Unit 03 Lecture Notes & Learning Standards

8/3/2019 Unit 03 Lecture Notes & Learning Standards

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Anderson - MCHS States of Matter AP Chem.Unit 03

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Unit Objectives

Chapter 10 Gases

AP10-1,2-01 Describe the general characteristics of gases and compared to other states o

matter, and list the ways in which gases are distinct

AP10-1,2-02 Define atmosphere, torr, and pascal, the most important units in which

pressure is expressed. Also describe how a barometer and a manometer

work.

AP10-3,4,5-01 Describe how a gas responds to changes in pressure, volume, temperature an

quantity of gas.

AP10-3,4,5-02 Use the ideal gas equation to solve for one variable (P,V, n, or T), given the

other three variables of information from which they can be determined

AP10-3,4,5-03 Use the gas laws, including the combined gas law to calculate how one

variable of a gas (P,V,n,or T) responds to changes in one or more of the other

variables.AP10-3,4,5-04 Calculate the molar mass of a gas, given the gas density under specified

conditions of temperature and pressure. Also calculate gas density under

stated conditions, knowing the molar mass.

AP10-6-01 Calculate the partial pressure of any gas in a mixture, given the composition

that mixture.

AP10-6-02 Calculate the mole fraction of a gas mixture, given its partial pressure and

total pressure of the system.

AP10-7,8-01 Describe how the distribution of speeds and the average speed of gas

molecules changes with temperature.AP10-7,8-02 Describe how the relative rates of effusion and diffusion of two gases depend

on their molar masses. (Graham's Law)

AP10-7,8-03 Use the principles of the Kinetic Molecular Theory of gases to explain the

nature of gas pressure and temperature at the molecular level.

AP10-9-01 Explain the origin of deviations shown by real gases from the ideal gas

relationship.

AP10-9-02 Cite the general conditions of P and T under which real gases most closely

approximate ideal gas behavior.

AP10-9-03 Explain the origins of the correction terms to P and V that appears in the vander Waals equation.


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Unit Objectives

Chapter 11 Intermolecular forces, Liquids & Solids

AP11-1-01 Employ the kinetic-molecular model to explain the differences in motion of

particles in gases, liquids and solids and how these relate to their states.

AP11-1-02 Describe the various types of intermolecular attractive forces, and state the

kinds of intermolecular forces expected for a substance given its molecular

structure.

AP11-3,4,5-01 Explain the meaning of the terms viscosity, surface tension, critical

temperature, and critical pressure, and account for the variations in these

properties in terms of intermolecular forces and temperature.

AP11-3,4,5-02

Explain the way in which the vapor pressure of a substance changes with

intermolecular forces and temperature.

AP11-3,4,5-03 Describe the relationship between the pressure on the surface of a liquid andthe boiling point of the liquid.

AP11-3,4,5-04 Given the needed heat capacities and enthalpies for phase changes, calculate

the heat absorbed or evolved when a qiven quantity of a substance changes

from one condition of another.

AP11-6-01 Draw a phase diagram of a substance given appropriate data, and use a phas

diagram to predict which phases are present at any given temperature and

pressure.

AP11-7,8-01 Distinguish between crystalline and amorphous solids.

AP11-7,8-02 Determine the net contents in a cubic unit cell, given a drawing or verbaldescription of the cell. Use this information, together with the atomic weight

of the atoms in the cell and the cell dimensions, to calculate the density of te

substance.

AP11-7,8-03 Describe the most efficient packing patterns of equal-sized spheres.

AP11-7,8-04 Predict the type of solid (molecular, covalent network, ionic, of metallic)

formed by a substances, ,and predict its general properties.


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Schedule & Assignments

Text Readings, Notes & Practice --- Option 1Complete the readings before discussions in class. Carefully read through the sample

exercises and attempt the practice exercise in the text. Pay particular attention to vocabulary

& concept reviews. After class discussion practice a min. of 5 EOCs for each section.

(Correct answers are in the back of the book or on MCHS AP CHEMISTRY.COM) Do what

you need to MASTER each concept. You should be able to summarize the main concepts in

the section both conceptually and mathematically.

Assignments for Unit: Option 1

Chapter 10 Gases

Text Section EOCQsDate

completeTopic

10-1 Characteristics of gases Pg 394-395 11-24

10-2 Pressure Pg 395 - 398 11-24Chemistry & LifeBlood Pressure Pg 39810-3 The Gas Laws Pg 398 - 402 25-2810-4 The Ideal Gas Equation Pg 402 - 406 29-4410-5 Further Applications of the ideal gasequation

Pg 406 - 410 45-58

Chemistry put to WorkGas Pipelines Pg 40910-6 Gas Mixtures and Partial Pressures Pg 410 - 413 59-7010-7 Kinetic Molecular Theory Pg 414 - 416 71-82A Closer LookThe ideal gas equation*** Pg 416

10-8 Molecular Effusion and Diffusion Pg 417 - 420 71-82Chemistry Put to WorkGas Separations Pg 42010-9 Real Gas Deviations from Ideal Behavior Pg 420 - 424 83-88

