Catherine MacGowan Armstrong Atlantic State University Lecture Presentation Chapter 11 Liquids, Solids, and Intermolecular Forces © 2013 Pearson Education,

Catherine MacGowan

Armstrong Atlantic State University

Lecture Presentation

Chapter 11

Liquids, Solids, andIntermolecular Forces

© 2013 Pearson Education, Inc.


• SOLIDS– Have rigid shape, fixed volume. External shape can

reflect the atomic and molecular arrangement.• Reasonably well understood

• LIQUIDS– Have no fixed shape and may not fill a container

completely• Not well understood

• GASES– Expand to fill their container

• Good theoretical understanding

The States of Matter


Changing Phases

• Matter can be transformed from one phase to another by changing the:– temperature (addition or release of energy)– pressure


A Molecular Comparison of the Three Phases of Water


A Graphical Representation ofthe Phases of Water: A Phase Diagram


Strong Covalent Bond

Intramolecular Forces

Very Weak AttractionsVery Weak Attractions

Intermolecular ForcesIntermolecular Forces

Intermolecular Forces

O = OO = O


Why is water usually a liquid and not a gas?

Why does liquid water boil at such a high temperature for such a small molecule?

Why does ice float on water?

Why is I2 a solid whereas Cl2 is a gas?

Why are NaCl crystals little cubes?

The answer: Intermolecular Forces

Questions:


Intramolecular

• Attractive forces that hold atoms together to form a bond

• Types: – Ionic– Metallic– Covalent

• True• Polar

Intermolecular

• Attractive forces between molecules and/or ions

• Types– Ion-ion– Ion-dipole– Dipole-dipole

• Hydrogen bonding– Dipole-induced– Induced-induced

• Dispersion• Van der Waals

Intra- vs. Intermolecular Forces


• When comparing one on oneone on one,, intermolecular forces are WEAKER in strength than intramolecular forces.

• What makes intermolecular bonds APPEAR stronger than intramolecular forces is in the numbers.

– Have “many” intermolecular interactions in a solid or liquid, not just one

• Networking• Spider web

• Size of charge and distance between charges play a role in intermolecular attraction strength.

– Larger the charge, the stronger the attraction

– Greater the distance, the weaker the attraction

Intermolecular Forces


• Particles are attracted to each other by electrostatic attractive forces.

– It is a positive-negative type interaction.

• The strength of the attractive forces vary; some are small and some are large.

– Type of particles (e.g., molecule or ion)

– Size of ion

– Polarity strength

• The stronger the attractive forces between the particles, the more they resist moving or breaking apart.

– No material completely lacks particle motion.

Intermolecular Forces Are Attractive Forces


Types of Intermolecular Forces


• Usually associated with ionic Usually associated with ionic compounds.compounds.

– Cation to anionCation to anion

• Crystalline structureCrystalline structure

– Ordered arrangementOrdered arrangement

• Strongest of intermolecular Strongest of intermolecular forcesforces

– 400 to 4000 kJ/mole400 to 4000 kJ/mole

– Range due to ion sizeRange due to ion size

• Explains why ionic solids tend Explains why ionic solids tend to have HIGH melting points to have HIGH melting points temperaturestemperatures

Ion–Ion Attractive


• Intermolecular attractions Intermolecular attractions observed with a polar molecule observed with a polar molecule (dipole) and ion(dipole) and ion

– Salt (ionic compounds) Salt (ionic compounds) dissolving in waterdissolving in water

• Strength of attraction affected Strength of attraction affected by size of ion and polarity of by size of ion and polarity of dipoledipole

– 1/101/10thth or 10% strength of a or 10% strength of a covalent (intramolecular) covalent (intramolecular) attractionattraction

– 40 to 600 kJ/mole40 to 600 kJ/mole

Ion–Dipole (Permanent) Attractions


• Dipole-dipole forces bind molecules having permanent dipoles to one another.

– Polar molecule interactions

• 1 to 5% the strength of a covalent bond

– 5 to 40 kJ/mole

• Strength varies.

– Depends on polarizing strength

• Special dipole-dipole interaction is hydrogen bonding.

Dipole (Permanent)–Dipole (Permanent) Interactions


• It is a special type of dipole-dipole attraction, which enhances dipole-dipole attractions.

Types of Hydrogen Bonds [ X – H - - - :Y ]

N – H - - - :N - O – H - - - :N - F – H - - - :N -N – H - - - :O - O – H - - - :O - F – H - - - :O -N – H - - - :F - O – H - - - :F - F – H - - - :F-

H bonding is strongest when X and Y are N, O, or F.

Hydrogen Bonding


H bonding is especially strong in water H bonding is especially strong in water because:because:

– the O—H bond is very polarthe O—H bond is very polar

– there are 2 lone pairs on the O there are 2 lone pairs on the O atomatom

H bonding accounts for many of water’s H bonding accounts for many of water’s unique properties.unique properties.

