Intermolecular Forces, Liquids and Solids

LIQUIDS AND SOLIDS

Compared to gases, liquids and solids

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

They are held together by intermolecular forces

INTERMOLECULAR FORCES Inside molecules (intramolecular) the atoms

are bonded to each other. Intermolecular refers to the forces between the

molecules. The forces that hold one molecule to another

molecule. These forces arise from unequal distribution

of the electrons in the molecule and the electrostatic attraction between oppositely charged portions of molecules.

These forces cause changes of state by causing changes AMONG the molecules NOT within them.

INTERMOLECULAR FORCES Physical properties such as melting points,

boiling points, vapor pressures, etc. can be attributed to the strength of the intermolecular attractions present between molecules.

It works like this: the lower the boiling point (or vapor pressure or melting point), the weaker the intermolecular attractions; the higher the boiling point, the stronger the intermolecular attractions.

Hydrogen bonds are not true bonds—they are just forces of attraction that exist between a hydrogen atom on one molecule and the unshared electron pair on fluorine, oxygen or nitrogen atoms of a neighboring molecule.

INTERMOLECULAR FORCES Strong – covalent bonding and ionic bonding Weak – Dipole-dipole and London dispersion

forces During phase changes the molecules stay

intact. Energy is used to overcome forces.

DIPOLE - DIPOLE the force of attraction that enables two polar

molecules to attract one another. Polar molecules are those which have an uneven charge distribution since their dipole moments do not cancel.

Compounds exhibiting this type of IMF have higher melting and boiling points than those exhibiting weaker IMFs.

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.

HYDROGEN BONDING the force of attraction between the hydrogen atom of one

molecule and an unshared electron pair on F, O, or N of a neighboring molecule (a special case of dipole-dipole).

Effects boiling point. This is the strongest IMF. Never confuse hydrogen bonding with a bonded

hydrogen. Especially strong dipole-dipole forces when H is attached

to F, O, or N. These three because they have high electronegativity and they are small enough to get close.

The unique physical properties of water are due to the fact that it exhibits hydrogen bonding. As a result of these attractions, water has a high boiling point, high specific heat, and many other unusual properties.

HYDROGEN BONDING

(a) The Polar Water Molecule (b) Hydrogen Bonding Among Water Molecules

HYDROGEN BONDINGThe boiling points of the covalent hydrides of the elements in groups 4A, 5A, 6A, and 7A.

LONDON DISPERSION FORCES Also called induced dipole-induced dipole. the force of attraction between two non polar

molecules due to the fact that they can form temporary dipoles.

Nonpolar molecules have no natural attraction for each other. Without these forces, we could not liquefy covalent gases or solidify covalent liquids.

These forces are a function of the number of electrons in a given molecule and how tightly those electrons are held. Non - polar molecules also exert forces on each other. Electrons are not evenly distributed at every instant in time.

Have an instantaneous, temporary dipole. Induces a dipole in the atom next to it.

LONDON DISPERSION FORCES(a) An

instantaneous polarization can occur on atom A

(b) Nonpolar molecules such as H2 also can develop instantaneous and induced dipoles

LONDON DISPERSION FORCES Weak, short lived. Lasts longer at low temperature. Eventually molecules are long enough to make

liquids. More electrons, more polarizable. Bigger molecules, higher melting and boiling points. Much, much weaker than other forces. Also called van der Waal’s forces. Since all molecules have electrons, all molecules have

these forces. These forces range from 5 to 40 kJ/mol. The strength of this force increases as the number of

electrons increases due to increasing polarizability.

PHASES OF MATTER

The phase of a substance is determined by three factors:

temperature. pressure. strength of intermolecular forces.

LIQUIDS Many of the properties due to internal

attraction of atoms. Surface tension Beading Capillary action Viscosity

Stronger intermolecular forces cause each of these characteristics to increase.

SURFACE TENSION The resistance to an increase in its surface

area (polar molecules). High surface tension indicates strong IMFs. Molecules are attracted to each OTHER. A molecule in the interior of a liquid is

attracted by the molecules surrounding it, whereas a molecule at the surface of a liquid is attracted only by the molecules below it and on each side.

SURFACE TENSION Molecules at the

top are only pulled inside.

