1/21/2014 1 Intermolecular forces, Liquids and Solids Ch. 10.

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Intermolecular forces, Liquids and Solids

Ch. 10Ch. 10

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Some Characteristics of Gases, Liquids and Solids and the Microscopic Explanation for the Behavior

GASGAS LIQUIDLIQUID SOLIDSOLID

Assumes shape/ volume of container –particles move past one another

Assumes shape of container –particles move/ slide past one another

Retains fixed volume /shape –rigid particles locked in place

Compressible-lots of free space between particles

Not easily compressible -little free space between particles

Not easily compressible-little free space between particles

Flows easily-particles move past one another

Flows easily-particles move/ slide past one another

Does not flow easily-rigid particles cannot move/slide past one another

Diffusion within gas occurs rapidly-expand easily to fill available space

Diffusion within liquid occurs slowly-does not expand to fill container

Diffusion within solid occurs extremely slowly

Not dense More dense than gases More dense than liquid

No significant attractive forces between molecules

Significant attractive forces between molecules

Most significant forces between molecules

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Intermolecular ForcesIntermolecular Forces

• Ideal gas law used to describe gases– Volume of gas molecule is too small

• 99.9% of gas is empty space

– Gas molecules so far apart • No significant intermolecular attraction

– If gases truly ideal (zero volume/ attractive forces), couldn’t condense them to liquids

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• To condense real gases–Intermolecular attractive forces must

overcome KE of gas molecules

• Increasing attractive forces by–Decreasing distance between

molecules

–Increasing pressure which forces gas molecules closer together

• Decreasing temperature lowers average KE

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Three recognized types of Three recognized types of intermolecular attractionsintermolecular attractions

IntramolecularIntramolecular

IntramolecularIntramolecular

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Intermolecular forces Intermolecular forces (all electrostatic) van der Waals forces van der Waals forces (developed equation for

predicting deviation of gases from ideal behavior) Dipole-Dipole attractionDipole-Dipole attraction: dipole molecules orient themselves

so that +/ ends close to each other London Dispersion Forces London Dispersion Forces

Relatively weak forces that exist among noble gas atoms/ nonpolar molecules. (Ar, C8H18)

Caused by instantaneous dipole, in which electron distribution becomes asymmetrical

Ease with which electron “cloud” of atom can be distorted-polarizability

Hydrogen bondsHydrogen bonds: dipole-dipole attraction in which H is bound to highly electronegative atom (F/O/NF/O/N)

Ion-dipole force Ion-dipole force (solutions-also electrostatic)

Ion-Dipole ForcesIon-Dipole Forces

• Exists between ion and partial charge on end of polar molecule

• Important for solutions of ionic substance in polar liquids

– + ions surrounded by - ends of water molecules

– - ions surrounded by + ends of water molecule

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Dipole-Dipole ForcesDipole-Dipole Forces(about 1% strength of covalent bonds)

• Weaker than ion-dipole forces

• Neutral polar molecules (dipolesdipoles) attract each other with δ+ / δ- ends

• Responsible for mutual solubility of polar molecules, such as NH3 in H2O

• Explains how and why polar molecules may be condensed to liquid state

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London ForcesLondon Forces(dispersion forces, instantaneous dipole (dispersion forces, instantaneous dipole

forces, induced dipole forces)forces, induced dipole forces)

• Weakest forces of attraction• Main form of attraction in all nonpolar molecules

– No other type of van der Waals forces exist• Nonpolar He/Ar can be condensed, so some kind of

attractive interactions exist• Movement of electrons within electron cloud

cause instantaneousinstantaneous or momentarymomentary dipole moment (self-polarization)

– Resultant dipole induces polarization in neighboring molecules

– Mutual attraction of opposite ends of neighboring dipoles

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• As long as molecules are close together, this electron movement can occur over huge numbers of molecules

• Whole lattice of molecules can be held together in solid using van der Waals dispersion forces

• An instant later, distribution of electrons shifts

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Hydrogen BondingHydrogen Bondingspecial dipole-dipole attraction

• Hydrogen covalently bonded to highly electronegative elements (N, O, F)

• Bond strength higher than other dipole-dipole attractions

• Important in bonding of molecules (water/DNA)

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Strength of dispersion forces• Much weaker than covalent bonds

– Size of attraction varies considerably with size/shape of molecule

• In comparing intermolecular forces, molar mass is 1st consideration– When molar masses comparable, differences in

intermolecular attractions mainly due to differences in molecular polarity (strengths of dipole attraction)

– When molar masses differ greatly, differences in intermolecular attractions mainly due to differences in molar masses themselves (strength of London forces)