Visualizing Concepts MUST DO1-10

Additional Exercises MUST DO94,99,100,102,107

Integrative Exercises MUST DO112,113,116,120,121

Chapter 11 IMFs, Liquids & SolidsText Section EOCQs

Topic11-1 A Molecular Comparison of Gases,Liquids and Solids

Pg 436 - 439 9-12

11-2 Intermolecular Forces Pg. 439 - 447 13-2811-3 Some Properties of Liquids Pg 447 - 448 29-3211-4 Phase Changes Pg 449 - 452 33-42

Chemistry Put to Work.Supercritical fluid extraction

pg 453

11-5 Vapor Pressure Pg 453 - 455 43-50


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A Closer LookThe ClausiusClapeyronEquation

11-6 Phase Diagrams Pg 456 - 458 51-5611-7 Structure of Solids Pg 458 - 464 57-6811-8 Bonding in Solids Pg 464 - 469 69-78

A Closer Look X ray diffraction by CrystalsPg 465

A closer LookThe Third form of Carbon Pg468

Chapter Review, Key Skills & Key Equations Pg 469 -471

Visualizing Concepts Pg 471 MUST DO1-8

Additional Exercises Pg 477-478 MUST DO79,81,8385,89,90

Intergrative Exercises 478479 MUST DO100,101,102

Chapter 10. Gases

Common Student Misconceptions


Students need to be told toalways use temperature in Kelvin in gas problems. Students should always use units in gas-law problems to keep track of required

conversions. Due to several systems of units, students often use ideal gas constants with units

inconsistent with values. Students often confuse the standard conditions for gas behavior (STP) with the

standard conditions in thermodynamics. Ideal gas behavior should discussed as just that, ideal; students should be

reminded that real gases do not behave ideally, especially at high pressures and/orlow temperatures.

Students expect a change in the gas particle distribution upon temperature changesat constant V. Students commonly confuse effusion and diffusion. Students often think that all atoms of the same element must have the same

oxidation number and that this number is uniquely related to the atoms locationin the periodic table.

It is important that students know that titrations can be conducted not only usingacids and bases, but also in precipitation and oxidationreduction reactions.


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Lecture Outline

10.1 Characteristics of Gases(EOCQS 15,16,19,22,23)

All substances have three phases: solid, liquid, and gas.

Substances that are liquids or solids under ordinary conditions may also existas gases. These are often referred to as vapors. Many of the properties of gases differ from those of solids and liquids: Gases are highly compressible and occupy the full volume of their containers. When a gas is subjected to pressure, its volume decreases. Gases always form homogeneous mixtures with other gases. Gases only occupy a small fraction of the volume of their containers.

As a result, each molecule of gas behaves largely as though other moleculeswere absent.

FORWARD REFERENCES Thermodynamics of phase changes will be discussed in Ch. 19. Such important gaseous reactions as the Haber process or equilibria involvingnitrogen oxides will be covered in Ch. 15.

10.2 Pressure(EOCQS 15,16,19,22,23) Pressure is the force acting on an object per unit area:

Atmospheric Pressure and the Barometer

The SI unit of force is the newton (N). 1 N = 1 kg-m/s2

The SI unit of pressure is the pascal (Pa). 1 Pa = 1 N/m2

A related unit is the bar, which is equal to105 Pa. Gravity exerts a force on the Earths atmosphere.

A column of air 1 m2 in cross section extending to the upper atmosphereexerts a force of 105 N.

Thus, the pressure of a 1 m2

column of air extending to the upperatmosphere is 100 kPa. Atmospheric pressure at sea level is about 100 kPa or 1 bar.

The actual atmospheric pressure at a specific locationdepends on the altitude and weather conditions.

Atmospheric pressure is measured with a barometer.

A

FP


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If a tube is completely filled with mercury and then inverted into acontainer of mercury open to the atmosphere, the mercury will rise 760 mm upthe tube. Standard atmospheric pressure is the pressure required to support 760mm of Hg in a column. Important non-SI units used to express gas pressure include:

atmospheres (atm) millimeter of mercury (mm Hg) or torr 1 atm = 760 mm Hg = 760 torr = 1.01325 x 10 5 Pa = 101.325 kPa.

FORWARD REFERENCES Osmotic pressure (in atm) will be calculated in Ch. 13 (section 13.5). Kps and thermodynamic equilibrium constants in Ch. 15 will use pressure (in atm). Pressure and Le Chteliers principle will be discussed in Ch. 15 (section 15.7).

10.3 The Gas Laws (EOCQS 25-28) The equations that express the relationships among T(temperature), P (pressure),

V(volume), and n (number of moles of gas) are known as gas laws.The PressureVolume Relationship: Boyles Law

Weather balloons are used as a practical application of the relationship betweenpressure and volume of a gas.

As the weather balloon ascends, the volume increases. As the weather balloon gets further from Earths surface, theatmospheric pressure decreases.

Boyles law: The volume of a fixed quantity of gas, at constant temperature, isinversely proportional to its pressure.

Mathematically: A plot ofVversus P is a hyperbola. A plot ofVversus 1/P must be a straight line passing through the origin.

The working of the lungs illustrates Boyles law. As we breathe in, the diaphragm moves down, and the ribs expand;

therefore, the volume of the lungs increases. According to Boyles law, when the volume of the lungs increases, the

pressure decreases; therefore, the pressure inside the lungs is less thanatmospheric pressure.

Atmospheric pressure then forces air into the lungs until the pressureonce again equals atmospheric pressure.

As we breathe out, the diaphragm moves up and the ribs contract.Therefore, the volume of the lungs decreases.