Molecules in the middle are attracted in all directions.

BEADING If a polar substance is placed on a non-polar

surface, there are cohesive but no adhesive forces.

CAPILLARY ACTION Spontaneous rising of a liquid in a narrow tube. Adhesive forces between molecule and glass

overcome cohesive forces between molecules themselves.

The narrower the tube, the more surface area of glass, the higher the column of water climbs!

The weight of the column sets the limit for the height achieved.

IMFs are cohesive, connecting like materials. Adhesive forces connect to something else. Glass is polar so it attracts water molecules.

CAPILLARY ACTION Hg liquid behaves opposite to water. Water

has a higher attraction for glass than itself so its meniscus is inverted or concave, while Hg has a higher attraction for other Hg molecules! Its meniscus is convex.

VISCOSITY Resistance to flow (molecules with strong

intermolecular forces). Increases with molecular complexity (long C

chains get tangled and larger electron clouds are more polarizable due to the presence of additional electrons) and increased with increasing IMFs.

Large forces, more viscous. Cyclohexane has a lower viscosity than

hexane because it is a circle- more compact.

TYPES OF SOLIDSTwo major types: Crystalline - highly regular arrangement of

components in their structure; example – salts, metals

Amorphous - considerable disorder in their structure; example - glass

Interparticle interactions and the ability to pack the particles together provide the main criteria for the structures of solids.

BODY-CENTERED CUBIC

FACE-CENTERED CUBIC

REPRESENTATION OF COMPONENTS OF A CRYSTALLINE SOLID Lattice - a three dimensional grid that

describes the locations of the pieces in a crystalline solid and designates the centers of components (atoms, ions, or molecules) that makes up the substance.

Unit Cell - The smallest repeating unit in of the lattice.

Three common types.

REPRESENTATION OF COMPONENTS OF A CRYSTALLINE SOLIDA. Covalent network solidB. Ionic salt crystal latticeC. Ice – formed due to hydrogen bonding

REPRESENTATION OF COMPONENTS OF A CRYSTALLINE SOLID network covalent—carbon in diamond form

—here each molecule is covalently bonded to each neighboring C with a tetrahedral arrangement. Graphite, on the other hand, exists as sheets that slide and is much softer.

ionic salt crystal lattice - Coulomb’s Law dictates the strength of the lattice

ice - notice the “hole” in the hexagonal structure and all the hydrogen bonds. The “hole” is why ice floats—it is less dense in the solid state than in the liquid state.

X-RAY ANALYSIS OF SOLIDSX-ray diffraction — A bending or scattering

of light. The beams of light are scattered from a regular array of points in which the spacing between the components are comparable with the λ of the light. It is due to constructive interference when the waves of parallel beams are in phase and to destructive interference when the waves are out of phase.

The diffraction pattern can be used to determine the interatomic spacing.

TYPES OF CRYSTALLINE SOLIDSIonic Solidcontains ions at the points of the lattice that

describe the structure of the solid (think NaCl).VERY high MP’s.Hard. Ion-Ion Coulombic forces are the strongest of all

attractive forces. “IMF” usually implies covalently bonded

substances, but can apply to both types.

TYPES OF CRYSTALLINE SOLIDSMolecular Solid: discrete covalently bonded

molecules at each of its lattice points (sucrose, ice).

Atomic Solid: atoms of the substance are located at the lattice points. Carbon—diamond, graphite and the fullerenes. Boron, and silicon as well.

TYPES OF CRYSTALLINE SOLIDS

THE BOOK DRONES ON Using diffraction patterns to identify crystal

structures. Talks about metals and the closest packing

model. It is interesting, but trivial. We need to focus on metallic bonding.

Why do metal atoms stay together? How their bonding affects their properties.

METALLIC SOLIDS Metals are characterized by high thermal and

electrical conductivity, malleability, and ductility.

These properties are explained by the nondirectional covalent bonding found in metallic crystals.

METALLIC SOLIDSclosest packing a model that uses hard spheres to represent the

atoms of a metal. These atoms are packed together and bonded to

each other equally in all directions. It will be easiest for you to understand if you can

imagine taking a cubic box and pouring in marbles. The marbles will layer, perhaps directly on top of one another, but perhaps one layer slides into the “dimple” made by the first layer so that the two layers are offset a bit. Next, remove the marbles and place tennis balls into the box. They will fill the box differently since they are of a different size.

