Title: Liquids and Solids
1Chapter 10
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3Van Der Waals Forces
- These are intermolecular forces of attraction
between neutral molecules. - The Nobel Prize in Physics 1910 (Johannes van der
Waals) - "for his work on the equation of state for gases
and liquids"
4intER vs. intRA molecular forces
- Intramolecular forces are the forces within a
molecule or ionic compound - NaCl Ionic bond between atom of Na and atom of
Cl - Intermolecular forces are the forces between
molecules or ions and molecules - Example Solid liquid gas
-
- .
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6Intermolecular Force Model Basis of Attraction Energy (kJ/mol) Example
Ion-dipole Ion and polar molecule 600 40 Na H2O
Dipole-dipole Partial charges of polar molecules 25 5 HCl HCl
Hydrogen bond H bonded to N, O, or F, and another N, O, or F 40 10 H2O NH3
London dispersion Induced dipoles of polarizable molecules 40 0.05 Xe Xe
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83 Types of van der Waals Forces
- Dipole-Dipole forces
- London Dispersion forces
- Hydrogen bonding
9DIPOLE-DIPOLE FORCES
- These are forces of attraction that occur between
polar molecules. (big difference in electron
negativity) - These forces are effective only when polar
molecules are very close. As distance increase
strength of bond decreases. - For molecules of approximately equal mass and
size, the strength of force of attraction
increases as the polarity increases. - Radius have an effect on strength of dipole.
10DIPOLE-DIPOLE FORCES
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11DIPOLE-DIPOLE FORCES
- Molecules with larger dipole moments have higher
melting and boiling points (hard to break) than
those with small dipole moments. - Dipole attractions are relatively weak and tend
to be liquids or gas at room temperature.
12HYDROGEN BONDING
- A special type of dipole-dipole interaction
between the hydrogen atom in a polar bond and an
unshared electron pair of an element that is very
electronegative usually a F, O, or N atom on
another molecule - (note that all of these have very high ENs and
small atomic radii).
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14FYI
15HYDROGEN BONDING
- These types of bonds are super-humanly strong.
- (4X stronger that diopole dipole)
16HYDROGEN BONDING IN WATER
17HYDROGEN BONDING
18WHY HYDROGEN BONDING IS EFFECTIVE
- F, O, N are extremely small and very
electronegative atoms. - Hydrogen atoms have no inner core of electrons,
therefore, the positive side of the bond dipole
has the concentrated charge of the partially
exposed, nearly bare proton of the nucleus - in other words, the atoms have a large
difference in electronegativity and their nuclei
can get really close.
19IMPORTANCE OF HYDROGEN BONDING
- Are important biologically, in stabilizing
proteins and keeping DNA together. - Also explains why ice is less dense than water
(see text).
20LONDON DISPERSION FORCES
- Fritz London
- These are forces that arise as a result of
temporary dipoles induced in the atoms or
molecules.( its a temporary accident!) - All molecules have some degree of LD forces
21LONDON DISPERSION FORCES
22LONDON DISPERSION FORCES
- LD forces occur between neutral non-polar
molecules. (nobles gases and nonpolar compounds) - LD forces are weak
- The greater the number of electrons the greater
the LD force. (ie the greater the melting and
boiling pt.) - LD force molecules have Low melting and boiling
pts
23See Graphic on next slide
- The motion of electrons in an atom or molecule
can create an instantaneous dipole moment. - EX in a collection of He (g) the average
distribution of electrons about a nucleus is
spherical, the molecules are non-polar and there
is no attraction.
24INSTANTANEOUS AND INDUCED DIPOLES
Pg 454- 455 in text
25LONDON DISPERSION FORCES (CONT)
- These forces tend to increase in strength with an
increase in molecular weight (The size of the
molecule generally increases with mass and the
electrons are less tightly heldallows the
electron cloud to be more easily distorted. - These forces are stronger in linear molecules
than comparable bunched up molecules.
