Transcript (a) Solution 2013 General Chemistry I

Chapter 5. LIQUIDS AND SOLIDS
INTERMOLECULAR FORCES
5.1 The Origin of Intermolecular Forces
5.2 Ion-Dipole Forces
5.3 Dipole-Dipole Forces
5.4 London Forces
5.5 Hydrogen Bonding
5.6 Repulsions
LIQUID STRUCTURE
5.7 Order in Liquids
5.8 Viscosity and Surface Tension
2012 General Chemistry I
INTERMOLECULAR FORCES (Sections 5.1-5.6)
5.1 The Origin of Intermolecular Forces
 Phase: uniform in both chemical composition and
physical state
- Condensed phase: simply a solid or liquid phase
- Condensed phases form when attractive
intermolecular forces between molecules pull them
together; repulsions dominate at even shorter
separations.
- Intermolecular forces
are weak compared
with bonding forces,
but boiling points and
sublimation points
depend on
their strength.
- All interionic and almost all
intermolecular forces can be
traced to the coulombic
interaction between charges.
- Distance dependence of potential energy of interaction
1/r :
1/r2 :
1/r3 :
1/r6 :
between ions (ionic bonding)
between ions and dipoles
between stationary dipoles
between rotating dipoles
TABLE 5.1 Interionic and Intermolecular Interactions
5.2 Ion-Dipole Forces
 This is the attractive force between ions and polar
molecules in liquid or solid phase.
 Hydration: attachment of water molecules to
ionic solute particles is an example of iondipole interaction.
H2O
 The potential energy of ion-dipole interactions
(~15 kJ mol-1)
z = the charge number of the ion
m = the electric dipole moment of the
polar molecule
 Water of crystallization: smaller and highly
charged cations strongly attract polar water
molecules in the solid phase.
- Note hydrated salts of Li and Na vs.
anhydrous salts of K, Rb, Cs, and NH4+
E.g. Na2CO3.10H2O compared with K2CO3
(Na+; 102 pm, K+; 138 pm)
- Note BaCl2 · 2H2O vs. anhydrous KCl
(Ba2+; 136 pm, K+; 138 pm)
5.3 Dipole-Dipole Forces
 This is the attractive force between polar molecules.
 Between stationary polar
molecules in the liquid
phase (~2 kJ mol-1)
 Between rotating polar
molecules in the gas phase
(~0.3 kJ mol-1)
Self-Test 5.1B
Which will have the higher boiling point, 1,1-dichloroethene
or tr ans-1,2-dichloroethene?
Solution
1,1-dichloroethene is polar, whereas tr ans-1,2-dichloroethene
is nonpolar:
H
Cl
C
Cl
H
Cl
C
C
H
H
C
Cl
Hence, dipole-dipole (as well as London) forces exist
in 1,1-dichloroethene, giving it the higher boiling point.
5.4 London Forces
 London force (induce dipole-induced dipole force,
dispersion force) ~2 kJmol-1: it exists between all
molecules but is the only interaction between
nonpolar molecules.
- Attractive interactions
due to instantaneous
fleeting dipole moments
- Fluctuation of the electron
distribution in one molecule→
temporary dipole → second
temporary dipole in the other molecule
→ ···
Time
- Potential energy (strength) of the London interaction is
given by
 Polarizability a is the ease with which molecular electron clouds
can be distorted: a increases with number of electrons.
- A large linear molecule is
more likely to have stronger
London interactions (and hence
a higher boiling point) than a
smaller or nonlinear one.
Examples
Alkanes
C5H12; mobile liquid
C15H32; viscous liquid
C18H38; waxy solid
- Halogens: gases (F2, and Cl2); liquid (Br2); solid (I2)
- Rod-shaped (pentane; Tb = 36 oC) vs.
spherical (2,2-dimethylpropane; Tb = 10 oC)
TABLE 5.2 Melting and Boiling Points of Substances
Allied intermolecular interactions
 Dipole-induced dipole interaction between a
polar molecule and a nonpolar molecule (~2
kJ mol-1)
– Van der Waals interactions is the collective name
for dipole-dipole forces between rotating polar
molecules, London forces, and
dipole-induced dipole forces
EXAMPLE 5.2
Explain the trend in the boiling points of the hydrogen halides:
HCl, -85 oC; HBr, -67 oC; HI, -35 oC.
