STATES OF MATTER

CHEM 15

Table 12.1

A Macroscopic Comparison of Gases, Liquids, and Solids

State Shape and Volume Compressibility Ability to Flow

Gas Conforms to shape and volume of container

high high

Liquid Conforms to shape of container; volume limited by surface

very low moderate

Solid Maintains its own shape and volume

almost none almost none

Figure 12.3

A cooling curve for the conversion of gaseous water to ice.

Phase Changes

solid liquid gas

melting

freezing

vaporizing

condensing

sublimation

endothermic

exothermic

deposition

GASES(e.g. Air: 78% N2, 21% O2, vapors (N2, H2, F2, Cl2, He, Ne, Ar, Kr, Xe))coined by Jan van Helmont in 1624 from the Greek word chaos which means “confusion” or “empty space”

have the complete freedom to fill its container thus they assume the volume and shape of the vessel at which they are contained.

are highly compressible

forms homogeneous mixtures

molecules are far apart: ~0.1% of total volume is occupied by gas molecules compared to 70% for lquids

*Substances that are liquids or solids under ordinary conditions can usually also exist in the gaseous state and they are often referred to as vapors

Variables that affect gases:1. volume, V 3. pressure, P2. temperature, T 4. number of moles, n

*Pressure – a force that tends to move something in a given directionP = F/A F = force; A = area

*gases exert pressure on walls with which they are in contacte.g. balloons, tires

SI Unit: 1 kg/m.s2 = 1 N/m2 = 1 Pa1 atm = 760 mmHg = 760 Torr = 1.01325 x 105 Pa or 101.325kPa

* standard atmospheric pressure = typical pressure at sea level; pressure exerted by the atmosphere on the Earth’s surface* Standard Temperature and Pressure (STP): T = 273 K, P = 1 atm

barometer – device used for atmospheric pressure measurement– made from an inverted glass tubing of more than 760 mm in length. This tube is filled with Hg and the height of the mercury column is proportional to the atmospheric pressure.manometer – device used to measure the pressure of enclosed gases whose pressures are near the atmospheric pressure – if the levels in the two arms of the tube are the same, the enclosed gas has the same pressure as the atmospheric pressure. – if the pressure of the enclosed gas is less than the atmospheric pressure, the mercury is forced toward the direction of the enclosed gas. – if the pressure of the enclosed gas is greater than the atmospheric pressure, the mercury is push towards the atmosphere.

GAS LAWS - equations that express the relationships among P, T, V, and n of gases

1. Boyle’s Law (Robert Boyle) : P-V Relationship“At constant temperature, the volume of a quantity of gas is

inversely proportional to the pressure.”at constant T: V α (1/P) ↔ PV = constantpressure, ↓ volume

2. Charles’ Law (Jacques Charles): temperature-volume relationship

“At constant pressure, the volume occupied by a fixed amount of gas is directly proportional to the absolute temperature.”at constant P : T α V ↔ V/T = constant↑ temperature, ↑ volume*Note: T = absolute temperature (Kelvin, K)

TK = 273 + oC

3. Avogadro’s Law (Amadeo Avogadro): quantity – volume relationshipAvogadro’s Hypothesis: “Equal volume of gases at the same temperature and pressure contain equal number of particles (molecules).”

Avogadro’s Law: “The volume of gas maintained at constant temperature and pressure is directly proportional to the number of moles of the gas.”at constant T and P: V α n ↔ V = constant x nsame volume of any gas, say 10 mL, be it H2, He, O2, NH3 , and so on, at the same temperature and pressure, contain the same number of moles (or molecules)at STP, V = 22.4 L

4. Gay-Lusaac’s Law/Laws of Combining Volumes (Joseph Louis Gay-Lusaac)“At a given pressure and temperature, the volume of gases that react with one another

are in ratios of small whole numbers.”e.g. 2H2 + O2 → 2H2O

1omL 5 mL 10 mL

Amonton’s Law : P1

T1

=P2

T2

Boyle’s Law: P1V1 = P2V2

Charles’ Law :

