intermolecular forces: liquids, solids & phase … · liquids, solids & phase changes...
TRANSCRIPT
INTERMOLECULAR FORCES:
LIQUIDS, SOLIDS & PHASE CHANGES (Silberberg, Chapter 12)
Ideal gas
molecules act
independently
PV=nRT
Liquid
molecules “stick” to
one another
Real gas
molecules
attract/repulse one
another
nRTnbVV
anP ))((
2
2
Intermolecular interactions
Centrosymmetric formamide dimer at the crystal geometry.
Blue shows regions of the molecular surface with a partial
positive charge, red those regions with a partial negative charge
and white regions which are neutral.
Types of Intermolecular interactions (intermolecular forces)
1. London force (“dispersion force”)
Electron cloud gives instantaneous dipole moment (e- build up at one
end of molecule leaving the nucleus at the other end partially
exposed)
Molecules stick together because their partial charges attract one another
London forces act between all types of molecules (polar and non-polar).
• Strength increases with molar mass : heavier molecules have more electrons (further from the nuclei)
bigger fluctuations in electron movement (they are more polarizable)
Eg. F2 and Cl2 are gases, Br2 is a liquid, I2 a solid at room temperature.
Strength influenced by molecule shape
Rod-shaped molecules have greater London forces of
attraction because the instantaneous dipoles can get
closer
Note: “van der Waals forces” is the generic term for ALL intermolecular forces.
It is NOT interchangeable for “dispersion forces.”
2. Dipole-dipole interactions:
Polar molecules have permanent partial charges.
Polar molecules that are brought close to one another tend to orient their dipole moments so that the plus
end of one molecule faces the minus end of another. This is a dipole-dipole interaction:
•Strength depends on magnitude of bond dipoles and shape of molecule. If bond dipoles cancel one
another within a molecule, then the molecule itself has no dipole moment.
3. Hydrogen bonding:
Occurs when a hydrogen atom is bonded to a strongly electronegative atom with lone pairs (N, O or F)
R-X-H … Y-R’ (X, Y = F, O or N)
O
H H
O
H H
O
H H
O
H H
Eg. water
+
+
-
-
covalent bond
hydrogen bond
Hydrogen bonding is the strongest of the intermolecular interactions
CCH3
O
O
H
C CH3
O
O
Hacetic acid
NH H
NHH
NH Hamine
In hexagonal ice I, (natural form of ice on Earth), each water molecule is H-bonded to 4
neighbouring molecules in a tetrahedral arrangement. H-bonds are less ordered at 20°C but they still
account for the cohesiveness of liquid water where each molecule remains H-bonded to an average
of 3.5 neighbours (at that temperature water molecules exchange their positions about 1011 times per
second !). Finally, the bonds weaken at higher temperatures and water vaporizes because thermal
motion causes an uncertain random orientation of the water dipoles. Vapour is like a gas where
molecules are too distant and move too fast to be able to interact.
http://wwwarpe.snv.jussieu.fr/td_2_eng/lsh.html
Table of intermolecular forces and energies
Type of interaction
Typical
energy
(kJ mol-1)
Interacting species
Intermolecular
London (dispersion) 0.5-40 Polarizable e- clouds
(all atoms/molecules)
dipole-dipole 5-25 dipole-dipole
(polar molecules)
Hydrogen bonding 10-40 Polar bond to H – dipole
(molecules where H is bonded to
electronegative atom with lone
pairs)
Chemical bonding
Ionic 400-4000 Cation-anion
Covalent 150-1100 atoms sharing electrons in a bond
Properties of liquids; influence of intermolecular interactions:
Viscosity
• Viscosity = resistance to flow
• High viscosity slow flow rate
• Strong intermolecular forces molecules held together can’t
move past one another easily high viscosity
• Viscosity as temperature because molecules have higher Ek
O
H H
H
H
H
H
H
Hvs
But London forces can add up to quite significant degree:
CH3(CH2)4CH=CHCH2CH=CH(CH2)7COOH CH3(CH2)7CH=CH(CH2)7COOH
Linoleic acid Oleic acid
Properties of liquids:
Surface tension
In a liquid, molecules experience equal
forces from all directions except at the
surface (experience a net inward force)
Strong intermolecular forces surface tension
Water meniscus curves upward because forces
between water molecules and the oxygen atoms
and OH groups in glass are stronger than the
forces between water molecules.
Mercury meniscus curves downwards
because forces between mercury atoms
are stronger than between Hg atoms and
the glass
capillarity
intermolecular interactions
occur in 3 dimensions
int. interactions
occur across &
below surface
net vector
for attractive
forces is down
Phase transition Name Example
Gas to liquid Condensation or liquefication
Dew
Liquid to gas Vaporisation Boiling water steam
Gas to solid Condensation or deposition
Frost
Solid to gas Sublimation Evaporation of CO2
Liquid to solid Freezing water ice
Solid to liquid Melting or fusion ice water
Phase transitions
solid liquid gas
melting
freezing
vaporizing
condensing
sublimination exothermic
endothermic
Vapour Pressure
• At a given T, molecules in liquid have a distribution of speeds.
• Some move fast enough to escape liquid (enter gas phase)
• At higher T, average kinetic energy is higher so greater fraction
of molecules can escape liquid
Consider liquid in sealed container:
Some fast moving molecules escape into gas phase
These gas molecules exert pressure on the liquid surface
The pressure as more molecules enter gas phase
Some gas molecules collide with surface and stick to it, re-entering the liquid phase
Eventually, rate of molecules evaporating equals rate of molecules condensing (equilibrium is reached)
vapour pressure : partial pressure exerted by a vapour over a liquid at a fixed temperature
Vapour pressure is characteristic for a given liquid (or solid)
Function of the intermolecular forces
Vapour pressure as temperature for a given liquid
Example: Which would you expect to have a higher boiling point, p-dichlorobenzene or o-dichlorobenzene?
