Prince George’s Community College
CHM1010 (Shah)
Chapter 11A
Liquids, Solids, and IMF
Outlines
 Solids, Liquids, and Gases
 IMF: London Dispersion Forces, Dipole-Dipole Forces, Hydrogen
bonding, Ion Dipole Forces
 Define: Surface tension, Viscosity
 Vaporization and Vapor Pressure
 The Clausius-Clapeyron Equation: Vapor Pressure and
Temperature
 Sublimation, Fusion
 Crystalline solids
States of matter: Solid, liquid, and gas
State
Density
Shape
Volume
Solid
High
Definite
Definite
Strength
of IMF
Strong
Liquid High
Indefinite
Definite
Moderate
Gas
Indefinite
Indefinite Weak
Low
Crystalline solids
The atoms or molecules that
compose them are arranged in
well ordered three –dimensional
array example: salt
Compressible
Generally can’t
be compressed
Not easily
compressible
Easily
compressible
Amorphous solids
The atoms or molecules that
compose them have no long range
order (different from the normal
solid state). Example: glass
Intramolecular forces (covalent bond) are relatively strong bonding
forces that hold atoms together in the molecular unit.
Intermolecular forces: Inter molecular forces are the relatively
weak forces that hold one molecular unit to other molecular unit. IMF
as a whole are usually called van der Waals forces.
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Types of IMF:
Dipole -dipole forces
London dispersion forces
Hydrogen bonds: A hydrogen bond between molecules is an
intermolecular force in which a hydrogen atom covalently bonded to
another atom (N, F, O)
Ion- dipole forces:
Dipole - Dipole forces:
Molecules with dipole moments can attract each other electro statically
by lining up so that positive and negative ends are close to each other.
This is called a dipole - dipole attraction. The forces created by this
attraction are called Dipole - Dipole forces.
Dipole - Dipole forces are typically only about 1% as strong as
covalent or ionic bonds. Particularly strong dipole - dipole forces,
however, are seen among molecules in which Hydrogen is bound to a
highly electronegative atom such as N, O or F (because of great
polarity and close approach of the dipoles). This is also called a
hydrogen bonding.
Dipole forces affect melting points and boiling points. Hydrogen bonds
are stronger dipole forces.
More polar the molecules are, the stronger the dipole forces between
molecules.
Hydrogen bond is much stronger in HF than HBr, HI, and HCl
EN difference is  1.8, 0.8, 0.5, 1.0
London Disperse forces: The forces that exist among noble gas atoms
and nonpolar molecules are called London dispersion forces.
London Dispersion forces result from temporary dipoles.
The Importance of London dispersion forces increases greatly as the
size of the atom increases (as the # of electrons in the molecules
increases).
Note: London and dipole forces between molecules are usually much
weaker than ionic forces in crystals. The very strong forces of ionic
bonds lead to much higher melting points and boiling points for ionic
compound than for molecular compounds.
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Example: M.P. for NaCl is 801°C and for I2 is 114°C.
The magnitude of LDF depends on molar mass and the shape of the
molecule.
For the same family: as molar mass increases strength increases so
boiling point increases.
Among n-pentane, n-hexane, n-heptane: n-heptane has a higher BP.
For the same molar mass: shape determines the strength: n-pentane
and neopentane: n-pentane has a higher BP. (larger the molecule,
higher the BP is) Another example: Formaldehyde and ethane
For the same molar mass if H bond is present: The molecule with “H”
bond has higher BP, and MP than molecule with no “H” bond.
Ethanol and Dimethyl ether
Ion-Dipole Force:
Ion-Dipole Force occurs when an ionic compound is mixed with a polar
compound. It is especially important in aq. Solution of ionic
compounds. Example: Sodium chloride is mixed with water.
Surface tension: The surface tension of a liquid is the energy required
to increase the surface area by a unit amount.
Surface tension increases as IMF increases
Viscosity: The resistance of a liquid to flow. Stronger the IMF higher
the viscosity is.
Capillary action: The ability of a liquid to flow against gravity up a
narrow tube when taking a blood sample.
Vaporization and Vapor Pressure:
Conversion from liquid to gas is called vaporization. Opposite to
vaporization (evaporation) is called condensation.
Volatile: Liquids that vaporize easily is known as volatile and those
that don’t vaporize easily is called nonvolatile. The rate of vaporization
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increases with increasing temperature/surface area or decreasing
strength of IMF.
Heat of vaporization: The amount of heat required to vaporize one
mole of a liquid to gas is its heat of vaporization (∆Hvap.)
Heat of vaporization of water at its normal boiling point of 100°C is 40.7kJ/mol (this process is endothermic)
H2O(l)  H2O(g)
∆Hvap. = +40.7kJ/mol
Heat of vaporization of 1mol of water condenses at its boiling point of
100°C is -40.7kJ/mol (this process is exothermic)
H2O(g)  H2O(l)
∆Hvap. = - 40.7kJ/mol
Example: Calculate the mass of water in g that can be vaporized at its
boiling point with 155kJ of heat. (answer: 68.6g H2O)
Calculate the amount of heat in kJ required to vaporize 2.58kg of
water at its boiling point. (5830kJ)
Dynamic equilibrium: Dynamic equilibrium occurs when the rate of
condensation is equal to the rate of evaporation.
Vapor pressure: The pressure of a gas in dynamic equilibrium with its
liquid is called its vapor pressure. The vapor pressure of a particular
liquid depends on the IMF present in the liquid and temperature.
Weak IMF result in volatile substances with high vapor pressures
because IMF are easily overcome by thermal energy. Strong IMF result
in nonvolatile substances with low vapor pressures.
Boiling point: The boiling point of liquid is the temperature at which its
vapor pressure equals the external pressure.
The normal boiling point of a liquid is the temperature at which its
vapor pressure equals 1 atm.
The Clausius – Clapeyron equation:
ln Pvap. = -(∆Hvap/R)(1/T)+ lnẞ
slope = - ∆Hvap/R
∆Hvap = -slope x R
where R = 8.314 J/mol.K
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ln(P2/P1) = - (∆Hvap/R)(1/T2 - 1/T1)
1. The vapor pressure of CH2Cl2 was measured as a function of
temperature, and the following results were obtained.
Temperature (K)
200
220
240
260
280
300
Vapor Pressure (torr)
0.8
4.5
21
71
197
391
Determine the heat of vaporization of dichloromethane.
(answer:31.4kJ/mol)
2. Methanol (CH3OH) has a normal boiling point of 64.6°C and a
heat of vaporization 35.2kJ/mol. What is the vapor pressure of
methanol at 12.0°C? (answer: 73.4 torr) R = 8.314 J/mol.K
Sublimation: The process by which a substance goes directly
from the solid to the gaseous state without passing through the
liquid state. Example: Sublimation of NH4Cl, solid CO2 (dry iice)
Deposition: The opposition to sublimation is deposition, the
transition from directly gas to solid.
Fusion (melting): The increasing thermal energy causes the
water molecules to vibrate faster and faster. At the melting
point, the molecules have enough thermal energy to overcome
the IMF that hold them at their stationary points, and the solid
turns into liquid. This process is known as fusion or melting.
The opposite of melting is known as freezing.
Melting
H2O(s)  H2O(l) ∆Hfus. = +6.02kJ/mol at mp (0°C)
Freezing
H2O(l)  H2O(s)
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∆H = -∆Hfus. = -6.02kJ/mol