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– Magnitude of self-polarization increases with increasing numbers of electrons• Polarizability indicates ease electron

“cloud” of atom can be distorted to give dipolar charge distribution

– Large atoms with many electrons (electrons loosely held) exhibit higher polarizability than small atoms (electrons tightly held near nucleus)

• As atomic # increases, # electrons increases

– More electrons, more distance they can move– Increased chance of temporary dipole and

dipole interactions– Bigger the dispersion forces

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• Gaseous polar molecules show little attraction for each other (far apart)

– Rapidly become weaker as distance between dipoles increases

– Unimportant in gas phase due to distance between molecules

• In liquids and solids, molecules10x closer– Boiling point (condensation point) indicative of

attractive forces between molecules• Measure of how much KE must be increased so

that it can overcome attractive forces in liquid• Greater attractive forces/polarity, higher BP• Low boiling point indicates low attractive forces

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–BP/MP increase going down group• # e’s/atomic radius increase

–Why I is solid/Br liquid at room temperature

–Why all halogens have much higher BPs than their neighboring noble gases

• More electrons = more distance they move = bigger temporary dipoles (“stickier”) = bigger dispersion forces

–Because of greater temporary dipoles, xenon molecules are "stickier" than smaller neon molecules

–Neon molecules will break away from each other at much lower temperatures than xenon molecules – neon has lower boiling point

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• Shape of molecules– More linear molecules w/greater contact between

molecules have higher BP than more spherically shaped molecules

• Develop bigger temporary dipoles due to electron movement than short fat ones containing same numbers of electrons

• Long thin molecules lie closer together-attractions most effective if molecules really close

– Responsible for mutual solubility of nonpolar molecules such as Br2 in CCl4

• Large electron clouds easily polarized, so they have higher BPs than neighboring noble gases

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• The boiling point of Argon is -189.4oC. – Why is it so low?

• Argon does not interact with other substances because it is so small and has a complete octet of valence electrons. Argon must be made quite cool to allow liquefication via London dispersion forces.

– How does this boiling point prove that London dispersion forces exist?

• If these forces did not exist, Argon would never liquefy.

– The boiling point of Xenon is -119.9oC. Why is it higher than that of Argon?

• Xenon is bigger and has more electrons than Argon. The likelihood of momentary dipoles is thus greater. It has a greater polarizability than Argon.

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• Put the following substances in order from lowest to highest boiling points:– C2H6

– NH3

– F2

• F2 can only exhibit intermolecular London forces.

• C2H6 is not especially polar, but it does have a very slight electronegativity difference between the carbons and the hydrogens.

• NH3 exhibits hydrogen bonding, thus giving it a relatively high boiling point.

• F2, C2H6, NH3

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The Liquid StateThe Liquid State

Properties of Liquids

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Surface TensionSurface Tensionresistance of liquid to increase in surface area (polar molecules)

• Increase in attractive forces between molecules at surface compared to forces between molecules in center

– Interior molecules attracted to molecules from every direction

• No net force on molecule (pulled in all directions)

– Surface molecules don’t have molecules to attract it on top side

• Molecule on surface drawn into bulk of liquid• Same total attractive force divided between fewer adjacent

molecules, resulting in stronger attractive force of surface molecule toward bulk of liquid

• Decreases with temperature (reduces intermolecular attractions

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• Intermolecular attractive forces act to minimize surface area of liquid– Geometric shape w/smallest ratio of surface area to

volume is sphere– Surface molecules drawn inward and area of

surface of liquid reduced– Small quantities of liquid form spherical drops– As drops get larger, weight deforms them into

typical tear shape

• Bubbles (hollow drop)Bubbles (hollow drop)– Surface tension acts to minimize surface Surface tension acts to minimize surface

(radius of spherical shell of liquid), but (radius of spherical shell of liquid), but opposed by pressure of vapor within bubbleopposed by pressure of vapor within bubble

– Bubbles in pure water tend to collapseBubbles in pure water tend to collapse– Bubbles with surfactant are stabilizedBubbles with surfactant are stabilized

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• Surface film– Smaller surface area causes liquid to behave as

though it has skin on surface– Surface tension enables insects to walk on

surface of water– High intermolecular forces = greater surface

tension

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Cohesion-AdhesionCohesion-Adhesion• Liquids confined within

a container have two types of forces present on molecules

– Cohesive forces-intermolecular forces among like molecules

– Adhesive forces-forces between unlike molecules for one another (polar molecules, oxygen in glass attracted to hydrogen in water)

