Transcript Chapter 11 - Richsingiser.com

Daniel L. Reger
Scott R. Goode
David W. Ball
http://academic.cengage.com/chemistry/reger
Chapter 11
Liquids and Solids
Characteristic Properties of Gases, Liquids,
and Solids
• Intermolecular forces are the
attractions that hold molecules
together in the liquid and solid states.
State
Volume Shape of State
Density
Compressibility
gas
assumes shape and
volume of container
low
easily
compressed
liquid
definite volume, assumes
shape of container
high
nearly
incompressible
solid
both definite shape and
volume
high
nearly
incompressible
Kinetic Molecular Theory
• What was kinetic molecular theory?
• Intermolecular forces:
• Do not change with temperature
• However, kinetic energy does
• So…
Physical State
Relation Between Energy of Attraction
and Kinetic Energy of Molecules
solid
kinetic energy << energy of attraction
liquid
kinetic energy ≈ energy of attraction
gas
kinetic energy >> energy of attraction
Example
• At room temperature, chlorine is a gas, bromine is a liquid,
and iodine is a solid. Arrange the molecules in order of
increasing intermolecular forces.
Phase Changes
• Intermolecular forces determine phase, but
temperature and pressure can influence phase too.
• Definitions:
• Evaporation is the process by which molecules
escape from the liquid to the gas phase.
• Condensation is the process by which
molecules go from the gas phase to the liquid
phase.
Vapor Pressure
• The vapor pressure is the partial
pressure of the gas when the rate of
evaporation equals the rate of
condensation.
Dynamic Equilibrium
• A state of dynamic equilibrium is one in
which the two opposing changes occur at
equal rates, so no net change is apparent.
At constant
temperature
Factors that Affect Vapor Pressure
• As temperature increases, the vapor
pressure of a liquid increases.
• The stronger the intermolecular
forces, the lower the vapor pressure
of the liquid at any temperature.
Vapor Pressure Curves
Which
substance
has the
weakest
intermolecular
forces?
(a) diethyl ether (b) ethanol (c) water
Boiling Point
• The boiling point of a liquid is the
temperature at which the vapor pressure
is equal to the external pressure.
• The normal boiling point of a liquid is the
temperature at which its equilibrium vapor
pressure is equal to 1 atmosphere.
• At the boiling point, bubbles filled with
vapor form below the surface of the liquid.
Other Vaporization Properties
• Enthalpy of vaporization (DHvap) is the
enthalpy change that accompanies the
conversion of one mole of a substance from a
liquid to a gas at constant temperature.
• The critical temperature is the maximum
temperature at which a substance can exist in
the liquid state.
• The critical pressure is the minimum pressure
needed to maintain the liquid state up to the
critical temperature.
Vaporization and Intermolecular Forces
• As the strength of intermolecular
forces increase:
•
•
•
•
vapor pressure of the liquid decreases;
boiling point increases;
enthalpy of vaporization increases;
critical temperature increases.
Example Calculation
• Butane boils at a temperature of -0.6°C and has a
ΔHvap = 22.3 kJ/mol. How much energy is necessary to
boil 150 g of butane?
Liquid-Solid Equilibrium
• The changes of a substance from liquid
to solid (freezing) and from solid to liquid
(melting or fusion) are also opposing
changes that involve a dynamic
equilibrium.
Definitions
• The melting point of a substance is the
temperature at which the solid and liquid
phases are in equilibrium when the
pressure is one atmosphere.
• There is very little effect of pressure on the
melting point of a solid.
• The enthalpy of fusion (DHfus) is the
enthalpy change that accompanies the
change of one mole of solid into liquid at
constant temperature.
Heating Curves
• A heating curve is a graph of
temperature of a sample versus heat
added.
Only One Phase is Present
• When only one phase is present (up to A,
B to C, D and after), then q = mCsDT.
Phase Transitions
• During a phase transition (A to B, C to D),
the temperature remains constant and q
= DH of the transition.
Example: Heating Curve
• In the heating curve below, identify
the phase transition between A and
B.
Example: Heating Curve
• From the heating curve below, determine
which phase (solid, liquid, or gas) has the
largest specific heat.
Solid-Gas Equilibrium
• Sublimation is the direct conversion of a
substance from the solid to the gas phase.
