Title: Chapter 11 Intermolecular Forces
1Chapter 11 Intermolecular Forces
211.1 Intermolecular Forces (IMF)
- IMF lt intramolecular forces (covalent, metallic,
ionic bonds) - IMF strength solids gt liquids gt gases
- Boiling points and melting points are good
indicators of relative IMF strength.
311.2 Types of IMF
- Electrostatic forces act over larger distances
in accordance with Coulombs law - Ion-ion forces strongest found in ionic
crystals (i.e. lattice energy)
4- Ion-dipole between an ion and a dipole (a
neutral, polar molecule/has separated partial
charges)
- Increase with increasing polarity of molecule and
increasing ion charge.
Ex Compare IMF in Cl- (aq) and S2- (aq).
lt
5- Dipole-dipole weakest electrostatic force exist
between neutral polar molecules
- Increase with increasing polarity (dipole moment)
of molecule
Ex What IMF exist in NaCl (aq)?
6- Hydrogen bonds (or H-bonds)
- H is unique among the elements because it has a
single e- that is also a valence e-. - When this e- is hogged by a highly EN atom (a
very polar covalent bond), the H nucleus is
partially exposed and becomes attracted to an
e--rich atom nearby.
7- H-bonds form with H-XX', where X and X' have
high EN and X' possesses a lone pair of e- - X F, O, N (since most EN elements) on two
molecules
F-H O-H N-H
F O N
8- There is no strict cutoff for the ability to
form H-bonds (S forms a biologically important
hydrogen bond in proteins). - Hold DNA strands together in double-helix
Nucleotide pairs form H-bonds
DNA double helix
9- H-bonds explain why ice is less dense than water.
10Ex Boiling points of nonmetal hydrides
- Conclusions
- Polar molecules have higher BP than nonpolar
molecules - ? Polar molecules have stronger IMF
- BP increases with increasing MW
- ? Heavier molecules have stronger IMF
Boiling Points (ºC)
- NH3, H2O, and HF have unusually high BP.
- ? H-bonds are stronger than dipole-dipole IMF
11Inductive forces
- Arise from distortion of the e- cloud induced by
the electrical field produced by another particle
or molecule nearby. - London dispersion between polar or nonpolar
molecules or atoms - Proposed by Fritz London in 1930
- Must exist because nonpolar molecules form liquids
Fritz London(1900-1954)
12- How they form
- Motion of e- creates an instantaneous dipole
moment, making it temporarily polar.
- Instantaneous dipole moment induces a dipole in
an adjacent atom - Persist for about 10-14 or 10-15 second
- Ex two He atoms
13 Geckos!
- Geckos feet make use of London dispersion forces
to climb almost anything. - A gecko can hang on a glass surface using only
one toe. - Researchers at Stanford University recently
developed a gecko-like robot which uses synthetic
setae to climb walls
http//en.wikipedia.org/wiki/Van_der_Waals27_forc
e
14- London dispersion forces increase with
- Increasing MW, of e-, and of atoms
(increasing of e- orbitals to be distorted) - Boiling points
- Effect of MW Effect of atoms
- pentane 36ºC Ne 246C
- hexane 69ºC CH4 162C
- heptane 98ºC
-
- ??? effect
- H2O 100C
- D2O 101.4C
- Longer shapes (more likely to interact with
other molecules) - C5H12 isomers 2,2-dimethylpropane 10C
- pentane 36C
15Summary of IMF
Van der Waals forces
16Ex Identify all IMF present in a pure sample of
each substance, then explain the boiling points.
