Transcript Slide 1

Unit 11:
Liquids and Solids
Intermolecular Forces
Intramolecular forces hold atoms together
in a molecule. (Covalent bonds)
Intermolecular forces are attractive forces between molecules
in a substance’s condensed phases (solids/liquids).
Intermolecular vs Intramolecular strength
• 41 kJ to vaporize (boil) 1 mole of water (inter)
• 930 kJ to break all H-O-H bonds in 1 mole of water (intra)
Intermolecular forces are much weaker
than intramolecular forces.
Intermolecular Force Characteristics
•
Weak bonds are readily reversible (form & reform)
•
Vary in strength, geometry, and specificity
•
Greater freedom of interaction than covalent
•
Can be accumulated in groups
to form significantly strong
interactions
Phase changes
Energy is needed to overcome the
intermolecular forces to allow the
molecules to "break free" of each other.
Boiling and Melting point are relative
measures of intermolecular
forces.
Helium has very little intermolecular forces
Boiling point is thus very low (4.22 K, −452.07 °F)
Predict the relative polarity.
Br-Cl
1.85 D
0.52 D
1.90 D
1.60 D
1.15 D
0.0 D
4 Types of Intermolecular Forces
#1 Dipole-Dipole Forces
Attractive forces between polar molecules
(have permanent dipole)
1.86 D; polar
Tb = -78 °C
0.0 D; Non-polar
Tb = -130 °C
Intermolecular Forces
#2 Ion-Dipole Forces
Attractive forces between an ion and a polar molecule
Ion-Dipole Interaction
Acetone and KCl
K+
d-
d+
ClThis type can only be found in a mixture of substances
Hydration: Most common Ion-Dipole force
Na+
d
+
d
H2 O
Cl
+
d
d
H2O forms hydration shell around ions
Intermolecular Forces
#3 London Dispersion Forces: (weakest)
Attractive forces that result from induced temporary dipoles
in non-polar atoms or molecules
Non-polar
(no dipole)
Non-polar
molecules can
become "Polarized"
London Dispersion Forces (Continued)
Instantaneous Dipoles Interacting With Each Other
These interaction patterns exist only momentarily,
new arrangements are formed in the next instant.
Dispersion forces are the primary forces acting
between non-polar molecules.
Dispersion Forces Continued
Polarizability is the ease with which the electron
distribution in the atom or molecule can be distorted.
Polarizability increases with:
• greater number of electrons
• more diffuse electron cloud
(Resonance)
Dispersion forces usually
increase with molar mass
(more electrons).
Phase changes of simple hydrocarbons
Example
What is the dominant type of intermolecular force that
exists between the following pairs?
(a) HBr and H2S
A. Both HBr and H2S are polar molecules. Therefore,
are dipole-dipole forces, as well as London
dispersion forces.
(b) Cl2 and CBr4
B. Both Cl2 and CBr4 are nonpolar, so there are only
dispersion forces between these molecules.
(c) I2 and Na+
C. I2 is a homonuclear diatomic molecule and
therefore nonpolar, so the forces between it and the
Na+ ion are ion-induced dipole dispersion forces .
(d) NH3 and C6H14
D. NH3 is polar, and C6H6 is nonpolar. The force is
dipole-induced dipole forces (dispersion).
#4 Hydrogen Bond (strongest)
The hydrogen bond is a special dipole-dipole interaction
between the hydrogen atom in a polar N-H, O-H, or F-H
bond and a polar O, N, or F atom.
d- ••• d+
H—N—X
X—N
X—O ••• H—O—X
H—F
X—F
Can form between identical molecules
Some can form H-bonds with
other molecules, but not
themselves.
:
Can also form between different molecules
•••
H—F
Hydrogen Bonds
4 H-bonds per
molecule
The 3 most electronegative
elements that H-bond
H-bonding network of water
With Hydrogen-bonds: H2O boils at 100 °C
Regular Dipole-dipole: H2S boils at -60 °C
Only Dispersion forces: CH4 boils at -164 °C
*Similar masses
Why is the hydrogen bond considered a “special”
dipole-dipole interaction?
Do form
Hydrogen bonds
Decreasing molar mass
Decreasing boiling point
Group
6A
5A
7A
4A
Does not form
Hydrogen bonds
Crucial in both DNA and Protein Structure
Hydrogen bonds hold the two DNA strands
together in the double helix
Strand 1
Strand 2
Maintain protein folds
••• Hydrogen bonds
Example
Methylamine
Which can form hydrogen bonds? Pure Substances
dd+
Yes, has N: and polar N-H
Acetone
No, has O: but C-H not polar
enough
Hydrogen
Cyanide
Urea
No, has N: but C-H not polar
enough
dd+
Yes, has O: and polar N-H
Example #2 Which can form hydrogen bonds with H2O?
