Dr. Williamson’s notes on Liquids and Solids CHAPTER 12 and 13
Kinetic-Molecular Theory All gases are composed of tiny particles (atoms/molecules) ! Particles are so small and distance between them so large that the volume of the particles are negligible (zero). ! Particles are in constant motion, colliding with each other and with the walls of the container ! Particles do not attract or repel each other, and collisions are elastic ! The average kinetic energy of the particles is directly proportional to the Kelvin temperature !
LIQUIDS AND SOLIDS
Dr. V. M. Williamson
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Heating Curve
at Constant Pressure
Kinetic-Molecular Description of Phases !
Schematic representation of the three common states of matter.
gas
cool
liquid
heat
cool
solid
heat 3
Solids, Liquids and Gases
A Matter of Forces
Properties Definite shape
Solid Liquid Gas Yes No No Compressibility Nearly Slight High none Density High Intermediate Low Fluid No Yes Yes Very Diffusing Rate Moderate Rapid slow VibraSlow Particle Motion Rapid tional Restricted Medium Far Apart Particle Distance Close Force of Attraction Strong
Medium Negligible
1
!
Electrostatic in nature
!
Intramolecular forces: _________ particles (molecules, atoms or ions), relatively strong, responsible for chemical properties. (Ionic, Polar Covalent, and Nonpolar Covalent)
!
Intermolecular forces: __________ particles (molecules, atoms or ions), relatively weak, responsible for physical properties
Dr. Williamson’s notes on Liquids and Solids Intermolecular Forces (IMF): Secondary Bonding
Intramolecular Forces: Primary Bonding Force
Model
Ionic
+
+
Energy (kJ/mol) 400-4000
Example
Attractive forces responsible for existence of ____________ _________ ! Balance between these forces and kinetic energy of particles determine particular physical state or property ! Stronger forces lead to ________ boiling points, heats of vaporization, heats of fusion ! There are four important intermolecular attractions. !
NaCl
+ +
+
Covalent + (polar/nonpolar)
+
+ + +
Metallic
+ + +
150-1100
HCl O2
75-1000
Cu
+ + + +
Ion-Ion Forces
Intermolecular Attractions and Phase Changes
Predominant in ______ compounds ! Very strong, consequently lead to ____ melting points for ionic compounds !
!
! The
force of attraction between two oppositely charged ions- Governed by
Arrange the following ionic compounds in the expected order of increasing melting and boiling points.
_________ IMF = _________ bp or mp
Coulomb’s Law:
(q )(q ) -
+
F∝
NaF, CaO, CaF2
2
d q and q are the ion charges. d is the distance between the ions. +
IMF⇑ BP⇑MP⇑
You do it! What important points must you consider?
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Melting Points
of Some Ionic Compounds
Intermolecular Attractions and Phase Changes
Na + F- 〈 Ca 2+ F2 − 〈 Ca 2+ O 2-
Cpd NaF NaCl NaBr KCl BaO
Charge First Put in order of increasing melting and boiling points.
MP(oC) 993 801 747 770 1923
Cpd CaF2 Na2S K 2S MgO CaO
MP(oC) 1423 1180 840 2800 2580
NaCl___NaF___MgO charge, size
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Dr. Williamson’s notes on Liquids and Solids Dipole-Dipole Forces
Dipole-Dipole Forces !
Predominant in _____ ________ molecules
!
Result from attraction of δ + end of one particle with δ – end of other particle
!
Effective only at short distances (∝ 1/d4)
Weaker since involve only partial charges (approximately 4 kJ/mol of bonds); about 1% as strong as primary ionic or covalent bonds. ! Molecules tend to align themselves so that the opposite charges are near each other !
Copyright © 1995 by Saunders College Publishing
Dipole-Dipole Forces
Among Polar Molecules
Dipole-Dipole Interactions
The _____ polar the molecule, the ______ is its boiling point.