BP(C) IMF Explanation
HF 20
HCl -85
HBr -67
HI -35
Lowest MW/weakest London, but most polar/strongest dipole-dipole and has H-bonds
Low MW/weak London, moderate polarity/dipole-dipole and no H-bonds
Medium MW/medium London, moderate polarity/dipole-dipole and no H-bonds
Highest MW/strongest London, but least polar bond/weakest dipole-dipole and no H-bonds
London, dipole-dipole, H-bonds
London, dipole-dipole
London, dipole-dipole
London, dipole-dipole
1711.3 Properties resulting from IMF
- Viscosity resistance of a liquid to flow
- Viscosity depends on
- -the attractive forces between molecules
- -the tendency of molecules to become entangled
- -the temperature
1811.3 Properties resulting from IMF
- Surface tension energy required to increase the
surface area of a liquid
19- 3. Cohesion attraction of molecules for other
molecules of the same compound - 4. Adhesion attraction of molecules for a
surface
20- Meniscus curved upper surface of a liquid in a
container a relative measure of adhesive and
cohesive forces - Ex
Hg
H2O
(cohesion rules)
(adhesion rules)
21Phase Changes
- Surface molecules are only attracted inwards
towards the bulk molecules. - Sublimation solid ? gas.
- Vaporization liquid ? gas.
- Melting or fusion solid ? liquid.
- Deposition gas ? solid.
- Condensation gas ? liquid.
- Freezing liquid ? solid.
- Energy Changes Accompanying Phase Changes
- Energy change of the system for the above
processes are
22Intermolecular Forces Bulk and Surface
23Phase Changes
- Energy Changes Accompanying Phase Changes
- Sublimation ?Hsub gt 0 (endothermic).
- Vaporization ?Hvap gt 0 (endothermic).
- Melting or Fusion ?Hfus gt 0 (endothermic).
- Deposition ?Hdep lt 0 (exothermic).
- Condensation ?Hcon lt 0 (exothermic).
- Freezing ?Hfre lt 0 (exothermic).
- Generally heat of fusion (enthalpy of fusion) is
less than heat of vaporization - it takes more energy to completely separate
molecules, than partially separate them.
24Phase Changes
- Energy Changes Accompanying Phase Changes
- All phase changes are possible under the right
conditions (e.g. water sublimes when snow
disappears without forming puddles). - The sequence
- heat solid ? melt ? heat liquid ? boil ? heat gas
- is endothermic.
- The sequence
- cool gas ? condense ? cool liquid ? freeze ? cool
solid - is exothermic.
25Phase Changes
Energy Changes Accompanying Phase Changes
26Phase Changes
- Heating Curves
- Plot of temperature change versus heat added is a
heating curve. - During a phase change, adding heat causes no
temperature change. - These points are used to calculate ?Hfus and
?Hvap. - Supercooling When a liquid is cooled below its
melting point and it still remains a liquid. - Achieved by keeping the temperature low and
increasing kinetic energy to break intermolecular
forces.
27Phase Changes
Heating Curves
28Heating Curve Illustrated
29Phase Changes
- Critical Temperature and Pressure
- Gases liquefied by increasing pressure at some
temperature. - Critical temperature the minimum temperature for
liquefaction of a gas using pressure. - Critical pressure pressure required for
liquefaction.
30Critical Temperature, Tc
31Transition to Supercritical CO2
32Supercritical CO2 Used to Decaffeinate Coffee
33Vapor Pressure
- Explaining Vapor Pressure on the Molecular Level
- Some of the molecules on the surface of a liquid
have enough energy to escape the attraction of
the bulk liquid. - These molecules move into the gas phase.
- As the number of molecules in the gas phase
increases, some of the gas phase molecules strike
the surface and return to the liquid. - After some time the pressure of the gas will be
constant at the vapor pressure.
34Gas-Liquid Equilibration
35Vapor Pressure
- Explaining Vapor Pressure on the Molecular Level
- Dynamic Equilibrium the point when as many
molecules escape the surface as strike the
surface. - Vapor pressure is the pressure exerted when the
liquid and vapor are in dynamic equilibrium.
36Vapor Pressure
- Volatility, Vapor Pressure, and Temperature
- If equilibrium is never established then the
liquid evaporates. - Volatile substances evaporate rapidly.