V1

T1

V2

T2

=

Combined Gas Laws :

P1V1 = P2V2

T1 T2

A

B

C

The Ideal Gas Equation: Combined Gas Laws

*ideal gas – hypothetical gas whose P, v, and T behavior are completely described by the ideal gas equation

Ideal Gas Equation: PV = nRTwhere R = ideal gas constant = 0.08206 L · atm / mol · K

P = pressure (atm)T = temperature (K)n = mole of gas (mol)V = volume (L)

Exercises1. A sample of a gas occupies 360mL under a pressure of 0.750atm. If the T is held constant, what volume will the sample occupy under a pressure of 1.00atm?2. At 0oC and 5.00 atm, a given sample of a gas occupies 75.0 L. The gas is compressed to a final volume of 30.0L at 0oC. What is the final pressure of the gas?3. A sample of N2 exerts a pressure of 0.988atm and occupies 12.3L when its T 450K. Assuming constant pressure, what volume will the gas occupy at 300K?4. At what T will a 10.0 L gas at constant pressure be brought to if the same gas originally occupied 3000cm3 at 27oC?5. Two identical cylinders of N2 contain the same weight of N2. The T of one cylinder is 20oC and the other is 100oC. If the pressure the one at 20oC is 1520mmHg, what is the pressure of the other cylinder?6. In a laboratory, 200cm3 of O2 gas is collected at a pressure of 73 cmHg and temperature of 30oC, Compute the volume of the O2 gas under standard conditions.7. A sample of gas occupies 400mL at STP. What volume will the sample occupy at 71oC and 2.50 atm?8. Calculate the pressure of one mole of an ideal gas which occupies 12.0 L at 25oC.9. A gas contained in 50L under 8 atm of pressure and at 20oC. How many moles of gas are there in the container?

GAS MIXTURES AND PARTIAL PRESSURES

PARTIAL PRESSURE AND MOLE FRACTION

FURTHER APPLICATIONS OF THE IDEAL GAS EQUATION

Gas Densities and Molar Mass:the higher the molecular mass and pressure of a gas, the denser is the gas.more denser gases will lie below that of less dense gases.hot gases tend to be less dense than cool gases.

Volumes of Gases in Chemical ReactionThe coefficients in a balanced chemical equation tell us the

relative amounts (in moles) of reactants and products in a reaction.Collecting Gases Over Water*When a gas is collected over water, the total pressure or atmospheric pressure is equal to the sum of the pressures of the gas and the water vapor.

Patm = Pgas + PWATER

Exercises 1:

1. From data gathered by Voyager 1, scientist have estimated the composition of the atmosphere of Titan, Saturn’s largest moon. The total pressure on the surface of Titan is 1220 Torr. The atmosphere consist of 82 mole percent N2, 12 mol percent Ar, and 6.0 mol percent CH4. Calculate the partial pressure of these gases.

2. A sample of 5.0 g He gas, 2.0 g H2 gas and 10.0 g water vapor exerts a pressure of 0.5 atm at 25oC. Calculate the mole fraction of each gas, the partial pressure of each gas and the volume occupied by the mixture.

Exercises1. What is the density of CCl4 vapor at 714 torr and 125oC?

2. The safety air bags in automobile are inflated by nitrogen gas generated by the rapid decomposition of sodium azide, NaN3:

NaN3 (s) 2 Na (s) + 3 N2 (g)If an air bag has a volume of 36 L and is to be filled with nitrogen gas at a pressure of 1.15 atm at a temperature of 26.00C, how many grams of NaN3 must be decomposed?

