liquids, solids and changes of state
DESCRIPTION
Kinetic-molecular view of liquids and solids All real gases can be condensed to liquids by lowering the temperature and increasing the pressure. This decreases the average speed of the molecules. When moving slow enough, they will be attracted to each other and form a liquid. Increased P and TTRANSCRIPT
12 - 1
Liquids, Solids andLiquids, Solids andChanges of StateChanges of State
Kinetic-Molecular View of Liquids and SolidsKinetic-Molecular View of Liquids and SolidsIntermolecular AttractionsIntermolecular Attractions
Properties of LiquidsProperties of LiquidsVapor Pressure and Boiling PointVapor Pressure and Boiling Point
Melting Points and FreezingMelting Points and FreezingHeating and Cooling CurvesHeating and Cooling Curves
Phase DiagramsPhase DiagramsCrystalsCrystals
12 - 2
Kinetic-molecular viewKinetic-molecular viewof liquids and solidsof liquids and solids
All real gases can be condensed to liquids by lowering the temperature and increasing the pressure.
• This decreases the average speed of the molecules.
• When moving slow enough, they will be attracted to each other and form a liquid.
IncreasedP and T
12 - 3
Kinetic-molecular viewKinetic-molecular viewof liquids and solidsof liquids and solids
If the temperature is further decreased:• Molecules can no longer move about
freely.• Motion is limited to vibration.
Rapid temperature decrease results in a disorderly arrangement - amorphousamorphous.
A slow temperature decrease allows molecules to form a crystallinecrystalline solid.
12 - 4
Intermolecular forcesIntermolecular forcesFor molecules to form liquids and solids, there
must be attractions between the them.Intermolecular attractive forces
•dipole-dipole attraction including dipole-dipole attraction including hydrogen bondinghydrogen bonding
•London (dispersion) forcesLondon (dispersion) forces
Relative strengthRelative strengthhydrogen bonding > dipole-dipole > London
12 - 5
Dipole-dipole attractionsDipole-dipole attractions- When electrons that make up a bond are
not equally shared because of a difference in electronegativity.
+ and - ends are attracted to each other.
HCl+-
HCl+-
HCl+-
HCl+-
HCl+-
HCl+-
HCl+-
HCl+-
HCl+- HCl
+-
HCl+-
solid liquid
12 - 6
Hydrogen bondingHydrogen bonding
An unusually strong dipole-dipole attraction.
• Occurs when hydrogen is bound to fluorine, oxygen and nitrogen -- the most electronegative elements.
• The small sizes of the elements involved and the large electronegativity differences result in large + and - values.
• Hydrogen bonds are usually represented using a dashed line.
12 - 7
Hydrogen bondingHydrogen bonding
The hydrogens of one water moleculeinteract with theoxygen on otherwater molecules.
12 - 8
London forcesLondon forcesTemporary dipole attractions that exist
between molecules - also called the dispersivedispersive.
Results from random electron motion.Relatively weak force.
12 - 9
Properties of liquidsProperties of liquids
DiffusionDiffusionThis takes place in both liquids and gases. It is the spontaneous mixing of materials that results from the random motion of molecules.
12 - 10
ViscosityViscosityResistance to flow.This increases with increased intermolecular
attractions.
Also, liquids composed of long, flexible molecules can entwine, resulting in increased viscosity - motor oil.
Properties of liquidsProperties of liquids
CH3CH2CH2
OHCH3CH CH2
OHOHCH2CH CH2
OHOHOH
Increasing viscosity
12 - 11
Properties of liquidsProperties of liquidsSurface TensionSurface Tension
Force in the surface of a liquid that makes the area of the surface as small as possible.
Molecules at thesurface interactonly with neighborsinside the liquid.
12 - 12
Properties of liquidsProperties of liquids
Capillary actionCapillary actionIt is the competition between two forces.
Cohesive forcesCohesive forcesThe attractions between molecules of a substance.
Adhesive forcesAdhesive forcesAttractions between molecules of different substances.