METALLIC SOLIDS In the diagram in each layer, a given sphere is surrounded by six

others. a) exhibits aba packing—the second layer is like the first, but it

is displaced so that each sphere in the second layer occupies a dimple in the first layer. The spheres in the third layer occupy dimples in the second layer so that the spheres in the third layer lie directly over those in the first layer hence aba. aba has the hexagonal unit cell shown below and the resulting structure is hexagonal closest packed (hcp) structure. ababab….

METALLIC SOLIDS

b) exhibits abc packing—the spheres in the third layer occupy dimples in the second layer so that no spheres in the third layer lie above any in the first layer. The fourth layer is like the first. abc has a face-centered cubic unit cell and the resulting structure is cubic closest packed (ccp) structure. abcabc…

HEXAGONAL CLOSEST PACKED STRUCTURE

CUBIC CLOSEST PACKED STRUCTURE

BONDING MODELS FOR METALS Remember, metals conduct heat and

electricity, are malleable and ductile, and have high melting points.

These facts indicate that the bonding in most metals is both strong and non-directional.

It is difficult to separate metallic atoms, but easy to move them provided they stay in contact with each other!

BONDING MODELS FOR METALS

Electron Sea Model: A regular array of metals in a “sea” of electrons. I A & II A metals pictured at left.

Band (Molecular Orbital) Model: Electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms.

METALLIC BONDING

1s2s2p3s

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.

1s2s2p3s

3p

Magnesium Atoms


1s2s2p3s

3p

Magnesium Atoms

The 3s and 3p orbitals overlap and form molecular orbitals.


1s2s2p3s

3p

Magnesium Atoms

Electrons in these energy level can travel freely throughout the crystal.

ll


1s2s2p3s

3p

Magnesium Atoms

This makes metals conductors Malleable because the bonds are flexible.

ll

BONDING MODELS FOR METALSMetal alloys: a substance that has a mixture of elements and has metallic properties

substitution alloys - in brass, one third of the atoms in the host copper metal have been replaced by zinc atoms. Sterling silver - 93% silver and 7% copper. Pewter - 85% tin, 7% copper, 6% bismuth and 2% antimony. Plumber’s solder - 95% tin and 5% antimony.

interstitial alloy - formed when some of the interstices (holes) in the closest packed metal structure are occupied by small atoms. Steel - carbon is in the holes of an iron crystal. There are many different types of steels, all depend on the percentage of carbon in the iron crystal.

NETWORK COVALENT SOLID Also called Network Atomic Solid Composed of strong directional covalent

bonds that are best viewed as a giant molecule.

Both diamond and graphite are network solids.

The difference is that diamond bonds with neighbors in a tetrahedral 3-D fashion, while graphite only has weak bonding in the 3rd dimension.

NETWORK COVALENT SOLID brittle—diamond is the hardest substance on

the planet, but when a diamond is “cut” it is actually fractured to make the facets.

do not conduct heat or electricity carbon or silicon-based

CARBON- A SPECIAL ATOMIC SOLIDThere are three types of solid carbon.1. Amorphous- coal2. Diamond- hardest natural substance on earth;

colorless; insulates both heat and electricity. hard, colorless and an insulator.

3. Graphite- slippery; black; conducts electricity.

How the atoms in these network solids are connected explains why.

DIAMOND- EACH CARBON IS SP3

HYBRIDIZED, CONNECTED TO FOUR OTHER CARBONS.

Carbon atoms are locked into tetrahedral shape.

sp3 hybridization and 109.5° bond angles.

Strong s bonds give the huge molecule its hardness.

WHY IS DIAMOND AN INSULATOR?

The space between orbitals make it impossible for electrons to move around

Empty MOs

Filled MOs

E

GRAPHITE IS DIFFERENT

Each carbon is connectedto three other carbons and sp2 hybridized.

The molecule is flat with120º angles in fused six member rings.

The p bonds extend above and below the plane. This p bond overlap forms a huge p bonding

network Electrons are free to move throughout these delocalized orbitals.

