Title: States of Matter: Liquids and Solids
1States of Matter Liquids and Solids
2States of Matter
- Comparison of gases, liquids, and solids. (See
Figure 11.2)
- 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.
3Changes of State
- A change of state or phase transition is a change
of a substance from one state to another.
gas
liquid
solid
4Vapor Pressure
- Liquids are continuously vaporizing.
- If a liquid is in a closed vessel with space
above it, a partial pressure of the vapor state
builds up in this space. - The vapor pressure of a liquid is the partial
pressure of the vapor over the liquid, measured
at equilibrium at a given temperature. (See
Figure 11.4)
5Vapor Pressure
- The vapor pressure of a liquid depends on its
temperature. (See Figure 11.7)
- As the temperature increases, the kinetic energy
of the molecular motion becomes greater, and
vapor pressure increases. - Liquids and solids with relatively high vapor
pressures at normal temperatures are said to be
volatile.
6Henrys Law
- Look up Henrys Law in your textbook.
- Complete the activity keeping this law in mind.
7Boiling Point
- The temperature at which the vapor pressure of a
liquid equals the pressure exerted on the liquid
is called the boiling point.
- As the temperature of a liquid increases, the
vapor pressure increases until it reaches
atmospheric pressure. - At this point, stable bubbles of vapor form
within the liquid. This is called boiling. - The normal boiling point is the boiling point at
1 atm.
8Freezing Point
- The temperature at which a pure liquid changes to
a crystalline solid, or freezes, is called the
freezing point.
- The melting point is identical to the freezing
point and is defined as the temperature at which
a solid becomes a liquid. - Unlike boiling points, melting points are
affected significantly by only large pressure
changes.
9Heat of Phase Transition
- To melt a pure substance at its melting point
requires an extra boost of energy to overcome
lattice energies.
- The heat needed to melt 1 mol of a pure substance
is called the heat of fusion and denoted DHfus.
10Heat of Phase Transition
- To boil a pure substance at its melting point
requires an extra boost of energy to overcome
intermolecular forces.
- The heat needed to boil 1 mol of a pure substance
is called the heat of vaporization and denoted
DHvap. (see Figure 11.9)
11A Problem to Consider
- The heat of vaporization of ammonia is 23.4
kJ/mol. How much heat is required to vaporize
1.00 kg of ammonia?
- First, we must determine the number of moles of
ammonia in 1.00 kg (1000 g).
12A Problem to Consider
- The heat of vaporization of ammonia is 23.4
kJ/mol. How much heat is required to vaporize
1.00 kg of ammonia?
- Then we can determine the heat required for
vaporization.
13Clausius-Clapeyron Equation
- We noted earlier that vapor pressure was a
function of temperature.
- It has been demonstrated that the logarithm of
the vapor pressure of a liquid varies linearly
with absolute temperature.
14A Problem to Consider
- Carbon disulfide, CS2, has a normal boiling point
of 46C (vapor pressure 760 mmHg) and a heat of
vaporization of 26.8 kJ/mol. What is the vapor
pressure of carbon disulfide at 35C?
- Substituting into the Clausius-Clapeyron
equation, we obtain
15A Problem to Consider
- Carbon disulfide, CS2, has a normal boiling point
of 46C (vapor pressure 760 mmHg) and a heat of
vaporization of 26.8 kJ/mol. What is the vapor
pressure of carbon disulfide at 35C?
- Taking the antiln we obtain
16Phase Diagrams
- A phase diagram is a graphical way to summarize
the conditions under which the different states
of a substance are stable.
- The diagram is divided into three areas
representing each state of the substance. - The curves separating each area represent the
boundaries of phase changes.
17Phase Diagrams
- Below is a typical phase diagram. It consists of
three curves that divide the diagram into regions
labeled solid, liquid, and gas.
.
B
C
solid
liquid
pressure
.
gas
A
D
temperature
18Phase Diagrams
- Curve AB, dividing the solid region from the
liquid region, represents the conditions under
which the solid and liquid are in equilibrium.
.
