General Chemistry
Principles and Modern Applications
Petrucci • Harwood • Herring
8th Edition
Chapter 13 (12 in 9th Ed.): Liquids, Solids and
Intermolecular forces
Chem 1050
Chapter 13
Dr. Azra Ghumman
Contents
13.1
13-2
13-3
13.4
13-5
13-6
13-9
Chem 1050
Intermolecular Forces and some Properties of Liquids
Vaporization of Liquids: Vapor Pressure ( page 485 and 487
omitted)
Some Properties of Solids
Phase diagrams
Van der Waals Forces
Hydrogen Bonding (effect on viscosity is omitted)
Energy Changes in the Formation of Ionic Crystals.
Chapter 13
Dr. Azra Ghumman
Intermolecular forces
• Intermolecular forces- forces between molecules
– Nonbonding forces.
– Determine the physical properties of the substance e.g.
f.p., b.p., m.p. and vapour pressure.
• Intramolecular forces – forces within a molecule.
– Bonding forces.
– Determine chemical properties of the substance e.g.
heat of combustion, reactivity etc.
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1
Intermolecular Forces and Some Properties of
Liquids
• Cohesive Forces-Intermolecular forces between like
molecules.
• Adhesive Forces-Intermolecular forces between unlike
molecules.
• Surface Tension(Greek letter gamma γ)-Energy or work
required to increase the surface area of a liquid.
– Units : J m-2
– Effect of Temperature
• Wetting agents- Substances that reduce the surface
tension of of water and allow it to spread more easily.
• Viscosity-A liquid’s resistance to flow.
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Dr. Azra Ghumman
Intermolecular Forces
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Intermolecular Forces
Meniscus formation
•In water
Cohesive Forces < Adhesive Forces
•In Hg
Concave
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Convex
Cohesive Forces >Adhesive
Forces
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Vaporization of Liquids:
Vapor Pressure
•Molecules having K.E sufficiently high are able to overcome intermolecular
forces of attraction and escape from the surface of the liquid into gaseous state
(vaporization).
Dynamic equilibrium
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Factors effecting rate of vapourization
Evaporation or Vapourization: Passage of molecules from
the surface of liquid into gaseous state i.e. vapours.
(liquid →gas).
Factors effecting rate of vaporization
• Increased temperature
• Increased surface area of the liquid
• Weak intermolecular forces
Condensation: Conversion of vapours to a liquid.
– reverse of vapourization (gas→ liquid).
Chem 1050
Chapter 13
Dr. Azra Ghumman
Enthalpy of vapourization
Enthalpy of vapourization (∆
∆Hvap)-The quantity of heat that
must be absorbed if a certain quantity of liquid is
vapourized at a constant temperature.
– Units kJ mol-1
∆Hvap = Hvapor – Hliquid
= - ∆Hcondensation
– Vapourization is an endothermic process and ∆Hvap is
always positive.
– Is condensation an endothermic or exothermic
process?
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Enthalpy of Vapourization
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Chapter 13
Dr. Azra Ghumman
Vapour pressure
•
•
•
•
•
Dynamic equilibrium
liquid
vapour
Vapour pressure- The pressure exerted by molecules of a substance
in gaseous state (vapours) in dynamic equilibrium with its liquid at a
given temperature.
Volatile liquids- Liquids with high vapour pressure at room
temperature (low b.p)e.g. acetone, diethyl ether
– The weaker the intermolecular forces(IMF) the more volatile is the
liquid.
Non volatile liquid- Liquid with very low vapour pressure at room
temperature e.g. Hg. (strong IMF, high b.p)
Vapour pressure curve- A plot of vapour pressure vs temperature
– Vapour pressure increases with temperature.
Chem 1050
Chapter 13
Dr. Azra Ghumman
Vapour pressure curve
Pvap of water at 25°C = 23.8 mmHg
Diethyl ether
C6H6
water
toluene
aniline
•Normal Boiling Point
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Chapter 13
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Vapour pressure
Pvap of a liquid depends only on the particular liquid and its temperature.
Mercury
Pvap
manometer
of liquid
Chem 1050
Pvap
independent
of Vliq
Pvap
independent
Pvap dependent
on T
of Vgas
Chapter 13
Dr. Azra Ghumman
Boiling and Boiling point
Boiling: A state, at particular temperature, in which bubbles
of vapours form throughout a liquid and escape with a
pressure equal to the pressure over the surface of the liquid
(barometric pressure).
•Normal Boiling Point-Temperature at which the vapour
pressure of a liquid is equal to standard atmospheric
pressure (1 atm =760 mmHg).
•The critical point is the condition of temperature and
pressure at which a liquid and its vapour become
indistinguishable
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Dr. Azra Ghumman
Boiling point and Barometric pressure
• Effect of external pressure on boiling point;
– Boiling point of the liquid increases with increase in
external pressure.
• As the external pressure increases, the vapour pressure of the
liquid required for boiling increases.e.g.
• Tb of H2O =100°C at 1atm
• Tb of H2O > 100 °C at P > 1atm
– At lower barometric pressure (at high altitudes), b.p. of
water is low and it takes longer to cook the food.
