1 Ch 11 Intermolecular Attractions and The Properties of Liquids and Solids Brady & Senese, 5th Ed.
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Transcript of 1 Ch 11 Intermolecular Attractions and The Properties of Liquids and Solids Brady & Senese, 5th Ed.
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1
Ch 11 Intermolecular Attractions and The Properties of Liquids and Solids
Brady & Senese, 5th Ed.
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Index
11.1. Gases, liquids, and solids differ because intermolecular forces depend on the distances between molecules11.2. Intermolecular attractions involve electrical charges11.3. Intermolecular forces and tightness of packing affect the physical properties of liquids and solids11.4. Changes of state lead to dynamic equilibria11.5. Vapor pressures of liquids and solids are controlled by temperature and intermolecular attractions11.6. Boiling occurs when a liquid's vapor pressure equals atmospheric pressure11.7. Energy changes occur during changes of state11.8. Changes in a dynamic equilibrium can be analyzed using Le Châtelier's principle11.9 Crystalline solids have an ordered internal structure11.10 X-Ray diffraction is used to study crystal structures11.11 Physical properties of solids are related to their crystal types11.12. Phase diagrams graphically represent pressure-temperature relationships
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11.1. Gases, liquids, and solids differ because intermolecular forces depend on the distances between molecules
3
States Of Matter
• Condensed phases have greater interaction between particles
• Conversion to less condensed phases requires energy
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11.2. Intermolecular attractions involve electrical charges 4
Electronegativity Review
Electronegativity: A measure of the attractive force that one atom in a covalent bond has for the electrons of the bond
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11.2. Intermolecular attractions involve electrical charges 5
Bond Dipoles
• Two atoms with different electronegativity values will share electrons unequally
• Electron density is uneven, with a higher charge concentration around the more electronegative atom
• Bond dipoles are indicated with delta (δ) notation that indicates that a partial charge has arisen
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11.2. Intermolecular attractions involve electrical charges 6
Net Dipoles
• Symmetrical molecules, even if they have polar bonds, are non-polar because the bond dipoles cancel
• Asymmetrical molecules are polar because the bond dipoles do not cancel-- they have permanent, net dipoles
• Molecular dipoles cause molecules to interact, and the distance between the molecules increases the amount of interaction
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11.2. Intermolecular attractions involve electrical charges 7
Dipole - Dipole Attractions
Interaction between the net dipoles in polar molecules are called dipole-dipole attractions. About 1% as strong as a covalent bond Decrease as molecular distance increases
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11.2. Intermolecular attractions involve electrical charges 8
Dipole-Dipole Interactions (H-bonding)
• Especially strong attractions found in polar molecules with hydrogen atoms bonded to F,O, or N
• These strong dipole-dipole interactions are called hydrogen bonds
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11.2. Intermolecular attractions involve electrical charges 9
Case Study: Snowflakes
What accounts for the fact the that snowflakes are always 6-sided?
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Your Turn!
Which of the following is not likely to have hydrogen bonding? Hint- sketch the molecules!
A. CH3CO2H
B. CH3OH
C. CH3O2CH3
D. All of these exhibit hydrogen bonding
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11
Your Turn!
Which statements are true of a molecule that exhibits dipole-dipole attractions?
i. It is polar
ii. It is asymmetrical
iii. It must contain H attached to F, O, or N
A. i & ii only
B. ii & iii only
C. i only
D. ii only
E. iii only
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11.2. Intermolecular attractions involve electrical charges 12
London Forces
• When atoms near one another, their valence electrons interact
• Repulsion causes the electron clouds in each to distort and polarize
• Instantaneous, induced dipoles result from this distortion
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11.2. Intermolecular attractions involve electrical charges 13
London Forces (con.)
• effect enhanced with increased particle mass
• effect diminished by increased distance between particles
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11.2. Intermolecular attractions involve electrical charges 14
Shape & Attraction
• Longer atom chains exert greater the attraction.• Compact molecules have fewer sites for interaction• Compact molecules have lower attraction for one
another
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11.2. Intermolecular attractions involve electrical charges 15
Intermolecular Attractions
ionic
charged particles attract one another
polar covalent
dipoles attract one another
non-polar covalent
electrons synchronize
contain H-X bonds
X= F, O, or N
do not contain H-X bonds
X= F, O, or N
ionic attractions
and London forces
hydrogen bonding
and London forces
dipole-dipole attractions
and London forces
London forces
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11.2. Intermolecular attractions involve electrical charges 16
Compounds of Different Polarity Interact
• Ionic compounds interact with polar molecules using ion-dipole forces
• Ionic compounds can induce dipoles in non-polar molecules, resulting in ion-induced dipoles
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Your Turn!
