Transcript Chapter 11: Liquids & Solids

Chapter 11: Liquids & Solids
 The molecular compounds like water, ammonia, and
carbon dioxide have different physical properties because
of the intermolecular forces.
 Comparison of all three phases:
Liquids & Solids
Liquids & Solids
of r
th la
n g cu
re le
St o
rm ns
te o
In ti
ac
?
ttr
ow
A
Fl
it
e?
ill
bl
si
W
es
pr
om
C
State Shape Volume
Solid
fixed
fixed
Liquid indefinite fixed
Gas indefinite indefinite
Density
high
high
low
No No
No Yes
Yes Yes
very strong
intermediate
weak
Compressibility
Changes of State
 Changes in state can be induced by a change in
temperature or pressure.
Intermolecular Forces
 Forces between molecules.
 Always LESS in energy than actual bond.
 The attractive force between two HCl molecules is
about 16 kJ/mol.
 The bond dissociation energy of the HCl bond is
about 431 kJ/mol.
Intermolecular Forces
 One method to compare the strength of
intermolecular forces is to examine the substance’s
boiling point.
 When the forces are relatively weak, then the boiling
point is small.
 Ex) HCl, bp = -85oC.
 There are three main types of intermolecular forces
between neutral substances.
Intermolecular Forces
 In the LS packet, we
identified molecules
as being polar or
non-polar based on
shape, types of
atoms, etc.
 Polar molecules have
a dipole – that is a
positive and negative
end.
Intermolecular Forces
 Thus, the first type of force is called Dipole-Dipole
(or DD) forces and occurs for any polar molecule.
 The larger the dipole moment, the more DD forces.
Intermolecular Forces
 If you cool any non-polar molecule or atom to a low




enough temperature, then it will liquefy.
Yet, these have no reason to be attractive to each
other.
Fritz London first proposed a theory in 1930.
On average, electrons in an atom like He are evenly
distributed.
But, in one INSTANT, the two electrons may both be
on the same side.
Intermolecular Forces
 Thus, in that one INSTANT, a He atom would have
an instantaneous dipole.
 This is called the London Dispersion (or LD) force.
Intermolecular Forces
 Since all molecules have electrons, they all have a LD
force.
 The polarizability of an atom or molecules electrons
depends on two factors.
The number of electrons. More electrons = More LD
forces.
2. The shape of molecule. More spread out = more LD
forces.
1.
Intermolecular Forces
Non-polar Alkanes
Intermolecular Forces
 The formula C5H12 has three structural isomers.
CH3 – CH2 – CH2 – CH2 – CH3
CH3
CH3 – C – CH3
CH3
CH3
CH3 – CH2 – CH – CH3
Non-polar Branched Alkanes
Name
Molar Mass
Boiling Point
Pentane
72.15 g/mol
36.1oC
Methylbutane
72.15 g/mol
27.7oC
Dimethylpropane
72.15 g/mol
10oc
Butane
58.12 g/mol
-0.5oC
Methylpropane
58.12 g/mol
-11.7oC
 Comparison of the group 4A, 5A, 6A, and 7A
hydrides shows an interesting result.
 What type of forces do the group 4A have?
 Group 6A?
 What is the notable exception?
Intermolecular Forces
 The third type of force is a
special case of DD force and is
called Hydrogen Bonding (or
HB).
 The name “Hydrogen Bonding”
is a misnomer!
 HB can only occur when:
 H is bonded to either N, O, or
F.
 The N, O, or F atom has at least
one lone pair.
Intermolecular Forces
 The strength of Hydrogen Bonding varies from 5 kJ
to 40 kJ, which is still much weaker than a covalent
bond (200 – 1000 kJ).
 However, it is MUCH stronger than DD or LD forces.
 Thus, it can greatly increase the boiling point
temperatures of molecules.
Intermolecular Forces
 HB forces are very important in biochemistry.
 Proteins are made from the twenty amino acids.
 The structure of the amino acid has both an –OH
group and an –NH2 group that can HB.
R O