- Electronegativity differences: HCl > HBr > HI
- Number of electrons and London forces: HCl < HBr < HI
→ not by dipole-dipole forces, but by London forces
Self-Test 5.2A
Account for the trend in boiling points of the noble gases,
which increase from helium to xenon.
Solution
In the noble gases, only London forces need be considered:
These increase as the number of electrons increases (size
of the atom increases):
He (2)
B.Pt
Ne(10) Ar(18)
Kr(36)
Xe(54)
5.5 Hydrogen Bonding
Some compounds are characterized by exceptionally
high Tb due to hydrogen bonding: examples include
NH3, H2O, HF – see Fig. 5.9.
Hydrogen bonding
London forces
 Hydrogen bonding: an attraction in which a
hydrogen atom bonded to a small, strongly
electronegative atom, specifically N, O, or F, is
attracted to a lone pair of electrons on another N,
O, or F atom. Intermolecular and intramolecule
types exist.
- strong electrostatic interaction ~20 kJ mol-1
O…….H-O
linear but asymmetric (101 pm vs 175 pm)
- Hydrogen fluoride, (HF)n
- DNA
double
helix
- Acetic acid dimer (vapor)
- Protein
folding
Self-Test 5.3B
Which of the following molecules can take part in hydrogen
bonding with other molecules of the same compound:
(a) CH3OH; (b) PH3; (c) H-O-Cl?
Solution
(a) and (c), since these both have a H-O covalent bond:
CH3
H
O
H3C
H
Cl
H
O
O
H
O
Cl
(most likely)
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Example 5.1
Identify the kinds of intermolecular forces that might arise between
molecules of each of the following substances:
(a) NH2OH; (b) CBr4; (c) H2SeO4; (d) SO2
Solution
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Example 5.5
Suggest, giving reasons, which substance in each of the following
pairs is likely to have the higher normal melting point (Lewis
structures may help your arguments):
(a) HCl or NaCl; (b) C2H5OC2H5 (diethyl ether) or C4H9OH (butanol);
(c) CHI3 or CHF3; (d) C2H4 or CH3OH.
Solution
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5.6 Repulsions
-Intermolecular repulsions arise from the
overlap of orbitals on neighboring molecules
and the requirements of the Pauli exclusion
principle.
- They are important only at very short
distances:
1
Ep
r 12
LIQUID STRUCTURE (Sections 5.7-5.8)
5.7 Order in Liquids
- The liquid phase lies between the extremes of
the gas and solid phases.
gas phase: moving with almost complete freedom
minimal intermolecular forces
solid phase: locked in place by intermolecular forces
oscillate around an average location
- In the liquid phase,
molecules have short-range
order but not long-range
order.
- Water loses only 10% of
hydrogen bonds upon
melting and the rest are
continuously broken
and reformed.
5.8 Viscosity and Surface Tension
 Viscosity: resistance to flow, indication of
the intermolecular force strength
- Water and glycerol: very viscous due to
hydrogen bonding
- Hydrocarbon oils and grease: viscous due to
tangling long chains
- Viscosity usually decreases with temperature
due to higher energy of molecules.
Viscosities of
common liquids
Linear alkane chains in
Heavy hydrocarbon oil
 Surface tension: the net inward pull, an indication
of the intermolecular force strength
- Water: three times larger than many other
liquids, due to hydrogen bonds
- Mercury: more than six times that of water,
partially covalent
 Wetting: strong interactions of water with the
materials’ surface. Water maximizes its contact
with the materials by hydrogen bonding.
 Capillary action: adhesive
forces between a liquid and
surface vs. cohesive forces
within the liquid
– Meniscus: indication of the relative
strength of adhesion and cohesion
H2O
Hg
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Example 5.21
Predict how each of the following properties of a liquid varies as
the strength of intermolecular forces increases and explain your
reasoning: (a) boiling point; (b) viscosity; (c) surface tension.
Solution
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Example 5.23
Predict which liquid in each of the following pairs has the greater
surface tension: (a) cis-dichloroethene or trans-dichloroethene;
(b) benzene at 20 oC or benzene at 60 oC.
Solution
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Example 5.25
Rank the following molecules in order of increasing
viscosity at 50 oC: C6H5SH, C6H5OH, C6H6.
