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Chapter 12: Intermolecular Attractions and the Properties of Liquids and solids

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Title: Chapter 12: Intermolecular Attractions and the Properties of Liquids and solids


1
Chapter 12 Intermolecular Attractions and the
Properties of Liquids and solids
  • There are important differences between gases,
    solids, and liquids
  • Gases - expand to fill their container
  • Liquids - retain volume, but not shape
  • Solids retain volume and shape

2

Properties can be understood in terms of how
tightly the molecules are packed together and the
strength of the intermolecular attractions
between them.
3
  • Intermolecular forces are the attractions between
    molecules
  • Intramolecular forces are the chemical bonds
    within the molecule
  • Intramolecular forces are always stronger than
    intermolecular forces
  • Intermolecular forces control the physical
    properties of the substance

4
  • There are only a few important types of
    intermolecular forces

Strong intramolecular attractions exist between H
and Cl within HCl molecules. These attractions
control the chemical properties of HCl. Weaker
intermolecular attractions exist between
neighboring HCl molecules. Intermolecular
attractions control the physical properties of
this substance.
5
  • Dipole-dipole attractions
  • Polar molecules tend to align their partial
    charges
  • The attractive force is about 1 of a covalent
    bond and drops off as 1/d3 (ddistance between
    dipoles)

The net interaction of the imperfectly aligned
molecules is attractive.
6
  • Hydrogen bonds
  • Very strong dipole-dipole attraction that occur
    when H is covalently bonded to to a small, highly
    electronegative atom (usually F, O, or N)
  • Typically about ten times stronger than other
    dipole-dipole attractions
  • Are responsible for the expansion of water as it
    freezes

7

(a) Polar water molecule. (b) Hydrogen bonding
produces strong attractions in the liquid. (c)
Hydrogen bonding (dotted lines) between water
molecules in ice form a tetrahedral configuration.
8
  • London forces
  • The (very) weak attractions between nonpolar
    molecules
  • Arise from the interactions of instantaneous
    dipoles on neighboring molecules

An instantaneous dipole on one molecule can
produce and induced dipole on another. The net
interaction of these over time is attractive.
9
  • These instantaneous dipole-induced dipole
    attractions are called London dispersion forces,
    London forces, or dispersion forces
  • London forces decrease as 1/d6 (ddistance
    between molecules)
  • Strength depends on three factors
  • Polarizability is a measure of the ease with
    which the electron cloud on a particle is
    distorted
  • It tends to increase as the electron cloud volume
    increases

10
Large electron clouds are more easily deformed
than small ones. The magnitude of the resulting
partial charge is also larger. The larger
molecules experience larger London forces than
small molecules.
  • The boiling point of the halogens and noble gases
    demonstrate this

11
  • London forces depend on the number of atoms in
    the molecule
  • The boiling point of hydrocarbons demonstrates
    this trend

12
  • Hexane, C6H14, (right) has a BP of 68.7oC while
    the BP propane, C3H8, (left) is 42.1oC because
    hexane has more sites (marked with ) along its
    chain where attraction to other molecules can
    occur.

13
  • Molecular shape affects the strength of London
    forces
  • More compact molecules tend to have lower London
    forces than longer chain-like molecules
  • For example the more compact neopentane molecule
    (CH3)4C has a lower boiling point than n-pentane,
    CH3CH2CH2CH2CH3
  • Presumably this is because the hydrogens on
    neopentane cannot interact as well as those on
    n-pentane with neighboring molecules

14
  • Space filling models of two molecules with
    formula C5H12. The H atoms in the more compact
    neopentane cannot interact as well with
    neighboring molecules as the H atoms in the more
    chain-like n-pentane.

