molecules:
Ne, Ar, Kr, Xe, Rn
 diatomics
 H2,

N2, O2, Cl2, F2, I2, Br2
Pure Hydrocarbons (CxHy):
 CH4,

if both atoms are same: (7)
C2H6, C3H8
small symmetrical molecule
 CO2,
CF4, CCl4
Dispersion Forces and Size
 Dispersion
forces ↑ with molecule size
 larger the electron cloud, the greater the
fluctuations in charge can be
 Rn
> Xe > Kr > Ar > Ne > He
 I2 > Br2 > Cl2 > F2
 C8H18 > C5H12 > C3H8 > CH4
Boiling point of N2 is 77 K (-196˚C)
IMF are very weak dispersion forces
2.
•
•
•
Dipole-dipole forces:
intermediate IMF
occur between polar molecules (they have a
partial charge at each pole – one is typically
much larger than the other)
Click here for animation (slide 3 of 13)
Dipole-dipole Forces & Polar Molecules
Molecule shows
permanent
separation of
charge; has poles:
one end partly (-) &
one end partly (+)
Polar Molecules
Polar means molecule has poles: (+) & (-)
geometry and electron distribution are
not symmetrical
3.
•
•
Hydrogen bonds:
strongest IMF
occur between molecules that have:
H-F H-O or H-N bonds ONLY
Hydrogen Bonding
H-O
N-H
Occurs between molecules with H-F, H-O, or H-N
bonds
Hydrogen Bonding
 Hydrogen
bonding is extreme case of
dipole-dipole bonding
 F, O, and N are all small and electronegative
 strong
electrons attraction
 H has only 1 electron, so if being pulled away H
proton is almost “naked”
H
end is always positive &
F, O, or N end is always negative
Strength of Hydrogen Bonding
 Fluorine
most electronegative element, so
 H-F
bonds are most polar and exhibit
strongest hydrogen bonding
 H-F
> H-O > H-N
(H-bonding…sound like FON to me!!!)
Hydrogen bonding:
• strongest IMF
• influences physical props a great deal
H-F > H-O > H-N
Strongest
Intermolecular
Force
Hydrogen Bonding
Dipole-Dipole
Dispersion
Indicate type of IMF for each molecule:
 NH3
•
 Ar
•
 N2
•
 HCl
•
 HF
•
 Ne
•
 O2
•
 HBr
•
 CH3NH2
•
Hydrogen bonding
Dispersion forces
Dispersion forces
Dipole-dipole forces
Hydrogen bonding
Dispersion
Dispersion
Dipole-dipole
Hydrogen bonding
H
H
O H
O H
H-Bonding = strongest IMF
much harder to “pull” molecules apart
H
H
C
H
H
H
H
C
H
H
Dispersion Forces= weakest IMF
much easier to “pull” molecules apart
IMF
vs.
Physical Properties
 If
IMF  then:
point 
 Melting point 
 Heat of Fusion  Change from solid to liquid w/o changing temp
 Heat of Vaporization  Change from liquid to gas w/o changing temp
 Boiling
while:
 Evaporation
Rate 
Rate at which conc. will go from liquid to gas
Why do some substances exist
as gases, some as liquids, and
some as solids at room temp?
#1 reason = IMF
#2 reason =
temperature (avg. KE)
IMF vs. Temp
 IMF
more important as temp is lowered
temperature – low evaporation rate
 High temperature – high evaporation rate
 Low
Intermolecular forces determine phase
 “Competition”
between strength of IMF & KE
determines phase
 If
IMF are strong,
substance will be solid
or liquid at room temp
Particles want to
clump together
 If
IMF are weak,
substance will be gas at
room temp
Particles free to
spread apart
It’s a balancing act!
Kinetic Energy
Intermolecular
Forces
High (fast)
weak
[this substance = a gas at room temperature]
Intermolecular
Forces
strong
Kinetic
Energy
Low (slow)
[this substance = a condensed phase (solid/liquid)]
REMEMBER…
 Temp
= average KE
 If we change T we change KE
 Increase KE will help “pull” molecules
apart (overcome IMF)