Liquids and Solids

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Chapter 10

• Liquids and Solids

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Van 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"

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intER 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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Intermolecular 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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3 Types of van der Waals Forces

• Dipole-Dipole forces
• London Dispersion forces
• Hydrogen bonding

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DIPOLE-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.

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DIPOLE-DIPOLE FORCES
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_



_
_

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DIPOLE-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.

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HYDROGEN 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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FYI
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HYDROGEN BONDING

• These types of bonds are super-humanly strong.
• (4X stronger that diopole dipole)

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HYDROGEN BONDING IN WATER
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HYDROGEN BONDING
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WHY 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.

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IMPORTANCE OF HYDROGEN BONDING

• Are important biologically, in stabilizing
  proteins and keeping DNA together.
• Also explains why ice is less dense than water
  (see text).

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LONDON 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

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LONDON DISPERSION FORCES
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LONDON 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

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See 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.

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INSTANTANEOUS AND INDUCED DIPOLES
Pg 454- 455 in text
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LONDON 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.

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LONDON DISPERSION FORCES

• LD forces are generally the WEAKEST
  intermolecular forces.
• Molecules with more electrons will experience
  more LD forces

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n-pentane vs neopentane

• BP 309.4 K BP 282.7 K
• Same atomic masses different structure

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Generalizations 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.

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Because 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
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Which noble gas element has the lowest boiling
point?

• He
• Ne
• Ar
• Kr
• Xe

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The chemical forces between HCl is/are

• Dispersion
• Covalent bond
• Hydrogen bond
• Dipol-dipole
• Two of the above

All Molecules Have
Not symmetrical Polar
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Consider 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

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Which 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
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ION-DIPOLE FORCES

• Attraction between an ion and the partial charge
  on the end of a polar molecule.

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ANSWER
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• A.Identify the types of bonds in
• Glucose
• Cyclohexane
• B.Glucose is soluble in water but cyclohexane is
  not. Why?

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• 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.

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ION-DIPOLE FORCES
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ION-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)

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ION-DIPOLE FORCES AND THE SOLUTION PROCESS
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Homework

• Pg 504-505
• s 35, 36, 37, 39 (you may need to read 10.1
  for this part esp. LD portion)

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10.2 Properties of liquids

• Liquids are vital to our lives.
• Water is a means of food preparation
• Cooling machines n industrial processes
• Recreation
• Cleaning
• Transportation

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CHARACTERISTICS OF LIQUIDS

• Surface tension
• Capillary action
• Viscosity

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COHESIVE FORCES

• Intermolecular forces that bind like molecules to
  one another (e.g. hydrogen bonding).

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ADHESIVE FORCES

• Intermolecular forces that bind a substance to a
  surface.

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SURFACE 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.

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Surface 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.
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CAPILLARY ACTION

• Another way surface tension manifests.
• The rise of liquids up very narrow tubes. This
  is limited by adhesive and cohesive forces.

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Formation of meniscus

• Water adhesive forces are greater than cohesive
  forces
• Mercury Cohesive are greater than adhesive
  forces.

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VISCOSITY

• 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
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Viscosity 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.

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Homework

• Pg 505
• s 43-45 all

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10.3 Structure of Solids

• Two ways to categorize solids
• Crystalline
• Amorphous

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Crystalline Solid

• Ridged and long range order of its atoms.
• Solids have flat surfaces
• Sharp melting points
• EX Quartz, diamond, sodium Chloride.

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Amorphous Solid

• Lack a well defined arrangement
• No long range order
• IMF vary in strength
• DO NOT have sharp melting points.
• EX rubber, glass

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Unit 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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X-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
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X-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?

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Braggs Equation Demo
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Example

• 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

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The 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

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Homework

• Pg 505
• s 47, 63, 64,

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Changes of state

• Transformation from one state to another

Condensation
Vaporization AKA steam
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Changes in state

• Liquid ? Gas Vaporization Endothermic
• Gas ? Liquid Condensation Exothermic

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• Solid ? Gas Sublimation Endothermic
  
• Gas ? Solid Deposition Exothermic

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• Solid ? Liquid Melting Endothermic
• Liquid ? Solid Freezing Exothermic

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Changes 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)

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Energy 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

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Question

• 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

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• 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.

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Heating Curve

• A plot of the temperature versus time

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Heat of Vaporization
Heat of Fusion
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Example

• 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

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Water? vapor 100C q4
Vapor 100C
Water 0C q3
Ice? water 0C q2
ICE -25 C q1
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• 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)

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Critical 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.

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Which 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
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Vapor 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)

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• At higher temperature more molecules have enough
  energy
• Higher vapor pressure.

of Molecules
T2
Kinetic energy
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• 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
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Evaporation

• Molecules at the surface break away and become
  gas.
• Only those with enough KE
• escape
• Evaporation is a cooling process.
• It requires heat.
• Endothermic.

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Condensation

• Change from gas to liquid
• Achieves a dynamic equilibrium with vaporization
  in a closed system.
• What the heck is a dynamic equilibrium?

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Dynamic 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.

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Dynamic 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

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VP example
In a closed container the number of partials
changing from liquid ? vapor will eventually
equal the number changing form vapor ? liquid.
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Boiling 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.
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Boiling Pts. of H2O at Various Elevations
Altitude compared to Sea Level (m) Boiling Point (C)
1609 98.3
177 100.3
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How 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)

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Boiling Point

• ? boiling pt by
• ? in IMF
• Or
• ? VP
• At high altitudes (low air pressure) water boils
  at a lower temperature

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Normal Boiling Point

• Temperature at which something boils when the vp
  1 atm
• Note the lower the external pressure the lower
  the boiling point.

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Freezing 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.

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Phase 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

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Triple point three phase are in equilibrium
with each other at the same time
1 atm

Melting Point
Critical point
Boiling Point
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Critical Point The temp beyond which the
,molecules of a substance have to much kinetic
energy to stick together to form a liquid.
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Water
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)
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Homework

• Pg 508
• s 85, 87,89, 91,