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Chapter 12. Intermolecular Forces:. Liquids, Solids, and Phase Changes. If you are doing this lecture “online” then print the lecture notes available as a word document, go through this ppt lecture, and do all the example and practice assignments for discussion time.

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Chapter 12

Chapter 12

Intermolecular Forces:

Liquids, Solids, and Phase Changes

If you are doing this lecture “online” then print the lecture notes available as a word document, go through this ppt lecture, and do all the example and practice assignments for discussion time.


Why do gases differ from liquids solids

Why do gases differ from liquids & solids?

Gases are tiny particles far apart with no attraction for each other (ideally) and they are moving rapidly in random directions

Gases obey a set of laws: Ideal Gas Laws

But liquids & solids don't have a set of "laws" because....

Liquids: - condensed from gases

- not compressible therefore not much space

between molecules

-moving randomly but more slowly, are attracted

to each other!

Solids: - ordered fixed in place particles

- close together

- strong forces of attraction!


Interparticle forces

Interparticle Forces

Interparticle forces of attraction between "particles" that affect physical state & physical behavior

Note: particles can be ions, atoms or molecules

True chemical bonding forces are intra-particle forces: within a chemical substance

They are atom to atom or ion to ion, not molecule to molecule: ionic bonding, metallic bonding, covalent bonding, and (new to us) network covalent bonding

Interparticle forces are between particles, not within them

Special group of interparticle forces between molecules and atoms:

- called intermolecular forces

- affect behavior of covalent compounds

Next slide is table from your packet of handouts


Chapter 12

Properties

High MP, brittle solid

Low MP, soft when solid

Range of MPs, malleable

ADD:

Network CovalentAtom-Atom like cov bond500+Diamond

Very high MP, very hard solid


Network covalent bonding

Network Covalent Bonding

Make special note of the network covalent solids that I added to Table 12.2:

- covalent bonds are extended throughout the crystal solid

- diamond and SiC, also SiO2

Diamond is fcc with 4 C's in "holes" in unit cell, which we see later in chapter

Graphite has unique structure (see diagram)


Chapter 12

**

**

**

**“True” Intermolecular Forces in pure substances


Ion dipole attraction

Ion-Dipole Attraction:

Interparticle forces involved in dissolving ionic solids in water or other polar solvents:

Ion-dipole attraction has to overcome ion-ion attraction in a solid’s crystal lattice

This is why some compounds are soluble and some are not

Called Hydration of an ion: typically endothermic - takes energy to pull ions apart

Some energy is gained back as Heat of Hydration - polar water molecules orient themselves and surround the individual ions

More detail coming soon in Chapter 13


The major intermolecular im forces

The Major InterMolecular (IM) Forces

The three major IM Forces are very weak to moderate forces of attraction between molecules (or atoms)

- London (dispersion) forces (LDF)

- Dipole-dipole attractions

- Special case of enhanced dipole attraction called Hydrogen-bonding

- Dipole-induced dipole attraction, between different types of molecules or atoms


London dispersion forces ldf

London Dispersion Forces (LDF)

"Instantaneous dipole" causes neighboring electron clouds to also move to one side, inducing a dipole in them

Leads to a small force of attraction between slight positive and slight negative ends of two different particles – these attractions are called London dispersion forces

Strength of LDF depends on: (1) size of electron cloud in atom and/or (2) number of atoms in a molecule

Polarizability increases down a group

Polarizability decreases left to right across a period

Look at the elemental halogens: fluorine and chlorine are gases; bromine is liquid and iodine is solid because of size of e- cloud around each molecule


Chapter 12

Dispersion forces among nonpolar molecules.

Figure 12.14

instantaneous dipoles

separated Cl2 molecules

An instantaneous dipole in one Cl2 molecule will induce a dipole in a nearby Cl2 molecule. The partial charges attract the molecules together. This process takes place with the particles throughout the container.


Chapter 12

Figure 12.15

Molar mass and boiling point.