Mixture of substances
Acetone
Propane
Hydrogen
Sulfide
Hydrogen
Fluoride
Formic Acid
Boron trihydride
Na+
Example
C3H8 , H2S & BH3 do not have N, O, or F and can’t H-bond
The Na+ cation instead forms ion-dipole interactions
The rest do form hydrogen bonds.
Two modes with Formic Acid
Crash Course: Polar & Non-Polar Molecules (H-bonds)
www.youtube.com/watch?v=PVL24HAesnc
Dipole-dipole, dipole-induced dipole, and
dispersion forces are commonly referred
to as van der Waals Forces.
(Omitting H-bonds & forces with ions (electrostatic)
Distances generally need
to be small and requires
close contact of molecules.
van der Waals Forces
PBS NOVA: Making Things Smarter Gecko Adhesive
TedEd: How do geckos defy gravity?
www.youtube.com/watch?v=YeSuQm7KfaE
Forces Hierarchy
Covalent Bonds – Hold atoms together in molecules
300 - 400 kJ/mol
Ionic Bonds – Electrostatic force between charges
600 – 4K kJ/mol in crystal lattice; Drastically reduced in H2O
Hydrogen Bonding – H bonded to N, O, or F
20 – 40 kJ/mol
Dipole-Dipole Force – Polar molecules (2 – 8 kJ/mol)
London Dispersion (Induced-Dipole) – Ion and non-polar
London Dispersion (Instantaneous-Dipole) – non-polar
Less than 1 kJ/mol
How to assign forces and rank relative boiling points?
Highest boiling points
Hydrogen
Bond
Yes
Question 1
or
Is it Polar?
Dipole-Dipole
Question 2a
Yes
Does it have Hbonded to N, O, or F?
No
2nd highest boiling points
No
Larger mass =
higher boiling pts
or
Question 2b
London
Dispersion
Lowest boiling points
How big is it?
(polarizability)
Smaller mass =
lower boiling pts
Practice Problems
a) State the major intermolecular force involved in each
chemical when found in its liquid phase
b) Place the above chemicals in order of decreasing boiling
points based on their types of intermolecular forces
c) Draw a water molecule Hydrogen bonding to HCN
Properties of Liquids: Surface Tension
Surface tension is the energy required to stretch or
increase the surface area of a liquid.
Strong intermolecular forces
result in high surface tension
Properties of Liquids: Cohesion/Adhesion
Cohesion is the attraction between like molecules
Adhesion is an attraction between unlike molecules
Adhesion forming a meniscus
Adhesion >
Cohesion
"capillary action"
Cohesion > Adhesion
Mercury
Properties of Liquids: Density
Density =
mass
volume
Water density is unique:
Most solid phases are more dense
than their liquid phases
benzene
water
Ice is less dense than water,
resulting in rising to the top.
Water’s Max Density 40C
Density decreases
upon freezing
(abnormal)
TedEd: How polarity makes water behave strangely
www.youtube.com/watch?v=ASLUY2U1M-8
H2O density ↑ as
Temp ↓ (normal)
Properties of Liquids: Viscosity
Viscosity: a fluid’s resistance to flow.
As intermolecular forces ↑
the increase viscosity ↑
Lot’s of H-bonding
TedEd: Non-Newtonian Fluids
www.youtube.com/watch?v=KB43fM_ozKQ
Honey
Sucrose: mucho H-bonds
Motor oils are made to have a precise viscosity to
reduce engine friction, but still move the parts.
• Several variants exist for the various climates.
• As temperature ↑
Viscosity ↓
• More energy to break
intermolecular forces.
Crash Course: Liquids
www.youtube.com/watch?v=BqQJPCdmIp8
Properties of Liquids: Solvent ability
Liquids with similar intermolecular forces are likely
to dissolve to them as solutes them into solution.
“like dissolves like”
• Polar molecules are soluble in polar solvents
Sugars in H2O
(H-bonding or Dipole-Dipole)
• Ionic compounds are more soluble in polar solvents
Na+Cl- in H2O or NH3 (l)
(Ion-dipole forces)
• Non-polar molecules are soluble in non-polar solvents
Fats/Oils in C6H6
(London Dispersion forces)
The solute/solvent can both be liquids
Two liquids are said to be miscible if they are
completely soluble in each other.