B
C
solid
liquid
pressure
.
gas
A
D
temperature
19Phase Diagrams
- Usually, the melting point is only slightly
affected by pressure. For this reason, the
melting point curve, AB, is nearly vertical.
.
B
C
solid
liquid
pressure
.
gas
A
D
temperature
20Phase Diagrams
- If a liquid is more dense than its solid, the
curve leans slightly to the left, causing the
melting point to decrease with pressure.
.
B
C
solid
liquid
pressure
.
gas
A
D
temperature
21Phase Diagrams
- If a liquid is less dense than its solid, the
curve leans slightly to the right, causing the
melting point to increase with pressure.
.
B
C
solid
liquid
pressure
.
gas
A
D
temperature
22Phase Diagrams
- Curve AC, which divides the liquid region from
the gaseous region, represents the boiling points
of the liquid for various pressures.
.
B
C
solid
liquid
pressure
.
gas
A
D
temperature
23Phase Diagrams
- Curve AD, which divides the solid region from the
gaseous region, represents the vapor pressures of
the solid at various temperatures.
.
B
C
solid
liquid
pressure
.
gas
A
D
temperature
24Phase Diagrams
- The curves intersect at A, the triple point,
which is the temperature and pressure where three
phases of a substance exist in equilibrium.
.
B
C
solid
liquid
pressure
.
gas
A
D
temperature
25Phase Diagrams
- The curves intersect at A, the triple point,
which is the temperature and pressure where three
phases of a substance exist in equilibrium.
.
B
C
solid
liquid
(see Figures 11.11 and 11.12)
pressure
.
gas
A
D
temperature
26 Rhombic v. Monoclinic
27Phase Diagrams
- The temperature above which the liquid state of a
substance no longer exists regardless of pressure
is called the critical temperature.
.
B
C
solid
liquid
pressure
.
gas
A
D
Tcrit
temperature
28Phase Diagrams
- The vapor pressure at the critical temperature is
called the critical pressure. Note that curve AC
ends at the critical point, C.
.
B
Pcrit
C
solid
liquid
(see Figure 11.13)
pressure
.
gas
A
D
Tcrit
temperature
29Properties of Liquids Surface Tension and
Viscosity
- The molecular structure of a substance defines
the intermolecular forces holding it together.
- Many physical properties of substances are
attributed to their intermolecular forces. - These properties include vapor pressure and
boiling point. - Two additional properties shown in Table 11.2 are
surface tension and viscosity.
30Properties of Liquids Surface Tension and
Viscosity
- Surface tension is the energy required to
increase the surface area of a liquid by a unit
amount.
- A molecule within a liquid is pulled in all
directions, whereas a molecule on the surface is
only pulled to the interior. (See Figure 11.16). - As a result, there is a tendency for the surface
area of the liquid to be minimized (See Figure
11.18 ).
31Properties of Liquids Surface Tension and
Viscosity
- Surface tension is the energy required to
increase the surface area of a liquid by a unit
amount.
- This explains why falling raindrops are nearly
spherical, minimizing surface area. - In comparisons of substances, as intermolecular
forces increase, the apparent surface tension
also increases. - intermolecular forces surface tension
32Intermolecular Forces Explaining Liquid
Properties
- Viscosity is the resistance to flow exhibited by
all liquids and gases.
- Viscosity can be illustrated by measuring the
time required for a steel ball to fall through a
column of the liquid. (see Figure 11.20) - Even without such measurements, you know that
syrup has a greater viscosity than water. - In comparisons of substances, as intermolecular
forces increase, viscosity usually increases. - intermolecular forces viscosity
33Intermolecular Forces Explaining Liquid
Properties
- Many of the physical properties of liquids (and
certain solids) can be explained in terms of
intermolecular forces, the forces of attraction
between molecules.
- Three types of forces are known to exist between
neutral molecules. - Dipole-dipole forces
- London (or dispersion) forces
- Hydrogen bonding
34Intermolecular Forces Explaining Liquid
Properties
- The term van der Waals forces is a general term
including dipole-dipole and London forces.