The layers slide by each other.

SILICON AS A NETWORK SOLID Silicon is to geology what carbon is to biology! The

most significant silicon compounds involve chains with silicon-oxygen bonds.

silica - empirical formula SiO2—not at all like its cousin CO2. Quartz and some types of sand are silicon dioxide as opposed to a clear colorless gas such as carbon dioxide. Bonding is the difference.

Silicon cannot use its valence 3p orbitals to form strong π bonds with oxygen, mainly due to the larger size of the silicon atom and its orbitals—you get inefficient overlap. INSTEAD of forming π bonds, the silicon atom satisfies the octet rule by forming single σ bonds with FOUR OXYGEN atoms.

SILICON AS A NETWORK SOLIDEach silicon is in the center of a

tetrahedral arrangement of oxygen atoms. This means that although the empirical formula is SiO2, the structure is based on a network of SiO4 tetrahedral with shared oxygen atoms.

When silica is heated above its MP of about 1600ºC and cooled rapidly, an amorphous solid forms. We call it glass—it’s really a supercooled, ultra viscous liquid with a great deal of disorder.

MOLECULAR SOLIDS Molecules occupy the corners of the lattices

(not atoms). Different molecules have different forces

between them. These forces depend on the size of the

molecule and on the strength and nature of dipole moments.

Ice and dry ice are examples. Allotropes of sulfur and phosphorous are included. S8 or P4 occupy the lattice positions in these allotropes of these elements.

MOLECULAR SOLIDSCharacterized by strong covalent bonding

within the molecule yet weak forces between the molecules.

It takes 6 kJ of energy to melt one mole of solid water since you only have to overcome hydrogen bonding while it takes 470 kJ of energy to break one mole of O—H bonds.

Molecules such as CO2, I2, P4, and S8 have no dipole moment.

The melting pt and boiling pt increase since the molecules are MORE attracted to each other as a result of this polarizing of the electron cloud.

MOLECULAR SOLIDS

Those without dipoles: Most are gases at 25ºC. The only forces are London Dispersion Forces

which are dependent on size of atom. Large molecules (such as I2 ) can be solids even

without dipoles. As the size of the molecule increases (radius or

molar mass), the London dispersion forces increase because the larger the molecule the more electrons, the more polarizable its electron cloud. If it is more polarizable, temporary dipoles can easily form.

MOLECULAR SOLIDSThose with dipoles: Dipole-dipole forces are generally stronger than

London Dispersion forces Hydrogen bonding is stronger than Dipole-Dipole

forces. No matter how strong the IMF, it is always much,

much weaker than the forces in bonds. Stronger forces lead to higher melting and

freezing points.

WATER IS SPECIAL Each molecule has two polar O-H

bonds.

HO

H

d+

d+

d-

WATER IS SPECIAL

Each molecule has two polar O-H bonds.

Each molecule has two lone pairs of e- on its oxygen.

Each oxygen can interact with four hydrogen atoms.

HO

H

d+

d+

WATER IS SPECIAL

This gives water an especially high melting and boiling point.

HO

H

d+

d+

HO

H

d+

d+

HO

H

d+

d+

IONIC SOLIDS Stable, high-melting substances held together by

STRONG electrostatic forces that exist between oppositely charged ions.

Coulomb’s Law dictates the strength of the electrostatic force.

The extremes in dipole dipole forces - atoms are actually held together by opposite charges.

Huge melting and boiling points. Atoms are locked in lattice → hard and brittle. Every electron is accounted for so they are poor

conductors-good insulators.

SUMMARY

PRACTICE ONEUsing Table 10.7, classify each of the following substances according to the type of solid it forms.

a. Gold

b. Carbon dioxide

c. Lithium fluoride

d. Krypton

VAPOR PRESSURE Vaporization - change from liquid to gas at

boiling point Evaporation - change from liquid to gas

below boiling point. Boiling – change from liquid to gas at the

boiling point - – energy is added. ENDOTHERMIC – energy must be

absorbed. Heat (or Enthalpy) of Vaporization (∆Hvap) –

the energy required to vaporize one mole at 1 atm.