– High boiling point compared to High boiling point compared to others of similar molecular massothers of similar molecular mass

– Liquid’s density is greater than Liquid’s density is greater than its solid’s density.its solid’s density.

– High heat capacity: 4.184 J/gHigh heat capacity: 4.184 J/gooCC

Hydrogen Bonding and Water


Water Has a Relatively High Boiling Point Compared to Other Molecular Compounds


• The interaction between a The interaction between a polarized/charged particle and a nonpolar polarized/charged particle and a nonpolar particleparticle

• Strength varies from 2 to 10 kJ/mole.Strength varies from 2 to 10 kJ/mole.– Depends on:Depends on:

• Polarizability of the nonpolar Polarizability of the nonpolar particleparticle

• Charge and size of ionCharge and size of ion• Polarity of the polar moleculePolarity of the polar molecule

• Two ways to induce a dipole moment in a Two ways to induce a dipole moment in a nonpolar particlenonpolar particle– Ion to induced dipoleIon to induced dipole– Dipole to induced dipoleDipole to induced dipole

• This interaction explains how nonpolar This interaction explains how nonpolar molecules such as Omolecules such as O2 2 and Iand I22 can dissolve can dissolve

in water (polar).in water (polar).

Induced–Induced Dipoles: Dispersion Forces


• Formation of an induced dipole between two nonpolar molecules

• Interaction (contact) cause an instantaneous dipole to form.

• Associated with molecules having covalent bonding

• Organic (nonpolar substances)

• Non metals

• Weakest of the intermolecular forces

• 0.05–40 kJ/mole

• Polariziability of nonpolar molecule a factor in strength

• Number of contact interactions

• Size of molecule (molecular mass)

• Also referred to as:

• Dispersion forces

• Van der Waal forces No dipole No dipole

• Explains:

• Why butter melts but does not form “butter gas”

• Why gases can liquefy

Dispersion Forces

I - I I - I


Relationship between Molecular Shapeand Dispersion Forces

• Dispersion forces are the weakest of all the intermolecular forces.

– The more “contact” spots, the greater the effect dispersion forces have on a molecule’s physical properties such as boiling point.

• Example:

– n-pentane (36 oC) vs. neopentane (2,2-dimethyl propane )(10 oC)


Boiling Points of Hydrocarbons

Note linear relation between boiling point and molar mass of the compounds.


Intermolecular Forces and Physical Properties


Surface Tension

•It is a liquid’s resistance to increase its surface area.• Liquid spontaneously seeks to minimize its surface area.• To minimize surface area, liquids form drops that are spherical.

•It is the energy that is required to increase a liquid’s surface area.• Energy required is directly related to the type of intermolecular forces within the liquid.

• The greater/stronger the intermolecular forces, the greater energy required to increase the surface

area.

•Surface tension of H2O = 72.8 mJ/m2 at room temperature

•Surface tension of C6H6 = 28 mJ/m2

•Temperature affects a liquid’s surface tension.–Increase temperature, reduce surface tension

•An increase in average kinetic energy of molecules•More energy to “break” intermolecular bonds


Capillary Action

• It is the spontaneous rising of a liquid in a It is the spontaneous rising of a liquid in a narrow tube or movement up a piece of narrow tube or movement up a piece of paper against the pull of gravity.paper against the pull of gravity.

• It is the result of two forces working together: cohesive and adhesive forces. – Cohesive forces hold the liquid

molecules together.– Adhesive forces attract the outer liquid

molecules to the tube’s surface.• Shape of meniscus indicates whether

or not similar polarities


Capillary Action and Meniscus

• The curving of the liquid surface in a thin tube is due to the competition between adhesive and cohesive forces.

• The meniscus of water is concave in a glass tube because its adhesion to the glass is stronger than its cohesion to itself.

• The meniscus of mercury is convex in a glass tube because its cohesion to itself is stronger than its adhesion to the glass.

– Metallic bonds are stronger than intermolecular attractions.

Water MercuryWater Mercury


Capillary Action and Absorption

• Movement of water up a piece of paper depends Movement of water up a piece of paper depends on H bonds between Hon H bonds between H22O and the OH groups of O and the OH groups of

the cellulose in the paper.the cellulose in the paper.

• Chromatography separates the components of a mixture by their distinctive attraction to the mobile phase and the stationary phase..


Viscosity

• It is a measure of a liquid’s resistance to flow.

• It is the energy required to move an object through a fluid.

• Affected by temperature– As temperature increases ,the liquid’s viscosity

decreases.– More kinetic energy

• Size of liquid’s molecule– A more complex, larger molecule is more viscous.