- The higher the temperature, the higher the
average kinetic energy, the faster the liquid
evaporates.
37Liquid Evaporates when no Equilibrium is
Established
38Vapor Pressure
Volatility, Vapor Pressure, and Temperature
39Vapor Pressure
- Vapor Pressure and Boiling Point
- Liquids boil when the external pressure equals
the vapor pressure. - Temperature of boiling point increases as
pressure increases. - Two ways to get a liquid to boil increase
temperature or decrease pressure. - Pressure cookers operate at high pressure. At
high pressure the boiling point of water is
higher than at 1 atm. Therefore, there is a
higher temperature at which the food is cooked,
reducing the cooking time required. - Normal boiling point is the boiling point at 760
mmHg (1 atm).
40Phase Diagrams
- Phase diagram plot of pressure vs. Temperature
summarizing all equilibria between phases. - Given a temperature and pressure, phase diagrams
tell us which phase will exist. - Features of a phase diagram
- Triple point temperature and pressure at which
all three phases are in equilibrium. - Vapor-pressure curve generally as pressure
increases, temperature increases. - Critical point critical temperature and pressure
for the gas. - Melting point curve as pressure increases, the
solid phase is favored if the solid is more dense
than the liquid. - Normal melting point melting point at 1 atm.
41Phase Diagrams
- Any temperature and pressure combination not on a
curve represents a single phase.
42Phase Diagrams
- The Phase Diagrams of H2O and CO2
- Water
- The melting point curve slopes to the left
because ice is less dense than water. - Triple point occurs at 0.0098?C and 4.58 mmHg.
- Normal melting (freezing) point is 0?C.
- Normal boiling point is 100?C.
- Critical point is 374?C and 218 atm.
- Carbon Dioxide
- Triple point occurs at -56.4?C and 5.11 atm.
- Normal sublimation point is -78.5?C. (At 1 atm
CO2 sublimes it does not melt.) - Critical point occurs at 31.1?C and 73 atm.
43Phase Diagrams
The Phase Diagrams of H2O and CO2
4411.4 Phase Changes
- Processes
- Endothermic melting, vaporization, sublimation
- Exothermic condensation, freezing, deposition
I2 (s) and (g)
Microchip
45Water Enthalpy diagram or heating curve
4611.5 Vapor pressure
Pressure cooker 2 atm
Normal BP 1 atm
10,000 elev 0.7 atm
29,029 elev (Mt. Everest) 0.3 atm
- A liquid will boil when the vapor pressure equals
the atmospheric pressure, at any T above the
triple point.
4711.6 Phase diagrams CO2
- Lines 2 phases exist in equilibrium
- Triple point all 3 phases exist together in
equilibrium (X on graph) - Critical point, or critical temperature
pressure highest T and P at which a liquid can
exist (Z on graph)
Temp (ºC)
- For most substances, inc P will cause a gas to
condense (or deposit), a liquid to freeze, and a
solid to become more dense (to a limit.)
48Phase diagrams H2O
- For H2O, inc P will cause ice to melt.
49 50 5111.7-8 Structures of solids
- Amorphous without orderly structure
- Ex rubber, glass
- Crystalline repeating structure have many
different stacking patterns based on chemical
formula, atomic or ionic sizes, and bonding
52Cubic Unit Cells in Crystalline Solids
- Primitive-cubic shared atoms are located only at
each of the corners. 1 atom per unit cell. - Body-centered cubic 1 atom in center and the
corner atoms give a net of 2 atoms per unit cell. - Face-centered cubic corner atoms plus half-atoms
in each face give 4 atoms per unit cell.