Solution
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Chapter 5. LIQUIDS AND SOLIDS
SOLID STRUCTURES
5.9 Classification of Solids
5.10 Molecular Solids
5.11 Network Solids
5.12 Metallic Solids
5.13 Unit Cells
5.14 Ionic Structures
THE IMPACT ON MATERIALS
5.15 Liquid Crystals
5.16 Ionic Liquids
2012 General Chemistry I
SOLID STRUCTURES (Sections 5.9-5.14)
5.9 Classification of Solids
 Crystalline solid: a solid in which the atoms,
ions, or molecules lie in an orderly array with
crystal faces
 Amorphous solid: one in which the atoms, ions, or
molecules lie in a random jumble
quartz
amorphous silica
Classification of Crystalline Solids
According to the bonds that hold their atoms, ions, or
molecules in place:
Metallic: consisting of cations held together by
a sea of electrons
Ionic: built from the mutual attractions of cations
and anions
Molecular: assemblies of discrete molecules held
in place by intermolecular forces
Network: consisting of atoms covalently bonded
to their neighbors throughout the extent of the
solid
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Example 5.35
Classify each of the following solids as ionic, network, metallic, or
molecular: (a) quartz, SiO2; (b) limestone, CaCO3; (c) dry ice, CO2;
(d) sucrose, C12H22O11; (e) polyethylene, a polymer with molecules
consisting of chains of thousands of repeating –CH2CH2- units.
Solution
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5.10 Molecular Solids
Molecular solids consist of molecules held together
by intermolecular forces; physical properties depend
on the strengths of those forces.
Amorphous molecular solids: as soft as paraffin
wax
Crystalline molecular solids:
- sucrose: numerous hydrogen bonds between OH
groups account for high melting point at 184 oC
- ultrahigh-density polyethylene: smooth yet tough
Molecular Solids and Liquids: Melting and
Freezing
Most substances increase in density on freezing:
water is an important exception. Ice at 0 oC is
less dense than water at 0 oC due to a more
open hydrogen-bonded structure.
- ice
water
benzene
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Example 5.33
Glucose, benzophenone (C6H5COC6H5), and methane are examples of
compounds that form molecular solids. The structures of glucose and
benzophenone are given here. (a) What types of forces hold these
molecules in a molecular solid? (b) Place the solids in order of
increasing melting point.
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Solution
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5.11 Network Solids
Network solids are characterized by a strong covalent
bond network throughout the crystal: they are very
hard and rigid, with high Tm and Tb
E.g. Two common allotropes (forms of an element
that differ in the way in which the atoms are linked)
of carbon have very different network structures.
Diamond, with an sp3 hybrid s-bonding framework, is
one of the hardest substances
Graphite has flat sheets of sp2
hybrid s-bonds with weak
bonding between sheets. It
conducts electricity parallel to
the sheet, and is soft and
slippery
- Ceramic materials: noncrystalline inorganic oxides,
great strength
5.12 Metallic Solids
In metallic solids, the cations are bound together
by their interaction with the sea of the electrons
that they have lost.
 Close-packed structure: the spheres stack
together with the least waste of space
- Hexagonal close-packed structure (hcp):
packed with the sequence of ABABAB···
Coordination number = 12
(3 plane below + 6 own plane + 3 plane
above): this is the maximum.
 Coordination number: the number of
nearest neighbors of each atom
- Cubic close-packed structure (ccp):
packed with the sequence of ABCABC···
Coordination number of ccp = 12
Occupied space in a ccp:
 Holes: the gaps (interstices) between the atoms in a crystal
- Octahedral hole: a dip in a layer coincides
with a dip in the next layer
- Tetrahedral hole: a dip between three atoms
is directly covered by another atom
5.13 Unit Cells
 Unit cell: the smallest unit that, when stacked
together repeatedly without any gaps and without
rotations, can reproduce the entire crystal.
- Face centered cubic (fcc, cubic F)
- Body centered cubic (bcc, cubic I)
- Primitive cubic (cubic P)
Cubic F
Cubic I
Cubic P
 Bravais lattices: 14 basic patterns of unit cell in 3D crystalline
systems; P = primitive; I = body-centered; F = face-centered; C = with
lattice point on two opposite faces; R = rhombohedral
Unit cells are characterized by
lengths a, b, c and angles a, b, g
- Cubic unit cells
- Primitive cubic (cubic P)
- Body centered cubic (bcc, cubic I)
- Face centered cubic (fcc, cubic F)
Self-Test 5.4A
How many atoms are there in a primitive cubic cell?