15
  • Ion-dipole and ion-induced dipole attractions are
    the attractions between an ion and the dipole or
    induced dipole of neighboring molecules

(a) The negative ends of water dipoles surround a
cation. (b) The positive ends of water dipoles
surround an anion. The attractions can be quite
strong because the ions have full charges.
16
  • It is sometimes possible to predict physical
    properties (like BP and MP) by comparing the
    strengths of intermolecular attractions

Ion-dipole attractions hold water molecules in a
hydrate. Water molecules are found at the
vertices of an octahedron around the aluminum ion
in AlCl36H2O.
17
  • Dipole-dipole occur between molecules with
    permanent dipoles about 1 - 5 of a covalent
    bond.
  • Hydrogen bonding occur when molecules contain
    N-H and O-H bonds about 5 to 10 of a covalent
    bond.
  • London dispersion present in all substances are
    weak, but can lead to large net attractions.
  • Ion-dipole occur when ions interact with polar
    molecules can lead to large net attractions.
  • Ion-induced dipole occur when an ion induces a
    dipole on neighboring particle depend on ion
    charge and the polarizability of its neighbor

18
  • Intermolecular attractions determine how tightly
    liquids and solids pack
  • Two important properties that depend on packing
    are compressibility and diffusion
  • Compressibility is a measure of the ability of a
    substance to be forced into a smaller volume
  • Solids and liquids are nearly incompressible
    because they contain very little space between
    particles

19
  • Diffusion occurs more rapidly in gases than in
    liquids and solids

Diffusion in a gas (a) and liquid (b). Gas
molecules move a much greater distance than
liquid molecules between collisions. As a result,
diffusion occurs more rapidly in the gas.
20
  • The strength of intermolecular attractions
    determine many physical properties
  • Volume and shape
  • Attractions in gases are not strong enough to
    retain either volume or shape
  • Attractions in liquids and solids are strong
    enough so they retain their volume
  • Attractions in solids are stronger than for
    liquids so that solids also retain shape
  • Surface tension is the tendency of a liquid to
    take a shape with minimum surface area

21
  • Molecules at the surface have higher potential
    energy than those in the bulk of the liquid
  • The surface tension of a liquid is proportional
    to the energy needed to expand its surface area
  • In general, liquids with strong intermolecular
    attractions have large surface tensions

Surface tension holds moist particles of sand
together. Separation is resisted because the
surface area of the water would increase.
22
  • Wetting is the spreading of a liquid across a
    surface to form a thin film
  • For wetting to occur, the intermolecular
    attractive force between the surface and the
    liquid must be about as strong as within the
    liquid itself
  • Surfactants are added to detergents to lower the
    surface tension of water
  • The wetter water can then gets better access to
    the surface to be cleaned

23
  • Viscosity is the resistance to changing the form
    of a sample
  • Gases have viscosity, but respond almost
    instantly to form-changing forces
  • Solids, such as rocks, normally yield to forces
    acting to change their shape very slowly
  • Liquids are what most people associate with
    viscosity
  • Viscosity is also called internal friction
    because it depends on intermolecular attractions
    and molecular shape

24

Acetone is a polar molecule and experiences
dipole-dipole and London forces. Ethylene glycol,
which also has ten atoms, also participates is
hydrogen-bonding. The viscosity of ethylene
glycol is larger than the viscosity of acetone.
25
  • A change in state is called a phase change
  • Evaporation is the change in state from liquid to
    gas
  • Sublimation is the change from solid to gas
  • Both deal with the motion of molecules
  • You have also probably noticed that the
    evaporation of liquids produce a cooling effect

26
Molecules that are able to escape from the liquid
have kinetic energies larger than the average.
When they leave, the average kinetic energy of
the remaining molecules is less, so the
temperature is lower.
  • The rate of evaporation depends on the
    temperature, surface area, and strength of the
    intermolecular attractions

27
At higher temperature, the total fraction of
molecules with kinetic energy large enough to
escape is larger so the rate of evaporation is
larger.
  • For a given liquid, the rate of evaporation per
    unit surface area is greater at a higher
    temperature

28

Kinetic energy distribution in two different
liquids, A and B, at the same temperature. The
minimum kinetic energy required by molecules A to
escape is less than for B because the
intermolecular attractions in A are weaker than
in B. This causes A to evaporate faster than B.
29
  • As soon as a liquid is placed in an empty
    container, it begins to evaporate
  • Once in the gas phase, molecules can condense by
    striking the surface of the liquid and giving up
    some kinetic energy
  • The rate of evaporation equals the rate of
    condensation at equilibrium
  • This can occur in a system where the molecules
    are constrained to remain close to the liquid
    surface

30

(a) The liquid begins to evaporate in the closed
container. (b) Dynamic equilibrium is reached
when the rate of evaporation and condensation are
equal.
31
  • Similar equilibria are reached in melting and
    sublimation