53Common Lattice Structures
54Types of Crystalline Solids
Type Particles Forces Notable properties Examples
Atomic Atoms London dispersion Poor conductors Very low MP Ar (s),Kr (s)
A small (2 cm long) piece of rapidly melting
argon ice (the liquid is flowing off at the
bottom) which has been frozen by allowing a slow
stream of the gas to flow into a small graduated
cylinder which was immersed into a cup of liquid
nitrogen
55Molecular crystals Molecules (polar or non-polar) London dispersion, dipole-dipole, H-bonds Poor conductors Low to moderate MP SO2(s) CO2 (s), C12H22O11, H2O (s)
Sucrose (liq at room T)
Ice(liq at room T)
Carbon dioxide, dry ice(g at room T)
56Covalent (a.k.a. covalent network) Atoms bonded in a covalent network Covalent bonds Very hard Very high MP Generally insoluble Variable conductivity C (diamond graphite) SiO2 (quartz) Ge, Si, SiC, BN
Diamond
Graphite
SiO2
57Ionic Anions and cations Crystals shatter! Ion-ion (ionic bonding) High Lattice Energy Hard brittle High MP,BP Poor conductors Some solubility in H2O NaCl, Ca(NO3)2
58Metallic Metal cations in a diffuse, delocalized e- cloud Metallic bonds Usually face-centered or body centered Excellent conductors Malleable Ductile High but wide range of MP Cu, Al, Fe (hard) Alloys Pb, Au, Na (soft)
59Overall
- Physical properties depend on these forces. The
stronger the forces between the particles, - (a) the higher the melting point.
- (b) the higher the boiling point.
- (c) the lower the vapor pressure (partial
pressure of vapor in equilibrium with liquid or
solid in a closed container at a fixed
temperature). - (d) the higher the viscosity (resistance to
flow). - (e) the greater the surface tension (resistance
to an increase in surface area).
60Practice
- Determine the type of solid and the forces
holding the particles together - SiO2 Covalent Network Covalent Bonds
- NaNO3 Ionic Electrostatic Att.
- C2H6 Molecular Dispersion
- CH3OH Molecular Dispersion, Dipole-Dipole,
H-Bond - C(diamond) Covalent Network Covalent Bonds
- Al Metallic Metallic
- Kr Atomic (Molecular) Dispersion
- H2O Molecular Dispersion, Dipole-Dipole, H-Bond
61Extra Material
- The following pages contain some additional
material and review items
62Examples
63Ionic Solids
- stable, high melting points
- held together by strong electrostatic forces
between oppositely charged ions - larger ions are arranged in closest packing
arrangement - smaller ions fit in the holes created by the
larger ions
64Cubic Unit Cells in Crystalline Solids
- Primitive-cubic shared atoms are located only at
each of the corners. 1 atom per unit cell. - Body-centered cubic 1 atom in center and the
corner atoms give a net of 2 atoms per unit cell. - Face-centered cubic corner atoms plus half-atoms
in each face give 4 atoms per unit cell.
65Common Lattice Structures
66Calculations involving the Unit Cell
- The density of a metal can be calculated if we
know the length of the side of a unit cell. - The radius of an metal atom can be determined if
the unit cell type and the density of the metal
known - Relationship between length of side and radius of
atom - Primitive 2r l FCC BCC
- E.g. Polonium crystallizes according to the
primitive cubic structure. Determine its density
if the atomic radius is 167 pm. - E.g.2 Calculate the radius of potassium if its
density is 0.8560 g/cm3 and it has a BCC crystal
structure.
67Figure 11.31
- Length of sides a, b, and c as well as angles a,
b, g vary to give most of the unit cells. Return
to unit cells
68Unit Cells in Crystalline Solids
- Metal crystals made up of atoms in regular arrays
the smallest of repeating array of atoms is
called the unit cell. - There are 14 different unit cells that are
observed which vary in terms of the angles
between atoms some are 90, but others are not.