By Boyles law, the pressure increases and air is forced out.

constantort1

constan PVP

V


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The TemperatureVolume Relationship: Charless Law

We know that hot-air balloons expand when they are heated. Charless law: The volume of a fixed quantity of gas at constant pressure is

directly proportional to its absolute temperature. Mathematically:

Note that the value of the constant depends on the pressure and numberof moles of gas.

A plot ofVversus Tis a straight line. When Tis measured in C, the intercept on the temperature

axis is273.15 C. We define absolute zero, 0 K =273.15 C.

The QuantityVolume Relationship: Avogadros Law

Gay-Lussacs law of combining volumes: At a given temperature and pressure thevolumes of gases that react with one another are ratios of small whole numbers.

Avogadros hypothesis: Equal volumes of gases at the same temperature andpressure contain the same number of molecules.

Avogadros law: The volume of gas at a given temperature and pressure isdirectly proportional to the number of moles of gas.

Mathematically:V= constant n

We can show that 22.4 L of any gas at 0 C and 1 atmosphere contains6.02 x 1023 gas molecules.FORWARD REFERENCES

Vapor pressure vs. temperature will be discussed in Ch. 13 (section 13.5). Increasing entropy of gases with temperature as well as entropy of

gases vs. other states of matter will be discussed in Ch. 19 (section 19.3).

10.4 The Ideal-Gas Equation(EOCQS 29-44 split)

Summarizing the gas laws: Boyle: V 1/P (constant n, T) Charles: V T (constant n, P) Avogadro: V n (constant P, T) Combined: VnT/P

Ideal gas equation: PV = nRT

constantorconstant T

VTV


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constant nRT

PV

An ideal gas is a hypothetical gas whose P, V, and Tbehavior iscompletely described by the ideal-gas equation. R = gas constant = 0.08206 L-atm/mol-K

Other numerical values of R in various units are given in Table10.2.

Define STP (standard temperature and pressure) = 0 C, 273.15 K, 1 atm.

The molar volume of 1 mol of an ideal gas at STP is 22.41 L.Relating the Ideal-Gas Equation and the Gas Laws

IfPV= nRTand n and Tare constant, then PVis constant and we have Boyleslaw. Other laws can be generated similarly.

In general, if we have a gas under two sets of conditions, then

We often have a situation in which P, V, and Tall change for a fixed number ofmoles of gas. For this set of circumstances,

Which gives

FORWARD REFERENCES The ideal gas constant will be used in Ch. 14 in the Arrhenius equation (section

14.5). The ideal gas constant will be used in Ch. 15 in conversions between Kc and Kp

(section 15.2) and to relate Gibbs free energy with the equilibrium constant in Ch.19 (section 19.7) as well as with the cell EMF in Ch. 20 (section 20.6).

22

22

11

11

Tn

VP

Tn

VP

2

221

1

1

T

VP

T

VP


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10.5 Further Applications of the Ideal-Gas Equation (EOCQS

45-58 split)

Gas Densities and Molar Mass

Density has units of mass over volume.

Rearranging the ideal-gas equation withMas molar mass we get

The molar mass of a gas can be determined as follows:

Volumes of Gases in Chemical Reactions

The ideal-gas equation relates P, V, and Tto number of moles of gas. The n can then be used in stoichiometric calculations.FORWARD REFERENCES

Solubility of gases vs. temperature (Henrys law) will be covered in Ch. 13 (section13.3).

10.6 Gas Mixtures and Partial Pressures (EOCQS 59-70 split)

Since gas molecules are so far apart, we can assume they behaveindependently.

Dalton observed: The total pressure of a mixture of gases equals the sum of the pressuresthat each would exert if present alone. Partial pressure is the pressure exerted by a particular component of a

gas mixture. Daltons law of partial pressures: In a gas mixture the total pressure is given by

the sum of partial pressures of each component:Pt = P1 + P2 + P3+

RT

PMd

P

dRTM

RT

P

V

n

RT

PM

V

nM


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V

RTn

V

RTnnnP

tt )( 321

Each gas obeys the ideal gas equation. Thus,

Partial Pressures and Mole Fractions

Let n1 be the number of moles of gas 1 exerting a partial pressure P1, then

P1 = Pt Where is the mole fraction (n1/nt). Note that a mole fraction is a dimensionless number.

Collecting Gases over Water

It is common to synthesize gases and collect them by displacing a volume ofwater.

To calculate the amount of gas produced, we need to correct for the partial

pressure of the water:Ptotal= Pgas + Pwater

The vapor pressure of water varies with temperature. Values can be found in Appendix B.

FORWARD REFERENCES Vapor pressure, volatility and temperature relationships will be introduced in Ch. 11

(section 11.5) and further applied to Raoults Law in Ch. 13 (section 13.5). Aira mixture of gaseswill be discussed in Ch. 18 (section 18.1) and 22 (section

22.7).

10.7 Kinetic-Molecular Theory (EOCQS 71-82 all)

The kinetic-molecular theory was developed to explain gas behavior. It is a theory of moving molecules.

Summary: Gases consist of a large number of molecules in constant random motion.

The combined volume of all the molecules is negligible compared withthe volume of the container.

Intermolecular forces (forces between gas molecules) are negligible. Energy can be transferred between molecules during collisions, but the

average kinetic energy is constant at constant temperature. The collisions are perfectly elastic.