- Van der Waals forces are the weak attractive
forces in a large number of substances. - Hydrogen bonding occurs in substances containing
hydrogen atoms bonded to certain very
electronegative atoms. - Van der Waals forces 0.1 to 10 kJ/mol
- Hyderogen bonding 10 to 40 kJ/mol
35Dipole-Dipole Forces
- Polar molecules can attract one another through
dipole-dipole forces.
- The dipole-dipole force is an attractive
intermolecular force resulting from the tendency
of polar molecules to align themselves positive
end to negative end.
36London Forces
- London forces are the weak attractive forces
resulting from instantaneous dipoles that occur
due to the distortion of the electron cloud
surrounding a molecule.
- London forces increase with molecular weight. The
larger a molecule, the more easily it can be
distorted to give an instantaneous dipole. - All covalent molecules exhibit some London force.
37(No Transcript)
38Van der Waals Forces and the Properties of Liquids
- In summary, intermolecular forces play a large
role in many of the physical properties of
liquids and gases. These include
- vapor pressure
- boiling point
- surface tension
- viscosity
39Van der Waals Forces and the Properties of Liquids
- The vapor pressure of a liquid depends on
intermolecular forces. When the intermolecular
forces in a liquid are strong, you expect the
vapor pressure to be low.
- As intermolecular forces increase, vapor
pressures decrease.
40Van der Waals Forces and the Properties of Liquids
- The normal boiling point is related to vapor
pressure and is lowest for liquids with the
weakest intermolecular forces.
- When intermolecular forces are weak, little
energy is required to overcome them. - Consequently, boiling points are low for such
compounds.
41Van der Waals Forces and the Properties of Liquids
- Surface tension increases with increasing
intermolecular forces.
- Surface tension is the energy needed to reduce
the surface area of a liquid. - To increase surface area, it is necessary to pull
molecules apart against the intermolecular forces
of attraction.
42Van der Waals Forces and the Properties of Liquids
- Viscosity increases with increasing
intermolecular forces because increasing these
forces increases the resistance to flow.
- Other factors, such as the possibility of
molecules tangling together, affect viscosity. - Liquids with long molecules that tangle together
are expected to have high viscosities.
43Hydrogen Bonding
- Hydrogen bonding is a force that exists between a
hydrogen atom covalently bonded to a very
electronegative atom, X, and a lone pair of
electrons on a very electronegative atom, Y.
- To exhibit hydrogen bonding, one of the following
three structures must be present.
- Only N, O, and F are electronegative enough to
leave the hydrogen nucleus exposed.
44Hydrogen Bonding
- Molecules exhibiting hydrogen bonding have
abnormally high boiling points compared to
molecules with similar Van der Waals forces.
- For example, water has the highest boiling point
of the Group VI hydrides. (see Figure 11.24A) - Similar trends are seen in the Group V and VII
hydrides. (see Figure 11.24B)
45Hydrogen Bonding
- A hydrogen atom bonded to an electronegative atom
appears to be special.
- The electrons in the O-H bond are drawn to the O
atom, leaving the dense positive charge of the
hydrogen nucleus exposed. - Its the strong attraction of this exposed
nucleus for the lone pair on an adjacent molecule
that accounts for the strong attraction. - A similar mechanism explains the attractions in
HF and NH3.
46Hydrogen Bonding
47Solid State
- A solid is a nearly incompressible state of
matter with a well-defined shape. The units
making up the solid are in close contact and in
fixed positions.
- Solids are characterized by the type of force
holding the structural units together. - In some cases, these forces are intermolecular,
but in others they are chemical bonds (metallic,
ionic, or covalent).
48Solid State
- From this point of view, there are four types of
solids.
- Molecular (Van der Waals forces)
- Metallic (Metallic bond)
- Ionic (Ionic bond)
- Covalent (Covalent bond)
49Types of Solids
- A molecular solid is a solid that consists of
atoms or molecules held together by
intermolecular forces.
- Many solids are of this type.
- Examples include solid neon, solid water (ice),
and solid carbon dioxide (dry ice).
50Types of Solids
- A metallic solid is a solid that consists of
positive cores of atoms held together by a
surrounding sea of electrons (metallic bonding).