Immiscible liquids are not soluble in each other.
Olive oil
in water
Acetone
DIY Hydrophobic Coating:
www.youtube.com/watch?v=OzN8uI5USXM
The Cleansing Action of Soap (Detergents)
Surfactants are amphipathic in that they possess a
non-polar hydrophobic region and a polar
hydrophilic region.
Dissolve to form colloids
Solid Structure
Particles become locked in place
Solids are divided into 2 categories: crystalline & amorphous.
A crystalline solid:
• Possess rigid and long-range order.
• Particles occupy specific repeating positions.
• Maximizes bonding (most stable).
An amorphous solid:
• Lacks long-range order.
• Occurs when freezing is rapid
• Not enough time to find optimal positions
Glass is an amorphous solid
Glass is made by quickly cooling molten salts
Typically SiO2, Na2O, with CaO
The atoms don’t have time to form their preferred
crystalline pattern and are “locked” into random positions.
This is why glass shatters into irregular shapes.
SmarterEveryDay: Prince Rupert's Drop at 130,000 fps
www.youtube.com/watch?v=xe-f4gokRBs
A unit cell is the basic repeating structural unit of
a crystalline solid, which forms a crystal lattice.
Each lattice point represents an atom, molecule, or ion
lattice point
Unit Cell
Unit cells in 3 dimensions
Seven Basic Unit Cells: based on identity of solid
26
3-D Structure of solid Water
Rigid Hexagonal
Structure of ice.
Empty Space within
= lower density
At 4°C, there is H2O
molecules within the
hexagons
TedEd: Why does ice float in water?
https://www.youtube.com/watch?v=UukRgqzk-KE
NaCl is simple cubic: all 90°
Cl-
Na+
Quartz (SiO4)
is hexagonal
A polymorphous substance can form more
than one type of crystal lattice depending on
the conditions (temperature, pressure)
Both CaCO3
orthorhombic
rhombohedral
Allotropes are various forms of a pure element that
adopt different lattice types.
Carbon has the largest number of allotropes
Physical/chemical properties
can vary between allotropes.
Incendiary ammunition
White
Phosphorous
P4
S6
Types of Crystals: Ionic Crystals
•
•
•
•
•
Lattice points occupied by cations and anions
Held together by electrostatic attraction
Hard, brittle, high melting point
Poor conductor of heat and electricity
Ionic compounds form these types
CsCl
ZnS
CaF2
Superconductors
Meissner effect: Magnetic suspension
Can transfer electrical energy
with zero resistance at
critically low temperatures
(40 – 95K).
Royal Institute: Levitating Superconductor on a Möbius strip
www.youtube.com/watch?v=zPqEEZa2Gis
1986
< 91K
MgB2 :
Cheaper and easier to
fabricate superconductive
Japan & South Korea’s
material discovered in
Maglevs can reach 270 mph
2001.
Types of Crystals: Molecular Crystals
•
•
•
•
Lattice points occupied by molecules
Held together by intermolecular forces
Soft, low melting point (dependent on size & forces)
Poor conductor of heat and electricity
Water
H2O(s)
Benzene
C6H6(s)
Tm = 0 °C
Tm = 6 °C
Dry ice
CO2(s)
Table Sugar
C12H22O11
Tm = -78 °C
Tm = 186 °C
Types of Crystals: Covalent Crystals
•
•
•
•
•
Essentially a giant molecule formed by an atomic
network of shared electrons.
Lattice points occupied by atoms
Held together by covalent bonds
Hard, high melting point
Poor conductor of heat and electricity
Carbon Allotropes:
Diamond versus
Graphite
Crash Course: Network Solids and Carbon
www.youtube.com/watch?v=b_SXwfHQ774
ChemMatters: Graphene
www.youtube.com/watch?v=SXmVnHgwOZs
Types of Crystals: Metallic Crystals
•
•
•
•
Lattice points occupied by metal atoms
Held together by metallic bonds
Soft to hard, low to high melting point
Good conductors of heat and electricity
nucleus &
inner shell emobile “sea”
of delocalized e-
Metallic compounds
• Electron sea theory – low ionization energies, ecan easily move from atom to atom
– Explains the hardness of metals – difficult to separate
• Held together by delocalized electrons
– Act like “glue”
Types of Crystals
Crash Course: Solids
www.youtube.com/watch?v=bzr-byiSXlA
X-ray Diffraction Patterns of a Crystal can
reveal structure
X-rays have a wavelength comparable in
magnitude of lattice point lengths.