Cohesion causes water to form drops, surface

tension causes them to be nearly spherical, and

adhesion keeps the drops in place

WaterWater

• Adhesive forces greater than cohesive forces

• MeniscusMeniscus (curved upper surface) is concave (curved downward)

MercuryMercury

• Cohesive forces greater than adhesive

forces

• Meniscus is convex (curved upward)

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• Capillary actionCapillary action– Instantaneous rising of liquid in

narrow tube– Combination of cohesion/adhesion– Distance traveled dependent on

diameter of tube

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• Which would have a higher surface tension, H2O or C6H14? Why? Would the shape of the H2O meniscus in a glass tube be the same or different than C6H14?– Water, having large dipole moment, has relatively

large cohesive forces. Hexane is essentially nonpolar so it has low cohesive forces. Water would have higher surface tension.

– Water meniscus is concave because adhesive forces of water to polar constituents on surface of glass are stronger than its cohesive forces. Hexane would have a convex meniscus because it has very small adhesive forces, and the slightly larger cohesive forces would dominate.

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ViscosityViscositymeasure of liquid’s resistance to flow

• Increases with intermolecular forces– To flow, liquid’s molecules must move past each

other– Move more freely in solutions with relatively low

attractive forces– Nonpolar molecules attracted to each other by only

London forces have lower viscosities than polar liquids like water

• Increases with molecular size• Is temperature dependent

– Decrease in viscosity with increased temperature – Attributed to average molecular KE of liquid which

overcomes attractive forces between molecules

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Homework:

Read 10.1-10.2, pp. 449-455Read 10.1-10.2, pp. 449-455

Q pp. 500-501, #13, 15, 36, 38, 40, 41, 42Q pp. 500-501, #13, 15, 36, 38, 40, 41, 42

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Solids

Structure and TypesStructure and Types

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Amorphous solidsAmorphous solidsnoncrystalline solids

• Greek “without form” – No orderly structure (arrangement)– Lack well-defined faces/shapes– Many are mixtures of particles that don’ stack

together well

• Plastic, glass, rubber

• No distinct, sharp melting point, but soften gradually over large temperature range

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Crystalline SolidsCrystalline Solids• Atoms, ions, or molecules

ordered in well-defined 3-D arrangements (lattice)

• Unit cellUnit cell: smallest repeating unit of lattice

• LatticeLattice: unit cells repeated in space in all three dimensions, characteristic of crystalline solid

– Lattice point: Lattice point: part of atom in lattice

– Way spheres arranged in layers determines what type of unit cell we have

Primitive or Simple CubePrimitive or Simple Cube

ShapeShape• Atom in each corner

of cube• Atoms in contact

along cell edge

StackingStacking• Spheres in each layer

lay on spheres below/above them– Stacking pattern:

AAAAAAA. . .

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Coordination #• # # atoms/ions surrounding

atom/ion in crystal lattice• 6 (in red)-has 6 immediate

neighbors• Value gives measure of how

tightly spheres packed together– Larger coordination #, closer

spheres to each other

– Not close packed (least efficient method-52%)

– Very rare packing arrangement for metals (ex. Polonium)

Atoms/unit cellAtoms/unit cell

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Body Centered Cubic (BCC)Body Centered Cubic (BCC)

ShapeShape• Atom at each corner/in

center of cube– Atoms in contact along

body diagonal

StackingStacking• 2nd layer fit into

depressions of 1st layer/3rd layer into 2nd

– Stacking pattern: ABABABAB. . .

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Coordination #• 8 (each sphere in contact

with 4 spheres in layer above/4 below)

• Not close packed (less efficient packing-68%)

Atoms/unit cellAtoms/unit cell

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Face-centered Cube (FCC)

Shape• Atom at each

corner/atom in center of each face of cube

StackingStacking• Each layer diagonally

next to each other. Alternating layer in crevices between spheres.– Stacking pattern:

ABABABAB. . .

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Coordination #Coordination #• 12 (more efficient at 74%

packing)• Not closest packed (Fe,

alkali metals)

Atoms/unit cellAtoms/unit cell

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Determine net number of Na+ and Cl- ions in unit cell

• Na+ – (¼ per edge) x (12 edges) = 3– (1 per center) x (1 center) = 1

• Cl- – (1/8 per corner) x (8 corners) = 1– (½ per face) x (6 faces) = 3

• Correct since # Na = # Cl– (4 = 4)

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Bravais lattice Bravais lattice is infinite array of discrete points with arrangement and orientation that appears exactly the same, from whichever points array viewed