Which of the following is not a true statement
A. All molecules exert London Forces
B. Only non-polar molecules exert London Forces
C. The greater the mass the greater the London Force
D. Another name for London Forces in “induced dipole attraction”
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11.2. Intermolecular attractions involve electrical charges 18
Effects of Intermolecular Forces (IMF)
Attractive
Force
MP/ BP/ density/ surface tension/ viscosity
Vapor Pressure/ Rate of Evaporation
Increasing IMF Increasing Decreasing
Increasing molar mass
Increasing Decreasing
Increasing branching
Increasing Decreasing
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11.2. Intermolecular attractions involve electrical charges 19
Learning Check
Identify the kinds of intermolecular forces present in the following compounds and then rank them in order of increasing boiling point: H2S, CH3OH, CBr4, and Ne
C Br
Br
Br
Br
HS
HC O
H
HH
HNe
hydrogenbonding
London forces
London forces
dipole-dipole
MM=331.63 MM=20.18•CH3OH >H2S> CBr4> Ne
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Your Turn!
Arrange the following in terms of increasing strength of intermolecular forces: CH4, CO2, HF
A. CH4>CO2>HF
B. CO2>HF>CH4
C. HF>CH4>CO2
D. None of theseHF>CO2>CH4
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11.3. Intermolecular forces and tightness of packing affect the physical properties of liquids and solids
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Packing Effects
• Gases have high compressibility due to the spaces between particles
• Diffusion is impeded by collisions, hence improved in gases
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11.3. Intermolecular forces and tightness of packing affect the physical properties of liquids and solids
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Surface Tension
surface tension - resistance to spreading out and increasing surface area
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11.3. Intermolecular forces and tightness of packing affect the physical properties of liquids and solids
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Viscosity
• Viscosity - resistance to flow
• larger molecules collide and interact more often, impeding their flow
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11.3. Intermolecular forces and tightness of packing affect the physical properties of liquids and solids
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Wetting
• Wetting - the ability of a liquid to spread across a surface
• The greater the similarity in attractive forces between the liquid and the surface, the greater the wetting effect
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11.3. Intermolecular forces and tightness of packing affect the physical properties of liquids and solids
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Solubility
• “Like dissolve like”- the more similar the polarity of two substances, the greater their ability to interact with each other rather
• Surfactants substances that have polar and non-polar
characteristics improve a liquid’s wetting properties allow non-polar substances to dissolve in polar
solvents
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Your Turn!
Which of the following are not expected to be soluble in water?
A. HF
B. CH4
C. CH3OH
D. All are soluble
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11.3. Intermolecular forces and tightness of packing affect the physical properties of liquids and solids
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Changes of state involve IMF
Gas State
Liquid State
Solid State
sublimation
evaporation
meltingfusion
condensation
deposition
Overcoming or establishing intermolecular attractions results in a change in energy
Exothermic, releases heat Endothermic, absorbs heat
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11.3. Intermolecular forces and tightness of packing affect the physical properties of liquids and solids
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Rates Of Evaporation Vary With Temperature
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11.4. Changes of state lead to dynamic equilibria 29
Changes Of State Involve Equilibria
• The fraction of molecules in the condensed state is higher when the intermolecular attractions are higher
• Intermolecular attractions must be overcome to separate the particles, while the separated particles are simultaneously attracted to one another
separated phasecondensed phase
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11.4. Changes of state lead to dynamic equilibria 30
Boiling Involves Equilibrium
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11.4. Changes of state lead to dynamic equilibria 31
Melting Involves Equilibrium• As solid melts, particles separate and liquid forms• liquid particles simultaneously interact to reform solid• At equilibrium, the rates of both processes are equal• A similar equilibrium exists when liquid vaporizes and gas simultaneously
condenses
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11.5. Vapor pressures of liquids and solids are controlled by temperature and intermolecular attractions
32
Vapor Pressure
• vapor pressure-the pressure of a gas in a closed container in which the solid and liquid phases are at equilibrium
• liquid + heat of vaporization ↔ gas
• Varies with the temperature and substance
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11.6. Boiling occurs when a liquid's vapor pressure equals atmospheric pressure 33
Boiling Point
• The temperature at which vapor pressure equals the atmospheric pressure, is called the boiling point
• Normal boiling point specifies that the atmospheric pressure achieved is 1 atm
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11.6. Boiling occurs when a liquid's vapor pressure equals atmospheric pressure 34
Case Study: The “Love Meter”
A liquid is trapped in a closed container. When the bottom bulb is held in the hand, the liquid rises up the tube into the upper bulb and appears to boil.