26LONDON DISPERSION FORCES
- LD forces are generally the WEAKEST
intermolecular forces. - Molecules with more electrons will experience
more LD forces
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28n-pentane vs neopentane
- BP 309.4 K BP 282.7 K
- Same atomic masses different structure
29Generalizations Regarding Relative Strengths of
IM Forces
- If molecules have comparable molecular weights
and shapes, dispersion forces are approximately
equal. Any difference in attractive forces is
due to dipole-dipole attractions. - If molecules differ widely in molecular weight,
dispersion forces are the decisive factor. The
most massive molecule has the strongest
attractions.
30Because melting points (MPs) and boiling points
(BPs) of covalent molecules increase with the
strengths of the forces holding them together, it
is common to use MPs and BPs as a way to compare
the strengths of intermolecular forces. This is
shown below, with the molecular formulas, molar
masses and normal BPs of the first five
straight-chain hydrocarbons. Molecular Formula
Molar Mass
Normal BP (C) CH4
16
- 161.5 C2H6
30
- 88.5 C3H8 44
- 42.1 C4H10
58
- 0.5 C5H12
72
36.1
31Which noble gas element has the lowest boiling
point?
32The chemical forces between HCl is/are
- Dispersion
- Covalent bond
- Hydrogen bond
- Dipol-dipole
- Two of the above
All Molecules Have
Not symmetrical Polar
33Consider the following list of compounds. How
many of these have hydrogen bonding as their
principle IMF
Hydrogen Bonding is between H and highley EN
atoms such as N, O, F, and H
- HCl
- NH3
- CH3OH
- H2S
- CH4
- PH3
34Which of the following statements are false or
correct and why?
- O2 is dipole dipole
- HCl is hydrogen bonding
- CO2 is dipole dipole
- NH3 is hydrogen
FALSE London Dispersion symmetric/nonpolar
FALSE Dipole dipole not symmetric/polar
FALSE London Dispersion symmetric/nonpolar
TRUE H N,O, or F
35ION-DIPOLE FORCES
- Attraction between an ion and the partial charge
on the end of a polar molecule.
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37ANSWER
38- A.Identify the types of bonds in
- Glucose
- Cyclohexane
- B.Glucose is soluble in water but cyclohexane is
not. Why?
39- A.Identify the types of bonds in
- Glucose
- H, LD, VanderWal, Dip-dip
- 2. Cyclohexane
- LD only
- B.Glucose is soluble in water but cyclohexane is
not. Why? - Glucose is polar and cyclohexane is nonpolar.
Polar compounds are soluble in polar solvent and
visversa.
40ION-DIPOLE FORCES
41ION-DIPOLE FORCES (CONT)
- The magnitude of attraction increases as either
the charge of the ion increases or magnitude of
the dipole moment increases. - Ion-dipole forces are important in solutions of
ionic substances in polar liquids (e.g. water)
42ION-DIPOLE FORCES AND THE SOLUTION PROCESS
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44Homework
- Pg 504-505
- s 35, 36, 37, 39 (you may need to read 10.1
for this part esp. LD portion)
4510.2 Properties of liquids
- Liquids are vital to our lives.
- Water is a means of food preparation
- Cooling machines n industrial processes
- Recreation
- Cleaning
- Transportation
46CHARACTERISTICS OF LIQUIDS
- Surface tension
- Capillary action
- Viscosity
47COHESIVE FORCES
- Intermolecular forces that bind like molecules to
one another (e.g. hydrogen bonding).
48ADHESIVE FORCES
- Intermolecular forces that bind a substance to a
surface.
49SURFACE TENSION
- A measure of the inward forces that must be
overcome in order to expand the surface area of a
liquid. - The greater the forces of attraction between
molecules of the liquid, the greater the surface
tension.
50Surface Tension Cont.
- Surface tension of a liquid decreases with
increasing temperature. - The stronger the intermolecular forces the
stronger the surface tension.
Water has a high surface tension do to hydrogen
bonding.
51CAPILLARY ACTION
- Another way surface tension manifests.
- The rise of liquids up very narrow tubes. This
is limited by adhesive and cohesive forces.
52Formation of meniscus
- Water adhesive forces are greater than cohesive
forces - Mercury Cohesive are greater than adhesive
forces.
53VISCOSITY
- The resistance of a liquid to flow.
- The less tangled a molecule is expected to be,
the less viscous it is.
Water less Viscosity
syrup high Viscosity
54Viscosity Cont.