The average kinetic energy of the gas molecules is proportional tothe absolute temperature.

Kinetic-molecular theory gives us an understanding of pressure andtemperature on the molecular level. The pressure of a gas results from the collisions with the walls of the

container.


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The magnitude of the pressure is determined by how often and how hardthe molecules strike.

The absolute temperature of a gas is a measure of the average kinetic energy. Some molecules will have less kinetic energy or more kineticenergy than the average (distribution).

There is a spread of individual energies of gas molecules in any sample

of gas. As the temperature increases, the average kinetic energy of the gasmolecules increases.

As kinetic energy increases, the velocity of the gas molecules increases. Root-mean-square (rms) speed, u, is the speed of a gas molecule

having average kinetic energy. Average kinetic energy, , is related to rms speed:

= mu2 where m = mass of the molecule.

Application to the Gas Laws We can understand empirical observations of gas properties within the framework

of the kinetic-molecular theory. Effect of an increase in volume (at constant temperature):

As volume increases at constant temperature, the average kinetic of thegas remains constant.

Therefore, u is constant. However, volume increases, so the gas molecules have to travelfurther to hit the walls of the container.

Therefore, pressure decreases. Effect of an increase in temperature (at constant volume): If temperature increases at constant volume, the average kinetic energy

of the gas molecules increases. There are more collisions with the container walls. Therefore, u increases.

The change in momentum in each collision increases (moleculesstrike harder).

Therefore, pressure increases.FORWARD REFERENCES

The collision model in Ch. 14 (section 14.5) will be based on the kinetic-moleculartheory.


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1

2

2

1

M

M

r

r

M

RTu

3

10.8 Molecular Effusion and Diffusion (EOCQS 71-82all)

The average kinetic energy of a gas is related to its mass: = mu2

Consider two gases at the same temperature: the lighter gas has a higher rms speedthan the heavier gas.

Mathematically:

The lower the molar mass,M, the higher the rms speed for that gas at aconstant temperature. Two consequences of the dependence of molecular speeds on mass are:

Effusion is the escape of gas molecules through a tiny hole into anevacuated space.

Diffusion is the spread of one substance throughout a space orthroughout a second substance.

Grahams Law ofEffusion

The rate of effusion can be quantified. Consider two gases with molar masses,M1 andM2, and with effusion rates, r1and r2, respectively:

The relative rate of effusion is given by Grahams law:

Only those molecules which hit the small hole will escape throughit.

Therefore, the higher the rms speed the more likely that a gas moleculewill hit the hole.

We can show

1

21

2 2

1

M

M

u

u

r

r


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Diffusion and Mean Free Path

Diffusion is faster for light gas molecules. Diffusion is significantly slower than the rms speed.

Diffusion is slowed by collisions of gas molecules with one another. Consider someone opening a perfume bottle: It takes awhile to detect theodor, but the average speed of the molecules at 25 C is about 515 m/s (1150

mi/hr). The average distance traveled by a gas molecule between collisions is called

the mean free path.

At sea level, the mean free path for air molecules is about 6 x 106cm.FORWARD REFERENCES

Similar molar mass related issues (e.g., passing of particles of solute throughsemipermeable membranes) for solutions will be discussed in Ch. 13 (section 13.5).

10.9 Real Gases: Deviations from Ideal Behavior

(EOCQS 83,85,88) From the ideal gas equation:

For 1 mol of an ideal gas, PV/RT= 1 for all pressures. In a real gas, PV/RTvaries from 1 significantly. The higher the pressure the more the deviation

from ideal behavior. For 1 mol of an ideal gas, PV/RT= 1 for all temperatures.

As temperature increases, the gases behave more ideally. The assumptions in the kinetic-molecular theory show where ideal gasbehavior breaks down: The molecules of a gas have finite volume. Molecules of a gas do attract each other.

As the pressure on a gas increases, the molecules are forced closer together. As the molecules get closer together, the free space in which themolecules can move gets smaller. The smaller the container, the more of the total space the gas molecules

occupy. Therefore, the higher the pressure, the less the gas resembles an ideal

gas. As the gas molecules get closer together, the intermolecular distances

decrease. The smaller the distance between gas molecules, the more likely that

attractive forces will develop between the molecules. Therefore, the less the gas resembles an ideal gas.

nRT

PV


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2

2

V

an

As temperature increases, the gas molecules move faster and further apart. Also, higher temperatures mean more energy available to break

intermolecular forces. As temperature increases, the negative departure from ideal-gas behavior

disappears.

The van der Waals Equation

We add two terms to the ideal gas equation to correct for The volume of molecules:

For molecular attractions:

The correction terms generate the van der Waals equation:

where a and b are empirical constants that differ for each gas. van der Waals constants for some common gases can be

found in Table 10.3. To understand the effect of intermolecular forces on pressure, consider a molecule

that is about to strike the wall of the container. The striking molecule is attracted by neighboring molecules. Therefore, the impact on the wall is lessened.

Chapter 11. Intermolecular Forces,Liquids, and Solids


Students confuse intermolecular and intramolecular forces.

Students often do not appreciate how important information from earlier chapters

is for the understanding of concepts in this chapter.

Students have difficulty predicting the relative strength of intermolecular forces

involved in different materials. Students are unaware that there can be intramolecular hydrogen bonding.

Iondipole interactions are technically interparticular forces.