3. A sample of KClO3 is partially decomposed producing O2 that is collected over water. The volume of gas collected is 0.250 L at 26oC and 765 Torr of total pressure. How many moles of O2 are collected? How many grams of KClO3 were decomposed?PH2O at 26oC = 25 torr

KINETIC MOLECULAR THEORY*provides a model to explain the regularity observed in the behavior of ideal gases.*formulated by Rudolf Clausius

Summary:*Gases consist of large number of molecules that are in continuous, random motion.*These gaseous molecules are widely separated in space. Actual volume of gas molecules is negligible compared with empty space between them.*Attractive and repulsive forces between the molecules are negligible.*The collisions are perfectly elastic. Energy can be transferred from one molecule to the other during collisions, but the average kinetic energy of the molecules remain constant.*The average kinetic energy of the molecules is proportional to absolute temperature. At any given temperature, the molecules of all gases have the same kinetic energy.

APPLICATION OF KMT1. Effect of volume increase on pressure at constant temperature

the speed at which molecules travel remains unchanged when temperature is constant (kinetic energy is proportional to temperature)

an increase in volume also increases the distance at which the molecules travel. When this happens, there will be fewer collisions on the container walls per unit time. Fewer collisions result to decrease in pressure. Increase in volume, decrease in pressure

2. Effect of temperature increase on pressure at constant volumeNo change in volume means constant number of collisions against the container walls.

An increase in temperature means an increase in kinetic energy. The greater the kinetic energy, the more collisions occur. Hence, the pressure of the gas increases. Increase in temperature, increase in pressure

Molecular Effusion and Diffusionroot-mean-square-speed (rms speed),

3RT

*Note: Particles of lighter gases have a higher rms speed than particles of heavier one.

μ =

Exercises1.Calculate the rms speed of N2 molecule at

25oC.

3RT

rms = = (3)(8.3145 J.mol/K)(298.15K) (28.02 g/mol)(1 kg/1000g)

√

= 515 or 5.2 x 102 m/s

Note : 1 J = 1 kg (m2/s2)

2. Arrange the following in increasing rms speed: N2, O2, H2O, NO at 298 K

N2 = 28.02 g/molO2 = 32 g/molH2O = 18.02 g/molNO = 30.01 g/mol

Answer: O2 < NO< N2 < H2O

Effusion – the escape of gas molecules through a tiny hole.

Diffusion – the spread of one substance throughout a space or throughout a second substance.

Graham’s Law (Thomas Graham)“The effusion and diffusion rates of a gas are

inversely proportional to the square root of its molar mass.”

Rate α1

Exercises1. Arrange the following gases in order of increasing rate of diffusion: CO, NO2,

HF, H2S, CO2 at the same temperature

CO= 28.01 Answer: NO2 < CO2 < H2S < CO < HFNO2 = 46.01HF = 20.008H2S = 34.086CO2 = 44.01

= d1

d2

2. An unknown gas composed of homonuclear diatomic molecule effuses at a rate that is only 0.355 times that of O2 at the same temperature. What is the identity of the unknown gas?

3. Two gases, O2 and N2, are introduced into the two ends of a 100.0cm tube. At what distance from the point where O2 is introduced will the two gases meet?

Deviations from Ideal Gas Behavior (Real or nonideal)High Pressures – molecules are more crowded; volume of gas molecules cannot be considered negligible; observed volume of gas > ideal volume

Low Temperature – average KE decreases while intermolecular forces remain; cooling a gas deprives the molecules of the energy needed to overcome the intermolecular attractive forces; less number of collisions with walls; observed P < ideal pressure

Intermolecular Forces of Attraction

-are the forces that exist between molecules

-the strengths of intermolecular forces vary for each substance

-they are generally much weaker than ionic or covalent bond (intramolecular)

-properties of liquids, such as boiling points and melting points, depend on the strength of intermolecular forces of attraction (IMFA)

Intermolecular Forces:

INTRAMOLECULAR BONDING

INTERMOLECULAR BONDING

Types of Intermolecular Forces of Attraction

ion-dipole forcedipole-dipole forceHydrogen bondingLondon dispersion

Polar molecule (due to difference in electronegativity)

Figure 12.12 Polar molecules and dipole-dipole forces.

solid

liquid

Polarizability and Charged-Induced Dipole Forces

distortion of an electron cloud

•Polarizability increases down a group

size increases and the larger electron clouds are furtherfrom the nucleus

•Polarizability decreases left to right across a period

increasing Zeff shrinks atomic size and holds the electronsmore tightly

•Cations are less polarizable than their parent atom because they are smaller.