• Greater entanglement


Solubility and Intermolecular Forces

• Solubility depends, in part, on the attractive forces of the solute and solvent molecules.

– Similar polarities• “Like dissolves like.”

– Miscible liquids will always dissolve in each other.

• Polar substances dissolve in polar solvents.– They have hydrophilic groups (water loving).

• OH, CHO, C=O, COOH, NH2, Cl

• Nonpolar molecules dissolve in nonpolar solvents.– Their compounds have hydrophobic groups (fear of water).

• Hydrocarbons (e.g., C–H, C–C, C=C)

• Molecules with hydrophilic and hydrophobic parts– Solubility in polar substance becomes a competition between

• the attraction of the polar groups for other polar groups• the attraction of the nonpolar groups for their own kind


Immiscible Liquids

• Solubility depends on the attractive forces between the solute and solvent molecules.– Similar polarities are soluble

with each other.• “Like dissolves like.”

• Immiscible liquids have NOTHING NOTHING in COMMON in COMMON with each other.– Polar compounds mixed with

nonpolar compounds– Do not have anything to

interact with or attract to each other.

– They do not “mix” or are not soluble with each other.


Overview of Liquid and Solid Properties

Water (l) Water (s)ice


Kinetic–Molecular Theory of Liquids

• When the attractive forces are strong enough so the kinetic energy can only partially overcome them, the material will be a liquid.

• In a liquid, the particles are packed together with only very limited translational or rotational freedom.– Liquid unlike gases are almost

incompressible.


Liquids

•• Liquid particles have close contact Liquid particles have close contact with one another, which limits their with one another, which limits their movement.movement.

• They have appreciable They have appreciable intermolecular forces.intermolecular forces.– Polar substances Polar substances

• Hydrogen bondingHydrogen bonding– Nonpolar substancesNonpolar substances

• Dispersion forcesDispersion forces

•• Liquids volume is the capacity of the Liquids volume is the capacity of the container.container.


Explaining the Properties of Liquids

• Liquids have higher densities than gases and are incompressible because the particles are in contact.

• They have an indefinite shape because the limited translational freedom of the particles allows them to move around enough to get to the container walls.

• It also allows them to flow.

• However, they have a definite volume because the limit on their freedom keeps them from escaping the rest of the particles.


Trends in the Strength of Intermolecular Attractions

• The stronger the attractions between the atoms or molecules, the more energy more energy it will take to separate them.

• Boiling a liquid requires addition of enough energy to overcome all the INTERMOLECULARINTERMOLECULAR attractions between the particles.– The bonds that make up the

molecule--the INTRAMOCULARINTRAMOCULAR forces--DO NOT BREAK.DO NOT BREAK.

• The higher boiling point of the liquid, the stronger the intermolecular attractive forces.


Vaporization/Evaporation

• It takes energy for molecules to go from one phase to another.

• Vaporization requires input of energy to overcome the attractions between molecules.

– Only a small fraction of the molecules in a liquid have enough energy to escape.

– By increasing the temperature, the fraction of the molecules with “escape energy” or evaporation energy increases.

• The higher the temperature, the faster the rate of evaporation by the high-energy molecules at the surface of the liquid.

• The greater the liquid’s surface area:– the faster the rate of evaporation– the quicker a liquid because a gas (vapor)


Condensation

Molecules in the vapor state can lose energy through molecular collisions with other molecules.

• The result:– Molecules are captured back into the liquid when they collide with

the liquid’s surface.– Others may stick together to form droplets of liquid.

This process is condensation.condensation.

• Open vs. closed container– In an open container:

• The vapor molecules generally spread out faster than they can condense.

– In a closed container:• The vapor molecules have limited area to spread out, resulting

in the vapor molecules condensing.


Condensation vs. Vaporization

Condensation

• Condensation is an exothermic exothermic processprocess.

Vaporization

• Vaporization is an endothermicendothermic process.


Intermolecular Forces andEvaporation and Condensation

Both phase change processes are affected by the strength of the liquid’sintermolecular forces.

• Weak interactions:– Vaporization:

• The weaker the attractive forces between molecules, the less energy to vaporize.

• The weaker the attractive forces, the faster the rate of evaporation.– Condensation:

• The weaker the attractive forces, the more energy is needed for the vapor molecules to condense.

• Liquids that evaporate easily are said to be volatilevolatile.• E.g., gasoline, fingernail polish remover

• Liquids that do not evaporate easily are called nonvolatilenonvolatile.• E.g., motor oil, salt water


• At a given temperature, the vapor At a given temperature, the vapor

pressure of a liquid depends on its pressure of a liquid depends on its

intermolecular forces.intermolecular forces.

– The GREATER the strength of the The GREATER the strength of the

intermolecular forces, the more intermolecular forces, the more

energy is needed to go from liquid energy is needed to go from liquid

to gas phase, and the greater the to gas phase, and the greater the

boiling point temperature. boiling point temperature.