12 - 13
Properties of liquidsProperties of liquidsCapillary action.Capillary action.
MercuryCohesive is largerthan adhesive.
WaterAdhesive is largerthan cohesive.
Capillary tube
meniscus
12 - 14
Properties of liquidsProperties of liquidsVaporizationVaporization
The formation of a gas from a liquid.At any temperature, at least a few of the molecules in the liquid are moving fast enough to escape.
initially equilibrium
after sometime
molecules leave andrenter liquid atthe same rate
12 - 15
EquilibriumEquilibriumA state where the forward and reverse
conditions occur at the same rate.
DynamicEquilibrium
I’m in staticequilibrium.
12 - 16
EquilibriumEquilibrium
A point is ultimatelyreached where therates of the forwardand reverse changesare the same.
At this point, equilibrium is reached.
Rate
Time
12 - 17
Chemical equilibriumChemical equilibrium
A dynamic process on the molecular level achieved when concentration of reactants and products remain constant over time.
- for a physical process:
H2O(l) H2O(s)
(reactant) (product)
- the equilibrium process is indicated with an equilibrium arrows.
12 - 18
EquilibriumEquilibriumCo
ncen
tratio
n
Time
Kinetic EquilibriumRegion Region
12 - 19
Le Chatelier’s principleLe Chatelier’s principle
Any stress placed on an equilibrium system Any stress placed on an equilibrium system will cause the system to shift to minimize will cause the system to shift to minimize the effect of the stress.the effect of the stress.
You can put stress on a system by adding or removing something from one side of a reaction.
N2(g) + 3H2 (g) 2NH3 (g)
What effect will there be if you added moreammonia? How about more nitrogen?
12 - 20
Vapor Pressure and boiling pointVapor Pressure and boiling point
Equilibrium vapor pressureEquilibrium vapor pressureThe pressure of a vapor in equilibrium with a liquid.
It depends on:It depends on:• the intermolecular forces in the liquid.• temperature.
It is independent of:It is independent of:• the volume of the liquid or vapor• the surface area of the liquid
12 - 21
Boiling PointsBoiling Points
Boiling pointBoiling point - temperature where the vapor pressure equals atmospheric pressure.
Boiling point of water at various elevationsCity Elevation BP (oC) San Francisco Sea level 100.0 Salt Lake City 4,390 95.6 Denver 5,280 95.0 La Paz, Bolivia 12,795 91.4 Mount Everest 20,028 76.5
This is the reason that cake mixes includehigh altitude baking instructions.
12 - 22
Boiling pointBoiling pointBoiling points are dependent on pressure.
NormalNormalboiling pointboiling pointThe boilingpoint at standardatmosphericpressure(760 mmHg) 0 50 100
Temperature, oC
Vapo
r pre
ssur
em
mHg
1000
500
0 Norm
al B
P
Standard atmospheric pressure
Vapor pressure of H2O
12 - 23
Melting pointMelting point
Normal melting pointNormal melting pointTemperature at which a solid changes to a liquid at atmospheric pressure.
Freezing point.Freezing point.The temperature at which a liquid changes to a solid.
For the same substance, these will both be at the same temperature.
12 - 24
Changes in stateChanges in stateA substance can usually be converted to
different states by adding or removing energy from a system.If energy must be added, the change is- endothermicendothermic
If energy is given off, the change is- exothermicexothermic
The same concept can also be applied to chemical reactions.
12 - 25
EndothermicEndothermicchanges of statechanges of state
SublimationSublimationThe direct conversion of a solid to a gas.Example - dry ice (solid CO2)
Melting or fusionMelting or fusionThe conversion of a solid to a liquid.Example - melting of ice
Evaporization or vaporizationEvaporization or vaporizationConverting a liquid to a gas.Example - boiling water
Most materials first melt then vaporize as you raise the temperature.