Solution
In a cubic P cell there are only eight corner atoms,
hence the total number of atoms per unit cell is:
8 x 1/8 = 1.
EXAMPLE 5.3
The density of copper is 8.93 gcm-3 and its atomic radius is 128 pm.
Is the metal (a) close-packed or (b) body-centered cubic?
(a) Fcc (ccp) and hcp cannot be distinguished by density only.
For 4 atoms in a fcc cell,
(b) For 2 atoms in a bcc cell,
close-packed cubic (fcc)
Self-Test 5.5A
The atomic radius of silver is 144 pm and its density is
10.5 g cm -3. Is the stucture face-centered cubic (fcc;
close packed) or body-centered cubic (bcc)?
Solution
We begin by assuming fcc;
4 x (107.9 g mol-1)
4M
d=
8
3/2
3
N Ar
=
83/2 x (6.022 x 1023 mol-1) x (1.44 x 10-8 cm)3
= 10.6 g cm -3. Hence silver has fcc structure.
190s
Example 5.47
One form of silicon has density of 2.33 gcm-3 and crystallizes in a cubic
Lattice with a unit cell edge of 543 pm. (a) What is the mass of each unit
cell? (b) How many silicon atoms does one unit cell contain?
Solution
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Example 5.49
What percentage of space is occupied by close-packed cylinders of
length l and radius r ?
Solution
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5.14 Ionic Structures
- Ionic structures are (like metallic structures) close
packed, with anions forming a slightly expanded
close-packed structure with smaller cations
occupying some of the enlarged holes in the
expanded lattice.
- Smaller tetrahedral hole for small cations,
larger octahedral hole for somewhat bigger cations
 Coordination number: the number of ions of
opposite charge immediately surrounding a
specific ion
 Radius ratio (r)
 Rock-salt structure of NaCl
Has Cl– ions forming an fcc structure
(expanded ccp) and Na+ ions
occupying all the octahedral holes
Found for a number of minerals
having ions of the same charge
number, including NaCl, KBr, RbI,
MgO, CaO, AgCl
- Rock-salt (fcc) structure: (6,6) coordination, the
coordination numbers of the cations and the anions
are both 6.
- Common whenever the cations and the anions
have very different radii; the cations can fit into
the octahedral holes in a fcc array of anions; 0.4
< r < 0.7
 Cesium chloride structure
- Cl– ions form an expanded primitive
cubic array and Cs+ ions occupy
large cubic holes (bcc).
- The radii of the cations and anions
are similar with r > 0.7.
r(Cs+) = 167 pm, r(Cl-) = 181 pm; r = 0.923
- It has (8,8) coordination and is less common.
Examples: CsBr, CsI, TlCl, TlBr
 Zinc-blend (sphalerite) structure,
r < 0.4
- ZnS: expanded ccp array of S2– and
small Zn2+ in half of the tetrahedral
holes: (4,4)-coordination
- NiAs: strong Ni–As covalent character
hcp As with Ni in all octahedral holes
(6,6)-coordination
Self-Test 5.6A
Predict (a) the likely structure and (b) the coordination
type of ammonium chloride. Assume that the
ammonium ion can be approximated as a sphere with a
radius of 151 pm.
Solution
Radius ratio, r
=
=
Radius of smaller ion
Radius of larger ion
151 pm
= 0.834
181 pm
This indicates (a) a cesium chloride structure with (b) (8,8)
-coordination.
Self-Test 5.7B
Estimate the density of cesium iodide from its crystal
structure.
Solution
CsI has a cesium chloride (bcc) type structure.
r(Cs+) = 170 pm; r(I-) = 220 pm
Length of diagonal, b = 170 + 2(220) + 170 pm = 780 pm
1
2
Length of side a = 780/3 = 450 pm
Hence unit cell volume is = 9.11 x 107 pm3 or 9.11 x 10-23 cm3
(1 pm3 = 10-30 cm3)
Each bcc unit cell has one Cs+ ion and one I- ion,
Density = mass/volume =
(132.91 + 126.90) g mol-1
(6.022 x 10
23
-1
mol )
9.11 x 10-23 cm3 = 4.74 g cm-3
192s
Example 5.53
Calculate the number of cations, anions, and
formula units per unit cell in each of the
following solids:
(a) the cesium chloride unit cell
Solution
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(b) The rutile (TiO2) unit cell
(c) What are the coordination numbers
of the ions in rutile?