Liquids


Equilibrium Vapor Pressure• A liquid’s vapor pressure is a function of

temperature.

• Vapor pressure increases with temperature rise.

• Vapor pressure curves (the liquid-gas line on a phase diagram)

• Temperature and pressure are at equilibrium between the liquid and gas phase.

• When a liquid’s vapor pressure = external/atmospheric pressure liquid boils.

This means boiling points of liquids This means boiling points of liquids change with altitudechange with altitude

• When pressure is lowered, the vapor pressure can equal the external pressure at a lower temperature.


Vapor-Liquid Dynamic Equilibrium

• If the volume of a chamber is increased, then a decrease in the vapor pressure inside the chamber occurs.

– Result:

• Fewer vapor molecules present for the given volume, which causes the rate of condensation to slow.

– Over time:

• The rate of vaporization will be faster than the rate of condensation, so the amount of vapor increases.

• Enough vapor accumulates so that the rate of condensation increases to the point where it is once again as fast as evaporation.

– Equilibrium is reestablished.

• At this point, the vapor pressure will be the same as it was before.


• When molecules of liquid are in the vapor state, they exert a When molecules of liquid are in the vapor state, they exert a VAPOR VAPOR PRESSURE.PRESSURE.

• EQUILIBRIUM VAPOR PRESSURE EQUILIBRIUM VAPOR PRESSURE is the pressure exerted by a vapor is the pressure exerted by a vapor over a liquid in a closed container when the over a liquid in a closed container when the rate of evaporation = the rate of evaporation = the rate of condensation.rate of condensation.


Liquids: Vapor Pressure and Hvap

HHvap vap

•It is the:– amount of energy (q) required to vaporize one mole of

the liquid– heat required (at constant heat required (at constant PP) to vaporize a liquid to a ) to vaporize a liquid to a

gasgas• Somewhat temperature dependent

•It is always endothermic; therefore, Hvap is positive.

•Referred to as the:– heat of vaporization orheat of vaporization or– enthalpy of vaporizationenthalpy of vaporization

Hcondensation = − Hvaporization


HHvap vap

LIQUID + heat LIQUID + heat VAPOR VAPOR

CompoundCompound ∆∆HHvapvap (kJ/mol) (kJ/mol) Intermolecular Intermolecular forceforce

HH22OO 40.7 (100 40.7 (100 ooC)C) H bondsH bonds

SOSO22 26.8 (-47 26.8 (-47 ooC)C) dipoledipole

XeXe 12.6 (-107 12.6 (-107 ooC)C) induced dipole induced dipole

Liquids: Vapor Pressure and Hvap


Heating Curve of a Liquid

• As a liquid is heated:– its temperature increases

linearly until it reaches the boiling point

– q = mass × Cs × T

• At the boiling point:– all the added heat goes into

boiling the liquid– the temperature stays constant

• Once all the liquid has been turned into gas, the temperature can again start to rise.


Heating Curve of Water


Determine heat (q) for H2O(s) H2O(g)

Segment 1:

Supercooled Ice to Ice

Heating 1.00 mole of ice from −25.0 °C up to the melting point, 0.0 °C

• q = mass × Cs × T1.00 mole of ice = 18.0 g

Cs = 2.09 J/g°CT = (0.0 oC – (25.0 oC) = 25.0 oC

q = (18.0 g) x (2.09 J/g°C) x 25.0 oC

q = 941 J

q = 0.941 kJ

Segment 2:

Ice to Water

Melting 1.00 mole of ice (0.0 °C)

• q = n∙Hfusion

n = 1.00 mole of ice

Hfus = 6.02 kJ/mol

q = (1.0 mol) x (6.02 J/mol°C)

q = 6.02 kJ


Segment 3:

Water heating

Heating 1.00 mole of ice from 0.0 °C up to its boiling point, 100.0 °C


Cs = 4.184 J/g°CT = (100.0 oC – (0.0 oC) = 100.0 oC

q = (18.0 g) x (4.184 J/g°C) x 25.0 oC

q = 7.52 x 103J

q = 7.52 kJ

Segment 4:

Water vaporizing to steam

Boiling 1.00 mole of water (100.0 °C)

• q = n∙Hvap

n = 1.00 mole of ice

Hvap = 6.02 kJ/mol

q = (1.0 mol) x (40.7 J/mol°C)

q = 40.7 kJ



Segment 5:

Steam Heating

Heating 1.00 mole of steam from 100.0 °C up to its boiling point, 125.0 °C


Cs = 2.01 J/mol°CT = (125.0 oC – (1000.0 oC) = 25.0 oC

q = (18.0 g) x (2.01 J/g°C) x 25.0 oC

q = 904 J

q = 0.904 kJ



Equilibrium Vapor Pressure and the Clausius-Clapeyron Equation

Clausius-Clapeyron equation used to find ∆Clausius-Clapeyron equation used to find ∆HH˚̊vapvap

– Can be used:

• with just two measurements of vapor

pressure and temperature

• to predict the vapor pressure if the heat of

vaporization and the normal boiling point

are known

• The Clausius–Clapeyron equationThe Clausius–Clapeyron equation

– The logarithm of the vapor pressure The logarithm of the vapor pressure PP is is

proportional to ∆proportional to ∆HHvapvap and to 1/ and to 1/TT..