Introduction to structures and types of solids

X-ray analysis of solidsX-ray analysis of solids

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DiffractionDiffraction• Scattering of light from regular array of points or lines• Spacing between points/lines/atoms, related to wavelength

of light• X-rays used because their wavelengths similar to distances

between atomic nuclei

Reflection of X-rays of wavelength from pair of atoms in two different layers of crystal

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• If incident X-ray beam hits crystal lattice, general scattering occurs– Most scattering

interferes w/itself and is eliminated (destructive destructive interferenceinterference)

– (b) Incident rays in phase but reflected rays out of phase: d2 (difference in distances traveled by two rays) = odd # of half wavelengths

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• Diffraction occurs when scattering in certain direction is in phase with scattered rays from other atomic planes– Combine to form waves

that reinforce each other (constructive constructive interferenceinterference)

– (a) Incident/reflected rays in phase: d1 (difference in distances traveled by two rays ) = whole # of wavelengths

X-ray diffractionX-ray diffraction

• One of best methods to determine crystal's structureOne of best methods to determine crystal's structure• Intense X-ray beam strikes crystalIntense X-ray beam strikes crystal• Crystal diffracts X-ray beam differently, depending Crystal diffracts X-ray beam differently, depending

on structure/orientationon structure/orientation– Atoms in crystal interact w/x-ray waves to produce Atoms in crystal interact w/x-ray waves to produce

interferenceinterference

• Diffraction pattern consists of reflections of different Diffraction pattern consists of reflections of different intensity used to determine crystal’s structureintensity used to determine crystal’s structure

• However, many different orientations of crystal However, many different orientations of crystal collected before true structure of crystal determinedcollected before true structure of crystal determined

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Bragg EquationBragg Equationresolution of X-ray diffraction detector

• Used for analysis of crystal structures– Each crystalline material has

characteristic atomic structure, diffracts X-rays in unique characteristic pattern

n = 2d sin

d = distance between atoms (unique for each mineral)

n = an integer = wavelength of x-rays

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• X rays of wavelength 1.54 Å were used to analyze an aluminum crystal. A reflection was produced at Ø = 19.3 degrees. Assuming n = 1, calculate the distance d between the planes of atoms producing this reflection.

• 2dsinθ = nλ

• 2(d)(sin 19.3) = 1(1.54 Å)

• 2(d)(0.3305) = 1.54 Å

• d = 2.33 Å = 233 pm

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• A topaz crystal has a lattice spacing (d) of 1.36 Å (1 Å = 1 x 10-10 m). Calculate the wavelength of X-ray that should be used if Ø = 15.0o (assume n = 1)

• 2dsinθ = nλ

• 2 (1.36 Å)(.259) = 1(λ)

• 0.704 Å = 70.4 pm

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Types of Crystalline Solids

(a) Atomic solids (b) Ionic solids (c) Molecular solids

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Structural particles

Principal attractive forces between particles

Characteristics (physical behavior) Examples

Atoms Nondirectional covalent bond involving positive ions and mobile valence electrons delocalized throughout crystal (metallic bond)

•Pure elements•All solids at 25oC except Hg•Wide range hardness (electrons move freely from atom to atom, bond metal atoms together with widely varying degrees of force)•Wide range melting point (between ionic and covalent compounds)•Excellent thermal/electrical conductors (mobile electrons quickly carry charge throughout metal)•Malleable/ductile (nondirectional, so stress alters but not destroys crystal)•Lustrous (interaction of electrons w/light)

•K•Na•Fe•Mg•Ca•Zn

Structure and Bonding in Metals

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Closest packingClosest packingmost efficient arrangement of spheres

• Uniform/hard spheres most efficiently use available space• Coordination number = 12Coordination number = 12

– Each sphere in contact with Each sphere in contact with • 6 spheres in its own layer6 spheres in its own layer• 3 spheres in layer above 3 spheres in layer above • 3 spheres in layer below3 spheres in layer below

• Stacking – 2nd layer does not lie directly over those in 1st layer– 3rd layer occupies dimples of 2nd layer in two ways

• Each sphere in 3rd layer lies directly over sphere in 1st layer (aba)• Each sphere can occupy positions so that no sphere in 3rd layer

lies over one in 1st layer (abc)

Hexagonal closest packing (hcp)Hexagonal closest packing (hcp)

• aba arrangement– Each layer identical to

layer two below it– Ex.: Be, Mg, Ti, Co,

Zn, Cd, He (at low T)– 2 atom unit cell (8 x 1/8

+ 1)– Packing efficiency =

74%

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Cubic close-packed (ccp)Cubic close-packed (ccp)

• Corresponds to face-centered cube (abc)– Each layer identical to

layer three below it

• Ex.: Ca, Sr, Ni, Pd, Pt, Cu, Ag, Au, Pb, Pt, Ne/Ar/Kr/Xe (at low T)

• Packing efficiency = 74%

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Coordination Numbers for Common Crystal Structures

Structure  Coordination Number  Stacking Pattern

simple cubic   6  AAAAAAAA. . .body-centered cubic  8  ABABABAB. . .hexagonal closest-packed  12  ABABABAB. . .cubic closest-packed  12  ABCABCABC.