Solution
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Example 5.61
Graphite forms extended two-dimensional layers.
(a)Draw the smallest possible rectangular unit
cell for a layer of graphite. (b) How many carbon
atoms are in your unit cell? (c) What is the
coordination number of carbon in a single layer
of graphite?
Solution
(a)
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Example 5.63
If the edge length of a fcc unit cell of RbI is 732.6 pm, how long would
the edge of a cubic single crystal of RbI be that contains 1.00 mol RbI?
Solution
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THE IMPACT ON MATERIALS
(Sections 5.15-5.16)
5.15 Liquid Crystals
 Liquid crystals are substances that flow like viscous liquids, but their
molecules lie in a moderately orderly array.
- mesophase: an intermediate state of matter with the fluidity of
a liquid and some of the molecular order of a solid
- responsive to changes in temperature and electric fields
- isotropic vs. anisotropic (due to ordering of rodlike molecules)
- p-azoxyanisole, a long and rodlike liquid crystal molecule
 Three classes of liquid crystals according to structure
- Nematic phase: the molecules lie together, all in the same direction
but staggered.
- Smectic phase: the molecules line up like soldiers on parade and
form layers.
- Cholesteric phase: the molecules form ordered layers, but neighboring
layers have molecules at different angles and so the liquid crystal has a
helical arrangement of molecules.
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 Two classes of liquid crystals according to method of preparation
- Thermotropic LC: made by melting solids. They have long rod shaped
Molecules.
Example: p-azoxyanisole
Uses: watches, LCD, thermometers, ···
- Lyotropic LC: ordering effects induced by a solvent. The molecules are
amphiphiles (surfactants), with hydrophilic and hydrophobic parts in
one molecule
Example: sodium lauryl sulfate
- LCD
television or monitor
Layers of a
liquid crystal
in a nematic
phase lie
between the
surfaces of
two
glass or
plastic plates.
190s
Example 5.65
Why do long hydrocarbon molecules that do not have multiple bonds,
such as decane, CH3(CH2)8CH3, not form liquid crystals?
Solution
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190s
Example 5.67
When long surfactant molecule having a polar headgroup and a
nonpolar “tail” are placed into water, micelles are formed in which
the nonpolar tails aggregate, with the polar headgroups pointing
out toward the solvent. Inverse micelles are similar but have the
nonpolar regions pointing outward. How can inverse micelles be
placed?
Solution
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5.16 Ionic Liquids
- A liquid at room temperature is likely to be a molecular substance.
There will be nonionic, weak intermolecular interactions and it
will have a relatively high vapor pressure.
 Ionic liquids: These are characterized by a relatively small
anions (BF4–) + large organic cation (e.g. 1-butyl-3methylimidazolium ion), preventing crystallization.
Low vapor pressure, novel solvent properties: reducing pollution
X-RAY DIFFRACTION
 The Technique
- interference: When two or more waves pass through the same region,
interference is observed as an increase (constructive) or
a decrease (destructive) in the total amplitude of the wave.
- diffraction: interference between waves that arises when there is an
object in their path
- Regular layers of atoms in a crystal giving a diffraction pattern
- Why x-rays?
The separation between layers of atoms in a crystal ~ 100 pm
corresponding to the x-ray region
 Experimental Techniques
- Powder diffraction technique: a monochromatic (single-frequency)
beam of x-rays is directed at a powdered sample spread on a
support.
- Bragg equation 2d sinq = l
with the angles q, to the spacing d, for x-rays of wavelength l
- Single-crystal diffraction technique
1) Growing a perfect single crystal of the sample (~ 0.1 mm)
- very challenging!
2) Placing the crystal at the center of a four-circle diffractometer
- raw data including intensities and angles of the diffraction
3) Fourier synthesis (conversion) into the locations of the atoms
- description of the atomic locations, bond lengths, and angles
X-Ray
diff. pattern
of DNA
Helix with
a regular
pitch and
radius
big angle
diffraction
Narrow
spacing of
10 bases
per turn of
helix