• ln ln PP = (- ∆ = (- ∆HH˚̊vap vap //RTRT) + C) + C

• ln (ln (PP2/2/PP1) = {-(1) = {-(HHvapvap//RR) [(1/) [(1/TT2) –(1/2) –(1/TT1)]}1)]}


Problem: Determine the Hvap for propyl amine from the following data.

P1 = 40.0 mm Hg T1 = 257 K R = 8.314 x 10-3 kJ/K mol

P2 = 100.0 mm Hg T2 = 274 K

ln (ln (PP2/2/PP1) = {-(1) = {-(HHvapvap//RR) [(1/) [(1/TT2) – (1/2) – (1/TT1)]}1)]}

ln (ln (100.0 mm Hg100.0 mm Hg) = {- () = {- (HHvapvap/8.314 x 10/8.314 x 10-3-3 kJ) [(1/274) – (1/257)]} kJ) [(1/274) – (1/257)]}

(40.0 mm Hg)(40.0 mm Hg)

9.16 x 109.16 x 10-1-1 = {- ( = {- (HHvapvap/8.314 x 10/8.314 x 10-3-3 kJ) [(3.65 x 10 kJ) [(3.65 x 10-3-3 ) – (3.89 10 ) – (3.89 10-3-3)]})]}

9.16 x 109.16 x 10-1-1 = {- ( = {- (HHvapvap/8.314 x 10/8.314 x 10-3-3 kJ) [(-2.42 x 10 kJ) [(-2.42 x 10-4-4)]})]}

7.62 x 107.62 x 10-3-3 kJ = {- ( kJ = {- (HHvapvap) [(-2.42 x 10) [(-2.42 x 10-4-4)]})]}

- 3.15 x 10- 3.15 x 1011 kJ = - ( kJ = - (HHvapvap) )

3.15 x 103.15 x 1011 kJ = kJ = HHvapvap


Kinetic–Molecular Theory of Solids

• When the attractive forces are strong enough so the kinetic energy cannot overcome them at all, the material will be a solid.

• In a solid, the particles are packed together without any translational or rotational motion.– The only freedom they have is vibrational

motion.


Solid Classifications

© 2013 Pearson Education, Inc.54

Solids: An Overview

• In solids:– The particles are packed tightly together.

• Results in solids being incompressible and having high density

– The particles have no translational or rotational freedom of motion.• Results in solids retaining their shape and volume when placed in a new

container

• It also prevents the particles from flowing.

• Types:

– Orderly geometric patternOrderly geometric pattern• Crystalline solidsCrystalline solids• Example: salt and diamonds

– Nonregular geometric pattern Nonregular geometric pattern over a long range• Amorphous solidsAmorphous solids• Example: plastic and glass


Classifying Crystalline Solids

Crystalline solid classifications can be based upon:

1) particle makeup or

2) the types of attractive forces holding the particles together

• Molecular solids:– Solids whose composite particles are molecules

– Example: Ice, which is composed of H2O molecules

• Ionic solids: – Solids whose composite particles are ions– Example: Table salt (NaCl), which is composed of Na+ and Cl-

• Atomic solids: – Solids whose composite particles are atoms

• Nonbonding atomic solids are held together by dispersion forces.• Metallic atomic solids are held together by metallic bonds.• Network covalent atomic solids are held together by covalent bonds.

– Example: Graphite is composed of carbon atoms.


Molecular Solids

• The lattice sites are occupied by molecules.

– CO2, H2O, C12H22O11

• The molecules are held together by intermolecular attractive forces.– dispersion forces, dipole–dipole

attractions, and H bonds

• Because the attractive forces are weak, they tend to have low melting points.– generally <300 °C


Water: An Extraordinary Substance

Water:• Is a liquid at room temperature

– Most molecular substances with similar molar masses are gases at room temperature.

• e.g., NH3, CH4

– Why?• due to H bonding between molecules

• Is an excellent solvent– Dissolves many ionic and polar molecular

substances because of its large dipole moment– Even many small nonpolar molecules have

some solubility in water.• e.g., O2, CO2

• Has a very high specific heat for a molecular substance

• Expands when it freezes so its solid (ice) is less dense than liquid (water)


Ionic Solids

Ionic solids • Have lattice sites occupied by ions, which are held together by

attractions between oppositely charged ions– Nondirectional

• Every cation attracts all anions around it and every anion attracts all cations around it.