- Viscosity decreases with increasing temperature
(molecules gain kinetic energy and can more
easily overcome forces of attraction). - Viscosity Increases as pressure increases.
- Liquids with strong IMF have a higher viscosity.
55Homework
5610.3 Structure of Solids
- Two ways to categorize solids
- Crystalline
- Amorphous
57Crystalline Solid
- Ridged and long range order of its atoms.
- Solids have flat surfaces
- Sharp melting points
- EX Quartz, diamond, sodium Chloride.
58Amorphous Solid
- Lack a well defined arrangement
- No long range order
- IMF vary in strength
- DO NOT have sharp melting points.
- EX rubber, glass
59Unit Cell
- The smallest part of a crystal that will
reproduce the crystal when repeated in a three
dimensions. - Three types
- Simple /primitive cubic cell
- Face centered cubic cell
- Body Centered
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61X-Ray Diffraction
- X-Ray crystallography lead to the discovery of
DNA
Beams of light shot at DNA and scattered to
reveal the double helix structure.
Watson, Franklin, Crick
62X-ray Diffraction
- Derived by the English physicists Sir W.H. Bragg
and his son Sir W.L. Bragg in 1913 to explain why
the faces of crystals appear to reflect X-ray
beams at certain angles of incidence (theta, ?). - d is the distance (?) between atomic layers in a
crystal - lambda ? is the wavelength (?) of the incident
X-ray beam - n is the order of the diffraction
- Braggs Equation n? 2d sin?
63Braggs Equation Demo
64Example
- The first order diffraction of x-rays from
crystal planes separated by 2.81 ? occurs at
11.8. - a. what is the wavelength of the x-ray
65The work
- n 1 (first order)
- d 2.81 ?
- T 11.8
- ? ?
- n? 2d sin?
- 1(?) 2(2.81 ?) sin 11.8
- Sin-1 24.14
- ? 24.14
66Homework
67 Changes of state
- Transformation from one state to another
Condensation
Vaporization AKA steam
68Changes in state
- Liquid ? Gas Vaporization Endothermic
-
- Gas ? Liquid Condensation Exothermic
69- Solid ? Gas Sublimation Endothermic
- Gas ? Solid Deposition Exothermic
70- Solid ? Liquid Melting Endothermic
- Liquid ? Solid Freezing Exothermic
71Changes of state
- The energy involved it phase changes is
calculated using - Heat of fusion (solid ? liquid or liquid? solid)
- Heat of vaporization (liquid? gas or gas? liquid)
-
72Energy Changes and Phase Changes
- Heat of Vaporization Vaporization is an
endothermic process ( it requires heat). Energy
is required to overcome intermolecular forces to
turn liq to gas. - Hvap is an Indicator of strength of IMF
- CH4 9.2 kJ/mol C3H8 18.1 kJ/mol
- Larger moleculegreater IMFgreater Hvap
73Question
- How much energy does it take to vaporizer 111 g
of water? - Given Hvap water 40.67 kJ/mol
- 111 g H2O 1 mol x 40.6kJ 250kJ
- 18g 1mol
74- Heat of Fusion the enthalpy change associated
with melting. (Solid to liquid.) - Hfusion water 6.01 kJ/mol
- NOTE heat of fusion is always smaller than heat
of vaporization. This makes sense think about the
level of order in the molecules in these
phases.
75Heating Curve
- A plot of the temperature versus time
76Heat of Vaporization
Heat of Fusion
77Example
- Calculate the enthalpy change associated with
converting 1.00 mole of ice -25ºC to water 150ºC
at 1 atm. Specific heat of ice, water, and steam
are 2.09 J/g ºC and 4.184 J/g ºC, 1.84 J/g ºC .