The strength of LDF increases with the number of electrons, which correlates with molar mass. Therefore, LDF increases down a group in the Periodic Table, which can be verified by the increase in boiling points.


Chapter 12

Molecular shape and boiling point.

Figure 12.16

fewer points for dispersion forces to act

more points for dispersion forces to act

Spherical molecular shapes make less contact with each other than do cyllindrical shapes, so they have a lower boiling point.

(It’s just Pentane.)

2,2-dimethylpropane

I have corrected the organic names – compare to your textbook, which is using old naming rules.


Chapter 12

liquid

Polar molecules and dipole-dipole forces.

Figure 12.11

solid

In a solid or a liquid, polar molecules are close enough for the attraction to hold. Orientation is more orderly in a solid because the average KE of the particles is lower.


Dipole dipole attraction

Dipole-dipole Attraction

Dipoles involve polar molecules which are attracted to each other because of the slight positive and slight negative "poles" to the molecules

Compare molecules of F2, HF, HCl, HBr, HI

Boiling points: F2<HCl<HBr<HI<HF


Data for im forces

Data for IM Forces

F2HFHClHBrHI

Tot # e-s1810183654

MM382036.581128

DEN01.81.00.80.5

Dip Mom01.41.10.80.4

% Disp100Low 81.494.599.5

% Dipole0High18.65.50.5

BP, K85 291188206238

DHvap6.86High161820


Enhanced dipole dipole or hydrogen bonding

Enhanced Dipole-Dipole or Hydrogen-Bonding?

Dipole forces are decreasing down the hydrohalogen group because DEN is decreasing

WHY is HF so very different in boiling point?

HF represents a special case of dipole-dipole attraction called Hydrogen-bonding

Occurs when H is bonded to a highly EN atom that is also very small: H to F, O or N

Size of EN atom is important also, allows H to get very close, as seen with radius in pm below:

N-70 O-73 F-72 Cl-100 (no H-bond)


Chapter 12

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.

(b)

(a)

(c)

PLAN:

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

(b)

SAMPLE PROBLEM 12.3

Drawing Hydrogen Bonds Between Molecules

of a Substance

SOLUTION:

(a) C2H6 has no H bonding sites.

(c)

The N-H is also attracted to the N-H.


Chapter 12

Hydrogen bonding and boiling point.

Figure 12.13

Boiling points of the binary covalent hydrides of Groups 14 – 17 plotted agains Period digit. Shows that H2O, HF and NH3 do not follow the downward trend, as shown by the dashed line for group 16.


Dipole induced dipole between two different compounds particles

Dipole-Induced Dipole: between two different compounds’ particles

Dipole-induced dipole forces account for limited solubility of oxygen in water

Ability to do this = function of polarizability of molecule

Compare H2 to I2: bigger molecule polarizes, soluble in water, which is demonstrated by the much greater solubility of I2 in water


Chapter 12

PROBLEM:

For each pair of substances, identify the dominant interparticle forces affecting the physical properties of each substance, and then select the substance with the higher boiling point.

(d) Hexane (CH3CH2CH2CH2CH2CH3)

or 2,2-dimethylbutane

SAMPLE PROBLEM 12.4

Predicting the Type and Relative Strength of

Intermolecular Forces

(a) MgCl2 or PCl3

(b) CH3NH2 or CH3F

(c) CH3OH or CH3CH2OH

  • 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.


Chapter 12

SAMPLE PROBLEM 12.4

Predicting the Type and Relative Strength of

Intermolecular Forces

continued

SOLUTION:

(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 CH3F is just dipole-dipole. 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 London dispersion force interaction. Therefore CH3CH2OH has the higher boiling point.

(d) Hexane and 2,2-dimethylbutane are both nonpolar with only London dispersion forces to hold the molecules together. Hexane has the larger surface area, thereby the greater dispersion forces and the higher boiling point.