Students confuse cohesion and adhesion.

Students do not realize that, under the right set of conditions, water also sublimes.

The term volatile is often used incorrectly, especially in the media.

Students often think that more viscous necessarily means more dense.

nRTnbVV

anP

2

2

nbV


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Lecture Outline

11.1 A Molecular Comparison of Gases, Liquids, and Solids

(EOC QS 9,11)

Physical properties of substances are understood in terms of kinetic-moleculartheory: Gases are highly compressible and assume the shape and volume of their

container. Gas molecules are far apart and do not interact much with one another.

Liquids are almost incompressible, assume the shape but not the volume ofthe container.

Liquids molecules are held together more closely than gas molecules butnot so rigidly that the molecules cannot slide past each other. Solids are incompressible and have a definite shape and volume. Solid molecules are packed closely together.

The molecules are so rigidly packed that they cannot easily slidepast each other.

Solids and liquids are condensed phases. Solids with highly ordered structures are said to be crystalline.

Converting a gas into a liquid or solid requires the molecules to get closer toeach other. We can accomplish this by cooling or compressing the gas.

Converting a solid into a liquid or gas requires the molecules to move furtherapart. We can accomplish this by heating or reducing the pressure on the gas.

The forces holding solids and liquids together are called intermolecular forces. Physical properties of liquids and solids are due to intermolecular forces.

These are forces between molecules.FORWARD REFERENCES

A comparison of phases in terms of entropy will be performed in Ch. 19 (section

19.3).


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11.2 Intermolecular Forces (EOC QS 13 27 split)

The attraction between molecules is an intermolecular force. Intermolecular forces are much weaker than ionic or covalent bonds.

When a substance melts or boils, intermolecular forces are broken. When a substances condenses, intermolecular forces are formed.

Boiling points reflect intermolecular force strength. A high boiling point indicates strong attractive forces.

Melting points also reflect the strength of attractive forces. A high melting point indicates strong attractive forces.

van der Waals forces are the intermolecular forces that exist between neutralmolecules. These include London dispersion forces, dipoledipole forces, and

hydrogen-bonding forces. Iondipole interactions are important in solutions.

These are all weak (


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London-Dispersion Forces

These are the weakest of all intermolecular forces. It is possible for two adjacent neutral molecules to affect each other.

The nucleus of one molecule (or atom) attracts the electrons of theadjacent molecule (or atom).

For an instant, the electron clouds become distorted. In that instant a dipole is formed (called an instantaneous dipole). One instantaneous dipole can induce another instantaneous dipolein an adjacent molecule (or atom).

These two temporary dipoles attract each other. The attraction is called the London dispersion force, or simply adispersion force. London dispersion forces exist between all molecules.

What affects the strength of a dispersion force?

Molecules must be very close together for these attractive forces tooccur.

Polarizability is the ease with which an electron distribution can bedeformed. The larger the molecule (the greater the number of electrons) the

more polarizable it is.

London dispersion forces increase as molecular weight increases. London dispersion forces depend on the shape of the molecule.

The greater the surface area available for contact, the greater thedispersion forces.

London dispersion forces between spherical molecules are smallerthan those between more cylindrically shaped molecules. Example: n-pentane vs. neopentane.


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Hydrogen Bonding

Experiments show that the boiling points of compounds with HF, HO, andHN bonds are abnormally high. Their intermolecular forces are abnormally strong.

Hydrogen bonding is a special type of intermolecular attraction.

This is a special case of dipoledipole interactions. H-bonding requires: H bonded to a small electronegative element (most important for

compounds of F, O, and N). an unshared electron pair on a nearby small electronegative ion or

atom (usually F, O, or N on another molecule). lement) lie

much closer to X than H.

bare proton to the X-. Bond energies of hydrogen bonds vary from about 4 kJ/mol to 25

kJ/mol. They are much weaker than ordinary chemical bonds.

Intermolecular and intramolecular hydrogen bonds have exceedinglyimportant biological significance.

They are important in stabilizing protein structure, in DNA structure andfunction, etc.

An interesting consequence of H-bonding is that ice floats. The molecules in solids are usually more closely packed than those in

liquids. Therefore, solids are usually more dense than liquids.

Ice is ordered with an open structure to optimize H-bonding. Water molecules in ice are arranged in an open, regular hexagon.

Each + H points towards a lone pair on O. Therefore, ice is less dense than water.

Ice floats, so it forms an insulating layer on top of lakes, rivers, etc.Therefore, aquatic life can survive in winter.

Water expands when it freezes. Frozen water in pipes may cause them to break in cold weather.

Comparing Intermolecular Forces

Dispersion forces are found in all substances. Their strength depends on molecular shapes and molecular weights.


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Dipoledipole forces add to the effect of dispersion forces. They are found only in polar molecules.

H-bonding is a special case of dipoledipole interactions. It is the strongest of the intermolecular forces involving neutral species. H-bonding is most important for H compounds of N, O, and F.

If ions are involved, iondipole (if a dipole is present) and ionic bonding arepossible. Iondipole interactions are stronger than H-bonds.

Keep in mind that ordinary ionic or covalent bonds are much stronger thanthese interactions!

FORWARD REFERENCES

Molecular materials held together by intermolecular forces will be called soft

materials in Ch. 12.

Intermolecular forces between polymer chains and in liquid crystals will be

mentioned in Ch. 12 (sections 12.6 and 12.8).