• Deposition is the reverse of the sublimation
process.
• Enthalpy of sublimation (DHsub) is the
enthalpy change for the conversion of one
mole of substance from solid to gas.
• DHsub = DHfus + DHvap
Enthalpy Diagram for Phase Changes
A Phase Diagram
• A phase diagram is a graph of pressure
versus temperature that shows the region
of stability for each phase.
Triple Point
• There is a unique combination of
pressure and temperature, called the
triple point (T), at which all three phases
(solid, liquid, gas) are at equilibrium.
Melting Point and Pressure
• The melting point of a substance
changes very little with pressure.
• The effect of pressure on the melting
point of a substance depends on the
relative density of the two phases.
Melting Point and Pressure
• If the solid is denser
than the liquid (which
is the more common
case), the melting point
increases with
increasing pressure.
• If the liquid is denser
than the solid (as in
H2O), the melting point
decreases with
pressure.
Intermolecular Attractions
• Electrostatic forces account for all
types of intermolecular attractions.
There are three types of attractions:
• Dipole-dipole attractions
• London dispersion forces
• Hydrogen bonding
Dipole-Dipole Attractions
• Dipole-dipole attractions result from
electronic forces between molecular
dipoles:
London Dispersion Forces
• London dispersion forces arise from the
attractions between instantaneous dipoles
and induced dipoles.
Dispersion Forces and Periodic Trends
• Polarizability is the ease with which a
charge distorts the electron cloud in a
molecule.
• Polarizability generally increases with the
number of electrons in the molecule.
• For related series of molecules,
London dispersion forces increase
going down any group in the periodic
table.
Boiling Points of Some Nonpolar Substances
Substance
Molar Mass
Boiling Point (C)
CH4
SiH4
GeH4
SnH4
16
32
77
123
-184
-112
-90
-52
F2
Cl2
Br2
I2
38
71
160
254
-188
-35
59
184
Hydrogen Bonding
• The unexpectedly high boiling points of
water, ammonia, and hydrogen fluoride
requires another kind of intermolecular
force.
Hydrogen Bonding
• Hydrogen bonding
occurs between a
hydrogen atom bonded
to N, O, or F, and a lone
pair of electrons on a
second N, O, or F.
• Hydrogen bonds are
sometimes shown as
dotted lines.
Structure of Solid Water
• Hydrogen bonding
causes ice to have
a lower density than
liquid water.
Example: Intermolecular Forces
• Identify the kind of intermolecular
forces, and predict which substance in
each pair has the stronger forces of
attraction.
(a) BF3, BBr3 (b) C2H5OH, C2H5Cl
Liquids: Surface Tension
• Surface tension is the energy needed
to increase the surface area of a liquid.
• Surface tension results from intermolecular
interactions.
Liquids: Capillary Action
• Capillary action causes water to rise in
a small diameter glass tube.
• Capillary action is the result of a
competition between:
• cohesion: the attraction of molecules for
other molecules of the same substance.
• adhesion: the attraction of molecules for
other molecules of a different substance.
Capillary Action
• Water rises because adhesion is stronger
than cohesion.
• Mercury is lowered because cohesion is
stronger than adhesion.
Liquids: Viscosity
• Viscosity is the resistance of a fluid to
flow.
• The stronger the intermolecular forces of
attraction, the greater the viscosity.
• Other factors contribute to viscosity as well,
like structure, size, and shape of
molecules.
Solids
• A crystalline solid: the units that
make up the solid are arranged in a
very regular, repeating pattern.
• Ionic compounds, metals, and solids of
small molecules are usually crystalline.
• An amorphous solid lacks the long
range order of a crystalline solid.
• Most plastics are amorphous solids.
(they are polymers)
Crystalline Solids
• Crystalline solids can be classified by
the nature of the forces that hold the
units together in a regular
arrangement.
• These forces are usually referred to
as crystal forces.
Molecular Solids
• Molecular solids consist of atoms or
small molecules held together by van
der Waals forces and/or hydrogen
bonding.
• Because these crystal forces are fairly
weak, molecular solids are generally soft
and low-melting.
• Examples are CO, Ar, I2, and most
organic molecules.
Covalent Network Solids
• In a covalent network solid, all of the
atoms in a crystal are held together by
covalent bonds.