The heat of fusion of ice is 6.01 kJ/mol and heat
of vaporization of water is 40.67 kJ/mol
78Water? vapor 100C q4
Vapor 100C
Water 0C q3
Ice? water 0C q2
ICE -25 C q1
79- 1 mol ice ? 1 mol ice ?1 mol water ? 1
mol water ? 1 mol steam -
- T -25ºC 0ºC 0ºC
100ºC 100ºC - qtotal q1 q2 q3
q4 -
- 1.)q 2.09(18g)(-25-0)
- 2.) q 6.02 KJ/mol (convert heat of fusion)
- 3.) q 4.184(18g) (100-0)
- 4.) q 40.7 KJ/mol (convert heat of
vaporization) - 5. ) q 1.84(18 g) (100-0)
80Critical Stuff
- Critical Temperature The temperature above which
it is impossible to liquefy the gas under study
no matter how high the applied pressure. - Critical Pressure The pressure required to
liquefy a gas as at its critical temperature - NOTE the critical temp of a gas gives an
indication of the strength of the IMF of that
gas. A substance with weak attractive forces
would have a low critical temp.
81Which gas can be liquefied at 25ºC
Gas Critical Temp ºC Critical Pressure atm
Ammonia 132 112
Ethanol 158 78
Argon -186 6
Critical Temp above 25ºC
Critical Temp under 25ºC
82Vapor Pressure (vp)
- Vapor Pressure Pressure exerted by molecules
that have enough energy to escape the surface. - As T ? VP ?evaporation ?
-
- Liquids with high VP are volatile (alcohol
evaporates easily) -
- Liquids that have strong IMF have low vapor
pressures. - (take a lot of energy to overcome IMF so it can
evaporate)
83- At higher temperature more molecules have enough
energy - Higher vapor pressure.
of Molecules
T2
Kinetic energy
84- Liquids with high VP are volatile (alcohol
evaporates easily) -
- Liquids that have strong IMF have low vapor
pressures. - (take a lot of energy to overcome IMF so it can
evaporate)
substance vapor pressure at 25oC
diethyl ether C4H10O 0.7 atm
Bromine Br2 0.3 atm
ethyl alcohol C2H5OH 0.08 atm
Water H2O 0.03 atm
85Evaporation
- Molecules at the surface break away and become
gas. - Only those with enough KE
- escape
- Evaporation is a cooling process.
- It requires heat.
- Endothermic.
86Condensation
- Change from gas to liquid
- Achieves a dynamic equilibrium with vaporization
in a closed system. - What the heck is a dynamic equilibrium?
87Dynamic equilibrium
- When first sealed the molecules gradually escape
the surface of the liquid - As the molecules build up above the liquid some
condense back to a liquid.
88Dynamic equilibrium
- As time goes by the rate of vaporization remains
constant - but the rate of condensation increases because
there are more molecules to condense. - Equilibrium is reached when
- Rate of Vaporization Rate of Condensation
89VP example
In a closed container the number of partials
changing from liquid ? vapor will eventually
equal the number changing form vapor ? liquid.
90Boiling Point
The vapor pressure of the liquid air pressure
above the liquid
Note The normal boiling point of water is 100oC.
The term normal refers to standard pressure or 1
atm, or also 101.3 kPa.
91Boiling Pts. of H2O at Various Elevations
Altitude compared to Sea Level (m) Boiling Point (C)
1609 98.3
177 100.3
92How to make something boil
- Increase the VP of the liquid (heat it) so that
the VP of the liquid is gt that of the atmosphere.
- 2. Lower the atmospheric pressure (pressure
above the liquid)
93Boiling Point
- ? boiling pt by
- ? in IMF
- Or
- ? VP
- At high altitudes (low air pressure) water boils
at a lower temperature
94Normal Boiling Point
- Temperature at which something boils when the vp
1 atm - Note the lower the external pressure the lower
the boiling point.
95Freezing point/melting point
- They are the same but in opposite directions.
- When heated the particles vibrate more rapidly
until they shake themselves free of each other. - Ionic solids have strong intermolecular forces so
a high mp. - Covalent/molecular solids have weak
intermolecular forces so a low mp.
96Phase Diagram
- A graphical way to summarize the conditions under
which equilibrium exists between different states
of matter. - Allows you to predict the phase of a substance
that is stable at a given temperature and pressure
97Triple point three phase are in equilibrium
with each other at the same time
1 atm
Melting Point
Critical point
Boiling Point
98Critical Point The temp beyond which the
,molecules of a substance have to much kinetic
energy to stick together to form a liquid.
99Water
Not Water
Phase diagrams of substances other than water the
slope of the solid liquid line slopes forward.
(positive) In water the slope of the
solid-liquid lines slopes downward. (negative)
100Homework