Transcript Liquids & Solids Chapter 13 Heat of Fusion/Vaporization H2O(s) ----> H2O(l) Hfusion = 6.02 kj/mol H2O(l) ----> H2O(g)  Hvaporization = 40.6 kj/mol From.

Liquids & Solids
Chapter 13
Heat of Fusion/Vaporization
H2O(s) ----> H2O(l) Hfusion = 6.02 kj/mol
H2O(l) ----> H2O(g)  Hvaporization = 40.6 kj/mol
From the  Ho values above, which two states
are most similar?
How do the attractive forces between the
molecules compare in these two states to the
third state?
Liquids & Solids
The liquid and solid states are considered
to be the condensed states.
Liquids & solids have much higher
densities than gases and are not
compressible.
Three States of Matter
Some Properties of a Liquid
Surface Tension: The resistance to an
increase in its surface area (polar molecules).
A sphere has the maximum volume for the
minimum surface area.
Some Properties of a Liquid
Capillary Action: Spontaneous rising of a liquid
in a narrow tube.
Viscosity: Resistance to flow (molecules with
large intermolecular forces).
Some Properties of a Liquid
Cohesive forces exist between molecules of a
liquid. Adhesive forces exist between the
liquid and its container.
Water & Its Phase Changes
Water is the most common liquid--covering
about 70% of the earth’s surface.
Water is necessary for life, moderates the
earth’s climate, a means of transportation,
and cools many industrial processes.
Pure water is colorless, odorless, tasteless,
freezes at 0 oC and boils at 100 oC at 1
atm.
Normal Melting/Freezing Point
The temperature at which a solid melts or freezes
at 1 atm pressure--0 oC for water.
Molecules break loose from lattice points and
solid changes to liquid. (Temperature is
constant as melting occurs.)
vapor pressure of solid = vapor pressure of liquid
Normal Boiling Point
The boiling temperature of a liquid at one
atmosphere pressure--100 oC for water.
Constant temperature when added energy is used
to vaporize the liquid.
vapor pressure of liquid = pressure of
surrounding atmosphere
Figure 13.3:
Both liquid
water and
gaseous water
contain H2O
molecules
Boiling Point
What effect does altitude have on boiling
point? Higher altitude--lower b.p.
Where on earth would have the highest
boiling point? Dead Sea
Where on earth would have the lowest boiling
oC
Top
of
Mt.
Everest--70
point?
Freezing of Water
When water freezes, it expands about 1/9th in
volume. This causes water pipes and
engine blocks to burst when frozen.
The density of ice is less than water and,
therefore, ice floats. If ice were more
dense than water, lakes and rivers would
freeze from the bottom up and aquatic life
could not survive.
Physical Changes & Energy
Changes
Endothermic
Exothermic
Melting
Condensing
Boiling
Freezing
Changes of state (melting, freezing, boiling,
& condensing) are constant temperature
processes!!!!!
Heats of Fusion & Vaporization
The molar heat of fusion of ice is 6.02 kJ/mol.
Hfusion = 6.02 kJ/mol
The molar heat of vaporization of water is
40.6 kJ/mol at 100 oC.
Hvaporization = 40.6 kJ/mol
Heating curve for water.
Q = (ms t)ice + m Hf + (ms t) water + m Hv +
(mst)steam
Q = KE & PE + PE + KE & PE
+ PE + PE & KE
Calculating Energy Changes
Calculate the amount of energy required to melt 8.5 g
of ice at 0 oC.
Q = mHfusion
Q = (8.5g HOH)(1mol/18.02g HOH)(6.02kJ/mol)
Q = 2.8 kJ
Calculating Energy Changes
Liquid to Gas
Calculate the energy (in kJ) required to heat 25g of
liquid water from 25 oC to 100 oC and change it
to steam at 100 oC. The specific heat (s) of water
is 4.18 J/gCo.
Q = mst + mHvaporization
Q = (25g)(4.18 J/gCo)(75 Co)(1kJ/1000J) +
(25g)(1mol/18.02g)(40.6kJ/mol)
Q = 7.8 kJ + 57 kJ
Q = 65 kJ
Calculating Energy Changes
Solid to Gas
Calculate the energy (in kJ) required to melt 15g of
ice at 0 oC, heat the water to 100 oC, and
vaporize it to steam at 100 oC.
Q = mHfusion + mst + mHvaporization
Q = (15g)(1mol/18.02g)(6.02kJ/mol) + (15g)(4.18
J/gCo)(100 Co)(1kJ/1000J) +
(15g)(1mol/18.02g)(40.6kJ/mol)
Q = 5.0 kJ + 6.3 kJ + 34 kJ
Q = 45 kJ
Types of Bonding
Intramolecular
Intermolecular
• within the molecule
•between molecules
•covalent bonding
•dipole-dipole forces
•ionic bonding
•hydrogen bonding
•London Dispersion
Forces
When ice changes to liquid and then to vapor, the
intramolecular forces (covalent bonds) stay intact, only
the weaker hydrogen bonds between molecules weaken
and break. These are, therefore, physical changes.