• The coordination number represents the number of close cation–anion interactions in the crystal.

– The higher the coordination number, the more stable the solid. • The lower the potential energy of the solid

– The coordination number depends on the relative sizes of the cations and anions that maintain charge balance.

• Anions are larger than cations.• The number of anions that can surround the cation is limited by

the size of the cation.– The closer in size the ions are, the higher the coordination number

is.


Problem: Match the compound with the type of solid.

Compound Solid Classification

a. KCl

b. C(s, graphite) S8

c. Kr

d. SrCl2

e. SiO2(s, quartz)

f. H2O

g. S8

h. Cu

i. Na


Problem: Match the compound with the type of solid.

Compound Solid Classification

a. KCl ionic

b. C(s, graphite) network covalent (atomic)

c. Kr atomic

d. SrCl2 ionic

e. SiO2(s, quartz) network covalent (molecular)

f. H2O molecular

g. S8 molecular (atomic)

h. Cu metallic

i. Na metallic

j. PBr3 molecular


Crystal Lattice

• When some liquids are allowed to cool slowly, their particles will arrange themselves so as to give the maximum attractive forces, which minimize the solid overall energy.

• The result will generally be a crystalline solid.

• The arrangement of the particles in a crystalline solid is called the crystal lattice.crystal lattice.

• The smallest unit that shows the pattern of arrangement for all the particles is called the unit unit cellcell.


Unit Cells

• Unit cells are three-dimensional. – Usually containing two or three layers of particles repeated over and over to

give the macroscopic crystal structure of the solid• Starting anywhere within the crystal results in the same unit cell.

• The number of other particles each particle is in contact with is called its coordination number.

– The higher the coordination number, the more interaction and the stronger the attractive forces holding the crystal together.

• The packing efficiency is the percentage of volume in the unit cell occupied by particles.

– The higher the coordination number, the more efficiently the particles are packing together.

• Lattice point:– Each particle in the unit cell is called a lattice point.

• Lattice planes:– Planes connecting equivalent points in unit cells throughout the lattice


Unit Cell Types


Simple Cubic Unit Cells

• All 90° angles between corners of the unit cell• Lengths of all the edges are equal• If the unit cell is made of spherical particles:

– ⅛ of each corner particle is within the cube– ½ of each particle on a face is within the cube– ¼ of each particle on an edge is within the cube


Body-Centered Cubic


Face-Centered Cubic


Layer Arrangements in Crystalline Structures: Closed Packed

• Closest-Packed Structure First Layer

• The second-layer atoms can sit directly over the atoms in the first layer (called an AA pattern), or

• The second layer can sit over the holes in the first layer (called an AB pattern).


Closest-Packed:Third Layer, with Offsetting 2nd Layer

• The third-layer atoms can align directly over the atoms in the first (called an ABA pattern).

Face-Centered Cubic


• Or the third layer can sit over the uncovered holes in the first (called an ABC pattern).

Hexagonal Closest-Packed

Closest-Packed:Third Layer, with Offsetting 2nd Layer


Ionic Crystals

CsClcoordination number = 8

Cs+ = 167 pmCl─ = 181 pm

NaClcoordination number = 6

Na+ = 97 pmCl─ = 181 pm


Cesium Chloride Structures

Statistic:• Coordination number = 8

• ⅛ of each Cl─ (184 pm) inside the unit cell

• Whole Cs+ (167 pm) inside the unit cell– Cubic hole Cubic hole = hole in simple

cubic arrangement of Cl─ ions

• Cs:Cl = 1: (8 × ⅛)

• The formula is CsCl.


NaCl: Rock Salt Structures

Statistics:• Coordination number = 6

• Cl─ ions (181 pm) in a face-centered cubic arrangement– ⅛ of each corner Cl─ inside the unit cell– ½ of each face Cl─ inside the unit cell

• Each Na+ (97 pm) in holes between Cl─

– Octahedral holesOctahedral holes– 1 in center of unit cell– ¼ of each edge Na+ inside the unit cell

• Na:Cl = (¼ × 12) + 1: (⅛ × 8) + (½ × 6) = 4:4 = 1:1

• The formula is NaCl.


Nonbonding Atomic Solids

Facts about atomic solids:

• Noble gases in solid form are classified as atomic solids.

• Held together by weak dispersion forces– Very low melting point

• Tend to arrange atoms in closest-packed structure.– Either hexagonal closest-packed or cubic closest-

packed• Maximizes attractive forces and minimizes energy


Metallic Atomic SolidsMetallic Atomic Solids

Facts about metallic atomic solids:

• Held together by metallic bonds

– Strength varies with sizes and charges of cations

• Coulombic attractions.