Calculate density of LiF

• 4.02 Å on edge

• Same arrangement of ions as NaCl

• 4(6.94 amu) + 4(19.0 amu) = 103.8 amu

D = 103.8 amu 1 g (1 Å)3 = 2.65 g/cm3

4.02 Å 6.02 x 1023 amu (10-8)-3 cm

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Calculating density of closest packed solid

• Silver crystallizes in a cubic closest packed structure. The radius of a silver atom is 144 pm. Calculate the desnity of solid silver.

• Textbook-length of edge of cube: d = r8 = 1448 = 407 pm• V = d3 = (407 pm)3 = 6.74 x 107 pm3 • 6.74 x 107 pm3 x (1 x 10-10cm)3/1pm = 6.74 x 10-23 cm3

• D = m = (4 atoms)(107.9 g/mol)(1 mol/6.022 x 1023 atoms) V 6.74 x 10-23 cm3

• 10.6 g/cm3

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• The radius of nickel atom is 1.24 Å (1Å = 1 x 10-8 cm). Nickel crystallizes with a cubic closest packed structure. Calculate the density of solid nickel.

• d = r8 = 1.24 Å8 = 3.51 Å

• V = d3 = (3.51 Å)3 = 43.2 Å3

• 43.2 Å3 x (1 x 10-8cm)3/1pm = 4.32 x 10-23 cm3

• D = m = (4 atoms)(58.69 g/mol)(1 mol/6.022 x 1023 atoms)

V 4.32 x 10-23 cm3

• 9.04 g/cm3

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Bonding Models for MetalsBonding Models for Metals•Electron Sea ModelElectron Sea Model: regular array of metal atoms in “sea” of electrons

•Band (Molecular Orbital) Band (Molecular Orbital) ModelModel: electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms

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Metal AlloysMetal Alloys

• Substance containing mixture of elements/has metallic properties

• Substitutional alloySubstitutional alloy– Host metal alloy atoms replaced by other

atoms of similar size – Brass-Cu/Zn, sterling silver-Ag/Cu,

pewter-Sn/Cu/Bi/Sb, plumber’s solder-Sn/Sb

• Interstitial alloyInterstitial alloy– Holes in closest packed metal structure

occupied by small atoms – Steel

Network Crystal

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Structural particles Principal attractive forces between particles

Characteristics (physical behavior) Examples

Nonmetal atoms (pure elements)

Directional covalent bonds leading to giant molecules (metallic bonds)

•Very hard (brittle)•Insoluble in most ordinary liquids•Sublime or melt at high temperatures (high MP)•Poor thermal and electrical conductors (most insulators)

•C (diamond, graphite)•SiC•BN•SiO2 (sand,

quartz)

DiamondDiamond GraphiteGraphite Quartz

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Semiconductors• Material w/electrical conductivity between

conductor and insulator• Conductivity is enhanced by dopingdoping

– In lattice, all atoms bond to 4 neighbors, leaving no free electrons to conduct current (insulator)

– Intentionally introducing impurities into extremely pure semiconductor which allows electrons in bonds to move

• Make electrons available for conduction (VA)• Form holes which can conduct current (IIIA)

http://solarhorizon.com.au/Flash/HoleMovementMovie.html

http://solarhorizon.com.au/Flash/DopingMovie.html

Polar Molecular CrystalsPolar Molecular Crystals

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Structural particles Principal attractive forces between particles

Characteristics (physical behavior) Examples

Discrete polar molecules occupy lattice point

Dipole-dipole attractions

•Low to moderate MP•Soluble in some polar liquids•Poor thermal/electrical conductors

•HCl•CHCl3•H2S

•C6H12O6

Molecules with H bonded to O, N, F

Hydrogen bonds •Low to moderate MP•Soluble in some hydrogen-bonded and some polar liquids•Poor thermal/electrical conductors

•H2O

•NH3

•HF• CH3OH

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Nonpolar Molecular CrystalsStructural particles Principal attractive forces

between particles

Characteristics (physical behavior) Examples

Atoms (8A), nonpolar molecules

Dispersion forces •Extremely low to moderate MP•Soluble in nonpolar solvents•Poor thermal/electrical conductors