Intermolecular Forces
Forces between (rather than within) molecules.
-
dipole-dipole attraction: molecules with
dipoles orient themselves so that “+” and
“” ends of the dipoles are close to each
other. (1 % as strong as covalent or ionic.)
-
hydrogen bonds: dipole-dipole attraction in
which hydrogen is bound to a highly
electronegative atom. (F, O, N)
10_208
–
–
+
+
(a)
–
+
+
–
–
+
–
+
+
–
–
+
–
–
+
+
Attraction
Repulsion
(b)
Electrostatic interaction of two polar molecules.
The polar water molecule and hydrogen bonds
among water molecules.
10_210
100
H2 O
Group 6A
Boiling point ( °C)
HF
0
H 2Te
SbH3
Group 7A
H 2Se
NH 3
H 2S
– 100
AsH 3
HCl
Group 5A
SnH4
HBr
GeH4
PH3
Group 4A
HI
SiH 4
CH 4
– 200
2
3
4
5
Period
The boiling points of the covalent hydrides of the
elements in Groups 4A, 5A, 6A, & 7A. The high
boiling points of HOH, NH3, HF is due to hydrogen
bonding.
10_211
+
+
Atom A
Atom B
H
–
+
+
Atom A
Atom B
–
H
H
Molecule A
H
Molecule B
Instantaneous dipole on molecule A
induces a dipole on molecule B
+
–
+
H
+
+
Atom A
Atom B
(a)
Molecule B
+
H
Instantaneous dipole on atom A
induces a dipole on atom B
+
H
No polarization
+
–
H
Molecule A
No polarization
–
H
H
Molecule A
–
+
H
H
Molecule B
(b)
Instantaneous and induced dipole moments
between nonpolar molecules -- London
Dispersion Forces. LDF forces are both weak
and short-lived.
London Dispersion Forces
Also Known As Induced Dipoles
Size of the London Dispersion Force depends on the
number of electrons and shapes of molecules
– the larger the molar mass, the larger the induced
dipole
-
- +
- -
-
-
- +
- -
-
- -- - -+
- - + -+
- - - -- +
- -- - -+
- - + -+
- - - -- +
London Dispersion Forces
-
relatively weak forces that exist among
noble gas atoms and nonpolar molecules.
(Ar, C8H18)
-
caused by instantaneous dipole, in which
electron distribution becomes asymmetrical.
-
the ease with which electron “cloud” of an
atom can be distorted is called
polarizability.
Vapor Pressure
Equilibrium
Liquid just
poured into
open
container,
little vapor
Evaporation faster
than Condensation
Evaporation as fast
as Condensation
Vapor Pressure
. . . is the pressure of the vapor present at
equilibrium with its liquid.
. . . is determined principally by the size of
the intermolecular forces in the liquid.
. . . increases significantly with temperature.
Volatile liquids have high vapor pressures.
Vapor Pressure
Low boiling point
•
high vapor pressure.
•
weak intermolecular forces.
Low vapor pressure
•
high molar masses.
•
strong intermolecular forces.
T1
(a)
Energy needed
to overcome
intermolecular
forces in liquid
Kinetic energy
Number of molecules
with a given energy
Number of molecules
with a given energy
10_245
Energy needed
to overcome
intermolecular
forces in liquid
T2
(b)
Kinetic energy
Boltzman Distribution -- number of molecules in
a liquid with a given energy versus kinetic energy
at two different temperatures.
Rate of Evaporation & High
Vapor Pressure
Liquids evaporate more rapidly and have a
high vapor pressure when the liquid:
1. has weak intermolecular forces.
2. is made up of lighter molecules.
3. is at a high temperature.
Vapor Pressure
Which of the following pairs have the highest
vapor pressure?
1. HOH(l) or CH3OH(l)
2. CH3OH(l) or CH3CH2CH2CH2OH(l)
Why?
Sublimation
•Change of a solid
directly to a vapor
without passing through
the liquid state.
•Iodine
•Dry Ice
•Moth Balls
Types of Solids
Crystalline Solids: highly regular
arrangement of their components [table salt
(NaCl), pyrite (FeS2)].
Amorphous solids: considerable disorder in
their structures (glass).
Representation of Components
in a Crystalline Solid
Lattice: A 3-dimensional system of
points designating the centers of
components (atoms, ions, or molecules)
that make up the substance.
Representation of Components
in a Crystalline Solid
Unit Cell: The smallest repeating unit of
the lattice.