Cl
Cl
Cl
Cl
Boiling point and melting point
In open container, atmosphere exerts a pressure on the liquid
surface. As T liquid molecules move more quickly & leave
surface more often. At some temperature, KE is great enough for
bubbles of vapour to appear in the liquid (“the liquid boils”).
Vapour pressure as function of
temperature and intermolecular
forces
The boiling point is the temperature at which the vapour
pressure equals the external pressure
The normal boiling point, Tb, of a liquid is defined as the boiling
point at 1 atm external pressure
The normal freezing point, Tf, of a liquid is the temperature at
which it freezes at 1 atm pressure
(Tf is only slightly dependent on pressure)
Strong intermolecular forces low vapour pressure
high boiling point
The normal boiling point, Tb, is defined as the boiling point at
1 atm external pressure
Heat of phase transitions
1. Gas cools:
2. Gas condenses:
3. Liquid cools:
5. Solid cools:
4. Liquid freezes:
Note:
• Within a phase, a change in heat is accompanied by a change in temperature (because of change in Ek). Heat lost or
gained depends on amount of substance, molar heat capacity for that phase and change in temperature.
•During a phase change, a change in heat occurs at constant temperature (energy needed to cause change in phase)
•Melting or boiling is endothermic; Condensation or freezing is exothermic.
Phase diagrams
• summarise phases of a substance under different conditions of temperature and pressure
Phase diagram of CO2:
solid liquid
gas
triple
point
critical
point
Amorphous and crystalline silicon dioxide.
When molten silica cools quickly it becomes a
glass. The atoms (red=O; black=Si) are
arranged in a disorderly fashion.
Quartz is a crystalline form of silica, SiO2. The
atoms (red=O; black=Si) are arranged in an
orderly network.
The Solid State (Silberberg, Chapter 12)
Amorphous: atoms/molecules in random arrangement (“frozen liquid”)
Crystalline: atoms/molecules in ordered pattern. Often have flat “faces”and definite angles at edges
(formed by orderly stacks of atoms)
Class Examples Characteristics
metallic s- and d- block elements malleable, ductile, lustrous, electrically and
thermally conducting
ionic NaCl, CsCl, KNO3, CuSO4 .
5 H2O
hard, rigid, brittle,
high melting points,
those soluble in water conduct electrical current;
when molten they are electrical conductors
network
(covalent
bonding)
B, C, BN, SiO2
hard, rigid, brittle,
very high melting points, insoluble in water
molecular
(covalent
bonding;
intermolecular
interactions)
H2O (ice), S8, I2, glucose,
sucrose, naphthalene
relatively low melting and boiling points, brittle
if pure
Crystalline solids can be classified in terms of their bonding:
portion of a 3-D
lattice
lattice point
unit
cell
The crystal lattice and unit cell
• All particles in a crystal are
packed in an orderly way
• Imagine placing a point at each
repeating part of the pattern –
the points form a regular array
called a crystal lattice
• All lattice points are identical
The smallest part of the crystal that, if repeated in 3-D gives the whole crystal, is called the unit cell
Seven crystal systems (different shapes of unit cells) and fourteen kinds of unit cells exist in nature
We consider only the cubic crystal system (all sides are equal length (called a) and all angles 90)
In the cubic system there are 3 possible packing arrangements:
Simple cubic /
Primitive cubic (P)
atom at each of the 8 corners
of cube
Body-centred cubic (I)
atom at each corner
& atom at centre of cube
Face-centred cubic (F)
atom at each corner
& atom in centre of each face of
cube
1 atom
at center
Coordination number = 8 Coordination number = 12 Coordination number = 6
atom
at 8 corners
1 8 –
Atoms/unit cell = x 8
= 1
1 8 –
atom
at 8 corners
1 8 – atom
at 8 corners
1 8 –
– atom
at 6 faces
1 2
Atoms/unit cell =
( x 8) + ( x 6)=4 1 8 – 1 –
2
Atoms/unit cell =
( x 8) + 1 = 2 1 8 –
Primitive cubic Body-centred cubic Face-centred cubic
Metals
• cations (identical spheres) packed tightly together and surrounded by a “sea” of electrons
• denser than most solids
• malleable and ductile
• conduct heat and electricity
• many pack in cubic unit cells
Copper: face-centred cubic
Figure 12.28 Packing of spheres.
simple cubic
(52% packing efficiency)
body-centered cubic
(68% packing efficiency)
hexagonal
unit cell
Figure 12.28 (continued)
closest packing of first
and second layers
layer a
layer a
layer b
layer c
hexagonal
closest
packing cubic closest
packing
abab… (74%) abcabc… (74%)
expanded
side views
face-centered
unit cell
expanded view space-filling
Ionic solids
•formed from oppositely charged ions (overall crystal is neutral)
•held together by attraction of cations and anions
•hard, rigid, brittle
•conduct electricity when molten or in solution; not as solids
•anions close-pack and cations fit into interstitial spaces
Allotropes (polymorphs) of Carbon
Diamond and graphite are examples of Network solids
• atoms covalently bonded
• entire crystal is one network
• hard, brittle, high melting points
• individual molecules, held in crystal structure by
intermolecular forces only
• physical properties depend on types & strengths
of intermolecular interactions
• relatively low melting points (100 - 500C)
Molecular crystal structures of carbon
C60 Face centred cubic