Practice with intermolecular forces explain the forces behind this data

Practice with Intermolecular Forces: explain the forces behind this data:

1. Butane (CH3CH2CH2CH3) melts at -138oC and boils at 0.5oC, while acetone (CH3C=OCH3) melts at -95oc and boils at +56oC, yet both weigh 58 g/mol. Draw Lewis structures and explain the differences in MPs and BPs.

2. Guess BP order for CCl4, N2, Cl2, ClNO (chlorine-nitrogen-oxygen).


Chapter 12

H bonded to

N, O, or F

Figure 12.17 modified

Summary diagram for analyzing the interparticle forces in a sample.

Metal atoms only

INTERACTING PARTICLES

(atoms, molecules, ions)

METALLIC BONDING

ions present

ions not present

ions only

IONIC BONDING

(Section 9.2)

nonpolar

molecules or atoms

Only, LONDON

DISPERSION

FORCES

polar molecules only

DIPOLE-DIPOLE

FORCES

ion + polar molecule

ION-DIPOLE FORCES

polar + nonpolar

molecules

DIPOLE-

INDUCED DIPOLE

FORCES

HYDROGEN

BONDING

LONDON DISPERSION FORCES ALSO PRESENT IN ALL OF ABOVE.

NETWORK COVALENT BONDING POSSIBLE FOR VERY FEW ATOMS.


Practice

Practice

See handouts

Also chapter problems: 2, 29, 31, 33, 37, 39, 43, 45

4th ed. #119: What forces are overcome when the following events occur: (a) NaCl dissolves in water, (b) krypton boils, (c) water boils, (d) CO2 sublimes?


Chapter 12

QUESTIONS TO ASK IN PREDICTING THE KINDS OF INTERPARTICLE FORCES THAT WILL BE PRESENT IN A SOLID OR A LIQUID

 Start at the top with the first question, “Is it metallic?”. When you can answer yes, you are done. If the answer is no, keep going down the list. The “it” refers to whatever substance you are working with.

QuestionIf yes, this force is present* MP & BPExamples (MP, BP)

 Is it metallic? (ONLY metal present) Metallic bonding HighIron (1555, 3000)

Is it ionic? (cation & anion present) Ionic bonding HighNaCl(804, not defined)

Is it network covalent compound?Network covalent bonding HighDiamond (3550, not def), SiC, SiO2

Is the substance molecular? (covalent bonds present)

In the molecule, is H attached by a covalent Hydrogen bonding MediumWater (0.100)

bond to F, O or N?

Does the molecule have a dipole moment? Dipole-dipole attraction LowHCl (-114, -85)

Is it a molecule with no dipole moment? Only London Forces Very lowHydrogen (-257, -253)

Iodine (114, 183)

Does the substance consist of atoms with no Only London ForcesExtremely low Neon (-249, -246)

covalent bonds between them?

*Remember – London forces are present in all liquids and solids.

Practice with these, supposing all to be in liquid or solid phase: methane, ethanol, sucrose (look up structure), NaOH, SiC, F2O, Cl2O, octane, radon, uranium, hydrobromic acid.


Properties of liquids interparticle forces

PROPERTIES OF LIQUIDS & INTERPARTICLE FORCES:

Why would a metal object with higher density than water float on water?

Why can we fill a glass of water above its rim?

Surface tension is related to strength of attractive forces in liquid: the stronger the attractive forces the greater the surface tension

Surface tension is the energy required to increase surface area by a unit amount; units are J/m2


Chapter 12

The molecular basis of surface tension.

Figure 12.18

hydrogen bonding

occurs across the surface

and below the surface

the net vector

for attractive

forces is downward

Molecules in the interior of a liquid experience IM forces in all directions. Molecules at the surface experience a net attraction downward, causing the liquid to minimize the number of molecules at the surface, ergo surface tension.

hydrogen bonding

occurs in three

dimensions


Chapter 12

Table 12.3

Surface Tension and Forces Between Particles

Surface Tension

(J/m2) at 200C

Substance

Formula

Major Force(s)

diethyl ether

CH3CH2OCH2CH3

dipole-dipole; dispersion

1.7x10-2

ethanol

CH3CH2OH

2.3x10-2

H bonding

1-butanol

CH3CH2CH2CH2OH

2.5x10-2

H bonding; dispersion

water

H2O

7.3x10-2

H bonding

mercury

Hg

48x10-2

metallic bonding


Chapter 12

stronger cohesive forces

adhesive forces

Shape of water or mercury meniscus in glass.

