Ch. 11– Liquids and Intermolecular Forces
• The forces holding solids and liquids together are
called intermolecular forces.
–Intermolecular forces are much weaker than ionic or
covalent bonds.
Example: 16 kJ/mol to vaporize HCl compared to 431 kJ/mol
to break HCl into its elements.
–When a substance melts or boils, the intermolecular
forces are broken (not the covalent bonds).
Intermolecular Forces
• Boiling point reflects intermolecular force strength.
- A high boiling point indicates strong attractive forces.
- A high melting point also reflects strong attractive forces.
• There are 4 basic types of intermolecular forces…
1) ion-dipole forces
2) dipole-dipole forces
3) London dispersion forces
4) hydrogen bonding…(a special type of dipole force)
Ion-Dipole Forces
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/molvie1.swf
• Interaction between an ion and a dipole.
• Strongest of all intermolecular forces.
Na+ and Clions dissolved
in water.
Dipole-Dipole Forces
• Exist between neutral polar molecules.
• Weaker than ion-dipole forces.
• If two molecules have about the same mass and size, then
dipole-dipole forces increase with increasing polarity.
London Dispersion Forces
http://chemsite.lsrhs.net/FlashMedia/html/dipoleVsLondon.html
• Weakest of all intermolecular forces… London dispersion forces exist
between all molecules!
• A temporary, instantaneous dipole may be created by an uneven distribution
of electrons around the nucleus (nuclei) of an atom (molecule).
• London dispersion forces arise due to the Coulombic interaction of the
temporary dipole with the electron distribution in neighboring atoms and
molecules…(One temporary dipole can create another one!)
“instantaneous dipoles”
London Dispersion Forces
• The larger the molecule (the greater the number of electrons) the
“more polarizable” or the easier it is to create instantaneous dipoles.
• London dispersion forces increase as molecular weight increases.
• The differences in LD strength explains why some halogens are
gases at room temperature while others are liquid or solid.
London Dispersion Forces
• London dispersion forces also depend on the shape of the molecule.
•The greater the surface area available for contact, the greater the
dispersion forces.
• London dispersion forces between spherical nonpolar molecules are
lower than the forces between long nonpolar molecules.
less surface area, less force
Hydrogen Bonding
• A special case of dipole-dipole forces.
•
By experiments, the boiling pts. of compounds with H-F, H-O,
and H-N bonds are abnormally high. The intermolecular forces
are therefore abnormally strong.
•
H-bonding requires…
1) H bonded to a small, highly electronegative element (most
important for compounds of F, O, and N)
2) an unshared electron pair on a nearby small highly
electronegative ion or atom (usually F, O, or N on another
molecule).
Hydrogen Bonding Examples
abnormally high B.P.
1)
polar molecules
2)
nonpolar molecules
3)
Hydrogen Bonding in DNA
Dissolving and IMF’s
“Like Dissolves Like”
• Polar solutes dissolve in
polar solvents (or nonpolar
dissolves in nonpolar.)
Properties of Liquids
• Liquids vary in viscosity.
- Viscosity is the resistance of a liquid to flow.
-The “thicker” the liquid, the more viscosity.
-The stronger the intermolecular forces, the higher the viscosity.
-Typically, as temperature increases, viscosity decreases.
Properties of Liquids
Surface Tension
-Acts as a thin skin.
-The stronger the intermolecular forces,
the greater the surface tension.
Fun Facts: (1) “Soap makes water wetter” by reducing the surface tension.
(2) Soap doesn’t kill germs. Instead, the germs are attracted to the soap suds
and are washed down the drain.
• Normally, oil and water don't mix, so they separate into two different
layers. Soap breaks up the oil into smaller drops, which can mix with
the water.
• It works because soap is made up of molecules with two very different
ends. One end of the soap molecules is polar –
they are hydrophilic. The other end of the soap
molecues is nonpolar - they are hydrophobic.
• The hydrophobic ends of soap molecules attach
to the oil. The hydrophilic ends stick out into
the water. This causes a drop of oil to be
suspended in the water.
• This is how soap cleans your hands - it causes drops of grease and dirt
to be pulled off your hands and suspended in water. These drops are
washed away when you rinse your hands.