At the melting point a solid begins to change
into a liquid as heat is added. As long no heat
is added or removed melting (red arrows) and
freezing (black arrows) occur at the same rate an
the number of particles in the solid remains
constant.
32

At equilibrium, molecules evaporate from the
solid at the same rate as molecules condense from
the vapor.
33
  • When molecules evaporate, the molecules that
    enter the vapor phase exert a pressure called the
    vapor pressure
  • The equilibrium vapor pressure is the vapor
    pressure once dynamic equilibrium has been
    reached
  • The equilibrium vapor pressure is usually
    referred to as simply the vapor pressure
  • Vapor pressures can be measured using a manometer

34
  • Measuring the (equilibrium) vapor pressure of a
    liquid at a specific temperature

35

Variation of vapor pressure with temperature.
Ether is said to be volatile because it has a
high vapor pressure near room temperature.
36
  • Volume changes can effect vapor pressure

(a) Equilibrium exists between liquid and vapor.
(b) The volume is increased, the pressure drops,
and the rate of condensation drops. (c) Once more
liquid evaporates, equilibrium is re-established
and the vapor pressure returns to its initial
value.
37
  • Solids also have vapor pressures
  • At a given temperature, some of the particles at
    the solid have enough kinetic energy and escape
    into the vapor phase
  • When vapor particle collide with the surface,
    they can be captured
  • The pressure of the vapor that is in equilibrium
    with the solid is called the equilibrium vapor
    pressure of the solid

38
  • The boiling point of a liquid can be defined as
    the temperature at which the vapor pressure of
    the liquid is equal to the prevailing atmospheric
    pressure
  • The normal boiling point is the temperature at
    which the vapor pressure is 1 atm
  • Molecules with higher intermolecular forces have
    higher boiling points

39

Boiling points of the hydrogen compounds of
elements of Groups IVA, VA, VIA, and VIIA of the
periodic table. The boiling points of molecules
with hydrogen bonding are higher that expected.
40
  • Heating and cooling curves can be used to
    determine melting and boiling points

(a) A heating curve observed when heat is added
at a constant rate. (b) A cooling curve observed
when heat is removed at a constant rate. The
flat regions of the curves identify the melting
and boiling points. Supercooling is seen hear as
the temperature of the liquid dips below its
freezing point.
41
  • The energy associated with the phase changes can
    be expressed per mole
  • The molar heat of fusion is the heat absorbed by
    one mole of solid when it melts to give a liquid
    at the same temperature and pressure
  • The molar heat of vaporization is the heat
    absorbed when one mole of the liquid is changed
    to one mole of vapor at constant temperature and
    pressure

42
  • The molar heat of sublimation is the heat
    absorbed by one mole of a solid when it sublimes
    to give one mole of vapor at constant temperature
    and pressure
  • All of these quantities tend to increase with
    increasing intermolecular forces
  • The concept of equilibrium is important and will
    be encountered again
  • Equilibria are often disturbed or upset

43
  • According to Le Chateliers Principle
  • When a dynamic equilibrium in a system is upset
    by a disturbance, the system responds in a
    direction that tends to counteract the
    disturbance and, if possible, restore equilibrium
  • The term position of equilibrium is used to refer
    to the relative amounts of the substance on each
    side of the double (equilibrium) arrows
  • Consider the liquid vapor equilibrium

44
  • Increasing the temperature increases the amount
    of vapor and decreases the amount of liquid
  • We say that the equilibrium has shifted
  • This can also be referred to as a right shift
    because more vapor is produced at the expense of
    the liquid
  • Temperature-pressure relationships can be
    represented using a phase diagram

45

The phase diagram of water. The line AB is the
vapor pressure curve for ice BD the vapor
pressure curve for liquid water BC the melting
point line point B the triple point (the
temperature where all three phases are in
equilibrium) and point D labels the critical
point (and the critical temperature and
pressure). Above the critical temperature a
distinct liquid phase does not exist, regardless
of pressure.
46
  • A substance that has a temperature above its
    critical temperature and a density near its
    liquid density is called a supercritical fluid
  • Supercritical fluids have some unique properties
    that make them excellent solvents
  • The values of the critical temperature tends to
    increase with increased intermolecular
    attractions between the particles
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