• Solids of this kind are high melting and
often very hard because strong covalent
bonds hold the atoms together.
• Some examples of covalent network
solids are diamond (C), boron nitride
(BN), and silicon dioxide (SiO2).
Allotropes
• Allotropes are two or more molecular
or crystalline forms of an element in
the same physical state.
• O2 (oxygen) and O3 (ozone) are
examples of gas-phase allotropes.
• Many elements have two or more
allotropes in the solid phase: C, S, P,
Sn, among others.
Allotropes of Carbon
• Graphite and diamond are allotropes
of carbon that have different covalent
network structures.
Ionic Solids
• An ionic solid consists of oppositely
charged ions, held together by strong
electrostatic interactions.
• Ionic solids are high melting and usually
brittle – they tend to shatter under
impact.
• Binary compounds made up of a metal
and a nonmetal are in this category.
Metallic Solids
• Metallic solids are formed from metal
atoms, and are characterized by high
thermal and electrical conductivity,
metallic luster, and malleability.
• A special kind of bonding, metallic
bonding, is needed to account for
these unique properties.
Metallic Bonding
• The electron sea model for metallic
bonding views the solid as metal ions
in a “sea” of electrons formed from the
valence shell electrons.
• The electrons are very mobile and
adequately account for the conductivity
and malleability of metals.
• Another model for metallic bonding will
be discussed in Chapter 20.
Properties of Solids - Summary
Type of Solid Molecular
Covalent
Network
Ionic
Metallic
Structural
unit
Atoms or
molecules
Atoms
Ions
Atoms
Attractive
forces
Intermolecular
forces
Localized
covalent bonds
Ionic bonds
Metallic bonds
(delocalized)
Melting
points
Low melting,
often gases or
liquids at room
temperature
High melting
High melting
Variable, from
low to very high
Character
Soft
Hard and brittle
Brittle
Malleable
Electrical
conductivity
Poor
Variable,
depending on
structure
Poor in solid,
but good when
molten
Very high
The Bragg Equation
• The distances between layers of atoms
in a crystal, as measured by x ray
diffraction, are given by the Bragg
equation:
nl  2d sin q
where l = wavelength of x rays, d =
distance between layers of atoms, q =
angle of x ray diffraction, and n is a
whole number called the order.
Example: The Bragg Equation
• X rays (l = 154 pm) are diffracted by a
crystal at an angle of 18.5. Assuming n
= 1, calculate the distance between the
layers of atoms that cause this
diffraction.
Crystal Structure
• The arrangement of the units (atoms or
molecules) is described by the unit cell
– a small regular geometric figure that
defines the repeating pattern in the
crystal.
• The location of every particle in the crystal
can be determined from the size and shape
of the unit cell.
The Unit Cell
Single unit cell
Crystal lattice
The Unit Cell
• Each unit cell is defined by the length of
the edges (a, b, and c) and the angles
between them (a, b, and g).
The Cubic Unit Cells
• For simplicity, only the three cubit unit
cells are considered (a = b = c, a = b = g
= 90).
Crystalline Solids
There are several types of basic arrangements
in crystals, like the ones depicted above.
Example: Density from Crystal Data
• Calculate the density of nickel, which
crystallizes in a face-centered cubic cell
with an edge length of 351 pm.
Closest Packing Structures
• Closest packing is the arrangement of
spheres in the most efficient manner,
and results in the smallest empty space.
• There are two closest packing
arrangements, called hexagonal close
packing (HCP) and cubic close packing
(CCP).
• In both of these arrangements, each atom
has twelve nearest neighbors.
Close Packing Structures
Ionic Crystal Structures
• Cations and anions alternate in ionic
crystals, to maximize the attractive
interactions and minimize the repulsions.
• In an ionic crystal lattice, the composition of
the unit cell must correspond to the formula
of the compound. For example, in NaCl the
ratio of cations to anions is 1:1; in CaBr2,
the ratio of cations to anions is 1:2.
Ionic Unit Cell of NaCl
• NaCl has an FCC arrangement of Clions. What about the Na+?
Example: Ionic Unit Cell
• Calcium fluoride has a unit cell with a
face-centered cubic arrangement of the
Ca2+ ions. How many F- ions are
present in the unit cell?