- In this kind of bonding, positively charged
atomic cores are surrounded by delocalized
electrons. - Examples include iron, copper, and silver.
51Types of Solids
- An ionic solid is a solid that consists of
cations and anions held together by electrical
attraction of opposite charges (ionic bond).
- Examples include cesium chloride, sodium
chloride, and zinc sulfide (but ZnS has
considerable covalent character).
52Types of Solids
- A covalent network solid is a solid that consists
of atoms held together in large networks or
chains by covalent bonds.
- Examples include carbon, in its forms as diamond
or graphite (see Figure 11.27), asbestos, and
silicon carbide. - Table 11.5 summarizes these four types of solids.
53Physical Properties
- Many physical properties of a solid can be
attributed to its structure.
- Melting Point and Structure
- For a solid to melt, the forces holding the
structural units together must be overcome. - For a molecular solid, these are weak
intermolecular attractions. - Thus, molecular solids tend to have low melting
points (below 300oC).
54Physical Properties
- Many physical properties of a solid can be
attributed to its structure.
- Melting Point and Structure
- For ionic solids and covalent network solids to
melt, chemical bonds must be broken. - For that reason, their melting points are
relatively high. - See Table 11.1.
55Physical Properties
- Many physical properties of a solid can be
attributed to its structure.
- Melting Point and Structure
- Note that for ionic solids, melting points
increase with the strength of the ionic bond. - Ionic bonds are stronger when
- The magnitude of charge is high.
- The ions are small (higher charge density).
56Physical Properties
- Many physical properties of a solid can be
attributed to its structure.
- Melting Point and Structure
- Metals often have high melting points, but there
is considerable variability. - Melting points are low for Groups IA and IIA but
increase as you move into the transition metals. - The elements in the middle of the transition
metals have the highest melting points.
57Physical Properties
- Many physical properties of a solid can be
attributed to its structure.
- Hardness depends on how easily structural units
can be moved relative to one another. - Molecular solids with weak intermolecular
attractions are rather soft compared with ionic
compounds, where forces are much stronger.
58Physical Properties
- Many physical properties of a solid can be
attributed to its structure.
- Covalent network solids are quite hard because of
the rigidity of the covalent network structure. - Diamond and silicon carbide (SiC),
three-dimensional covalent network solids, are
among the hardest substances known.
59Physical Properties
- Many physical properties of a solid can be
attributed to its structure.
- Molecular and ionic crystals are generally
brittle because they fracture easily along
crystal planes. - Metallic solids, by contrast, are malleable.
60Physical Properties
- Many physical properties of a solid can be
attributed to its structure.
- Electrical Conductivity and Structure
- Molecular and ionic solids are generally
considered nonconductors. - Ionic compounds conduct in their molten state, as
ions are then free to move. - Metals are all considered conductors.
61Physical Properties
- Many physical properties of a solid can be
attributed to its structure.
- Electrical Conductivity and Structure
- Of the covalent network solids, only graphite
conducts electricity. - This is due to the delocalization of the resonant
p electrons in graphites sp2 hybridization.
62Crystalline Solids Crystal Lattices and Unit
Cells
- Solids can be crystalline or amorphous.
- A crystalline solid is composed of one or more
crystals each crystal has a well-defined,
ordered structure in three dimensions. - Examples include sodium chloride and sucrose.
- An amorphous solid has a disordered structure. It
lacks the well-defined arrangement of basic units
found in a crystal. - Glass is an amorphous solid.
63Shapes of Crystals
64Crystal Lattices
- A crystal lattice is the geometric arrangement of
lattice points in a crystal.
- A unit cell is the smallest boxlike unit from
which you can construct a crystal by stacking the
units in three dimensions (see Figure 11.29). - There are seven basic shapes possible for unit
cells, which give rise to seven crystal systems
used to classify crystals (see Figure 11.31 and
Table 11.6).
65Crystal Lattices
- A crystal lattice is the geometric arrangement of
lattice points in a crystal.
- These seven systems can have more than one
possible crystal lattice. - A primitive lattice has lattice points only at
the corners of each cell.