• Melting point varies

• Mostly closest-packed arrangements of the lattice points

– Cations

• Metal atoms release their valence electrons.

– Metal cation “islands” fixed in a “sea” of mobile electrons


Metallic Atomic Solids


Network Covalent Solids

Network covalent solids:• Atoms are attached to the nearest neighbors by covalent

bonds.– Due to the directionality of the covalent bonds, these solids do not tend

to form closest-packed arrangements in the crystal.

• The strength of the covalent bonds within the solids is why they tend to have very high melting points.– Generally >1000 °C

• The physical properties of network covalent solids are related to the dimensionality of the network.


Diamond: A Three-Dimensional Network Structure

• The carbon atoms in a diamond each have four covalent bonds to surrounding atoms.– sp3

– Tetrahedral geometry

• This arrangement effectively makes each crystal one giant molecule held together by covalent bonds.• A path of covalent bonds from any atom to every other atom

• Explains why diamond is one of the hardest naturally occurring materials found on earth


Graphite Structure:a Two-Dimensional Network

Graphite:• The carbon atoms in a sheet are covalently bonded together.

– Forming six-membered flat rings fused together• Similar to benzene• Bond length = 142 pm

– sp2 • Each C has three sigma bonds and one pi bond.

– Trigonal-planar geometry– Each sheet a giant molecule

• The sheets are then stacked and held together by dispersion forces.– Sheets are 341 pm apart.


Diamond vs. Graphite

Diamond

• Very high melting, ~3800 °C

– need to overcome some covalent bonds

• Very rigid

– due to the directionality of the covalent bonds

• Very hard

– due to the strong covalent bonds holding the atoms in position

– used as abrasives

• Electrical insulator

• Thermal conductor

– best known

• Chemically very nonreactive

Graphite

• Hexagonal crystals

• High melting, ~3800 °C

– need to overcome some covalent bonding

• Slippery feel

– Because there are only dispersion forces holding the sheets together, they can slide past each other.

• glide planes

– lubricants

• Electrical conductor

– parallel to sheets

• Thermal insulator

• Chemically very nonreactive


Diamond vs. Graphite


Silicates

Facts about silicates:

• ~90% of Earth’s crust

• Their solids tend to be extended arrays of SiO.

– Sometimes with Al substituted for Si— aluminosilicates

• Glass and quartz (sand) are types of silicate compounds.– Glass is the amorphous form.– Quartz

• SiO2 in pure form is clear, but impurities (metal atoms) add color.– Three-dimensional array of Si covalently bonded to four O– Tetrahedral– Properties:

• Melts at ~1600 °C• Very hard


Structure of Quartz and Glass

Quartz Glass


Band Theory: What Is it About?

• Metals and covalent network structures:– The solids result in every atom’s orbitals being

shared by the entire structure.

• This means there is a large number of molecular orbitals that have approximately the same energy.

• This is referred to as an energy bandenergy band.


Band Theory

• When two atomic orbitals combine, they produce both a bonding and an antibonding molecular orbital.

• When many atomic orbitals combine, they produce a band of bonding molecular orbitals and a band of antibonding molecular orbitals.

• The band of bonding molecular orbitals is called the valence valence band.band.

• The band of antibonding molecular orbitals is called the conduction bandconduction band.


Molecular Orbitals of Poly lithium


Band Gap

• At absolute zero, all the electrons will occupy the valence band.

– As the temperature rises, some of the electrons may acquire enough energy to jump to the conduction band.

• The difference in energy between the valence band and conduction band is called the band gapband gap.

– The larger the band gap, the fewer electrons there are with enough energy to make the jump.

• Types of band gaps and their conductivity:


Band Gap and Conductivity

• Conductor:Conductor:– The more electrons at any one time that a substance has in the conduction

band, the better conductor of electricity it is.• Band gap is ~0:

» Electrons will be almost as likely to be in the conduction band as the valence band, and the material will be a conductorconductor.

– Example: metals• The conductivity of a metal decreases with temperature.

• Semiconductor:Semiconductor:– Band gap is small:

• Significant number of the electrons will be in the conduction band at normal temperatures, and the material will be a semiconductor.semiconductor.

– Example: graphite• The conductivity of a semiconductor increases with temperature.

• Insulator:Insulator:– Band gap is large:

• Effectively no electrons will be in the conduction band at normal temperatures, and the material will be an insulatorinsulator. .


Doping Semiconductors

• DopingDoping is adding impurities to the semiconductor’s crystal to increase its conductivity.– The goal is to increase the number of electrons in the

conduction band.