( = 0.01 - 10 nm: ~500x smaller than visible)
Electrons scatter X-ray waves
Bragg’s law allows us to map electron-density
Named after William H. Bragg and William L. Bragg (father and son, shared Nobel prize)
Reflection of
X-rays from two
layers of Atoms
2d sin θ = n
(Bragg’s law 1913)
Reactions: How Can You
See an Atom?
https://www.youtube.com/w
atch?v=ipzFnGRfsfE
X-ray diffraction is used primarily to determine
biomolecular structures
Highly ordered crystalline samples are needed.
Kendrew and Perutz shared the Nobel Prize for
chemistry in 1962 for determining the first
protein structure using X-ray crystallography
(myoglobin: oxygen binding protein in muscle)
Royal Institute: Celebrating Crystallography
www.youtube.com/watch?v=uqQlwYv8VQI
A phase is a state of matter with definite
physical properties.
1 Substance - 3 Phases
Solid phase - ice
Liquid phase - water
Gas phase - vapor
Phase Changes
Least Order
Random fast movements
of particles
Medium Order
Particles held together,
but can slide past one
another
Greatest Order
Rigid and tightly packed
particles
Effect of Temperature on Kinetic Energy
Low
Evaporation
Raise Temp.
Higher
Evaporation rates
As temp. increases, a larger number of molecules
have enough kinetic energy to enter the gas phase.
This produces vapor. Evaporization occurs below the
boiling point and results in vapor pressure, exerted
force over a liquid of evaporated molecules.
The equilibrium vapor pressure is pressure
condensation and evaporation rates are equal
H2O (l)
H2O (g)
Rate of = Rate of
condensation evaporation
No net
change
Measurement of Vapor Pressure
Before Evaporation
(no vapor pressure)
Equilibrium reached;
Vapor pressure =
Eternal Pressure
Vapor of Water and Temperature
Boiling occurs when vapor pressure
reaches the external pressure
Vaporation rate > Condensation rate
1 atmosphere of pressure
The "Fizz Keeper"
Pumping in air above the liquid
favors CO2 staying dissolved in
solution (b.p. increase)
The Boiling Point is the temperature at which the vapor
pressure is equal to the external pressure.
Normal boiling point is the b.p. temperature at 1 atm of pressure
"Heat of vaporization"
Phase change is dependent on energy input (DH)
Boiling points are pressure-dependent
Cooking is about energy input, not temperature
Higher temps = faster energy transfer
Boiling points rise as pressure increase
Boiling points decrease as pressure decrease
Water boils at ~
212°F at sea level (1 atm)
160°F on top of Mt. Everest
250°F in a pressure cooker/autoclave
Molar heat of fusion (DHfus) is the energy required to
melt 1 mole of a solid substance at its freezing point.
Heats (DH) of fusion are smaller than vaporization
(Takes more energy to boil than melt)
Solid-Gas Equilibrium
CO2 (s) → CO2 (g)
H2O(g) → H2O (s)
SciShow: Snowflakes;
www.youtube.com/watch?v=q
bEw9SobgkA
A phase diagram summarizes the conditions at
which a substance exists as a solid, liquid, or gas.
Melting
curve
Boiling
curve
Freeze-drying
(Lyophilizing):
freezing, then put
under vacuum to
sublime off solvent
Sublimation
Triple point: all three
phases in equilibrium
Phase Diagram of H2O
Effect of Increase in Pressure
Phase Diagram of Carbon Dioxide
At 1 atm
CO2 (s)
CO2 (g)
Naturally sublimes at normal conditions
because triple point is above 1 atm
Sample of Things to Study
• Inter- vs Intra-molecular forces
• Effects on phase changes
• Dipole-Dipole examples
• Ion-Dipole examples (hydration)
• London Dispersion forces examples
• Polarizability effects on L. D.
• Hydrogen bond (drawing them)
• Able to rank forces by strength
• van der Waals forces
• Surface tension
• Cohesion/adhesion/capillary action
• Density (& water’s uniqueness)
• Viscosity (and temp effects)
• Solubility/Miscibility/Detergents
• Crystalline vs amorphous
• Unit cell/ lattice point
• Polymorphous and Allotropes
• Superconductors
• 4 types of crystals and properties
• X-ray diffraction
• Phases – solids/liquids/gas
• 6 types of phase changes
• Evaporation vs boiling
• vapor pressure
• Boiling point definition
• Pressure effects on boiling
• Heats of Vaporization
/Fusion/Sublimation
• Phase diagrams/triple points