Figure 12.19

capillarity

H2O

Hg

See note in box below or look in textbook.


Properties of liquids

Properties of Liquids

Capillary action: rising of a liquid through a narrow space against the force of gravity

Viscosity: resistance to flow, units in Newton-seconds/m2


Chapter 12

Table 12.4 Viscosity of Water at Several Temperatures

viscosity - resistance to flow

Viscosity (N*s/m2)*

Temperature(0C)

20

1.00x10-3

40

0.65x10-3

0.47x10-3

60

80

0.35x10-3

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


Why water is special

Why water is special:

Water molecules are 80% H-bonded at normal conditions

Molecules are so close together that you cannot tell which H's belong to which O in each molecule

This is important to life on earth (& possibly elsewhere)

Ice floats on liquid water because the solid (ice) is less dense than the liquid (good for fishies)

Ice forms a crystal structure in tetrahedral arrangement

Hydrogen-bonding also accounts for other physical properties:

Lower weight alcohols are very soluble in water because of the -OH functional group

Great solvent properties because water is so polar

Very high specific heat as noted back in Chapter 6

High surface tension


Chapter 12

The H-bonding ability of the water molecule.

Figure 12.20

hydrogen bond donor

hydrogen bond acceptor

Because it has two O-H bonds and two lone pairs, one water molecule can engage in as many as four hydrogen-bonding attractions to surrounding water molecules, which are arranged tetrahedrally.


Chapter 12

The hexagonal structure of ice.

Figure 12.21

A. The geometric arrangement of the hydrogen-bonding in water leads to open, hexagonally shaped crystal structure of ice. Thus, when water freezes, the volume increases.

B. The delicate six-pointed beauty of snowflakes reflects the hexagonal crystal structure of ice.


Solids crystal structures

SOLIDS & CRYSTAL STRUCTURES

SOLIDS: fixed particles that cannot move with velocity, but do vibrate and rotate in position, so they do have KE

Generally have long-range order - crystals have well-defined regular shapes, or if short-range order they are amorphous - no regular shape, like asphalt, wax, glass

Crystal structure includes the four types of solids

ionic (all cation-anion units)

metallic (Cu, Zn, U, etc.)

molecular/atomic (ice, I2, etc.)

network covalent (diamond, SiC, SiO2)


General properties of the four types of crystalline solids

General Properties of the Four Types of Crystalline Solids

1. Ionic (KNO3, MgO): high MP/BP; some water-solb, brittle, conduct only when molten or aqueous

2. Molecular (C10H8, I2): low MP/BP; more solb in nonpolar; nonconductors

3. Network Covalent (Cdiamond, SiC, SiO2): very high MP/BP; insolb, brittle, non- or semi-conductor

4. Metallic (Cu, Fe, U): wide range of MPs; insolb, malleable, ductile, elec conductor


Chapter 12

The striking beauty of crystalline solids.

Figure 12.22

celestite

pyrite

amethyst

halite

Figure 12.22 in current 2nd edition has wulfanite, barite, calcite, quartz as amethyst, and beryl (emerald)


Crystal structures

Crystal Structures

Crystals have a crystal lattice arrangement of which smallest pieces are unit cell in 3-D, containing > one formula unit

Seven basic types: cubic, tetragonal, orthorhombic, monoclinic, hexagonal, rhombohedral, triclinic

See packet of handouts and Dry Lab VI(?) in Lab Manual for “Crystal Structures and Characteristics”

Simplest are the cubic, of which there are three types

Simple cubic (sc): metals and ionic cmpds

Body-centered cubic (bcc): metals

Face-centered cubic (fcc): metals and ionic cmpds


Chapter 12

lattice point

unit cell

unit cell

portion of a 2-D lattice

portion of a 3-D lattice

The crystal lattice and the unit cell.