- simple cubic -- 1 atom/cell
- body-centered cubic -- 2 atoms/cell
- face-centered cubic -- 4 atoms/cell
10_213
Unit cell
Lattice
Example
(a)
Polonium
metal
Simple cubic
(b)
Uranium
metal
Body-centered
cubic
(c)
Gold
metal
Face-centered
cubic
Three cubic unit cells and the corresponding
lattices.
Types of Crystalline Solids
Atomic Solid: contains atoms at the lattice
points (diamond) -- variable melting points.
Ionic Solid: contains ions at the points of the
lattice that describe the structure of the solid
(NaCl) -- high melting points, strong
electrostatic forces between + & - ions.
Molecular Solid: discrete covalently bonded
molecules at each of its lattice points (sucrose,
ice) -- low melting points, weak attraction.
Crystalline Solids
Metallic
Molecular
malleable & ductile brittle & weak, or
soft & waxy solids
Usually high MP
MP < 300°C
High BP
Low BP
High Hvap
Low Hvap,
Hfusion
high density
low density
good conductor
insulator
soluble in other
metals
solubility varies
Ionic
hard & brittle
Atomic
Networks
very hard
MP > 300°C
MP > 1000°C
Very high BP
Very high BP
High Hvap, Hfusion Very high Hvap,
Hfusion
medium density
medium density
good electrical
very insulating
conductor when
very unreactive
molten or dissolved
in water
often soluble in
dissolve in very few
water
things
10_216
= Cl
= Na
Sodium chloride
=C
Diamond
(a)
(b)
= H2O
Ice
(c)
Three crystalline solids -- a) atomic solid, b) ionic
solid, and c) molecular solid.
The properties of solids are determined primarily by
the nature of the forces that hold the solid together.
Molecular Solids
•
molecular units at each lattice position.
•
strong covalent bonding within molecules.
•
relatively weak forces between molecules.
•
London Dispersion Forces -- CO2, I2, P4,
& S8.
•
Hydrogen Bonding -- H2O, NH3, & HF.
Figure 13.16: (Left) Sulfur crystals contain S8 molecules.
(Right) White phosphorus contains P4 molecules.
Network Solids
Composed of strong directional covalent
bonds that are best viewed as a “giant
molecule”.
- brittle
- do not conduct heat or electricity
- carbon, silicon-based
graphite, diamond, ceramics, glass
10_229
(a)
Diamond
Network solid structure of diamond.
Bonding Models for Metals
Metals are malleable, ductile, good
conductors of heat and electricity, have high
melting points, and are durable. The
bonding in metals is strong but
nondirectional.
Bonding Models for Metals
Electron Sea Model: A regular array of
metals in a “sea” of electrons. The cations
are mutually attracted to the valence
electrons--this attraction holds the metal
together.
The mobile electrons can conduct heat and
electricity and the cations are fairly easily
moved--making the metal malleable &
ductile.
Electron Sea Model
Atomic Solids that are made of metal atoms
– metal atoms release their valence electrons
– metal cations fixed in a “sea” of mobile electrons
– Leads to strong attractions that are non-directional
+
+
e-
+
+
e+
+
+
e-
e-
+
+
+
e-
e-
+
+
e-
+
e+
+
+
e-
e-
+
+
+
e-
e-
+
+
+
e-
e-
+
+
+
e-
e-
+
+
Metal Alloys
Substances that have a mixture of elements and
metallic properties.
1. Substitutional Alloy: some metal atoms
replaced by others of similar size.
brass = Cu/Zn
Metal Alloys
(continued)
2. Interstitial Alloy: Interstices (holes) in
closest packed metal structure are
occupied by small atoms.
steel = iron + carbon
3. Both types: Alloy steels contain a mix of
substitutional (Cr, Mo) and interstitial
(Carbon) alloys.
Substitutional
Alloy
Interstitial Alloy
Steels
Pure iron is relatively soft, ductile, and
malleable. The addition of carbon to iron
makes the bonds more directional.
Mild Steels - contain less than 0.2 % carbon.
Still ductile and malleable - used for nails,
cables, and chains.
Steels
Medium Steels - contain 0.2 - 0.6 % carbon.
Harder than mild steels - used for rails and
structural beams.
High-carbon steels - contain 0.6 - 1.5 % carbon.
Tough and hard - used for springs, tools, and
cutlery.
Alloy steels - mixed interstitial & substitutional
alloys. Stainless steel has cobalt and nickel
substituted for iron - resistant to corrosion.
Identifying Types of Solids
Name the type of solid formed by each of the following:
a. ammonia
a. molecular
b. iron
b. atomic
c. cesium fluoride
c. ionic
d. argon
d. atomic
e. sulfur (S8)
e. molecular
f. sulfur trioxide
f. molecular
g. barium oxide
g. ionic
h. gold
h. atomic