66Crystal Lattices
- A crystal lattice is the geometric arrangement of
lattice points in a crystal.
- Other lattices in the same crystal may have
lattice points on the faces of the unit cell. - Following is a description of the cubic crystal
system.
67Cubic Unit Cells
- A simple cubic unit cell is a cubic cell in which
the lattice points are situated only at the
corners (see Figure 11.30).
- A body-centered cubic unit cell is one in which
there is a lattice point in the center of the
cell as well as at the corners. - A face-centered cubic unit cell is one in which
there are lattice points at the center of each
face of the cell as well as at the corners, (see
Figures 11.32 and 11.33).
68Crystal Defects
- There are principally two kinds of defects that
occur in crystalline substances.
- Chemical impurities, such as in rubies, where the
crystal is mainly aluminum oxide with an
occasional Al3 ion replaced with Cr3, which
gives a red color. - Defects in the formation of the lattice. Crystal
planes may be misaligned, or sites in the crystal
lattice may remain vacant.
69Crystal Defects
- Point Defects
- Vacancies / Schottky defects
- Interstitial
- Antisite
- Topological
- Impurity / Substitution
- Line Defects
- Planar Defects
70Point Defects
- Vacancies / Schottky defects
71Point Defects
72Point Defects
- Antisite
- Na Cl Na Cl Na Cl
- Na Cl Na Cl Na Cl
- Na Cl Cl Cl Na Cl
- Na Cl Na Cl Na Cl
- Na Cl Na Cl Na Cl
- Na Cl Na Cl Na Cl
- Atom belongs in solid, but not in that place.
73Point Defects
- Topological
- chemical bonding is topologically different from
the surroundings - Ex 6 carbons in ring ? rings of 5 and 7, shile
total number of atoms remains constant.
74Point Defects
- Impurity / Substitution
- The technique of purposefully substituting, or
doping, a solid is used to produce microchips,
lasers, and in the amplification of light signals
through fiberopitic cable, to name a few.
75Line Defects
- Dislocations are linear defects around which some
of the atoms of the crystal lattice are
misaligned. - Edge dislocations are caused by the termination
of a plane of atoms in the middle of a crystal
76Line Defects
- a line defect
- a slip of the part of crystal over an atomic
plane relative to another part - A screw dislocation results when atomic planes
form a spiral ramp winding around the line of the
dislocation
77Planar Defects
- Grain boundaries
- Usually result when one crystal grows into another
78Planar Defects
- Antiphase
- Each side of the boundary has an opposite phase
For example if the ordering is usually ABABABAB,
an anti phase boundary takes the form of
ABABBABA.
79Planar Defects
- Stacking Fault
- a one or two layer interruption in the stacking
sequence when stacking one of the layers on top
of another, the atoms are not directly on top of
one another
80Calculations Involving Unit Cell Dimensions
- X-ray diffraction is a method for determining the
structure and dimensions of a unit cell in a
crystalline compound.
- Once the dimensions and structure are known, the
volume and mass of a single atom in the crystal
can be calculated. - The determination of the mass of a single atom
gave us one of the first accurate determinations
of Avogadros number.
81Determination of Crystal Lattice by X-Ray
Diffraction
- When x-rays are reflected from the planes of a
crystal, they show a diffraction pattern that can
be recorded on photographic film (see Figure
11.47).
- Analysis of these diffraction patterns allows the
determination of the positions of the atoms in
the unit cell of the solid.
82Sample Calculation
- Silver crystals are face-centered cubics, with a
cell edges of 4.086 angstroms. - What is the distance between center of the two
closest Ag atoms? - What is the atomic radius of silver in this
crystal? - How many nearest neighbors does each atom have?
83What is the distance between center of the two
closest Ag atoms?
- The hypotenuse is equal to twice the center to
center distance. - c 2(4.086)
- a b 4.086 angstroms
- (4.086)2 (4.086)2 c2
- c2 v33.3908
- c 5.7785 angstroms
- d c/2
- d 2.889 angstroms
c
a
b
4.086 angstroms
- Atoms are assumed to touch along face diagonals
84What is the atomic radius of silver in this
crystal?