•Anions are more polarizable than their parent atom because they are larger.

1. ION-DIPOLE FORCE

=exist between an ion and the partial charge on the end of a polar molecule (e.g. H2O)=polar molecules have dipoles, that is, they have partial positive and partial negative ends=specially important for solutions of ionic compounds in polar liquids, e.g. solution of NaCl in water

2. DIPOLE-DIPOLE FORCE-exist between neutral polar molecule-polar molecules attract each other when the negative end of molecule is in close proximity with the positive end of the other molecule

-effective only when polar molecules are close to each other-are generally weaker than ion-dipole force-for molecules with approximately equal mass and size, the strengths of dipole-dipole attraction increase with increasing polarity

3. Hydrogen Bonding (H-bonding)

-the strongest type of IMFA-a special case of dipole-dipole interaction occurring between molecules in which H is covalently bonded to small, electronegative atom (N, O, F)-the small size of the H atom makes the F, N, or O atom of one molecule approach the H atom of another molecule closely sufficient to produce an attraction strong enough to be called a bond

Illustration: H ― F ----------- H ― F

THE HYDROGEN BOND

a dipole-dipole intermolecular force

The elements which are so electronegative are N, O, and F.

A hydrogen bond may occur when an H atom in a molecule, bound to small highly electronegative atom with lone pairs of electrons, is attracted to the lone pairs in another molecule.

..F..

.. ..H O..

N.. FH

..

..

..

O.. ..

..NH

hydrogen bonddonor

hydrogen bondacceptor

hydrogen bondacceptor

hydrogen bonddonor

hydrogen bonddonor

hydrogen bondacceptor

Figure 12.21 The H-bonding ability of the water molecule.

hydrogen bond donor

hydrogen bond acceptor

SAMPLE PROBLEM 12.2 Drawing Hydrogen Bonds Between Molecules of a Substance

SOLUTION:

PROBLEM: Which of the following substances exhibits H bonding? For those that do, draw two molecules of the substance with the H bonds between them.

C2H6(a) CH3OH(b) CH3C NH2

O

(c)

PLAN: Find molecules in which H is bonded to N, O or F. Draw H bonds in the format -B: H-A-.

(a) C2H6 has no H bonding sites.

(c)(b)C O H

H

H

H

COH

H

H

H

CH3C N

O

H

H

CH3CN

O

H

H

CH3CN

O

H

H

CH3CN

O

H

H

4.LONDON DISPERSION FORCE

-forces formed from instantaneous (momentary) dipole moment created by the motion of electrons in an atom or molecule

-the movement of electrons in a molecule is influenced by repulsion of neighboring electrons

-IMFA that exist in nonpolar molecules ; the weakest of all known intermolecular forces; exist between all molecules but are generally overshadowed when stronger forces are present.

-the greater the polarizability of the electrons in a molecule, the more easily is its electron cloud can be distorted, hence, the stronger the London dispersion force.

-dispersion forces tend to increase in strength as the molecular weight increases (due to increasing size and polarizability)

Figure 12.15 Dispersion forces among nonpolar molecules.

separated Cl2

molecules

instantaneous dipoles

Figure 12.18

Summary diagram for analyzing the intermolecular forces in a sample.