Figure 12.23

A. The lattice is an array of points that defines the positions of the particles in a crystal structure. It is shown here as points connected by lines. One unit cell is highlighted.

A checkerboard is a two-dimensional analogy for a lattice.


Chapter 12

1/8 atom at 8 corners

The three cubic unit cells.

Figure 12.24 (1 of 3)

Simple Cubic

Atoms touch along edge of cube

Atoms/unit cell = 1/8 * 8 = 1

Cell length = 2r

coordination number = 6

See notes box below slide.


Chapter 12

1/8 atom at 8 corners

1 atom at center

coordination number = 8

The three cubic unit cells.

Figure 12.24 (2 of 3)

Body-centered Cubic

Atoms touch along main diagonal.

Atoms/unit cell = (1/8*8) + 1 = 2

Cell length = 4r/(3)1/2

See notes box below slide.


Chapter 12

1/8 atom at 8 corners

1/2 atom at 6 faces

coordination number = 12

The three cubic unit cells.

Figure 12.24 (3 of 3)

Face-centered Cubic

Atoms touch along face diagonal.

Atoms/unit cell = (1/8*8)+(1/2*6) = 4

Cell length = 4r/(2)1/2

See notes box below slide.


Cell length and cell volume

Cell length and cell volume

See figure 12.28 for derivation of cell length based on which cubic structure makes up the unit cell. For metals cell length is: sc length = 2r; bcc length = 4r/(3)1/2; fcc length = 4r/(2)1/2

Cell length determination is different for ionic compounds, which are simple cubic or face-centered cubic: sc length = 2(r+R)/(3)1/2; fcc length = 2(r+R)

Volume of any cube = (length)3


Crystal structures practice

Crystal Structures Practice

See handouts and practice problems

Also chapter problems: 61, 64, 67, 73, 75

4th ed. #98: Polonium is a rare radioactive metal that is the only element with a crystal structure based on the simple cubic unit cell. If its density is 9.142 g/cm3, calculate an atomic radius for a polonium atom.


Crystal structures practice1

Crystal Structures Practice

4th ed. #101: Tantalum, with D = 16.634 g/cm3, has a bcc structure with an edge length of 3.3058 Angstroms. Use its molar mass and this data to prove Avogadro’s number.


Chapter 12

exothermic

endothermic

Phase Changes

sublimination

vaporizing

melting

solid

liquid

gas

condensing

freezing

deposition


Vapor pressure

Vapor Pressure

Evaporation/vaporization: small fraction of molecules have high enough velocity to escape force of attraction at surface

RATE OF EVAPORATION: will increase with increasing T, since fraction of molecules with escape vel will increase

In a closed system, a dynamic equilibrium will be reached:

Rate of evaporation = rate of condensation

Vapor pressure: vapor molecules exert a partial pressure called vapor pressure


Chapter 12

Liquid-gas equilibrium.

Figure 12.4

A. In a closed flask at const T, with air removed, Pi = 0. As molecules escape surface to become vapor, P increases. B. At equilibrium, # of molecules escaping liquid = # of molecules condensing, P is constant. C. Plot of P vs. time shows P becomes constant.


Chapter 12

The effect of temperature on the distribution of molecular speed in a liquid.

Figure 12.5

With T1 lower than T2, most probable molecular speed, u1, is less than u2. Fraction of molecules with “escape velocity” is greater at the higher temperature. At higher T, equilibrium is reached with more molecules in the vapor phase, therefore at a higher P.


Vapor pressure practice

Vapor Pressure Practice

If 1.00 L of water is placed in 2.30x104 L closed room, will all the water evaporate? Given D = 0.997 g/mL, Vapor Pressure = 23.8 torr at 25.0oC.