Hydrogen Bonding
Hydrogen Bonding
Special case of _______________ interaction ! Requirements: !
✻ ✻
Consider H2O which is very polar molecule and has
hydrogen bonding (H bonded to an ___,___,or __)
Hydrogen covalently bonded to a small, highly electronegative element (__________) Electronegative element with lone pair available for hydrogen-bonding
Typically 15-20 kJ/mol (about 5x stronger than other dipole-dipole attractions) ! Responsible for anomalously high boiling points of some compounds !
! (note
covalent primary bond is still stronger that any IMF)
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Dr. Williamson’s notes on Liquids and Solids
Hydrogen Bonding in NH3
Hydrogen Bonding in C2H5OH
Hydrogen Bonding and Boiling Point
Hydrogen Bonding in DNA
Copyright © 1995 by Saunders College Publishing 22
London Forces, Dispersion Forces, Instantaneous Dipoles ! ! ! ! ! !
London Forces
Named for Fritz London (1900-1954) a German physicist-proposed in 1930 The only attractive force for ________ molecules London Forces are very _____ forces present in all particles, polar or nonpolar An uneven distribution of electrons form Depend on particle size or MM (bigger size, greater number of electrons, greater forces) Influenced by molecular shape (bigger surface area, greater forces)
Cpd MM He 4 Ne 20 Ar 40 Kr 84 Xn 131 Rn 222
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BP(K) 4.2 27 87 121 164 211
No dipole-dipole forces, yet these can be liquified, so some IM force exists ! Imagine cooling these down. Rn will liquefy first !
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Dr. Williamson’s notes on Liquids and Solids Factors Affecting London Forces London Dispersion Forces:
An Illustration
Copyright © 1995 by Saunders College Publishing
Factors Affecting London Forces
!
The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane).
!
This is due to the increased surface area in n-pentane.
ionic>hydrogen b.>dipole-dipole>London force ! Ionic- requires ions. Differentiate between Ionic by charge on ions and size ! Hydrogen bonding- requires H bonding to F, O or N (note molecule will be polar due to lone pairs on F,O,N). Differentiate by dipole moment or EN difference. ! Dipole-Dipole- requires polar molecules. Differentiate by dipole moment or EN difference. ! London Forces- requires nonpolar molecules. Differentiate by molar mass
The strength of dispersion forces tends to _______ with _________ molecular weight. ! Larger atoms have larger electron clouds, which are easier to polarize. !
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Ion-Dipole Interactions Ion-dipole interactions are an important force in solutions of ions. ! The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents. !
Dipole-Induced Dipole Forces
Copyright © 1995 by Saunders College Publishing
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Dr. Williamson’s notes on Liquids and Solids The Liquid State
The Liquid State-Viscosity !
!
An example of viscosity of two liquids.
Viscosity is the _________________. " For
example, compare how water pours out of a glass compared to molasses, syrup or honey.
!
IMF__ Vis__
Oil for your car is bought based on this property. "
10W30 or 5W30 describes the viscosity of the oil at high and low temperatures.
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The Liquid State-Surface Tension !
!
Surface Tension Demo
Surface tension is a measure of the unequal attractions that occur at the surface of a liquid.
IMF__ ST__
The molecules at the surface are attracted unevenly. 33
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Property of the Liquid State:
Capillary Action
The Liquid State-Capillary Action !
Capillary action is the ability of a liquid to rise (or fall) in a glass tube or other container
Drawing up of liquid inside narrow tube when adhesive forces between
the liquid and the tube walls exceed cohesive forces within the liquid ! Adhesive forces > cohesive forces: concave meniscus; e.g., water in glass ! Cohesive forces> adhesive forces: convex meniscus; e.g., mercury in glass ! Responsible for intake of water and nutrients by plants from the soil !
IMF__ CA__
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Dr. Williamson’s notes on Liquids and Solids The Liquid State
Evaporation
! Capillary
action also affects the meniscus of liquids.