INTERACTING PARTICLES(atoms, molecules, ions)

INTERACTING PARTICLES(atoms, molecules, ions)

ions onlyIONIC BONDING(Section 9.2)

ions onlyIONIC BONDING(Section 9.2)

ion + polar moleculeION-DIPOLE FORCESion + polar moleculeION-DIPOLE FORCES

ions present ions not present

polar molecules onlyDIPOLE-DIPOLE

FORCES

polar molecules onlyDIPOLE-DIPOLE

FORCES

HYDROGENBONDING

HYDROGENBONDING

polar + nonpolar moleculesDIPOLE-INDUCED DIPOLE FORCES

polar + nonpolar moleculesDIPOLE-INDUCED DIPOLE FORCES

nonpolar molecules onlyLONDON DISPERSIONFORCES only

nonpolar molecules onlyLONDON DISPERSIONFORCES only

DISPERSION FORCES ALSO PRESENT

H bonded toN, O, or F

Figure 13.1 The major types of intermolecular forces in solutions.

Strengths of IMFA

H-bonding > ion dipole > dipole-dipole > london dispersion

SAMPLE PROBLEM 12.3 Predicting the Type and Relative Strength of Intermolecular Forces

PROBLEM: For each pair of substances, identify the dominant intermolecular forces in each substance, and select the substance with the higher boiling point.

(a) MgCl2 or PCl3

(b) CH3NH2 or CH3F

(c) CH3OH or CH3CH2OH

(d) Hexane (CH3CH2CH2CH2CH2CH3)

or 2,2-dimethylbutaneCH3CCH2CH3

CH3

CH3PLAN: Use the formula, structure and Table 2.2 (button).

•Bonding forces are stronger than nonbonding(intermolecular) forces.

•Hydrogen bonding is a strong type of dipole-dipole force.

•Dispersion forces are decisive when the difference is molar mass or molecular shape.

SOLUTION:

SAMPLE PROBLEM 12.3 Predicting the Type and Relative Strength of Intermolecular Forces

continued

(a) Mg2+ and Cl- are held together by ionic bonds while PCl3 is covalently bonded and the molecules are held together by dipole-dipole interactions. Ionic bonds are stronger than dipole interactions and so MgCl2 has the higher boiling point.

(b) CH3NH2 and CH3F are both covalent compounds and have bonds which are polar. The dipole in CH3NH2 can H bond while that in CH3F cannot. Therefore CH3NH2 has the stronger interactions and the higher boiling point.

(c) Both CH3OH and CH3CH2OH can H bond but CH3CH2OH has more CH for more dispersion force interaction. Therefore CH3CH2OH has the higher boiling point.(d) Hexane and 2,2-dimethylbutane are both nonpolar with only dispersion forces to hold the molecules together. Hexane has the larger surface area, thereby the greater dispersion forces and the higher boiling point.

Examples1. Arrange the following types of interactions in order of

increasing stability: covalent bond, van der Waals force, hydrogen bonding, dipole interaction

2.Which has the highest boiling point: H2, He, Ne, Xe, CH4

3.Which is expected to have the highest melting point: PH3, NH3, (CH3)3N? Explain why.

Answer: van der Waals < dipole < hydrogen bonding < covalent

Answer: Xe, All are nonpolar molecules, but Xe has the greatest van der Waals forces because it has the most electrons

Answer: NH3 has the strongest intermolecular forces, thus it is expected to have the highest melting point.

II. Liquids

II. LIQUIDSintermolecular forces (IMF) are strong enough to hold the molecules close together; IMF are not strong enough to keep the molecules from moving past one another

Consequences:liquids flow and assume the shape of their containerliquids are denser than gases; a 70% of the volume of liquids are occupied by liquid moleculesliquids are incompressible; liquids do not expand to fill the containerdiffusion with a liquid occurs slowly

Properties of Liquids

1. Viscosity resistance of liquid to flow

related to the ease with which individual molecules of the liquid can move with respect to one another

it depends on the IMFA that exist in the liquid; the greater the liquid’s viscosity, the more slowly it flows

viscosity decreases with increasing temperature; at higher temperature, the average kinetic energy of the molecules is greater, hence, it more easily overcomes the IMFA