H–N–C–C –O –H

H H
Intermolecular Forces
 Predicting relative boiling points.
1. Determine the molecular weight.
2. Determine the type(s) of intermolecular forces




present.
If weights are similar, then LD < DD < HB
If weights are very dissimilar, then #2 probably
does not matter.
However, HB can really distort the bp’s!
Ex) H2O, bp = 100oC, MW = 18 g/mol versus CCl4,
bp = 76oC, MW = 154 g/mol.
Intermolecular Forces
 The strengths of
attractions between
the molecules may
affect a liquids
properties.
 Viscosity
 Surface Tension
Intermolecular Forces
 Viscosity is the resistance of a
liquid to flow.
 Liquids with low viscosity, like
water, will produce a “splash”
whereas liquids with high
viscosity, like corn syrup or
ketchup, will not.
Intermolecular Forces
 Viscosity tends to increase with more intermolecular




forces and molecular weight.
Many liquids, like water, have a consistent viscosity
over a wide range of temperatures.
Some liquids, like corn syrup, will decrease in
viscosity as the temperature increases.
Multi-weight motor oil actually increases with an
increase in temperature (ie. 5W – 30).
Non-Newtonian liquids (ie. “Slime”) have a variable
viscosity at the same temperature.
Surface Tension
 Surface Tension is the
“skin-like” appearance
of the surface.
 Results from surface
molecules seeking six
nearest neighbors like
interior molecules.
Surface Tension
Phase Changes
Phase Changes
 Energy when changing between solid and liquid
phase is called the Heat of Fusion and denoted as
DHfus.
 DHfus for water is 6.01 kJ/mol or 334 J/g.
 Energy when changing between liquid and gas is
called the Heat of Vaporization and denoted as
DHvap.
 DHvap for water is 40.67 kJ.mol or 2,260 J/g.
Heating Curve
Refrigeration
 The basics of
refrigeration.
 First law of
thermodynamics
at work again!
 Coolant is CF2Cl2
(old) or CF3CH2F
(new).
Vapor Pressure
 Above the surface of any liquid, some liquid molecules
will have enough energy to escape and become gas
molecules.
 In a closed system, an equilibrium will be achieved
between the gas molecules and the liquid.
 This is the vapor pressure.
Vapor Pressure
 As the temperature of the liquid increases, its vapor
pressure will increase.
Vapor Pressure
 When the vapor pressure
equals 1 atmosphere,
then the liquid
spontaneously becomes a
gas. You would call this
the boiling point.
 Does pure water always
boil at 100oC?
Clausius-Clapeyron Equation
 The graphs of vapor
pressure versus
temperature are
approximately an
exponential function.
 Mathematically, if you
take the natural
logarithm (ln key on
calculator) of the vapor
press versus 1/T, then
you get a linear
relationship.
Clausius-Clapeyron Equation
 P2  DH vap  1 1 
n  
  
