Liquids and solids Liquids and solids
Dec 26, 2015
Liquids and solidsLiquids and solids
compared to gases. They are incompressible. Their density doesn’t change
with temperature. These similarities are due
◦to the molecules being close together in solids and liquids
◦and far apart in gases What holds them close
together?
Inside molecules (intramolecular) the atoms are bonded to each other.
Intramolecular forces are covalent, ionic and metallic bonds
Intermolecular refers to the forces between the molecules...these may also involve those listed above.
These are what hold the molecules together in the condensed states.
Strong◦ covalent bonding◦ ionic bonding
Weak◦ Dipole-dipole◦ London dispersion forces
During phase changes the molecules stay intact.
Phase changes involve energy input or release…
Energy used to overcome I.M. forces.
Remember where the polar definition came from?
Molecules line up in the presence of a electric field. The opposite ends of the dipole can attract each other so the molecules stay close together.
1% as strong as covalent bonds. Weaker with greater distance. Small role in gases. (Which correction factor in the Van der
Waals Equation?)
Especially strong dipole-dipole forces when H is attached to N, O, or F
These three because-◦ They have high electronegativity.◦ They are small enough to allow close approach
of the dipoles. Affects boiling point, melting point, and
other colligative properties (freezing point
depression, boiling point elevation and osmotic pressure).
CH4
SiH4
GeH4SnH4
PH3
NH3 SbH3
AsH3
H2O
H2SH2Se
H2Te
HF
HI
HBrHCl
Boiling Points
0ºC
100
-100
200
+-
+
Ionic substances are soluble in water due to the high degree of polarity.
Ionic solubility in other solvents depends upon the solvent polarity.
Examine the solubility of an ionic compound K2SO4 in water vs. isopropyl alcohol.
70% isopropyl alcohol
Saturated K2SO4 soln.
Non-polar molecules also exert forces on each other.
Otherwise, nonpolar substances could not exist as solids or liquids.
Electrons are not evenly distributed at every instant in time.
Have an instantaneous dipole. Induces a dipole in the atom next to it. Induced dipole-induced dipole
interaction.
H H H HH H H H
+ +
H H H H
+ - +
Weak, short lived. More significant at lower temperatures. Eventually enough to change gases to liquids. More electrons on a particle = more
polarizable. Bigger molecules, higher melting and boiling
points. Much, much weaker than other forces when
compared…but…can be strong! Also called Van der Waal’s forces.
Many of the properties due to internal attraction of atoms.◦ Beading◦ Surface tension ◦ Capillary action
Stronger intermolecular forces cause each of these to increase.
Molecules in the middle are attracted in all directions.
Molecules at the the top are only pulled inside.
Minimizes surface area.
Liquids spontaneously rise in a narrow tube. Intermolecular forces are cohesive,
connecting like things. Adhesive forces connect to something else. Glass is polar. It adhesively attracts water molecules.
If a polar substance is placed on a non-polar surface (think of water beading up on a waxed car).◦ There are cohesive,◦ But no adhesive forces.
How much a liquid resists flowing. Large forces, more viscous. Large molecules can get tangled up, which
increases viscosity (corn syrup). Cyclohexane has a lower viscosity than
hexane (examine their structures). Because it is a circle-more compact.
◦ Hydrogen bonding◦ Polar bonding◦ LDF◦ Compare iodine, water and oil…
The phase of a substance is determined by three things.
The temperature. The pressure. The strength of intermolecular
forces.
Two major types.Amorphous- those with much disorder in their structure.
Crystalline- have a regular arrangement of components in their structure.
Lattice- a three dimensional grid that describes the locations of the pieces in a crystalline solid.
Unit Cell-The smallest repeating unit in of the lattice.
Three common types.
There are many amorphous solids.◦ Like glass.
We tend to focus on crystalline solids. Two Types.
◦ Ionic solids have ions at the lattice points.◦ Molecular solids have molecules.
Sugar vs. Salt.
Using diffraction patterns to identify crystal structures.
Talks about metals and the closest packing model.
It is interesting, but trivial…unless you are a crystallographer or geologist!
We need to focus on metallic bonding. Why do metal atoms stay together. How the bonding effects their properties.
1s
2s2p
3s
3pFilled Molecular Orbitals
Empty Molecular Orbitals
Magnesium Atoms
Filled Molecular OrbitalsEmpty Molecular Orbitals
The 1s, 2s, and 2p electrons are close to nucleus, so they are not able to move around.