WATER’S ∆H Water’s heat of vaporization is 40.7 kJ/mol.

This is a large value. Water makes life on this planet possible since

it acts as a coolant. The reason its Hvap is so large has everything

to do with hydrogen bonding. The IMFs in water are huge, thus a great deal of the sun’s energy is needed to evaporate the rivers, lakes, oceans, etc. of Earth.

Perspiration is a coolant for animals possessing sweat glands. Energy from your hot body is absorbed by the water solution to evaporate.

CONDENSATION Opposite of vaporization. Change from gas to liquid. EXOTHERMIC Achieves a dynamic equilibrium with vaporization

in a closed system.

DYNAMIC EQUILIBRIUMa. When first sealed, the molecules gradually

escape the surface of the liquid.b. As the gas molecules build up above the

liquid, some condense back to a liquid.

DYNAMIC EQUILIBRIUMAs time goes by the rate of

vaporization remains constant but the rate of condensation increases because there are more molecules to condense.

DYNAMIC EQUILIBRIUMEquilibrium 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

DYNAMIC EQUILIBRIUM

VAPOR PRESSURE The pressure caused by the gas above the liquid

at equilibrium. Liquids with high vapor pressures evaporate

easily. They are called volatile. They have weak IMFs.

Vapor pressure decreases with increasing IMFs. Bigger molecules (bigger dispersion forces) More polar molecules (dipole-dipole)

VAPOR PRESSURE Increases with increasing temperature. Easily measured in a barometer. Heat the particles up, speed the up, move

them out. Increasing the temperature increases the KE

which facilitates escape and the speed of the escapees.

They bang into the sides of the container with more frequency and more energy.

More molecules can attain the energy needed to overcome the IMFs in a liquid at a higher temperature since the KE increases.

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

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

TEMPERATURE EFFECT

Kinetic energy

# of

mol

ecul

es

T1

Energy needed to overcome intermolecular forces

Kinetic energy

# of

mol

ecul

es

T1

Energy needed to overcome intermolecular forcesT1

T2

At a higher temperature more molecules have enough energy - higher vapor pressure.

Energy needed to overcome intermolecular forces

MOLAR MASS AND VAPOR PRESSURE As Molar Mass increases, Vapor

Pressure decreases. Because as molecules increase in molar

mass, they also increase in the number of electrons.

As the number of electrons increase, the polarizability of the molecule increases so more induced dipole-induced dipole or dispersion forces exist, causing stronger attractions to form between molecules.

This decreases the number of molecules that escape and thus lowers the VP.

BIG EXCEPTION

Hydrogen bonding causes a major exception. Its presence greatly increases the IMFs of the

liquid. Water has an incredibly low VP for such a

light molecule - MM = 18.02g/mol.

MATHEMATICAL RELATIONSHIP

ln is the natural logarithmln = Log base ee = Euler’s number - an irrational number

like pR - universal gas constant, but use the “energy” R = 8.3145 J/K mol.C – a constant characteristic of the liquid.DHvap is the heat of vaporization in J/mol.

C+

D

T1

RH

=lnVP vap


Surprisingly this is the graph of a straight line.

PRACTICE TWO Using the plots in graph in notes, determine

whether water or diethyl ether has the larger enthalpy of vaporization.


If we know the values of ΔHvap and VP at one temperature, we can solve the expression for the constant, C and set a second expression for T2 equal to the first since the value of C is NOT dependent upon temperature:

This form is called the Clausius- Clapeyron equation.

D

12

vapvapT2

vapT1

T1-

T1

RH

= PPln

MATHEMATICAL RELATIONSHIP The Clausius-Clapeyron equation allows us to

estimate the vapor pressure at another temperature, if the vapor pressure is known at some temperature, and if the enthalpy of vaporization is known.

D

12

vapvapT2

vapT1

T1-

T1

RH

= PPln

PRACTICE THREEThe vapor pressure of water at 25°C is 23.8 torr, and the heat of vaporization of water at 25°C is 43.9 kJ/mol. Calculate the vapor pressure of water at 50°C.

SUBLIMATION solids also have vapor pressures. Some solids go directly to the vapor phase at

1atm, skipping the liquid phase all together! ENDOTHERMIC Examples: Iodine and dry ice (solid CO2)

Opposite of this is DEPOSITION

CHANGES OF STATE 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 DHfus.