• Tb of H2O =100°C at 1atm
• Tb of H2O = 95°C at 630 mmHg = 0.83atm
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Clausius-Clapeyron Equation for expressing
Vapour pressure data
ln P = -A (1/T) + B
(e)
(d) (c) (b)
(a)
Where A = ∆Hvap/R
y = mx + c
•A plot of ln P versus 1/T gives
a straight line with its slope
= -∆Hvap/R.
∆Hvap
P2
1
1
=
(
)
P1
T1 T2
R
Clausius-Clapeyron Equation
ln
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Dr. Azra Ghumman
An equation for expressing vapour pressure data
∆Hvap
P2
1
1
=
(
)
P1
T1 T2
R
Example: Vapour pressure of methyl alcohol is100 mmHg at
21.2 °C and its enthalpy of vapourization is 38.0 kJ mol-1.
What is its vapor pressure at 25.0 °C?
ln
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Phase changes
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13-3 Some Properties of Solids
• Melting, Melting point and heat of fusion:
– Melting and freezing points.
– The melting point of a solid and freezing point of its
liquid are identical.
• At this temperature, solid and liquid coexist in
equilibrium.
• Enthalpy of Fusion ∆Hfus – Heat required to melt a solid
at constant temperature is called enthalpy of fusion.
expressed in kJ mol-1.
• ∆Hfreezing= - ∆Hfus
Chem 1050
Chapter 13
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Cooling and Heating curves for water
melting point
freezing point
∆Hfus(H2O) = +6.01 kJ/mol
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Sublimation
•Sublimation: The direct passage of molecules from
solid to the vapor state.
Reverse process is called deposition.
•Enthalpy of sublimation ∆Hsub
∆Hsub = ∆Hfus + ∆Hvap
∆Hsub = -∆Hdeposition
Sublimation of Iodine, purple vap.is
produced at 70°C well below its mp 114°C
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Phase Diagrams (Iodine)
fusion curve
critical point
vapour pressure
curve
triple point
sublimation curve
OD represents the fusion curve. There is little effect of pressure on melting
point.
O is the triple point. A unique temperature and pressure at which three states
Chem 1050
Chapter 13
of matter
coexist.
Dr. Azra Ghumman
Phase Diagrams(CO2)
•Triple point is greater than one atmosphere,
so we do not form liquid. Sublimation occurs.
•Liquid present in fire extinguishers (P > 5.1 atm)
Chem 1050
Chapter 13
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The Critical Point
•The critical point is the condition of temperature and
pressure at which a liquid and its vapour become
indistinguishable
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Critical Temperatures and Pressures
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Chapter 13
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Water
Fusion curve (OD) has a negative slope. Unusual behavior. Ice skating.
Polymorphysm- The existence of a solid substance in more than one
form. This is common.
Chem 1050
Chapter 13
Dr. Azra Ghumman
Van der Waals Forces
•
Dispersion or London forces. (1928 Fritz London)
– Instantaneous dipoles.
• Asymmetric e- movement in an orbital to cause a polarization.
– Induced dipoles.
• Electrons move in response to an outside force e.g. nearby cation.
•
•
– Related to polarizability.
Polarizability- the tendency for charge separation to occur in a
molecule.
Polarizability increases with
– Increased # of electrons
Molecular mass
– # of atoms in a molecule
– Valence electrons in large molecules are easily polarizable.
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Dispersion forces and polarizability
•
•
Strength of Dispersion forces:
– Dispersion forces become stronger as polarizability increases.
Result: Strong IMF between molecules resulting in higher mp and bp
of covalent substances with increasing molecular mass.
e.g He atomic mass = 4 u
bp = 4K
Radan atomic mass = 222 u
bp = 211K
Halogens- the order of mp and bp. is
F2< Cl2 <Br2 <I2
Strength of dispersion forces also depend on the shape of the
molecules.
– Elongated molecules are more polarizable than small, compact and
symmetrical molecules
– Isomers-Two substances with identical number and kinds of atoms
but different molecular shapes may have different properties.
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Phenomenon of Induction
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Instantaneous and Induced Dipoles
a. Nonpolar molecule b. Instantaneous condition c. Induced dipole
Isomers- n-pentane and Neopentane
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Dipole Dipole Interactions
• Polar molecules- with permanent dipole.
• Dipoles tend to arrange themselves with the +ve end of one
dipole pointed toward the -ve end of a neighboring dipole.
•This tendency can affect mp
of solids and bp of liquids.
Chem 1050
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Comparing physical properties of polar and non
polar substances
• Example: Which of the following substances would you
expect to have the highest boiling point:
C3H8,CO2,CH3CN? Explain
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Chapter 13
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Summary of van der waals Forces
•
•
•
•
Dispersion (London) forces exist between all molecules
– Involve displacement of all the electrons in a molecule.
Forces associated with permanent dipoles involve displacement of the
electron pairs in a bond rather than in molecules as a whole.
– They exist in polar molecules.
– Adds to effect of dispersion forces.