Water will evap till room is at 23.8 torr partial pressure of water vapor.

See how many moles at that point:

n = PV/RT = 29.4 mol

How many moles in 1.0 L beaker? About 55 moles - won't all evaporate


Vapor pressure vs temperature

Vapor Pressure vs. Temperature

Boiling point: occurs when you see bubbles of gas forming in the liquid and coming to surface

Any pure liquid remains at constant T while boiling, since this is a change of state

Definition:

BP is the Temperature at which VP = barometric P

Why does water boil at 100oC in Fairfield and at 95oC in Denver?


Chapter 12

Figures 12.6 and 12.7

Figure 12.7

A linear plot of vapor pressure- temperature relationship.

Vapor pressure as a function of temperature and intermolecular forces.

The Clausius-Clapeyron equation comes from this graph: y = mx + b


You practice drawing and labelling a generic vp curve

You practice drawing and labelling a generic VP curve


Vapor pressure curves

Vapor pressure curves:

Initial "phase diagrams" incorporated into P/T diagrams that will include solid phase later

Why can NH3 be condensed from gas to liquid at Room T by compression, but N2 can't?


Relative humidity

Relative Humidity

NOT IN TEXT:

Relative humidity as reported by weather forecasters:

%water evap=actual partial P/equil vapor P * 100

If actual is 12.8 and VP for given T is 21.1, relative humidity is 61%


Vp and d h vap

VP and DHvap

DHvap is related to VP and T thru the Clausius-Clapeyron equation

ln P = (-DHvap/RT) + C (where R = 8.314 J/mol-K, T in K)

If plotted on a graph, the slope is:

=(ln p2 – ln p1)/(1/T2 – 1/T1)= -DHvap/R

Rearranges to Clausius-Clapeyron Equation (next slide)


Chapter 12

The Clausius-Clapeyron Equation

MEMORIZE!

Alternately, if you don’t want to use the negative sign:

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


Clausius clapeyron equation examples

Clausius-Clapeyron equation examples

Look at Sample Problem 12.2 in text.

My Example: hexane has DHvap = 30.1 kJ/mol and at 25.0oC, VP = 148 torr. What will VP be at 50.0oC?

ln (P2/148) = (-30.1x103J/8.314 J/mol-K)(1/323.15 – 1/298.15)

ln (P2/148) = 0.9394 (take antilog of both sides)

P2/148 = e0.9394 = 2.55

P2 = 379 torr


Practice1

Practice:

Chapter problems: 17 & 18

17: A liquid has DHovap of 35.5 kJ/mol and a BP of 122oC at 1.00 atm. What is its VP at 113oC?

18: What is the DHovap of a liquid that has a VP of 641 torr at 85.2oC and a BP of 95.6oC at 1.00 atm?


Chapter 12

test tube with ice

iodine solid

iodine vapor

iodine solid

Iodine subliming.

Figure 12.8 4th ed.,

not in principles


Simple phase diagrams

Simple Phase Diagrams

Are P vs. T diagrams showing three phases for pure elements or compounds, incorporates the VP curve

Critical Point is where liquid and gas cannot be distinguished from each other

Triple Point is where solid, liquid and gas phases meet and all three are present

Example: For water the Triple Point is 0.01oC and 4.58 torr

For CO2 Triple Point is at -56.7oC and 5.1 atm


Chapter 12

H2O

CO2

Phase diagrams for CO2 and H2O.

Figure 12.8

Each region depicts the T & P under which the phase is stable. Lines between regions show conditions at which two phases exist in equilibrium. The Critical Point shows conditions beyond which liquid and gas cannot be distinguished from each other. At the triple point, three phases exist is equilibrium. CO2 phase diagram is typical with forward sloping solid-liquid line; solid is more dense than liquid. H2O phase diagram is sloping backward; solid is less dense than liquid.


Phase diagrams

Phase Diagrams

You must draw and label phase diagrams based on data given to you and determine the physical state of a substance from its placement on a phase diagram.

Work on problems 20 & 22


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