Phase Changes & Energy
• Generally, it will take more heat to vaporize a liquid than
to melt a solid… (∆H(vap) > ∆H(fusion) ) Why?
- Every intermolecular bond is broken when vaporizing, but only
some of the intermolecular forces break when melting solids.
Liquefying Gases
• Gases can be liquefied by increasing pressure at some temperature.
- Critical temperature: the highest temperature at which a
substance can remain a liquid regardless of the pressure applied.
- Critical pressure: the pressure needed at the critical temperature.
Notice: As intermolecular attractions increase, critical temp, & pressure increase.
Vapor Pressure and Boiling Liquids
• Some of the molecules on the surface of a liquid have enough
energy to escape the attraction of the liquid. These molecules move
into the gas phase.
• As the number of molecules in the gas phase increases, some of the
gas phase molecules strike the surface and return to the liquid.
•After some time the pressure of the gas will be constant at the vapor
pressure.
• Dynamic Equilibrium: the
point when as many
molecules escape the surface
as strike the surface.
• Vapor pressure is the
pressure exerted when the
liquid and vapor are in
dynamic equilibrium.
Vapor Pressure and Boiling Liquids
• Liquids boil when the external pressure equals the vapor pressure.
• Temperature of the boiling point increases as pressure increases.
• Normal boiling point
is the boiling point at
760 mmHg (1 atm).
• A substance with a
high vapor pressure is
said to be volatile. It
readily evaporates.
•A substance with more
intermolecular forces
has a lower vapor
pressure and therefore a
higher B.P.
(a) Measuring vapor of a liquid by using a simple barometer.
(b) The water vapor pushed the mercury level down.
(c) Diethyl ether shows a higher vapor pressure than water.
Vapor Pressure and Boiling Liquids
• Two ways to get a liquid to boil:
1) increase temperature 2) decrease pressure (vacuum pump)
• Pressure cookers operate at high pressure.
- At high pressure the boiling point of water is higher than
at 1 atm.
- Therefore, there is a higher temperature at which the
food is cooked, reducing the cooking time required.
•Boiling points increase with molecular weight as long (as the
intermolecular forces are similar.)
- Example:
CH4 (m.w. = 16) < C2H6 (m.w. = 30) < C3H8 (m.w. = 44)
[ lowest B.P.]
[highest B.P.]
Pressure
Cooker
Volcanic Vent
Autoclave: Steam Sterilizer Using High Pressure
Phase Diagrams
• Shows the relationship between the 3 phases of matter at various
temperatures and pressures.
Triple Point: All 3 phases of matter at equilibrium.
Critical Point: The highest temperature at which the liquid phase can exist.
Phase Diagrams of H2O and CO2
• Notice the slope of the solid–liquid equilibrium line. This indicates that
water expands when it freezes and CO2 contracts when it freezes.
Phase Diagram of H2O
Regular Temp. & Pressure
Ch. 12.1-12.7—Solids and Modern Materials
• There are 2 general classifications of solids:
-Amorphous: No pattern to the arrangement of particles. Their
melting point is over a wide range of temperatures. They just get
softer and softer when heated. (Examples- glass, plastic, wax)
- Crystalline: Well-ordered, definite arrangement of atoms.
Crystals have a repeated structure and a melting point at a very
narrow range of temperatures. (Examples- metals, H2O, diamond,
table salt)
Crystalline Solids
There are four types of crystalline solid:
- Molecular (formed from molecules) - usually soft
with low melting points and poor conductivity.
- Covalent network (formed from atoms) - very hard
with very high melting points and poor conductivity.
- Ionic (formed form ions) - hard, brittle, high melting
points and poor conductivity.
- Metallic (formed from metal atoms) - soft or hard,
high melting points, good conductivity, malleable and
ductile.
Bonding in Crystalline Solids
•
Metallic bonds are formed from metal nuclei floating in a sea of electrons.
• The valence electrons from the metal atoms are considered to be delocalized
and not associated with any individual atom.
Crystalline Solids
Molecular
Covalent Network
Ionic
Metallic
Structure of Solids
• Unit Cell: The smallest repeating unit in a crystal is a unit cell.
