Chapter 11
States of Matter; Liquids and Solids
8–1
John A. Schreifels
Chemistry 211
Chapter 11-1
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
John A. Schreifels
Chemistry 211
Chapter 11-2
8–2
Comparison 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.
8–3
Figure 11.2 States of Matter
John A. Schreifels
Chemistry 211
Chapter 11-3
Phase Transitions
• Melting: change of a solid to a liquid. H2O(s)  H2O(l)
• Freezing: change a liquid to a solid. H2O(l)  H2O(s)
• Vaporization: change of a solid or
liquid to a gas. Change of solid to
vapor often called sublimation.
H2O(l)  H2O(g)
or
H2O(s)  H2O(g)
• Condensation: change of a gas to a H2O(g)  H2O(l)
liquid or solid. Change of a gas to a or
H2O(g)  H2O(s)
solid often called deposition.
8–4
John A. Schreifels
Chemistry 211
Chapter 11-4
Vapor 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
8–5
John A. Schreifels
Chemistry 211
Chapter 11-5
Temperature 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:
Hvap
1
ln P  
 C
RT
T
8–6
John A. Schreifels
Chemistry 211
Chapter 11-6
Clausius – 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.
Hvap 1
ln P  
 C
RT
T
Write the equation for each and
combine to get:
Hvap  1
P
1
ln 2 
   
P1
RT
 T1 T2 
8–7
John A. Schreifels
Chemistry 211
Chapter 11-7
Using 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 25.0°C.
E.g.2. The normal boiling point of benzene is 80.1°C; at 26.1°C it
has a vapor pressure of 100.0 mmHg. What is the heat of
vaporization?
John A. Schreifels
Chemistry 211
Chapter 11-8
8–8
Energy of Heat and Phase Change
• Heat of vaporization: heat
needed for the vaporization of a
liquid.
H2O(l) H2O(g) H = 40.7 kJ
• Heat of fusion: heat needed
for the melting of a solid.
H2O(s) H2O(l) H = 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 0°C. Determine the heat required to heat this mixture to
8–9
100.0°C and evaporate half of the water.
John A. Schreifels
Chemistry 211
Chapter 11-9
Phase 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.
John A. Schreifels
Chemistry 211
8–10
Chapter 11-10
Properties of Liquids
• Surface tension: the energy
required to increase the surface area
of a liquid by a unit amount.
• Viscosity: a measure of a liquid’s
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.
John A. Schreifels
Chemistry 211
8–11
Chapter 11-11
Intermolecular 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.
8–12
John A. Schreifels
Chemistry 211
Chapter 11-12
Boiling 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.
8–13
John A. Schreifels
Chemistry 211
Chapter 11-13
Comparisonof 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
8–14
John A. Schreifels
Chemistry 211
Chapter 11-14
Structure 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 exists8–15
in the solid.
John A. Schreifels
Chemistry 211
Chapter 11-15
Unit 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
8–16
John A. Schreifels
Chemistry 211
Chapter 11-16
Packing 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.
8–17
John A. Schreifels
Chemistry 211
Chapter 11-17
Packing 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.
8–18
John A. Schreifels
Chemistry 211
Chapter 11-18
Cubic 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.
8–19
John A. Schreifels
Chemistry 211
Chapter 11-19
Calculations 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:
r
2
l
4
BCC
r
3
l
4
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.
8–20
John A. Schreifels
Chemistry 211
Chapter 11-20
Figure 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
8–21
John A. Schreifels
Chemistry 211
Chapter 11-21