R  T1 T2 
 P1 
•R is molar gas constant = 8.314 J/K mol and the T is
the temperature in Kelvin
•Heat of vaporization must be in J/mol.
•Pressures can be in either atm or mmHg (must agree).
Phase Diagrams
 Display a single’s substances states of matter over a
wide range of P and T.
Carbon Dioxide
 The phase diagram of
CO2 shows that the
liquid phase can only be
found above a pressure
of 5.11 atm.
 As the temperature of
solid CO2 increases, it
undergoes sublimation.
Water
 The phase diagram of
water has one very
important difference.
What is it?
Solids
 Solids can be either
amorphous (random) or
crystalline (repeating pattern).
 Unit cell is the smallest
repeating pattern for the
crystalline structure.
 Analogy: a hotel with many
floors.
 Structure of unit cell can have
various lengths and angles.
Solids
 While many types of unit cells are possible, a few are
seen many times in structures of metals, molecular,
and ionic compounds.
 Cubic unit cells – two main versions.
 Body-centered cubic (BCC) – has atoms at each corner
and an atom in the body-center.
 Face-centered cubic (FCC) – has atoms at each corner
and an atom on each face.
 Important – just like a hotel room shares walls,
floors, and ceilings with other rooms, so does a unit
cell share atoms with other unit cells.
Solids
Solids
Solids
 Can also have atoms on edges in larger unit cells –
namely for ionic compounds.
 Thus, the following are the contributions for
locations on or in a unit cell:
Solids
 Unit cell calculations will follow the formula:
MW  C
Vc 
D  NA
 Where Vc is the volume of the cubic unit cell, MW is
the molar mass, C is the number of atoms per unit
cell, D is the density (m/V), and Na is Avogadro’s
Number.
Solids
 Another view – Closest
Packing Model.
 Assumes that atoms are
hard spheres.
 Maximize the density,
minimize the empty spaces.
Solids
 First layer – what is the most efficient method of
arrangement?
Solids
 Second layer is placed so that spheres sit in gaps from
previous row.
 Third layer can either repeat first layer yielding an
ABABAB… pattern.
 OR, the third layer is offset from the first two producing
an ABCABCABC… pattern.
Solids
 The ABABAB…
pattern produces a
unit cell called
hexagonal closest
packing or HCP.
 This is NOT a
cubic unit cell!
Solids
 The ABCABCABC… pattern produces a unit cell
called cubic closest packing or CCP.
 However, CCP is the same as FCC!
Solids
H
He
Hcp
Li
Bcc
Be
Hcp
B
C
N
O
F
Ne
Fcc
Na
Bcc
Mg
Hcp
Al
Fcc
Si
P
S
Cl
Ar
Fcc
K
Bcc
Ca
Fcc
Sc
Hcp
Ti
Hcp
V
Bcc
Cr
Bcc
Mn
Bcc
Fe
Bcc
Co
Hcp
Ni
Fcc
Cu
Fcc
Zn
Hcp
Ga
Ge
As
Se
Br
Kr
Fcc
Rb
Bcc
Sr
Fcc
Y
Hcp
Zr
Hcp
Nb
Bcc
Mo
Bcc
Tc
Hcp
Ru
Hcp
Rh
Fcc
Pd
Fcc
Ag
Fcc
Cd
Hcp
In
Sn
Sb
Te
I
Xe
Fcc
Cs
Bcc
Ba
Bcc
Hf
Hcp
Ta
Bcc
W
Bcc
Re
Hcp
Os
Hcp
Ir
Fcc
Pt
Fcc
Au
Fcc
Hg
Tl
Hcp
Pb
Fcc
Bi
Po
At
Rn
Solids
 All crystalline solids can be catagerized into one of
four types.
 Type 1: Molecular Solids
 Consist of atoms or molecules like Ne, CH4, and H2O.
 Are held together by relatively weak intermolecular
forces.
 Are soft and have low melting points (unless they have
a high MW).
 Poor conductors of heat and electricity.
Solids
 Type 2: Ionic Solids
 Consist of ions held together by their electrostatic
attractions.
 Unit cells are always larger since the smallest repeating
pattern must include two ions.
 When cation and anion are of similar sizes, get BCC type
arrangement. When anion is much larger, get a CCP
arrangement of anions with cations stuck in the “holes.”
 Hard and brittle and have high melting points.
 Poor electrical conductors as solids, but excellent when
melted.
Solids
(a) CsCl
(b) ZnS
(c) CaF2
Solids
 View of NaCl
Solids
 Type 3: Metallic Solids
 Atoms are held together by
a “sea of valence
electrons.”
 Can be soft (Na, Au) or
very hard (Fe, Co) with low
to very high melting
points.
 Excellent conductors of
both heat and electricity.
 Malleable and ductile.
Solids
 Type 4: Covalent Network Solids
 Consist of atoms held together in large networks of
covalent bonds.
 There are not many of these – C(diamond), SiO2,
quartz, SiC, and BN.
 Very hard with very high melting points.
 Poor conductors.
Solids
 Two forms of carbon, diamond and graphite.
Solids
 Comparing metal points of solids.
 First – determine the type of solid.
 Molecular is always the lowest of the four types.
 Second – if both are the same type of solid, then:
 Molecular is like bp’s. LD < DD < HB.

Ex) CH4 (-182 C) < COCl2 (-118 C) < H2O (0 C)
 Ionic depends on lattice energy – the larger the lattice
energy, the higher the mp.

Ex) NaCl (801 C) < MgO (2852 C)
Solids
 Metallic depends on the number of unpaired electrons.
More unpaired electrons = higher melting point.




K, 1 unpaired electron, mp = 64 C
Ti, 2 unpaired electrons, mp = 1668 C
Cr, 6 unpaired electrons, mp = 1907 C
Cu, 1 unpaired electron, mp = 1065 C
 Covalent network are always very high.


Quartz, mp = 1670 to 1710 C
Diamond, mp = 3550 C (highest of any naturally occurring
substance)