- A unit cell is the smallest unit with all the symmetry of the entire crystal.
•Three-dimensional stacking of unit cells is the crystal lattice.
Unit Cell
Delocalized π- bonding in Molecular Solids
Here’s a random fact… A long-standing assumption in chemistry is that
molecular solids will always be insulators. But in 1970, some molecular
solids that conduct electricity were discovered. Conduction was the
result of delocalized pi bonding in the molecular solids.
Molecular Solids: Polymers
• Molecular solids are usually composed of very large molecules, or
polymers, with important commercial and biological applications.
Commercial Polymers
Biological Polymers
Carbohydrates – sugars, starches, cellulose
Lipids - include fats, phospholipids and steroids.
Biological Polymers
Proteins - composed of amino acid monomers
Biological Polymers
Nucleic Acids - include DNA and RNA.
•
Covalent Network Solids: Carbon
Carbon has 3 allotropes…different crystal structures:
(1) Graphite: “flat layers” of atoms. Each atom of carbon forms three
covalent bonds, leaving one electron in each outer orbital
delocalized, creating multiple "free electrons" within each plane of
carbon. So graphite is electrical conductive within each plane.
(2) Diamond: “tetrahedral blocks” of interlocking atoms
(3) Buckyball: a “soccer ball-shaped sphere” of atoms
buckminsterfullerene
buckyball
diamond
graphite
Covalent Network Solids: Silicon
• Silicon is a covalent network solid and a semiconductor.
• Silicon forms a three-dimensional network similar in geometry to a
diamond.
• Silicon’s conductivity increases as temperature increases.
Semiconductor Doping
• Periodicity can be used to understand why doping with an element
with one extra valence electron converts silicon into an n-type
semiconducting (negative charge carrying) material, while doping
with an element with one less valence electron converts silicon into
a p-type semiconducting (positive charge carrying) material.
• Junctions between n-doped and p-doped materials can be used to
control electron flow, and thereby are the basis of modern
electronics. (Transistors!!)
http://www.youtube.com/watch?v=IcrBqCFLHIY&feature=youtu.be
Transistor
Metallic Solids: Alloys
• Metallic solids are often pure substances, but may also be mixtures
called alloys.
o Nearly all the common uses of iron involve alloy compositions.
o Pure gold is too soft to be used in jewelry, whereas alloys of
gold and copper are quite hard.
• Some properties of alloys can be understood in terms of the size of
the component atoms:
— Interstitial alloys form between atoms of
different radius, where the smaller atoms fill
the interstitial spaces between the larger atoms.
o Steel is an example in which carbon
occupies the spaces between the iron.
o The interstitial atoms make the lattice
more rigid, decreasing malleability and ductility.
• In some cases, alloy formation alters the chemistry of the surface.
o An example is the formation of a chemically inert oxide layer in
stainless steel. (Stainless steel contains 0.4 % C, 18 % Cr, and 1 % Ni.)
Metallic Solids: Alloys
— Substitutional alloys form between atoms of
comparable radius, where one atom substitutes
for the other in the lattice.
o Example: Brass…some copper atoms
are substituted with a different
element, usually zinc.
o The density of the alloy typically lies between those of the
component metals, and the alloy remains malleable and ductile.
• Alloys typically retain a sea of mobile electrons, and so they remain
good conductors.
• As mentioned before, metallic bonding can be represented as an
array of positive metal ions with valence electrons drawn among
them, as if the electrons were moving (i.e., a sea of electrons).
• The electron sea model can be used to explain several properties of
metals, including electrical conductivity, malleability, ductility, and
low volatility.
Interstitial Holes in Ionic Crystals
In a face-centered cubic crystal’s unit cell there is more than one type of “hole”.
If the octahedral holes are filled, the structure above results, with a 1:1 count for the
two types of ions in the salt.
If the tetrahedral holes are filled, a different structure exits. It will have a 2:1 count
for the ions in the salt.
In the figure below, The left shows the structure of NaCl and the right that of CaF2.
4 blue
&
4 gray
8 blue
&
4 gray
(By the way, not all of the “holes” need to be filled up. Ion size determines this.)