Breaking of solutesolute and solventsolvent intermolecular forces and replacing

them with solutesolvent interactions will take place in the solution process (Ch.

13).

The binding between the substrate and the active site in the enzyme action thanks to

the intermolecular forces will be discussed in Ch. 14 (section 14.7).

Hydrogen bonding and the formation of hydrated hydronium ions will be discussed

in Ch. 16 (section 16.2).

Hydrogen bonding will be partially responsible for the relative weakness of HF

compared to the strength of other binary acids involving halides (section 16.10).

Hydrogen bonding and high heat capacity, high melting, and boiling points of water

will be mentioned again in Ch. 18 (section 18.5). Intermolecular attractions in ice will be discussed in Ch. 19 (section 19.3).

Hydrogen bonding in alcohols will be discussed in Ch. 25 (section 25.5).

Hydrogen bonding in the helix of a protein will be discussed in

Ch. 25 (section 25.9).

11.3 Some Properties of Liquids (EOC QS 29,31)

Viscosity

Viscosity is the resistance of a liquid to flow.

A liquid flows by sliding molecules over one another. Viscosity depends on: the attractive forces between molecules.

The stronger the intermolecular forces, the higher the viscosity. the tendency of molecules to become entangled.

Viscosity increases as molecules become entangled with oneanother.


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the temperature. Viscosity usually decreases with an increase in temperature.

Surface Tension

Bulk molecules (those in the liquid) are equally attracted to their neighbors. Surface molecules are only attracted inward towards the bulk molecules.

Therefore, surface molecules are packed more closely than bulkmolecules. This causes the liquid to behave as if it had a skin.

Surface tension is the amount of energy required to increase the surface area ofa liquid by a unit amount.

Stronger intermolecular forces cause higher surface tension. Water has a high surface tension (H-bonding)

Hg(l) has an even higher surface tension (there are very strongmetallic bonds between Hg atoms).

Cohesive and adhesive forces are at play. Cohesive forces are intermolecular forces that bind molecules to one

another. Adhesive forces are intermolecular forces that bind molecules to a

surface. Illustrate this by looking at the meniscus in a tube filled with liquid.

The meniscus is the shape of the liquid surface. If adhesive forces are greater than cohesive forces, the

liquid surface is attracted to its container more than the bulk molecules.Therefore, the meniscus is U-shaped (e.g., water in glass). If cohesive forces are greater than adhesive forces, the meniscus iscurved downwards (e.g., Hg(l) in glass)

Capillary action is the rise of liquids up very narrow tubes. The liquid climbs until adhesive and cohesive forces are balanced by

gravity.FORWARD REFERENCES

Viscosity of liquid crystals will be discussed in Ch. 12 (section 12.8).

Viscosity of organic compounds, such as polyhydroxyl alcohols,

will be mentioned in Ch. 25 (section 25.5).


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11.4 Phase Changes (EOC QS 33 36 ALL, 37,39,41)

Phase changes are changes of state.

Matter in one state is converted into another state.

Sublimation: solid gas. Melting or fusion: solid liquid. Vaporization: liquid gas. Deposition: gas solid. Condensation: gas liquid. Freezing: liquid solid.

Energy Changes Accompanying Phase Changes

Energy changes of the system for the above processes are:melting or fusion: Hfus > 0 (endothermic). The enthalpy of fusion is known as the heat of fusion.

vaporization: Hvap > 0 (endothermic). The enthalpy of vaporization is known as the heat of vaporization.

sublimation: Hsub > 0 (endothermic). The enthalpy of sublimation is called the heat of sublimation.

deposition: Hdep < 0 (exothermic). condensation: Hcon < 0 (exothermic). freezing: Hfre < 0 (exothermic).

Generally the heat of fusion (enthalpy of fusion) is less than heat ofvaporization. It takes more energy to completely separate molecules than to partially

separate them. All phase changes are possible under the right conditions (e.g., water sublimes

when snow disappears without forming puddles). The following sequence is endothermic:heat solid melt heat liquid boil heat gas

The following sequence is exothermic:cool gas condense cool liquid freeze cool solid

Heating Curves

Plot of temperature change versus heat added is a heating curve. During a phase change adding heat causes no temperature change.


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The added energy is used to break intermolecular bonds rather thancause a temperature change.

These points are used to calculate Hfus and Hvap. Supercooling: When a liquid is cooled below its freezing point and it still

remains a liquid.

Critical Temperature and Pressure Gases may be liquefied by increasing the pressure at a suitable temperature. Critical temperature is the highest temperature at which a substance can exist

as a liquid. Critical pressure is the pressure required for liquefaction at this critical

temperature. The greater the intermolecular forces, the easier it is to liquefy a

substance. Thus the higher the critical temperature.

FORWARD REFERENCES

Supercritical fluids in green chemistry will be discussed in Ch. 18 (section 18.7).

Thermodynamics of phase changes will be further discussed throughout Ch. 19.

11.5 Vapor Pressure (EOC QS 43-49 ODD)

Explaining Vapor Pressure on the Molecular Level

Some of the molecules on the surface of a liquid have enough energy to escapethe attraction of the bulk liquid. These molecules move into the gas phase.

As the number of molecules in the gas phase increases, some of the gas phasemolecules strike the surface and return to the liquid.

After some time the pressure of the gas will be constant. A dynamic equilibrium has been established.