↑ T, ↓ viscosity↑ P, ↑ viscosity↑ complexity of molecule, ↑ viscosity↑ IMF, ↑ viscosity e.g. H2O > CH3CH2OH; oil > H2O

Table 12.4 Viscosity of Water at Several Temperatures

Temperature(0C)Viscosity (N*s/m2)*

20

40

60

80

1.00x10-3

0.65x10-3

0.47x10-3

0.35x10-3

*The units of viscosity are newton-seconds per square meter.

viscosity - resistance to flow

2. Surface tension•energy required to increase the surface area of a liquid by a unit amount•measure of the inward force in liquids•the stronger the IMFA, the larger the surface tension

Consequences of Surface Tension:Cohesion/Adhesion Cohesive force – IMFA that binds similar molecules Adhesive force – IMFA that binds different molecules/surface Cohesion > Adhesion → meniscus concave down Adhesion > Cohesion → meniscus concave up

Capillary Action – the rise of a liquid up a narrow tube. The surface area is increased when H2O rise up a narrow tube. The H2O levels stop rising until it is balanced by the gravitational pull

Figure 12.19 The molecular basis of surface tension.

hydrogen bondingoccurs in three

dimensions

hydrogen bondingoccurs across the surface

and below the surfacethe net vectorfor attractive

forces is downward

Table 12.3 Surface Tension and Forces Between Particles

Substance FormulaSurface Tension

(J/m2) at 200C Major Force(s)

diethyl ether

ethanol

butanol

water

mercury

dipole-dipole; dispersion

H bonding

H bonding; dispersion

H bonding

metallic bonding

1.7x10-2

2.3x10-2

2.5x10-2

7.3x10-2

48x10-2

CH3CH2OCH2CH3

CH3CH2OH

CH3CH2CH2CH2OH

H2O

Hg

Figure 12.20 Shape of water or mercury meniscus in glass.

adhesive forcesstronger

cohesive forces

H2O

capillarity

Hg

3. Vaporization/Evaporation•passage of molecules from liquid to a gaseous state•rate of vaporization increases with increased T and decreased IMFA

4. Vapor pressurethe pressure exerted by a vapor over the liquid when the liquid and vapor state are in dynamic equilibrium (equal rates of evaporation and condensation)

escaping tendency of the liquid

the liquids with weak IMFA have high vapor pressure, hence, they evaporate easily and are said to be volatile; vapor pressure increases with temperature. The higher the temperature, the higher also is the vapor pressure →

↑ T, ↑ VP ↑ IMF, ↓ VP↑ VP, ↓ boiling point

5. Boiling Point

temperature at which the VP of a liquid equals the external pressure acting on its surface

during boiling, every heat absorbed is used to convert liquid to gas and T remains constant until all the liquid has been converted

normal boiling point – boiling point at which the external pressure is equal to 1atm, e.g. H2O 100oC

Tc, critical temperature – the highest temperature at which a liquid can exist

Pc, critical pressure – the minimum pressure required to bring about liquefaction; greater IMFA → more readily a gas is liquefied → higher Tc

Figure 12.16

Molar mass and boiling point.

Figure 12.17 Molecular shape and boiling point.

more points for dispersion

forces to act

fewer points for dispersion

forces to act

Figure 12.14 Hydrogen bonding and boiling point.

Figure 12.13 Dipole moment and boiling point.

Problem: Compare the two substances base on the factors listed. Put X to the appropriate box.