Go to Figure 11.31
69Packing of Spheres and the Structures of Metals
- Arrays of atoms act as if they are spheres. Two
or more layers produce 3-D structure. - Angles between groups of atoms can be 90 or can
be in a more compact arrangement such as the
hexagonal closest pack (see below) where the
spheres form hexagons. - Two cubic arrays one directly on top of the other
produces simple cubic (primitive) structure. - Each atom has 6 nearest neighbors (coordination
number of 6) nearest neighbor is where an atom
touches another atom. - 54 of the space in a cube is used.
- Offset layers produces a-b-a-b arrangement since
it takes two layers to define arrangement of
atoms. - BCC structure an example.
- Coordination is 8.
70Packing of Spheres and the Structures of Metals
- FCC structure has a-b-c-a-b-c stacking. It takes
three layers to establish the repeating pattern
and has 4 atoms per unit cell and the
coordination number is 12.
71Metallic Crystals
- can be viewed as metals atoms (spheres) packed
together in the closest arrangement possible - closest packing- when each sphere has 12
neighbors - 6 on the same plane
- 3 in the plane above
- 3 in the plane below
72Bonding of Metals
- the highest energy level for most metal atoms
does not contain many electrons - these vacant overlapping orbitals allow outer
electrons to roam freely around the entire metal
73Bonding of Metals
- these roaming electrons
- form a sea of electrons
- around the metal atoms
- malleability and ductility
- bonding is the same in every direction
- one layer of atoms can slide past another without
friction - conductivity
- from the freedom of electrons to move around the
atoms
74Metal Alloys
- substance that is a mixture of elements and has
metallic properties - substitutional alloy
- host metal atoms are replaced by other metal
atoms - happens when they have similar sizes
- interstitial alloy
- metal atoms occupy spaces created between host
metal atoms - happens when metal atoms have large difference in
size
75Examples
- Brass
- substitutional
- 1/3 of Cu atoms replaced by Zn
- Steel
- interstitial
- Fe with C atoms in between
- makes harder and less malleable
76Chapter 11 Overview
- Changes of State
- Phase transitions
- Phase Diagrams
- Liquid State
- Properties of Liquids Surface tension and
viscosity - Intermolecular forces explaining liquid
properties - Solid State
- Classification of Solids by Type of Attraction
between Units - Crystalline solids crystal lattices and unit
cells - Structures of some crystalline solids
- Calculations Involving Unit-Cell Dimensions
- Determining the Crystal Structure by X-ray
Diffraction
Exam on Friday We will begin Chp 14 Thursday
77Comparison of Gases, Liquids and Solids
- Gases are compressible fluids. Their molecules
are widely separated. - Liquids are relatively incompressible fluids.
Their molecules are more tightly packed. - Solids are nearly incompressible and rigid. Their
molecules or ions are in close contact and do not
move.
78Phase Transitions
- Melting change of a solid to a liquid.
- Freezing change a liquid to a solid.
- Vaporization change of a solid or liquid to a
gas. Change of solid to vapor often called
sublimation. - Condensation change of a gas to a liquid or
solid. Change of a gas to a solid often called
deposition.
H2O(s) ? H2O(l) H2O(l) ? H2O(s) H2O(l) ? H2O(g)
or H2O(s) ? H2O(g) H2O(g) ? H2O(l) or H2O(g) ?
H2O(s)
79Vapor Pressure
- In a sealed container, some of a liquid
evaporates to establish a pressure in the vapor
phase. - Vapor pressure partial pressure of the vapor
over the liquid measured at equilibrium and at
some temperature. - Dynamic equilibrium
80Temperature Dependence of Vapor Pressures
- The vapor pressure above the liquid varies
exponentially with changes in the temperature. - The Clausius-Clapeyron equation shows how the
vapor pressure and temperature are related. It
can be written as
81Clausius Clapeyron Equation
- A straight line plot results when ln P vs. 1/T is
plotted and has a slope of ?Hvap/R. - Clausius Clapeyron equation is true for any two
pairs of points. - Write the equation for each and combine to get
82Using the Clausius Clapeyron Equation
- Boiling point the temperature at which the vapor
pressure of a liquid is equal to the pressure of
the external atmosphere. - Normal boiling point the temperature at which the
vapor pressure of a liquid is equal to
atmospheric pressure (1 atm).