1s
2s2p
3s
3p
Magnesium Atoms
Filled Molecular OrbitalsEmpty Molecular Orbitals
1s
2s2p
3s
3p
Magnesium Atoms
The 3s and 3p orbitals overlap and form molecular orbitals.
Filled Molecular OrbitalsEmpty Molecular Orbitals
1s
2s2p
3s
3p
Magnesium Atoms
Electrons in these energy levels can travel freely throughout the crystal.
Filled Molecular OrbitalsEmpty Molecular Orbitals
1s
2s2p
3s
3p
Magnesium Atoms
This makes metals thermal and electrical conductors…
and malleable because the bonds are flexible.
There are three types of solid carbon molecules (called allotropes)
Amorphous- coal…uninteresting. Diamond- hardest natural substance on
earth; electrical and thermal insulator.
Graphite- slippery, conducts electricity.
How the atoms in these network solids are connected explains these phenomena.
Carbon atoms are locked into tetrahedral shape.
Strong bonds give the huge molecule its hardness.
The space between orbitals make it impossible for electrons to move around
Empty MOs
Filled MOs
E
Each carbon is connected to three other carbons and
sp2 hybridized. The molecule is flat
with 120º angles in fused 6 member rings.
The bonds extend above and below the plane.
Electrons are free to move throughout these delocalized orbitals.
The layers slide by each other. However, he bonding within the layer is very
strong CARBON FIBER!!!
Molecules occupy the corners of the lattices. Different molecules have different forces
between them. These forces depend on the size of the
molecule. They also depend on the strength and nature
of dipole moments. Molecular solids with dipoles are typically in
a condensed state of matter at room temp.
Dipole-dipole forces are generally stronger than L.D.F.
Hydrogen bonding is stronger than Dipole-dipole forces.
No matter how strong the intermolecular force, it is always much, much weaker than the forces in bonds.
Stronger forces lead to higher melting and boiling points.
What about vapor pressure? Examples?
Most are gases at 25ºC. The only forces are London Dispersion
Forces. These depend on the size of the particle. Large molecules (such as I2 ) can be solids
even without dipoles. Examine the phases of the halogens!
Each molecule has two polar O-H bonds.
HO
H
-
Each molecule has two polar O-H bonds.
Each molecule has two lone pairs of electrons on its oxygen.
HO
H
Each molecule has two polar O-H bonds.
Each molecule has two lone pairs on its oxygen.
Each oxygen can interact with 4 hydrogen atoms.
UNIQUE!
HO
H
This gives water an especially high melting and boiling point.
HO
H
HO
H
HO
H
The extremes in dipole-dipole forces…atoms are actually held together by opposite charges.
VERY HIGH melting and boiling points. Atoms are locked in lattice, making the
substance hard and brittle. Every electron is accounted for so they are
poor conductors-good insulators. Molten ionic substances will conduct (free
moving ions or e- = conductors)
Vaporization - change from liquid to gas at boiling point.
Evaporation - change from liquid to gas below boiling point
Heat (or Enthalpy) of Vaporization
(Hvap )- the energy required to
vaporize 1 mol at 1 atm.
Vaporization is an endothermic process - it requires heat.
Energy is required to overcome intermolecular forces.
Responsible for cool earth. Why we sweat (boys) or perspire (girls)!
Change from gas to liquid. Always exothermic! Achieves a dynamic equilibrium
with vaporization in a closed system. A closed system means matter can’t
go in or out. What the heck is a “dynamic
equilibrium?”
When first sealed the molecules gradually escape the surface of the liquid.
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.
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.
As time goes by the rate of vaporization remains constant but the rate of condensation increases because there are more molecules to condense.
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.
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
Molecules are constantly changing phase
“Dynamic” The total amount of liquid and vapor
remains constant “Equilibrium”
The pressure above the liquid at equilibrium.
Liquids with high vapor pressures evaporate easily. They are called volatile.
Decreases with increasing intermolecular forces. ◦ Bigger molecules (bigger LDF)◦ More polar molecules (dipole-dipole)
Increases with increasing temperature. Easily measured in a barometer. Diagramed in the next series of slides…
Dish of Hg
Vacuum
Patm=
760 torr
A barometer will hold a column of mercury 760 mm high at one atm
Dish of Hg
Vacuum
Patm=
760 torr
A barometer will hold a column of mercury 760 mm high at one atm.