Factors Alcohols,CH3OH Water, H2O

IMFA

Rate of evaporation

Vapor pressure

Boiling Point

Viscosity

X

X

X

XX

III. Solids

IMF’s are strong enough to hold the molecules close together and virtually lock them in place

Consequences:•solids retain their shape and have fixed volume•solids are incompressible•diffusion within a solid is extremely slow•solids do not flow

Table 12.5 Characteristics of the Major Types of Crystalline Solids

ParticlesInterparticle Forces

Physical Behavior Examples (mp,0C)

Atomic

Molecular

Ionic

Metallic

Network

Group 8A(18)[Ne-249 to Rn-71]

Molecules

Positive & negative ions

Atoms

Atoms

Soft, very low mp, poor thermal & electrical conductors

DispersionAtoms

Dispersion, dipole-dipole, H bonds

Fairly soft, low to moderate mp, poor thermal & electrical conductors

Nonpolar - O2[-219], C4H10[-138], Cl2

[-101], C6H14[-95]

Polar - SO2[-73], CHCl3[-64], HNO3[-42], H2O[0.0]

Covalent bond

Metallic bond

Ion-ion attraction

Very hard, very high mp, usually poor thermal and electrical conductors

Soft to hard, low to very high mp, excellent thermal and electrical conductors, malleable and ductile

Hard & brittle, high mp, good thermal & electrical conductors when molten

NaCl [801]CaF2 [1423]

MgO [2852]

Na [97.8]Zn [420]Fe [1535]

Bonding in SolidsTypes of Solids Based on Bonding1. molecular solids

•consist of atoms/molecules held together by IMF•soft; low melting pointe.g. Ar, H2O, CH4, C12H22O11, CO2

2. covalent•network solids•consist of atoms held together in large networks or chains by covalent bonds•hard; high melting pointe.g. diamond (mp = 3550oC), quartz (SiO2), graphite (interconnected hexagonal planes; used as lubricants), silicon carbide (SiC), boron nitride (BN)3. ionic solidsions held together by ionic solids (electrostatic attraction)structure depends on size and charge of ionse.g. NaCl, MgO, ZnS, CaF2, Ca(NO3)24. metallic solidsconsist of metal ionse.g. Na, Cr

Structure of Solids

a. crystalline solidsb. amorphous solid

ions or molecules are in well-defined arrangement

no orderly structure

have regular shapes lack well defined shapes

have specific melting point no specific melting temperature

e.g. diamond, quartz, NaCl

PHASE DIAGRAMS – graphical way to summarize the conditions under which equilibria exist between different states of matter. This allows us to predict the phase of a substance that is stable at any given T and P

Features of a phase diagram includes:Vapor-pressure curve: generally as temperature increases, vapor pressure increases.Critical Point: critical temperature and pressure for the gas

Normal melting point: melting point at 1 atmTriple point: temperature and pressure at which all three phases are in equilibriumAny temperature and pressure combination not on the curve represents a single phase

*Note: Gases may be liquefied by increasing the pressure at a suitable temperature

Critical Temperature – the highest temperature at which a substance can exist as a liquidCritical Pressure – the minimum pressure required for liquefaction at this critical temperature

The greater the IMF, the easier it is to liquefy a substance, thus, the higher the critical temperature

Critical Point – point at which the liquid and vapor becomes indistinguishableSupercritical Fluid – have high densities like liquids and have low viscosity like gases

Phase Diagrams of H2O and CO2

1. WaterIn general, an increase in pressure favors the more compact phase of the material; this is usually the solidWater is one of the few substances whose solid form is less dense than the liquid form → the melting point curve for water slopes to the leftThe triple point occurs at 0.0098oC and 4.58 TorrThe normal melting point is 0oCThe normal boiling point is 100oCThe critical point is 374oC and 218 atm CO2

2. Carbon dioxide

The triple point occurs at -56.4oC and 5.11 atmThe normal sublimation point is -78.5oC. (At 1 atm, CO2 sublimes, it does not melt)The critical point occurs at 31.1oC and 73 atmFreeze drying: Frozen food is placed in a low pressure (<4.58 Torr) chamber → the ice sublimes

H2O