When comparing substances of roughly comparable molecular masses,
dipole forces can produce significant differences in properties such as
mp, bp and enthalpy of vaporization.
When comparing substances of widely different molecular masses,
dispersion forces are usually more significant than dipole forces.
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Effect of IMF,s on physical properties of
substances
•
Van derWaals Forces
Substance M (u)
b.p (K) ∆Hvap
F2
HCl
HBr
HI
85.01
188.11
206.43
237.80
38.00
36.00
80.92
127.91
6.86
16.15
17.61
19.77
%Dispersion
100
81.4
94.5
99.5
%Dipole
0
18.6
5.5
0.5
HCl and HBr are more polar compared to HI. This is not
sufficient to reverse trends produced by molecular masses.
Dispersion forces are predominant intermolecular forces in
the series of halides.
Chem 1050
Chapter 13
Dr. Azra Ghumman
13-6 Hydrogen Bonding
•
•
Hydrogen bonding- occurs when a H atom is covalently bonded to a
small highly electronegative atom such as F, O, N and is
simultaneously attracted to a highly electronegative atom in a
neighboring molecule e.g. H-bonding in H2O.
– Hydrogen bonds are possible only with certain H- containing
compounds because the nuclei of H atom is not shielded by inner
shell electron from attraction of lone pair of nearby atoms.
– e.g. H-bonding in HF
– Hydrogen bonds are relatively stronger than other intermolecular
forces.
– Energies of hydrogen bonding are of the order of 15-40kJ/mol (
covalent bond energies > 150kJ mol-1).
Weak H-bonds- with Cl and S (3rd period).
» E.g. H-Cl---H-Cl
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Hydrogen Bonding
Comparison of boiling points of some hydrides of the
elements of group 14, 15, 16 and 17
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Comparison of hydrides
• Group 16 hydrides
Intermolecular forces:
• Dispersion forces.
– All are polar. Type: AX2E2
– Polarizablity of the molecules increases as the total
number of electron increases, resulting in stronger
dispersion forces.
• Hydrogen bonding only in H2O.
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Hydrogen Bonding in HF(g)
In gaseous HF, molecules are associated in cyclic HF6 structure.
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Hydrogen Bonding in Water
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Intermolecular and intramolecular H-Bonding
Electron density contour showing intramolecular H-bonding
in salicylic acid.
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Chapter 13
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Properties effected by H-bonding
• Unusual freezing point behaviour of water and its effect on
aquatic life.
– Max. density of water at 3.98°C.
– When temp. falls below 4 °C, denser water sinks down
and ice at the top tends to insulate the water below from
further heat loss.
• Hydrogen bonding in living matter.
Chem 1050
Chapter 13
Dr. Azra Ghumman
Interionic Forces
Ionic Solids: consist of cations and anions held together by electrical
attraction of opposite charges (ionic bonds).
• Properties: nonvolatile and high mp solids (600-2000°C)
– Non-conducting as solids, but conducting in molten state.
• Lattice energy- The energy given off when separated, oppositely
charged, gaseous ions come together to form one mole of a solid ionic
compound.
– Lattice energies can be useful in making predictions about the m.p
and water solubilities of ionic compounds.
• Rough rule- The lower the lattice energy, the greater the quantity of
an ionic solid that can be dissolved in given quantity of solvent.
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Interionic Forces
•
The attractive force between a pair of oppositely charged ions
increases with increased charge on the ions and with decreased ionic
sizes (based on coulomb’s law).
Chem 1050
Chapter 13
Dr. Azra Ghumman
Energy Changes in the Formation of Ionic
crystals of NaCl(s)
Lattice energy is commonly
determined indirectly through
an application of Hess’s Law
known as Boron-Fajans- Haber cycle.
•Sublime one mol of solid Na
•Dissociate 0.5 mol of Cl2 (g)into 1 mol
of Cl(g).
•Ionize 1 mole of Na (g) to Na+ ion(g)
•Convert 1mole of Cl(g) to Cl- ion(g)
•Allow the Na+ (g) and Cl-(g) to form 1
mole of NaCl(s)
The overall ∆Hoverall = ∆H°f [NaCl(s)]
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Chapter 13
Dr. Azra Ghumman
Practice Example A
• The enthalpy of sublimation of cesium is 78.2kJ/mol, and
∆Hf°[CsCl(s)] = -442.8 kJ/mol. Use these values, together
with other data given below, to calculate the lattice energy
of CsCl(s).
Cs(s) →Cs(g)
∆H1= 78.2kJ/mol
1/2Cl2 (g) →Cl(g)
∆H2= + 122kJ/mol
Cs(g) →Cs+(g)+ e∆H3 (I1) = 375.7kJ/ mol
Cl(g) + e- →Cl- (g)
∆H4 (EA of Cl) =-349.0 kJ/mol
Cs+(g) + Cl- (g) →CsCl(s) ∆H5 (lattice energy) =? kJ/mol
Overall Cs(s) + 1/2Cl2(g) →CsCl(s) ∆H°f = -442.8kJ/mol
• ∆Hoverall = ∆H1 + ∆H2 + ∆H3 + ∆H4 + ∆H5
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