If we inject a volatile liquid in the barometer it will rise to the top of the mercury.
Dish of Hg
Patm=
760 torr
A barometer will hold a column of mercury 760 mm high at one atm.
If we inject a volatile liquid in the barometer it will rise to the top of the mercury.
There it will vaporize and push the column of mercury down.
Water
Dish of Hg
736 mm Hg
Water Vapor
The mercury is pushed down by the vapor pressure.
Patm = PHg + Pvap
Patm - PHg = Pvap
760 - 736 = 24 torr
Kinetic energy
# of
mol
ecu
les
T1
Energy needed to overcome intermolecular forces
Kinetic energy
# of
mol
ecu
les
T1
Energy needed to overcome intermolecular forces
T1
T2
At higher temperature more molecules have enough energy - higher vapor pressure.
Energy needed to overcome intermolecular forces
The graph of temperature versus heat applied is called a heating curve.
The temperature a solid turns to a liquid is the melting point.
The energy required to accomplish this change is called the Heat (or Enthalpy) of Fusion Hfus
-40
-20
0
20
40
60
80
100
120
140
0 40 120 220 760 800
Heating Curve for Water
IceWater and Ice
Water
Water and Steam
Steam
-40
-20
0
20
40
60
80
100
120
140
0 40 120 220 760 800
Heating Curve for Water
Heat of Fusion
Heat ofVaporization
Slope is Heat Capacity
Melting point is determined by the vapor pressure of the solid and the liquid.
At the melting point the vapor pressure of the solid = vapor pressure of the liquid
Solid Water
Liquid Water
Water Vapor Vapor
Solid Water
Liquid Water
Water Vapor Vapor
If the vapor pressure of the solid is higher than that of the liquid the solid will release molecules to achieve equilibrium.
Solid Water
Liquid Water
Water Vapor Vapor
While the molecules of condense to a liquid.
This can only happen if the temperature is above the freezing point since solid is turning to liquid.
Solid Water
Liquid Water
Water Vapor Vapor
If the vapor pressure of the liquid is higher than that of the solid, the liquid will release molecules to achieve equilibrium.
Solid Water
Liquid Water
Water Vapor Vapor
Solid Water
Liquid Water
Water Vapor Vapor
While the molecules condense to a solid.
The temperature must be below the freezing point since the liquid is turning to a solid.
Solid Water
Liquid Water
Water Vapor Vapor
If the vapor pressure of the solid and liquid are equal, the solid and liquid are vaporizing and condensing at the same rate…called The Melting point!
Solid Water
Liquid Water
Water Vapor Vapor
Reached when the vapor pressure equals the external pressure.
Normal boiling point is the boiling point at 1 atm pressure.
Super heating - Heating above the boiling point.
Supercooling - Cooling below the freezing point.
A plot of temperature versus pressure for a closed system, with lines to indicate where there is a phase change.
Water’s triple point diagram is truly an anomaly.
Unfortunately, it is the most widely know substances and familiar to us!
Carbon dioxide’s triple point is much more like that of a “normal” pure substance.
Make sure you are familiar with this fact!
Temperature
SolidLiquid
Gas
1 Atm
AA
BB
CCD
D D
Pre
ssur
e
D
SolidLiquid
Gas
Triple Point
Critical Point
Temperature
Pre
ssur
e
Solid Liquid
Gas
This is the phase diagram for water. The density of liquid water is higher
than solid water.
Temperature
Pre
ssur
e
Solid Liquid
Gas
1 Atm
This is the phase diagram for CO2
The solid is more dense than the liquid The solid sublimes at 1 atm.
Temperature
Pre
ssur
e
For carbon, the space between orbitals makes it impossible for electrons to move around…thus a great insulator! (not graphite)
Empty MOs
Filled MOs
E
Consider Carbon (in comparison to silicon)
In silicon, the gap between orbitals is smaller…making it possible for electrons to cross the gap.
Empty MOs
Filled MOs
EIn fact…like other semiconductors, the conductivity of Si increases with increased temperatures!
(unlike metals)
If you replace some of the Si atoms with some As atoms…you create a n-type semiconductor
As has one more e- than Si…which becomes available for the conduction band.
n-type means a negative type (extra neg. e-) p-type occurs when you replace some Si
atoms with B which has one less valence e- This creates an electron hole (positive area)