12 - 26
EndothermicEndothermicchanges of statechanges of state
Gas
Solid Liquid
subli
mation
evaporation or
vaporization
melting or fusion
12 - 27
ExothermicExothermicchanges of statechanges of state
Condensation or liquifactionCondensation or liquifactionThe conversion of a gas to a liquid or solid.Example - steam becoming water
Freezing or crystallizationFreezing or crystallizationWhen a liquid becomes a solid.Examples - formation of ice from water
Substances usually first condense to liquids and then become solids.
12 - 28
ExothermicExothermicchanges of statechanges of state
Gas
Solid Liquid
depo
sition
liquification or
condensation
freezing orcrystallization
12 - 29
Changes in state andChanges in state andattractive forcesattractive forces
As the attractive forces between molecules become larger, more energy is needed to separate them.
Vapor pressures become smaller, boiling points and melting points become larger.
Chemical MW Polarity Mp Bp N2 28 Nonpolar -210 -196 O2 32 Nonpolar -219 -183 NH3 17 Polar -78 -33 H2O 18 Polar 0 100NaCl 58 Ionic 801 1465
12 - 30
Heating and CoolingHeating and Cooling
Changes in state involve several steps.
Example.Example.Producing 150 Producing 150 ooC steam from -20 C steam from -20 ooC C
ice.ice.1. Heat ice up to 0 oC.2. Convert the ice to water.3. Heat the water from 0 oC to 100
oC.4. Convert the water to steam.5. Heat the steam to 150 oC.
12 - 31
Heating and CoolingHeating and Cooling
Heat of fusion, Heat of fusion, HHfusfusThe amount of thermal energy necessary to melt one mole of a substance at its melting point.
Heat of vaporization, Heat of vaporization, HHvapvapThe amount of thermal energy necessary to boil one mole of a substance at its boiling point.
12 - 32
Heating and CoolingHeating and Coolingmp Hfus bp Hvap
Substance oC kJ/mol oC kJ/mol
Br2 -7.3 10.5759.2 29.5
CH3CH2OH -117.0 4.6079.0 43.5
CH3(CH2)6CH3 -56.8 20.65 125.7 38.6
H2O 0.0 6.01 100.0 40.7
Na 97.8 2.60 883 98.0
12 - 33
Specific heatsSpecific heatsEach substance requires a different amount of
energy to increase its temperature.Specific heatSpecific heat - amount of energy needed to
increase a substance’s temperature by 1oC.It also depends on the state of the substance.
Substance J/g Substance J/gAluminum, solid 1.0 Ice 2.1Copper, solid 0.4 Water 4.2Hydrogen, gas 14.2 Steam 2.0Mercury, liquid 0.1
12 - 34
Phase DiagramsPhase DiagramsGraphs that show the states of a substance as
a function of both pressure and temperature.
solidliquid
gas
temperature
pres
sure
12 - 35
Phase DiagramsPhase DiagramsPartial phase diagram for water.Partial phase diagram for water.
ice
water
steam
-20 0 20Temperature, oC
Pres
sure
, mm
Hg
30
20
10
0
Triple point
12 - 36
Phase DiagramsPhase Diagrams
Triple pointTriple pointAll three phases are in equilibrium. Temperature and pressure are fixed.
The triple point for water is at 0.01 oC and 4.58 mmHg.
The triple point for water, 273.16 K is used to define the Kelvin temperature scale.
12 - 37
Phase DiagramsPhase Diagramssupercriticalfluid region
Tc
solidliquid
gas
Pc
temperature
pres
sure Critical
point
12 - 38
Phase DiagramsPhase DiagramsCritical pointCritical point
The end of the vapor pressure curve.
Critical temperature, Critical temperature, TTccThe temperature at the critical point.
Critical pressure, Critical pressure, PPccThe pressure at the critical point.
At temperatures above Tc, liquefying a gas is impossible, no matter what the pressure.
12 - 39
Phase DiagramsPhase DiagramsAt pressures and temperatures above the
critical point, a supercritical fluidsupercritical fluid is formed.
A supercritical fluid:• is a gas.• has a density similar to a liquid.• has a viscosity similar to a gas.