!
!
glass has polar
S-O bonds that the polar water is attracted to more strongly that to other water molecules
Evaporation is the process in which molecules escape from the surface of a liquid and become a gas. Evaporation is temperature dependent.
! The
Water
Hg 37
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Evaporation of Liquids
Evaporation of Liquids:
Dynamic Equilibrium
Liquid
IMF__ Evap__
Vapor
The rate of condensation is equal to the rate of evaporation
Copyright © 1995 by Saunders College Publishing
Molecules leave liquid surface and enter the gas phase ! In closed system, equilibrium is established !
Copyright © 1995 by Saunders College Publishing
Vapor Pressure
More About Vapor Pressure
Pressure exerted by the vapor in equilibrium with its liquid at a particular temperature ! Constant with constant temperature as long as both liquid and vapor are present ! Independent of container volume ! Obeys Dalton’s Law of Partial Pressures ! Increases with increasing temperature !
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!
Depends on nature of the substance
!
Compounds with weak intermolecular forces have high vapor pressures and low boiling points (volatile)
!
Boiling Point: Temperature at which ________ pressure = _________ pressure
!
Normal Boiling Point: Temperature at which vapor pressure = Pext = __________
Dr. Williamson’s notes on Liquids and Solids Vapor Pressure:
Vapor Pressure Curves
at Different Temperatures
A Butane Lighter
Compare at one temperature IMF⇑ VP⇓
Copyright © 1995 by Saunders College Publishing
The Liquid State !
diethyl ether ethanol water
! !
o
Compound MW(amu) B.P.( C)
Vapor Pressure (torr) and boiling point for three liquids at different temperatures. 0oC
!
The Liquid State
VP⇑ BP⇓
185 12 5
20oC
442 44 18
30oC
647 74 32
CH4
16
-161
36oC
C2H6
30
-88
78oC 100oC
C3H8
44
-42
n-C4H10
58
-0.6
n-C5H12
72
+36
normal boiling point
760 torr = 1 atm, so water has lowest vp What are the intermolecular forces in each of these compounds? Hydrogen bonding > london forces 45
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The Liquid State !
Elevations and Boiling Points
Arrange the following substances in order of increasing boiling points. Ne , NH3, Ar, NaCl, AsH3 You do it!
Boiling at different elevations Boiling is when vapor pressure = external pressure. ! What is the atmospheric pressure on the mountain top? It is less than at sea level. ! So, less vapor pressure needed, less heat needed to make vp, so lower boiling point ! (note: you will have to boil the eggs longer or add salt to get them cooked at this lower bp.) ! !
IMF__ BP__
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Dr. Williamson’s notes on Liquids and Solids The Solid State Normal Melting Point
!
!
Heat of Fusion
The ___________ melting point is the temperature at which the solid melts (liquid and solid in equilibrium) at exactly 1.00 atm of pressure.
!
Heat of fusion is the amount of heat required to melt ___________ of a solid at its melting point at constant temperature. + 334 J !! ! !! → 1.00 g H O at 0o C 1.00 g H 2 O (s) at 0o C ← ! 2 (ℓ) -334 J
The melting point increases as the strength of the intermolecular attractions increase.
• Heat of crystallization is the reverse of the heat of fusion.
IMF__ MP__
IMF__ ∆Hfus__ 49
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Heat of Vaporization
Molar Heat of Fusion or ΔHfusion
Heat of vaporization (ΔHvap): heat needed to convert one gram of liquid at its boiling point to vapor at the same temperature ! Molar heat of vaporization (ΔHvap): heat needed to convert one mole of liquid at its boiling point to vapor with the same temperature ! Specific heat: heat needed to raise temperature of 1 g of substance by one degree Celsius IMF__ ∆H __ !
The molar heat of fusion is the amount of heat required to melt ___________of a substance at its melting point. ! The molar heat of crystallization is the ____________of molar heat of fusion !