-40-20

020406080

100120140

0 40 120 220 760 800

HEATING CURVE FOR WATER

IceWater and Ice

Water

Water and Vapor

Vapor

HEAT OF FUSION DHFUS

the enthalpy change that occurs at the melting point (or freezing point).

This energy is clearly going into increasing the PE of the molecules since the temperature or average KE of the molecules has plateaued or is staying the same.

vapor pressure of solid = vapor pressure of liquid (EQ is established).

On the plateaus, calculate the energy change using q = ΔH[vap or fus]m

On the slants, calculate the E change using q = mc[solid, liquid or gas]ΔT

-40-20

020406080

100120140

0 40 120 220 760 800

HEATING CURVE FOR WATER

Heat of Fusion

Heat ofVaporization

Slope is Heat Capacity

MELTING POINT 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. Molecules break loose from lattice points and

solid changes to liquid. Temperature remains constant during

ANY phase change. Potential energy (PE) is changing like crazy while KE remains constant!

MELTING POINT The melting and boiling

points of water are determined by the vapor pressures of the solid and liquid states.

Below zero - VP of ice has a larger temperature dependence. This means the VP of ice increases more rapidly than the liquid’s VP for each increase in temperature.

A point is eventually reached where the VP solid = VP liquid.

We call this temperature the melting pt.

This figure shows VP of solid and liquid water as a function of temperatures near zero.

“NORMAL”normal melting point—the

temperature at which the VP solid = VP liquid AND P total = 1atm

normal boiling point—the temperature at which the VP liquid = exactly 1 atm

BOILING POINT 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.

SUPERCOOLING the substance is at a temperature below its

freezing point, yet it remains a liquid. Usually happens when the cooling has been

gradual and the degree of organization needed to form crystals has not happened.

At some point, organization happens and the solid is readily formed, the temperature rises back to the MP as the heat of fusion is released.

SUPERCOOLING OF WATER

SUPERHEATED the substance is at a temperature above its

boiling pt, yet it remains a liquid. Usually happens when heated very rapidly

(like in a microwave oven) and bubbles form in the interior with high internal pressures.

They often burst before reaching the surface making quite a mess.

This is called bumping in the lab. Prevent it by adding boiling chips to the flask. These chips are porous and have air trapped in them, upon heating tiny air bubbles form and act as “starters” for vapor bubble formation.

Solid Water

Liquid Water

Water Vapor Vapor

SOLID WATER AND LIQUID WATER EUILIBRIUM

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 (vaporize) to achieve equilibrium.

Solid Water

Liquid Water

Water Vapor Vapor

While the molecules will condense to a liquid.

This can only happen if the temperature is above the freezing point since the 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 = melting point.

Solid Water

Liquid Water

Water Vapor Vapor

PHASE DIAGRAMSRepresent phases of matter as a function of

temperature and pressure.

Critical temperature: temperature above which the vapor cannot be liquefied.

Critical pressure: pressure required to liquefy AT the critical temperature.

Critical point: critical temperature and pressure coordinates (for water, Tc = 374°C and 218 atm).

Triple point: the point where all three phases exist in equilibrium.

Solid Liquid

Gas

Triple Point

Critical Point

Temperature

Pres

sure

Solid Liquid

Gas

This is the phase diagram for water. The density of liquid water is higher than solid

water.

Temperature

Pres

sure

Solid1 Atm

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

Pres

sure

FOUR EXPERIMENTS

EFFECT OF PRESSURE

The wire will exert pressure on the block, melt it and begin a journey downward through the block due to the force of gravity acting on the weights. After the wire has left thesurface, the block refreezes! Most substances have a solid-liquid line that has a positive slope since their solid phase is more dense than the liquid.

REMEMBER!! EACH PHASE BOUNDARY REPRESENTS

AN EQUILIBRIUM SET OF PRESSURE ANDTEMPERATURE CONDITIONS!!

Be sure and use the word equilibrium in your writing as you explain phase diagrams.

Intermolecular Forces, Liquids and Solids

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Transcript of Intermolecular Forces, Liquids and Solids