Supercritical fluids have a number of uses. One example is their use for extractions - removal of caffeine from coffee.
12 - 40
The solid stateThe solid state
At room temperature, solids:•are not compressible•commonly have regular repeating units
Two types are observedCrystallineCrystalline solids have a definite melting
point. - ionic- ionic - covalent- covalent - molecular- molecular - metallic- metallicAmorphousAmorphous solids do not have a definite
melting point or regular repeating units.
12 - 41
Ionic solidsIonic solids
Ions make up the repeating units.
NaClNaCl
12 - 42
Covalent solidsCovalent solids
Repeating units of covalently bound atoms.
GraphiteGraphite
12 - 43
Molecular solidsMolecular solidsRepeating units are made up of molecules.
IceIce
12 - 44
Metallic solidsMetallic solidsRepeating units are made up of metal atoms,Valence electrons are free to jump from one
atom to another,
++ + +++ + +
++ + +++ + +
++ + +
++ + +++ + +++ + +
++ + +
12 - 45
Arrangement of units in crystalsArrangement of units in crystals
MetalsMetals• All atoms are spherical.• In a crystal, they are packed to minimize
the space they occupy.Coordination numberCoordination numberThe number of nearest neighbors that surround an atom in a crystalUnit cellUnit cellThe smallest three dimensional unit that describes the arrangement of the atoms
12 - 46
Simple cubic crystalsSimple cubic crystalsThis is one of the simplest arrangements to visualize.• Each atom has a coordination number of six.• Only 52% of the space is occupied.
Single layer Expanded model
12 - 47
Close PackingClose PackingThe crystals of most metals are of this type.• Each atom is surrounded by 6 neighbors in its
layer and a total of 12 in three dimensions..• This results in a high percentage of the space
being occupied.
Singlelayer
12 - 48
Close PackingClose Packing
The atoms in a layer mustrest in holes of the twolayers that touch it.
Two types of crystalscan result.
12 - 49
Cubic close packingCubic close packing
Face-centered cubic cellFace-centered cubic cell
Each face has fiveatoms the maximumamount of space isoccupied by theatoms - 74%.
12 - 50
Body-centered cubic unit cellBody-centered cubic unit cell
This unit cell is observed for all metals that do not crystallize in one of the two close-packed arrangements.
The exception is polonium.
The coordinationis eight.
68 % of the space isoccupied.
12 - 51
Body-centered cubic, GaAsBody-centered cubic, GaAs
12 - 52
Crystal structuresCrystal structuresCoordination % of space
Name number occupied Example
Face-centered 12 74 Alcubic
Body-centered 8 68 Nacubic
Simplecubic 6 52 Po
12 - 53
Ionic compoundsIonic compoundsCrystal structures for these compounds are
complicated by the following:• Two or more kinds of particles are involved.• The particles are usually differ in size and
often in charge.• Not all ions are spherical.
The major attractive force is electrostatic and crystals should allow the largest number of oppositely charged particles to touch.
12 - 54
Ionic compoundsIonic compoundsMany ionic compounds will
assume a close-packed arrangement of anions.
Small cations are placed in the holes.
Because each is surrounded by four spheres, the smaller holes are called tetrahedral holes.
12 - 55
Ionic compoundsIonic compounds
Many compounds will have this type of structure including LiCl, NaCl, NaBr, MgO, NiO, and NH4I.
NaClNaCl
12 - 56
Ionic compoundsIonic compounds
CsClCsCl
12 - 57
Ionic compoundsIonic compounds
AlAl22OO33
12 - 58
X-ray diffractionX-ray diffraction
This method is used to find the dimensions and shape of a crystal unit.
It provides a ‘fingerprint’ of a material which can be used:
•To deduce the structure of a material•To identify a substance•To tell structure of a polymer•For elemental analysis
12 - 59
X-ray diffractionX-ray diffraction
Film can be used fordetection of the patternsIt is now more commonto rotate the crystaland detect the x-rayswith a fixed positiondetector.This way, you havedata that can be processed by a computer