+6012 J
!! !! → 1.00 mole H O at 0o C 1.00 mole H 2 O (s) at 0o C ! ←! 2 (ℓ) -6012 J
IMF__ ∆Hfus__
vap
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Phase Diagrams
The Solid State !
Which requires more energy?
Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.
→ NaCl NaCl(s ) ← (ℓ ) or →H O H 2 O (s ) ← 2 (ℓ ) IMF__ ∆Hfus__ ionic IMF> hydrogen bonding IMF So ∆Hfus of NaCl is _____ than ∆Hfus of H2O 53
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Dr. Williamson’s notes on Liquids and Solids Phase Diagrams
Phase Diagrams It ends at the critical point (B); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.
The AB line is the liquid-vapor interface. ! It starts at the ______ point (A), the point at which all three states are in equilibrium. !
Phase Diagrams
Phase Diagrams The AD line is the interface between liquid and solid. ! The ______ _____ at each pressure can be found along this line. !
Each point along this line is the ________ point of the substance at that pressure.
Phase Diagrams
Phase Diagrams (P versus T) !
Below A the substance cannot exist in the __________ state. ! Along the AC line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line. !
Compare water’s phase diagram to carbon dioxide’s phase diagram.
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Dr. Williamson’s notes on Liquids and Solids Amorphous vs Crystalline Solids
SiO2
Types of Solids !
!
Classification based on morphology:
Crystalline and amorphous
crystalline
Classification based on bonding forces:
Ionic, molecular, covalent and metallic amorphous
Structure of Crystals
Amorphous vs Crystalline Solids Crystalline ! Have periodic, orderly arrangement of atoms ! Sharp melting points ! May be anisotropic; cleavage gives planar faces ! display X-ray diffraction patterns which reflect the molecular structure
E.g., sugar, salt, ice
!
Amorphous ! No well-defined, long-range ordered structure ! Broad melting point range ! Properties generally isotropic
E.g., glass, rubber, plastic, waxes
" As
an analogy, bricks are repeating units for buildings.
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Ionic Solids
Structure of Crystals
Held by ______ bonding forces ! Crystalline arrangement of _____ o ! High melting points (400-3000 C) ! Poor electrical and thermal conductors (become good conductors in molten state) ! Examples: CsCl, NaCl, ZnS !
There are three basic variations of the cubic crystal system. ! Simple cubic unit cells. !
" The
balls represent the positions of atoms, ions, or molecules in a simple cubic unit cell.
! !
Unit cells are the smallest repeating unit of a crystal.
Body-centered unit cells Face-centered unit cells 65
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Dr. Williamson’s notes on Liquids and Solids Molecular Solids
Molecular Solids
● Held
by _____ intermolecular forces ● Molecular Solids have __________ in each of the positions of the unit cell. ● Soft o ● Low melting points (–272 to 400 C) ● Poor electrical and heat conductors ● Examples: CO2, C6H6, C60
Copyright © 1995 by Saunders College Publishing
Covalent Solids
Covalent Bonding
referred to as “_______ ______” ● Atoms held by covalent bonds to each other ● Very hard; high-melting (1200 to 4000oC) ● Poor heat and electrical conductors ● Also
!
Some examples of covalent solids are: ! Diamond,
graphite, SiO2 (sand)
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Metallic Solids Metallic Solids may be thought of as positively charged nuclei surrounded by a sea of electrons. ! The positive ions occupy the crystal lattice positions. ! Soft to very hard ! Wide range of melting points
(-39 to 4000oC) ! High heat and electrical conductivity ! Examples: Li, Cu, Ni, Ag .. (metals)
Crystalline Solids
!
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Dr. Williamson’s notes on Liquids and Solids
ID Type of Bonding in Solids Fe(s) Ar(s) ! Diamond, C(s) ! Ice(s) ! CO2(s) ! NaCl(s) ! NH4NO3(s) ! !
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