• n-Type semiconductors n-Type semiconductors do not have enough electrons themselves to add to the conduction band, so they are doped by adding electron-rich impurities.

• p-Type semiconductorsp-Type semiconductors are doped with an electron-deficient impurity, resulting in electron “holes” in the valence band. – Electrons can jump between these holes in the valence

band, allowing conduction of electricity.


Diodes

• When a p-type semiconductor adjoins an n-type semiconductor, the result is a p-n junctiona p-n junction.

• Electricity can flow across the p-n junction p-n junction in only one direction.– This is called a diodediode.

• This allows the accumulation of electrical energy.– Called an amplifier amplifier


Thermodynamics of Melting

Melting requires input of energy to overcome the attractions between

molecules.

When a solid is heated:

– Its temperature rises and the molecules vibrate more vigorously.

• Once the melting point is reached:– The molecules have sufficient energy to overcome some of the attractions that hold them

in position and the solid melts (fuses).

Solids can melt when:

– There is a loss of the high-energy molecules from a solid, resulting in:

• a lowering of solid’s average kinetic energy

• a decrease in the substance’s temperature

• Melting is an endothermic process.

• Freezing, the reverse process of melting, is an exothermic process.


Heat of Fusion

Heat of fusion, Heat of fusion, HHfus fus or enthalpy of fusionor enthalpy of fusion

It is the amount of heat energy required to melt one mole of the solid.– Always endothermic; therefore, Hfus is positive (+).– Can be temperature dependent– Usually its value is less than its Hvap.

Hcrystallization = − Hfusion

Hsublimation = Hfusion + Hvaporization


Sublimation and Deposition

• Molecules or atoms in the solid have thermal energy that allows them to vibrate.

• Molecules at the surface with sufficient energy may break free from the surface and become a gas. This process is called sublimationsublimation.

• The reverse of sublimation is referred to as deposition.deposition.– It is the capturing of vapor molecules into a solid.

• Molecular solids have a vapor pressure.– Solid and vapor phases can exist in dynamic

equilibrium in a closed container at temperatures below the melting point.


The How and Why of Phase Changes

• Attractive forces between the molecules are fixed.– To change the material’s state (phase) requires changing the amount of kinetic energy

the particles have, or limiting their freedom.

• Solids melt when heated because the particles gain enough kinetic energy to partially overcome the attractive forces.

• Liquids boil when heated because the particles gain enough kinetic energy to completely overcome the attractive forces.– The stronger the attractive forces, the higher you will need to raise

the temperature.

• Gases can be condensed by decreasing their temperature and/or increasing the pressure.– Pressure can be increased by decreasing the gas volume.– Reducing the volume reduces the amount of translational freedom

the particles have.


• A phase diagram is a representation of a substance’s phase according to pressure and temperature.– It describes the different states and state changes that occur at

various temperature/pressure conditions.

• Regions represent states (phases).

• Lines and points represent state changes.• The liquid/gas line is the vapor pressure curve.

– At any point on a line, both phases are in equilibrium at the temperature and pressure.

– Points:• The critical point is the farthest point on the vapor pressure curve.• The triple point is the temperature/pressure condition where all

three states exist simultaneously.

• For most substances, freezing point increases as pressure increases.

What Is a Phase Diagram?


Phase Diagram of Water


The Critical Point

• Critical temperature:– The temperature required to produce a supercritical

fluid

• Critical pressure:– The pressure at the critical temperature

• At the critical temperature or higher temperatures, the gas cannot be condensed to a liquid, no matter how high the pressure gets.


Supercritical Fluid: A 4th Phase of Matter

• As a liquid is heated in a sealed container, more vapor collects, causing: – The pressure inside the container to rise– The density of the vapor to increase– The density of the liquid to decrease

• At some temperature, the meniscus between the liquid and vapor disappears and the states commingle to form a supercritical fluidsupercritical fluid.

• Supercritical fluids have properties of both gas and liquid states.


a. 20.0 °C, 72.9 atm

b. −56.7 °C, 5.1 atm

c. 10.0 °C, 1.0 atm

d. −78.5 °C, 1.0 atm

e. 50.0 °C, 80.0 atm

Problem: Determine the phases of matter for the following points on the phase diagram of CO2 .

following conditions?


a. 20.0 °C, 72.9 atm liquid

b. −56.7 °C, 5.1 atm solid, liquid, gas (triple point)

c. 10.0 °C, 1.0 atm gas

d. −78.5 °C, 1.0 atm solid, gas

e. 50.0 °C, 80.0 atm supercritical fluid

Problem: Determine the phases of matter for the following points on the phase diagram of CO2 .

following conditions?

Documents

Catherine MacGowan Armstrong Atlantic State University Lecture Presentation Chapter 11 Liquids, Solids, and Intermolecular Forces © 2013 Pearson Education,