E.g. Determine normal boiling point of chloroform
if its heat of vaporization is 31.4 kJ/mol and it
has a vapor pressure of 190.0 mmHg at
25.0C. E.g.2. The normal boiling point of
benzene is 80.1C at 26.1C it has a vapor
pressure of 100.0 mmHg. What is the heat of
vaporization?
83Energy of Heat and Phase Change
- Heat of vaporization heat needed for the
vaporization of a liquid. - H2O(l) ?H2O(g) DH 40.7 kJ
- Heat of fusion heat needed for the melting of a
solid. - H2O(s) ?H2O(l) DH 6.01 kJ
- Temperature does not change during the change
from one phase to another.
E.g. Start with a solution consisting of 50.0 g
of H2O(s) and 50.0 g of H2O(l) at 0C. Determine
the heat required to heat this mixture to 100.0C
and evaporate half of the water.
84Phase Diagrams
- Graph of pressure-temperature relationship
describes when 1,2,3 or more phases are present
and/or in equilibrium with each other. - Lines indicate equilibrium state two phases.
- Triple point- Temp. and press. where all three
phases co-exist in equilibrium. - Critical temp.- Temp. where substance must always
be gas, no matter what pressure.
- Critical pressure- vapor pressure at critical
temp. - Critical point- point where system is at its
critical pressure and temp.
85Properties of Liquids
- Surface tension the energy required to increase
the surface area of a liquid by a unit amount. - Viscosity a measure of a liquids resistance to
flow. - Surface tension The net pull toward the interior
of the liquid makes the surface tend to as small
a surface area as possible and a substance does
not penetrate it easily. - Viscosity Related to mobility of a molecule
(proportional to the size and types of
interactions in the liquid).
- Viscosity decreases as the temperature increases
since increased temperatures tend to cause
increased mobility of the molecule.
86Intermolecular Forces
- Intermolecular forces attractions and repulsions
between molecules that hold them together. - Intermolecular forces (van der Waals forces) hold
molecules together in liquid and solid phases. - Ion-dipole force interaction between an ion and
partial charges in a polar molecule. - Dipole-dipole force attractive force between
polar molecules with positive end of one molecule
is aligned with negative side of other. - London dispersion Forces interactions between
instantaneously formed electric dipoles on
neighboring polar or nonpolar molecules. - Polarizability ease with which electron cloud of
some substance can be distorted by presence of
some electric field (such as another dipolar
substance). Related to size of atom or molecule.
Small atoms and molecules less easily polarized.
87Boiling Points vs. Molecular Weight
- Hydrogen bonds the interaction between hydrogen
bound to an electronegative element (N, O, or F)
and an electron pair from another electronegative
element. Hydrogen bonding is the dominate force
holding the two DNA molecules together to form
the double helix configuration of DNA.
88Comparisonof Energies for Intermolecular Forces
Interaction Forces Approximate Energy
Intermolecular
London 1 10 kJ
Dipole-dipole 3 4 kJ
Ion-dipole 5 50 kJ
Hydrogen bonding 10 40 kJ
Chemical bonding
Ionic 100 1000 kJ
Covalent 100 1000 kJ
89Structure of Solids
- Types of solids
- Crystalline a well defined arrangement of
atoms this arrangement is often seen on a
macroscopic level. - Ionic solids ionic bonds hold the solids in a
regular three dimensional arrangement. - Molecular solid solids like ice that are held
together by intermolecular forces. - Covalent network a solid consists of atoms held
together in large networks or chains by covalent
networks. - Metallic similar to covalent network except
with metals. Provides high conductivity. - Amorphous atoms are randomly arranged. No
order exists in the solid. Example glass