Dynamic equilibrium is a condition in which two opposingprocesses occur simultaneously at equal rates.

In this case, it is the point when as many molecules escape thesurface as strike the surface.

Vapor pressure of a liquid is the pressure exerted by its vapor when theliquid and vapor are in dynamic equilibrium. The pressure of the vapor at this point is called the equilibrium

vapor pressure.


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Volatility, Vapor Pressure, and Temperature

If equilibrium is never established the vapor continues to form. Eventually, the liquid evaporates to dryness.

Liquids that evaporate easily are said to be volatile. The higher the temperature, the higher the average kinetic energy, the

faster the liquid evaporates.Vapor Pressure and Boiling Point Liquids boil when the external pressure at the liquid surface equals the vapor

pressure. The normal boiling point is the boiling point at 760 mm Hg (1 atm).

The temperature of the boiling point increases as the external pressureincreases.

Two ways to get a liquid to boil increase temperature or decrease pressure. Pressure cookers operate at high pressure. At high pressure the boiling point of water is higher than at 1 atm. Therefore, food is cooked at a higher temperature.

FORWARD REFERENCES

Vapor pressure reduction of the solvent in a solutiona colligative propertywill

be discussed in Ch. 13 (section 13.5).

11.6 Phase Diagrams (EOC QS 51 56 ALL)

A phase diagram is a plot of pressure vs. temperature summarizing allequilibria between phases. Phase diagrams tell us which phase will exist at a given temperature and

pressure. Features of a phase diagram include:

vapor-pressure curve: generally as temperature increases, vapor pressureincreases.

critical point: critical temperature and pressure for the gas. normal melting point: melting point at 1 atm.

triple point: temperature and pressure at which all three phases arein equilibrium.

Any temperature and pressure combination not on a curve represents asingle phase.


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Phase Diagrams of H2O and CO2

Water: In general, an increase in pressure favors the more compact phase of the

material. This is usually the solid.

Water is one of the few substances whose solid form is less dense thanthe liquid form.

The melting point curve for water slopes to the left. The triple point occurs at 0.0098 C and 4.58 mm Hg. The normal melting (freezing) point is 0 C. The normal boiling point is 100 C. The critical point is 374 C and 218 atm.

Carbon Dioxide:

The normal sublimation psublimes, it does not melt.)

The critical point occurs at 31.1 C and 73 atm. Freeze drying: Frozen food is placed in a low pressure (< 4.58torr) chamber. The ice sublimes.

FORWARD REFERENCES

Phase diagrams for a pure solvent and for a solution of a nonvolatile solute will be

discussed in Ch. 13 (section 13.5).

Phase equilibria at melting and boiling points will be further analyzed in Ch. 19.

11.7 Structures of Solids(EOC QS 57,59)

A crystalline solid has a well-ordered, definite arrangements of molecules,atoms, or ions.

Examples are quartz, diamond, salt, and sugar. The intermolecular forces are similar in strength.

Thus they tend to melt at specific temperatures.

In an amorphous solid, molecules, atoms or ions do not have an orderlyarrangement.

Examples are rubber and glass. Amorphous solids have intermolecular forces that vary in strength.

Thus they tend to melt over a range of temperatures.


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Unit Cells

Crystalline solids have an ordered, repeating structure. The smallest repeating unit in a crystal is a unit cell.

The unit cell is the smallest unit with all the symmetry of the entirecrystal.

The three-dimensional stacking of unit cells is the crystal lattice. Each point in the lattice is a lattice point which represents anidentical environment within the solid.

There are three types of cubic unit cells.

Primitive cubic.

The lattice points are at the corners of a simple cube with eachatom shared by eight unit cells.

Body-centered cubic (bcc).

Lattice points occur at the corners of a cube and in addition there

is a lattice point in the center of the body of the cube. The cornerlattice points are shared by eight unit cells, and the center atom iscompletely enclosed in one unit cell.

Face-centered cubic (fcc).

There are lattice points at the corners of a cube plus one latticepoint in the center of each face of the cube. Eight unit cells sharethe corner lattice points and two unit cells share the face latticepoints.

The Crystal Structure of Sodium Chloride It has a face-centered cubic lattice. There are two equivalent ways of defining this unit cell:

Cl (larger) ions at the corners of the cell, or Na+ (smaller) ions at the corners of the cell.

The cation to anion ratio in a unit cell is the same for the crystal. In NaCl each unit cell contains the same number of Na+ and Cl ions.

Note that the unit cell for CaCl2 needs twice as many Cl ions as Ca2+ ions.

Close Packing of Spheres

Crystalline solids have structures that maximize the attractive forces betweenparticles. Their particles can be modeled by spheres.

Each atom or ion is represented by a sphere. Molecular crystals are formed by close packing of the molecules.


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Maximum intermolecular forces in crystals are achieved by the close packingof spheres. A crystal is built up by placing close packed layers of spheres on top of

each other. There is only one place for the second layer of spheres.

There are two choices for the third layer of spheres: The third layer eclipses the first (ABAB arrangement). This is called hexagonal close packing (hcp).

The third layer is in a different position relative to the first(ABCABC arrangement). This is called cubic close packing (ccp).

Note that the unit cell of a ccp crystal is face-centered cubic.