ArCl2H2

CH4

I2

CO2

CCl4

Ionic CrystalsIonic Crystals

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Structural particles Principal attractive forces between particles

Characteristics (physical behavior) Examples

Positive and negative ions that give strongest attractive forces in chemistry

Ion-ion attraction

•Almost all have rigid lattice•Moderate to high MP/BP (large lattice energy required to separate ions)•Hard/brittle (when hit, atoms shift position which breaks +/- attraction)•Nonconductors as solids but conductors as liquids•Many dissolve in water

KClCaF2

CsBrMgOBaCl2NaCl

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Packing arrangement done to minimize anion-anion and cation-cation repulsions

• Structure of most binary ionic solids explained by closest packing of spheres

• Anions usually larger than cations packed as hcp or ccp arrangements with cations filling holes

• Nature of holes depends on anion : cation size– Trigonal holes formed by 3 spheres in same layer– Tetrahedral holes formed by sphere sitting in dimple of

three spheres in adjacent layer– Octahedral holes formed between 2 sets of 3 spheres in

adjoining layers of closest packed structures

Trigonal holes < tetrahedral holes < octahedral holes

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• Would AlP have a closest packed structure Would AlP have a closest packed structure which is more like NaCl or ZnS? Ionic radii: which is more like NaCl or ZnS? Ionic radii: – AlAl3+3+ = 50 pm, P = 50 pm, P3-3- = 212 pm = 212 pm– ZnZn2+2+ = 74 pm, S = 74 pm, S2-2- = 184 pm = 184 pm– NaNa++ = 95 pm. Cl = 95 pm. Cl-- = 181 pm = 181 pm

• Take anion to cation ratio in each case:Take anion to cation ratio in each case:– SS2-2-/Zn/Zn2+2+ = 2.49 (tetrahedral holes) = 2.49 (tetrahedral holes)– ClCl--/Na/Na+ + = 1.91 (octahedral holes)= 1.91 (octahedral holes)– PP3-3-/Al/Al3+3+ = 4.24 (?) = 4.24 (?)

• Aluminum ions very small compared to Aluminum ions very small compared to phosphorus ion, so not much room is phosphorus ion, so not much room is needed. Tetrahedral holes are adequate. needed. Tetrahedral holes are adequate. So AlP is more like ZnS than NaCl.So AlP is more like ZnS than NaCl.

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• Using table 10.7, and based on their properties, Using table 10.7, and based on their properties, classify each of the following substances as to classify each of the following substances as to the type of solid it forms.the type of solid it forms.

• FeFe• Atomic solid with metallic propertiesAtomic solid with metallic properties• CC22HH66

• Contains nonpolar molecules-molecular solidContains nonpolar molecules-molecular solid• CaClCaCl22• Contains CaContains Ca2+2+ and Cl and Cl- - ions-ionic solidions-ionic solid• GraphiteGraphite• Made up of nonpolar carbon atoms covalently Made up of nonpolar carbon atoms covalently

bonded in directional planes-network solidbonded in directional planes-network solid• FF22

• Nonpolar fluorine molecules-molecular solidNonpolar fluorine molecules-molecular solid

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Homework:

Read 10.3-10.7, pp. 456-483Read 10.3-10.7, pp. 456-483

Q pp. 501-504, #46, 48, 60, 68, 72Q pp. 501-504, #46, 48, 60, 68, 72

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Vapor Pressure

And change of stateAnd change of state

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Vaporization (Evaporation)Vaporization (Evaporation)

• Escape of surface liquid molecules to form gas• Rate depends on

– Nature of liquid– Surface area (↑ SA, ↑ evaporation)– Temperature (↑ T, ↑ proportion of molecules w/ KE

above escape energy)

• Always endothermic (positive)• Heat of vaporization (Enthalpy of vaporization, Heat of vaporization (Enthalpy of vaporization,

ΔΔHHvapvap) ) --energy required to vaporize one mole of liquid at 1 atm

Vapor pressureVapor pressure

• Develops in gas phase Develops in gas phase above liquid when liquid above liquid when liquid placed in closed containerplaced in closed container

• When evaporation occurs in closed container, gas molecules cannot escape to surroundings

• As more molecules enter gas phase, pressure increases, finally stopping at level (vapor pressurevapor pressure) dependent only on temperature

Equilibrium vapor pressureEquilibrium vapor pressure

• Pressure of vapor present Pressure of vapor present at equilibriumat equilibrium

• Gas molecules collide w/container walls/liquid– Most re-condense when

collide with liquid

• Initially, evaporation rate greater condensation rate

• As # gas molecules increase, collisions w/liquid surface increase

• Rate of evaporation eventually equals to rate of condensation-equilibrium

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At given temperature, not all molecules moving w/same KE. Small # molecules moving very slow (low KE), while few moving very fast (high KE). Vast majority somewhere in between.