• Is the liquid boiling?
• What causes the change?
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11.6. Boiling occurs when a liquid's vapor pressure equals atmospheric pressure 35
Case Study : Love On The Rocks
When the love meter is placed onto a table and a cube of ice is applied to the top bulb, the liquid appears to boil as before.
• Is this boiling?
• Why does this happen?
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11.6. Boiling occurs when a liquid's vapor pressure equals atmospheric pressure 36
Intermolecular Attractions Affect Boiling Point
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37
Your Turn!
Which of the following will affect the boiling point of a substance?
A. Molecular mass of the material
B. Intermolecular attractions
C. The external pressure on the material
D. All of these
E. None of these
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11.7. Energy changes occur during changes of state 38
Heating Curve
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11.7. Energy changes occur during changes of state 39
Cooling Curve
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11.7. Energy changes occur during changes of state 40
Changes Of State Involve Energy
Gas State
Liquid State
Solid State
Enthalpy of sublimationEnthalpy
of fusion
Enthalpy of Vaporization
• Phase changes involve specific energies• Energy is released when a more condensed phase
is formed
(-)Exothermic, releases heat (+)Endothermic, absorbs heat
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11.8. Changes in a dynamic equilibrium can be analyzed using Le Châtelier's principle 41
LeChâtelier’s Principle
• Any change to the quantities or conditions may cause the system to “shift” to re-establish equilibrium
• Increasing heat causes all endothermic processes to shift to form more product
• Increases in pressure results in a shift toward the more condensed phase
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11.9 Crystalline solids have an ordered internal structure 42
Crystal Structures Have Regular Patterns
• The smallest segment that repeats regularly is a unit cell• The unit cell repeats to form a regular highly symmetrical lattice• Square lattices have symmetry in 2 dimensions• Cubic lattices have symmetry is 3 dimensions
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11.9 Crystalline solids have an ordered internal structure 43
Three Types Of Unit Cells
• Crystals take common unit cell shapes: Simple cubic Body-centered cubic (BCC) Face- centered cubic (FCC)
• Type of unit cell depends on size of atoms involved
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11.9 Crystalline solids have an ordered internal structure 44
Simple Cubic
• Has one host atom at each corner edge length (a) = 2r where r is the radius of the atom or ion.
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11.9 Crystalline solids have an ordered internal structure 45
Body-centered Cubic (BCC)
• has one atom at each corner and one in the center edge length
3
4ra
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11.9 Crystalline solids have an ordered internal structure 46
Face- centered Cubic (FCC)
has one atom centered in each face, and one at each corner
ra 22
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11.9 Crystalline solids have an ordered internal structure 47
Spaces In Ionic Solids Are Filled With Counter Ions• In NaCl, the Cl- ions form a
unit cell that is face centered cubic
• Na+ ions, being smaller, fill the spaces between the Cl- ions
• If we count the atoms in the unit cell we have 6 of each, thus a 1:1 Na+:Cl- ratio
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11.9 Crystalline solids have an ordered internal structure 48
Learning Check: Identify the Formula for Each, Based on the Unit Cell.
CaF2CsF ZnS
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11.9 Crystalline solids have an ordered internal structure 49
Your Turn!
Which formula would be consistent with the face centered cubic unit cell shown?
A. GB4
B. GB
C. G4B16
D. None of these
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11.9 Crystalline solids have an ordered internal structure 50
Some Factors affecting crystalline structure
• the size of the atoms or ions involved• the stoichiometry of the salt.• the materials involved (some substances do not
form crystalline solids)
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11.9 Crystalline solids have an ordered internal structure 51
Unit Cells Stack To Form Crystal Lattice
• Square lattice- 2-dimensional arrays• Close packed structures: atoms are layered so
that the holes in one layer are covered by the atoms in the next layer cubic closest packed -3 alternating layers hexagonal close packed-2 alternating layers
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11.9 Crystalline solids have an ordered internal structure 52
Amorphous Solids (Glass)
• Have little order, thus referred to as “super cooled liquids”
• Edges are not clean, but ragged due to the lack of order
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11.9 Crystalline solids have an ordered internal structure 53
Your Turn!
Which factor does not affect the type of crystal
A. the radius of the ions
B. the ratio of the numbers of anions : cations
C. the molar mass of the ions
D. all affect the crystal
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11.10 X-ray diffraction is used to study crystal structures 54
X-Ray Crystallography
• X-rays are passed through a crystalline solid.
