John A. Schreifels

1
Chapter 11

• States of Matter Liquids and Solids

2
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

3
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.

4
Phase 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)
5
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

6
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

7
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.
• Write the equation for each and combine to get

8
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
25.0C. 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?
9
Energy 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.
10
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.

11
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 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.

12
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.

13
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.

14
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
15
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 exists in the solid.

16
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

17
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.

18
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.

19
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.

20
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 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.

21
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