- r d 2.889 angstroms
- 2 2
- d 1.445 angstroms
c
a
b
4.086 angstroms
- The hypotenuse is of the unit cell face is four
times the radius of the atom.
85How many nearest neighbors does each atom have?
- The central atom has
- 4 nearest neighbors in the x-y plane
- 4 nearest neighbors in the x-z plane, and
- 4 nearest neighbors in the y-z plane
- a total of 12 nearest neighbors.
86Sample Calculation
- Nickel has a face-centered unit cell with an edge
length of 352.4 pm. The density of nickel is
8.91 g/cm3. - From the atomic weight, calculate Avagadros
number.
87What is the mass of a single nickel atom?
c
a
b
352.4 pm
- pm to cm
- 352.4pm 1 x 10-12 m 100 cm
- 1 pm 1m
- 3.524 x 10-8 cm
88What is the mass of a single nickel atom?
- Change pm to cm
- Use the density to calculate mass
- 3.524 x 10-8 cm
- D m
- v
- 8.91 g m
- cm3 (3.524 x 10-8cm )3
- m 3.90 x 10-22 g/ unit cell
89- Determine atoms per unit cell to calculate g /
atom
- 8 vertices 1/8 atom 1 atom
- vertex
- 6 faces 1/2 atom 3 atoms
- face
- 4 total atoms
- 3.90 x 10-22 g/ unit cell
- 4 atoms
- 9.75 x 10-23 g / atom
90- Use atomic weight to calculate Avagadros number
- Atomic wt of Nickel is 59 g/ mol
- 59 g / mol
- 9.75 x 10-23 g / atom
- 6.05 x 1023 atoms / mol
91Operational Skills
- Calculating the heat required for a phase change
of a given mass of substance. - Calculating vapor pressures and heats of
vaporization. - Relating the conditions for the liquification of
a gas to the critical temperature. - Identifying intermolecular forces.
- Determining relative vapor pressure on the basis
of intermolecular attraction. - Identifying types of solids.
92Operational Skills
- Determining the relative melting points based on
types of solids. - Determining the number of atoms per unit cell.
- Calculating atomic mass from unit-cell dimension
and density. - Calculating unit-cell dimensions from unit-cell
type and density.
93Figure 11.2 Representation of the States of
Matter
Return to Slide 2
94Figure 11.4 Measurement of the vapor pressure
of water.
Return to Slide 4
95Figure 11.7 Variation of vapor pressure with
temperature.
Return to Slide 5
96Figure 11.9 Heating curve for water.
Return to Slide 9
97Figure 11.11 Phase diagram for water (not to
scale).
Return to Slide 24
98Figure 11.12 Phase diagrams for carbon dioxide
and sulfur (not to scale).
Return to Slide 24
99Figure 11.13 Observing the critical phenomenon.
Return to Slide 26
100Figure 11.16 Explaining Surface Tension
Return to Slide 28
101Figure 11.18 Demonstration of Surface Tension of
Water
Return to Slide 28
102Figure 11.20Comparison of the viscosities of
two liquids. Photo courtesy of James Scherer.
Return to Slide 30
103Figure 11.24 Boiling point versus molecular
weight for hydrides.
Return to Slide 41
104Figure 11.24 Boiling point versus molecular
weight for hydrides.
Return to Slide 41
105Figure 11.27 Structures of diamond and graphite.
Return to Slide 49
106Return to Slide 51
107Figure 11.29 A two-dimensional pattern.
Return to Slide 60
108Figure 11.31 Unit-cell shapes of the different
crystal systems.
Return to Slide 60
109Return to Slide 60
110Figure 11.30 Crystal structure and crystal
lattice of copper.
Return to Slide 63
111Figure 11.32 Cubic unit cells.
Return to Slide 63
112Figure 11.33 Space-filling representation of
cubic unit cells.
Return to Slide 63
113Figure 11.47 A crystal diffraction pattern.From
Preston, Proceedings of the Royal Society, A,
Volume 172, plate 4, figure 5A
Return to Slide 66