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Variations in Vapor PressureVariations in Vapor Pressure

• Related to intermolecular attractive forces– Liquids w/high intermolecular attraction have relatively

low vapor pressures (less volatilevolatile-liquids that evaporate readily)

– Liquids w/low intermolecular attraction have relatively high vapor pressures (more volatile)

• For similar-size molecules– Hydrogen-bonded substances have largest ΔHvap

values (less volatile/lower vapor pressure)– Polar substances have higher values than

similar-size nonpolar substances

• Vapor pressure increases w/temperature

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Boiling PointBoiling Pointvapor pressure of liquid = prevailing atmospheric

pressure above that liquid

• Increasing temperature increases KE increases molecular motion

• Forces of attraction between molecules (H bonding) disrupted

• Molecules break free of liquid and become gas

• At boiling point, liquid turns into gas

• Normal boiling pointNormal boiling point-BP of liquid at 1 atm

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Clausius-Clapeyron equationClausius-Clapeyron equation• Mathematical relationship between heat of vaporization and vapor

pressure as measures of intermolecular forces that attract molecules together in liquid state

– Relationship between vapor pressure and temperatureRelationship between vapor pressure and temperature

– P = vapor pressure– ΔHvap = heat of vaporization– R = universal gas law constant– T = Kelvin temperature– C = constant (eliminated in 2nd equation)

• Based on following assumptions (fail at high P, near critical point)– Volume of vaporized liquid negligible compared to volume of vapor– Vapor behaves as ideal gas– ΔHvao is constant over temperature interval of data– External pressure doesn’t affect vapor pressure

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• (a) Vapor pressure as function of temperature

• Quantitative nature of temperature dependence of vapor pressure

• Nonlinear

• (b) Plots of In(Pvap) versus 1/T (Kelvin temperature)

• Linear

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Clausius-Clapeyron

• Pg. 487, Sample 10.6• The vapor pressure of water at 25oC is

23.8 torr, and the heat of vaporization of water is 43.9 kJ/mol. Calculate the vapor pressure of water at 50oC.

• ln (23.8 torr) = -43,900 J/mol ( 1 – 1 ) PvapT2 torr 8.3145 J/K mol 323 K 298 K

• ln (23.8) = -1.37 PvapT2 torr• Take the antilog of both sides to get 23.8 =0.254 = 93.7 torr

PvapT2 torr

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• Water has a vapor pressure of 24 mmHg at 25oC and a heat of vaporization of 40.7 kJ/mol. What is the vapor pressure of water at 67oC?

• Solution: Simply use the Clausius-Clapeyron Equation to figure out the vapor pressure. We have to be a bit careful about the units of R: the units we're using are kJ, so R = 8.31x10-3 kJ/mol K.

ln(P2/P1) = -Hvap/R * (1/T2- 1/T1)

ln(P2/24) = -40.7 kJ/8.31x10-3 kJ/mol K *(1/340- 1/298)

ln(P2/24) = 2.03

P2/24 = 7.62

P2 = 182 mmHg

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• An unknown liquid has a vapor pressure of 88mmHg at 45oC and 39 mmHg at 25oC. What is its heat of vaporization? Solution: Again, use the Clausius-Clapeyron Equation. Here, the only thing we don't know is Hvap

ln(P2/P1) = -Hvap/R * (1/T2- 1/T1)

ln(88/39) = -Hvap/8.31x10-3*(1/318 - 1/298)

Hvap = 32.0 kJ

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• The vapor pressure of 1-propanol at 14.7oC is 10.0 torr. The heat of vaporization is 47.2 kJ/mol. Calculate the vapor pressure of 1-propanol at 52.8oC.