In both close-packed structures, each sphere is surrounded by 12 other s

pheres (6 in one plane, 3 above, and 3 below). Coordination number is the number of spheres directly surrounding a

central sphere.

When spheres are packed as closely as possible, there are small spacesbetween adjacent spheres (interstitial holes).

If unequally sized spheres are used, the smaller spheres are placed in theinterstitial holes. For example: Li2O

The larger O2 ions assume the cubic close-packed structurewith the smaller Li+ ions in the holes.

FORWARD REFERENCES

Coordination numbers will be used in Ch. 23 for transition metal complexes.

11.8 Bonding in Solids(EOC QS 69 78 ALL)

The physical properties of crystalline solids depend on the: attractive forces between particles and on the arrangement of the particles.

Molecular Solids

Molecular solids consist of atoms or molecules held together by intermolecularforces.

Weak intermolecular forces give rise to low melting points. Intermolecular forces include dipoledipole, London dispersion and H-

bonds.


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Molecular solids are usually soft. They are often gases or liquids are room temperature.

Efficient packing of molecules is important (since they are not regularspheres).

Molecular solids show poor thermal and electrical conductivity.

Examples: Ar(s), CH4(s), CO2(s), sucrose.

Covalent-Network Solids

Covalent-network solids consist of atoms held together, in large networks orchains, with covalent bonds.

They have much higher melting points and are much harder than molecularsolids. This is a consequence of the strong covalent bonds that connect the

atoms. Examples are diamond, graphite, quartz (SiO2), and silicon carbide (SiC). In diamond:

each C atom has a coordination number of 4. each C atom is tetrahedral. there is a three-dimensional array of atoms. Diamond is hard, and has a high melting point (3550C).

In graphite: each C atom is arranged in a planar hexagonal ring.

layers of interconnected rings are placed on top of each other. the distance between adjacent C atoms in the same layer is closeto that seen in benzene (1.42 vs. 1.395 in benzene).

electrons move in delocalized orbitals (good conductor). the distance between layers is large (3.41 ). the layers are held together by weak dispersion forces.

They slide easily past each other. Graphite is a good lubricant.

Ionic Solids Ionic solids consist of ions held together by ionic bonds.

They are hard, brittle and have high melting points. Ions (spherical) are held together by electrostatic forces of attraction. Recall: The larger the charges (Q1, Q2) and smaller the distance (d) between ions, the

stronger the ionic bond.


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The structure of the ionic solid depends on the charges on the ions and on therelative sizes of the atoms.

Examples of some ionic lattice types are: NaCl structure.

Each ion has a coordination number of six.

It has a face-centered cubic lattice. The cation to anion ratio is 1:1. Other similar examples: LiF, KCl, AgCl, and CaO.

CsCl structure. Cs+ has a coordination number of eight. It is different from the NaCl structure (Cs+ is larger than Na+). The cation to anion ratio is 1:1.

Zinc blende (ZnS) structure. S2 ions adopt a face-centered cubic arrangement. Zn2+ ions have a coordination number of four. The S2 ions are placed in a tetrahedron around the Zn2+ ions. Another example is CuCl.

Fluorite (CaF2) structure. Ca2+ ions are in a face-centered cubic arrangement. There are twice as many F ions as Ca2+ ions in each unit cell. Other examples are BaCl2 and PbF2.

Metallic Solids

Metallic solids consist entirely of metal atoms. Metallic solids are soft or hard. They have high melting points. They show good electrical and thermal conductivity. They are ductile and malleable.

Examples are all metallic elements (i.e., Al, Cu, Fe, Au) Metallic solids have metal atoms in hexagonal close-packed, face-centered

cubic or body-centered cubic arrangements.

Thus the coordination number for each atom is either 8 or 12. A problem that needs to be explained:

The bonding is too strong to be explained by London dispersion forcesand there are not enough electrons for covalent bonds.

Resolution: The metal nuclei float in a sea of delocalized valence electrons.


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Metals conduct heat and electricity because the valence electrons aredelocalized and are mobile.

X-Ray Diffraction When waves are passed through a narrow slit they are diffracted.

When waves are passed through a diffraction grating (many narrow slits inparallel) they interact to form a diffraction pattern (areas of light and darkbands).

Efficient diffraction occurs when the wavelength of light is close to the size ofthe slits.

The spacing between layers in a crystal is 220 , which is the wavelengthrange for X-rays.

X-ray diffraction (X-ray crystallography): X-rays are passed through the crystal and are detected on a photographic

plate. The photographic plate has one bright spot at the center (incident beam)

as well as a diffraction pattern. Each close-packing arrangement produces a different diffraction pattern.

Knowing the diffraction pattern, we can calculate the positions ofthe atoms required to produce that pattern.

We calculate the crystal structure based on a knowledge of thediffraction pattern.

FORWARD REFERENCES Classes of materials will be discussed throughout Ch. 12, [e.g., carbon nanotubes

(section 12.9)].

Perfectly crystalline solids will be mentioned in Ch. 19 (Third Law of

Thermodynamics).

Electrolysis of ionic salts will be further discussed in Ch. 20 (section 20.9).

Allotropic forms of carbon will be discussed in Ch. 22 (section 22.1).

Lubricating properties of silicate sheets will be mentioned in Ch. 22 (section 22.10).

Electron-sea model for metallic bonding as well as alloys will be

discussed in Ch. 23.

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Unit 03 Lecture Notes & Learning Standards