• Some x-rays are absorbed, most re-emitted in all directions
• Some emissions by the atoms are in phase, others out of phase
• Emission is recorded on film.
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11.10 X-ray diffraction is used to study crystal structures 55
X-ray Diffraction
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11.10 X-ray diffraction is used to study crystal structures 56
Interpreting Diffraction Data
• helps us estimate atomic and ionic radii
• nλ=2dsinθ n = an integer (1, 2, …) = wavelength of the
X–rays d = the interplane
spacing in the crystal = the angle of
incidence =the angle of reflectance of X–rays to the crystal planes
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11.10 X-ray diffraction is used to study crystal structures 57
Learning Check: Diffraction Data
The diffraction pattern of copper metal was measured with x-ray radiation of wavelength of 1.315Å. The first order (n=1) Bragg diffraction peak was found at an angle theta of 50.5 degrees. Calculate the spacing between the diffracting planes in the copper metal.
1(1.315 Ǻ)=2×d×sin(50.5°)nλ=2dsinθ
2.83 Ǻ =d
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11.10 X-ray diffraction is used to study crystal structures 58
Learning Check
Silver packs together in a faced center cubic fashion. The interplanar distance, d, corresponds to the length of a side of the unit cell, and is 4.07 angstroms. What is the radius of a silver atom?
ra 22r22A07.4
r = 0.536 angstroms
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11.11 Physical properties of solids are related to their crystal types 59
Ionic Crystals (ex. NaCl, NaNO3)
• Cations and anions found at lattice sites
• Relatively hard and brittle
• high melting points• Have strong attractive
forces between ions • Do not conduct
electricity in their solid states
• Conduct electricity well when molten
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11.11 Physical properties of solids are related to their crystal types 60
Molecular Crystals (ex. H2O, CO2)
• Lattice sites occupied either by atoms or by molecules
• If the molecules are relatively small, the crystals tend to: be soft have low melting points be held together by relatively weak intermolecular
attractions
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11.11 Physical properties of solids are related to their crystal types 61
Covalent Crystals
• Lattice positions occupied by atoms that are covalently bonded to other atoms at neighboring lattice sites
• Called network solids - interlocking network of covalent bonds extending all directions
• Covalent crystals tend to: be very hard have very high melting points have strong attractions between covalently bonded
atoms. • Examples: are quartz (SiO2) , silicon carbide (SiC)
and diamond
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13.3. Physical properties of solids are related to their crystal types 62
Metallic Crystals• Simplest models: lattice positions of a
metallic crystal occupied by positive ions
• Conduct heat and electricity well “cloud” of valence electrons surrounds
ions moving electrons transmit kinetic energy
rapidly through the solid
• lustrous When light shines on the metal, loosely
held electrons vibrate easily and re-emit the light with essentially the same frequency and intensity
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11.11 Physical properties of solids are related to their crystal types 63
Learning Check: Classify the following in terms of likely structure
ionic molecular covalent metallic
Sulfur, a substance that pulverizes when struck, is non-conductive of heat and electricity
Substance X. A white crystalline solid that conducts electrical current when molten or dissolved.
Z: shiny, conductive malleable with a high melting temperature.
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11.11 Physical properties of solids are related to their crystal types 64
Your Turn!
molecular crystals can contain all of the listed attraction forces except:
A. dipole-dipole attractions
B. electrostatic forces
C. London forces
D. hydrogen bonding
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11.12. Phase diagrams graphically represent pressure-temperature relationships 65
Phase Diagrams
• Show the effects of both pressure and temperature on phase changes
• Boundaries between phases indicate equilibrium.• Triple point: the temperature and pressure at
which s,l, and g are all at equilibrium• Critical point: the temperature and pressure at
which a gas can no longer be condensed
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11.12. Phase diagrams graphically represent pressure-temperature relationships 66
Phase Diagram of Water
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11.12. Phase diagrams graphically represent pressure-temperature relationships 67
Case Study: An Ice Necklace
• A cube of ice may be suspended on a string simply by pressing the string into the ice cube. As the string is pressed onto the surface, it becomes embedded into the ice.
• Why does this happen?
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11.12. Phase diagrams graphically represent pressure-temperature relationships 68
Phase Diagram – CO2
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11.12. Phase diagrams graphically represent pressure-temperature relationships 69
Your Turn!
At 89 °C and 760 mmHg, what physical state is present?
A. Solid
B. Liquid
C. Gas
D. Supercritical fluid
E. Not enough information is given