• ln(10.0/x) = - 47.2 kJ/mol/0.008314 kJ/K mol x (1/326 K – 1/287.9 K)• 100.2 torr

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Changes of StateChanges of State

Heating curvesHeating curves

Phase diagramsPhase diagrams

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•T/KE increases and PE constant•Only 1 phase present

•Ice starts to melt at 0oC•T/KE constant due to heat of fusion-80 calories/g (bonds not broken). PE increases until all ice melted•Solid/liquid in equilibrium

•T/KE increases and PE is constant•Only 1 phase present

•Water starts to boil at 100oC•KE constant due to heat of vaporization-539 cal/g PE increases until all water evaporated•Liquid/gas in equilibrium

•T/KE increases, PE is constant•Only 1 phase present

Heat capacity (Heat capacity (J/kg-oC-1)

• Amount of energy required to raise temperature of system by 1o

Heat of fusion (Enthalpy Heat of fusion (Enthalpy of fusion, of fusion, ΔΔHHfusfus))

• Energy required to convert mole of solid to mole of liquid at constant pressure

– Determined as length of first (solid-melting) plateau, which represents heat added, divided by # moles of sample

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Heat of vaporization Heat of vaporization (Enthalpy of (Enthalpy of vaporization, vaporization, ΔΔHHvapvap))

• Heat absorbed by one mole of liquid when it changes to gas at constant pressure

• Length of second (liquid-boiling) plateau divided by # moles of sample

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SublimationSublimationsubstance goes directly from solid to gaseous state

• Reasons for sublimation– Solids have vapor pressure, but it is normally very low– Solids with little intermolecular attraction may have

substantial vapor pressures and be able to sublime at room conditions.

• Enthalpy of sublimation (Δhsub)

– Energy in solid-gas transition

– State function, so ΔHsub = ΔHfus + ΔHvap

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Cooling CurveCooling Curve

Reverse-all features same except start with gas and condense to get solid as heat is removed

Condensation point/Crystallization point-used in place of BP/MP

ΔHcond requires gas to give off heat-always negative (exothermic)

ΔHvap = -ΔHcond

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Phase DiagramPhase Diagram

• relationship between pressure relationship between pressure and temperature and the three and temperature and the three states of matterstates of matter

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FusionFusion curvecurve(solid-liquid line(solid-liquid line Vapor pressure curveVapor pressure curve

Liquid-gas lineLiquid-gas line

SublimationSublimation curvecurveSolid-gas lineSolid-gas line

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Diagram for closed system

• Pressure plotted on y-axis/temperature on x-axis• Divided into three physical states by three lines

meeting at triple point– Solids-upper left– Liquids-upper right– Gases-lower part

• Along each line is equilibrium mixture of two phases on two sides of that line

– Liquid-solid line usually straight line-changes in pressure have very little effect on solids and liquids (only slightly compressible)

– Gas-solid and liquid-gas lines are curved upward

• Normal melting Normal melting pointpoint-temperature at which solid and liquid states have same vapor pressure under conditions where total pressure is 1 atmosphere

• Normal boiling Normal boiling pointpoint-temperature at which vapor pressure of liquid is exactly 1 atmosphere

• SupercoolingSupercooling -process of cooling liquid below its freezing point without its changing to solid

• SuperheatingSuperheating -process of heating liquid above its boiling point without its boiling

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Triple PointTriple Point• Each phase has

same temperature and vapor pressure

• All three phases exist together in equilibrium– Helium-only

substance that doesn’t have triple point since it has no solid phase

Critical PointCritical Point• Maximum temperature at which any

liquid can exist , defined by• Critical temperature-temp above

which substance can’t liquefy gas, regardless of how great pressure

• Critical pressure-pressure above which substance can no longer exist as gas, no matter how high temp– For water: 374°C and 218 atm– Depends on intermolecular attraction

(greater intermolecular forces, higher critical temperature)

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• Non-compressible/high density fluid

• Obtained by either• Heating gas above its critical T• Compressing liquid at higher P

than its critical pressure

• Above critical point, differences between gases and liquids disappear

• Density of gas can approach density of liquid phase

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Homework:Read 10.8-10.9, pp. 483-497Read 10.8-10.9, pp. 483-497Q pp. 500-505, #23, 29, 76 (have fun), 78, 88, 90Q pp. 500-505, #23, 29, 76 (have fun), 78, 88, 90Do 1 additional exercise and 1 challenge problemDo 1 additional exercise and 1 challenge problemSubmit quizzes by email to me:Submit quizzes by email to me:http://www.cengage.com/chemistry/book_content/0547125321_zumdahl/ace/launch_ace.html?folder_path=/chemistry/book_content/0547125321_zumdahl/ace&layer=act&src=ch10_ace1.xmlhttp://www.cengage.com/chemistry/book_content/0547125321_zumdahl/ace/launch_ace.html?folder_path=/chemistry/book_content/0547125321_zumdahl/ace&layer=act&src=ch10_ace2.xmlhttp://www.cengage.com/chemistry/book_content/0547125321_zumdahl/ace/launch_ace.html?folder_path=/chemistry/book_content/0547125321_zumdahl/ace&layer=act&src=ch10_ace3.xml