Chapter 11

Chapter 11

Liquids and solidsLiquids and solids

Almost all substances that are liquids are molecular, (held together by the covalent bonds within the molecule)

The physical properties of molecular liquids and solids is due to the intermolecular forces that hold them together

They are similar compared to gases. They are incompressible. Their density doesn’t change with

temperature. These similarities are due

• to the molecules being close together in solids and liquids

• and far apart in gases What holds them close together?

Intermolecular forces Inside molecules (intramolecular) the

atoms are bonded to each other. Intermolecular refers to the forces

between the molecules. These are what hold the molecules

together in the condensed states.

Intermolecular forces are much weaker than ionic or covalent bonds

Less energy is required to vaporize water, than to break the bonds between hydrogen and oxygen.

When a molecule changes state from solid-liquid-gas, the molecule itself stays together

Properties of liquids Many are related to the

intermolecular forces that hold them together

Boiling point- in order for a liquid to boil the molecules must overcome their attractive forces between them in order to separate. The stronger the intermolecular force, the higher the boiling point. (Melting point is based on the same relationship)

Ion dipole forces Refer to a solution, not a molecule. Force that exists between an ion and

the partial charge on the end of a polar molecules.

Like NaCl dissolved in H2O The strength of the force increases

as the charge on the ion increases, or the strength of the dipole increases.

Types of Intermolecular Forces3 Types for neutral molecules:

Dipole-dipole

London Dispersion

Hydrogen

All are based on electrostatic attractions

All called Van der Waals forces (man who predicted the deviation of gases from the ideal gas laws due to attractive forces between them)

Dipole - Dipole Molecules line up in the presence of a

electric field. The opposite ends of the dipole can attract each other so the molecules stay close together.

1% as strong as covalent bonds Only work when polar molecules are

very close together

For molecules of approximately equal mass and size, the strengths of intermolecular attractions increase with increasing polarity

The higher the dipole moment, the higher the boiling point

London Dispersion Forces Non - polar molecules also exert

forces on each other. Otherwise, no solids or liquids. Electrons are not evenly distributed at

every instant in time. Have an instantaneous dipole. Induces a dipole in the atom next to it. Induced dipole- induced dipole

interaction.

London Dispersion Forces Weak, short lived. Lasts longer at low temperature. Eventually long enough to make liquids. More electrons, more polarizable. Bigger molecules, higher melting and

boiling points. Much, much weaker than other forces. BP & MP of halogens and noble gases

increase as you go down

Hydrogen Bonding Especially strong dipole-dipole forces

when H is attached to F, O, or N These three because-

• They have high electronegativity.

• They are small enough to get close. Effects boiling point.

CH4

SiH4

GeH4SnH4

PH3

NH3 SbH3

AsH3

H2O

H2SH2Se

H2Te

HF

HI

HBrHCl

Boiling Points

0ºC

100

-100

200

Water

+-

+

Example

H H H HH H H H

- +

H H H H

- +

Liquids Many of the properties due to

internal attraction of atoms.

• Beading

• Surface tension

• Capillary action Stronger intermolecular forces cause

each of these to increase.

Viscosity Which will pour very easily, syrup or

lemonade? Both liquids so what accounts for the difference?

Viscosity- How much a liquid resists flowing.

The higher the viscosity, the slower it flows.

Related to how the molecules of the liquid move with respect to one another.

Viscosity decreases as temperature increases.

Surface tension

Molecules in the middle are attracted in all directions.

Molecules at the the top are only pulled inside.

Minimizes surface area.

Water has a high surface tension because of it’s hydrogen bonds

Cohesive- the intermolecular forces that bind like molecules together (very strong for water)

Adhesive- the intermolecular forces that bind a molecule to another surface.

Mercury/ water

Capillary Action Liquids spontaneously rise in a

narrow tube. Helps water move through plants Occurs because of the attraction

between the polar glass and the polar water molecules.

Beading If a polar substance is

placed on a non-polar surface. • There are cohesive,• But no adhesive forces.

When you wax

car.

Relation to intermolecular forces? The stronger the force, the bigger

the effect.

• Hydrogen bonding

• Polar bonding

• LDF

Phases

The phase of a substance is determined by three things.

The temperature. The pressure. The strength of intermolecular

forces.

Phase changes Directly related to the strength of the

intermolecular forces> Na has a boiling point of 883°C

(strong metallic bonds) While iodine readily sublimes at

room temperature (LDF) When your phase change is going to

a less ordered state (more entropy) energy must be supplied

Energy Requirement Fusion- to melt, the energy required

to change from the solid to liquid state is called the heat of fusion (Hfus)

Vaporization- to turn from a liquid to a gas, the energy required to change from a liquid to a gas is called the heat of vaporization (Hvap).

The heat of vaporization is usually much larger than the heat of fusion.

Fusion, vaporization, and sublimation are endothermic

Freezing, condensation, and deposition are exothermic

Steam burns, sweating cools

Changes of state The graph of temperature versus

heat applied is called a heating curve.

The temperature a solid turns to a liquid is the melting point.

-40

-20

0

20

40

60

80

100

120

140

0 40 120 220 760 800

Heating Curve for Water

IceWater and Ice

Water

Water and Steam

Steam

-40

-20

0

20

40

60

80

100

120

140

0 40 120 220 760 800

Heating Curve for Water

Heat of Fusion

Heat ofVaporization

Slope is Heat Capacity

Calculating Energy Changes Q=mcΔT C= specific heat (given in J/g-K) M=mass ΔT= change in T, no need to convert

To melt To vaporize molesHfus molesHvap

Calculate the enthalpy change of converting 1 mole of ice at -25°C to water vapor at 125°C. (The pressure is held constant at 1 atm). The specific heats of ice is 2.09 J/g-K, the specific heat of water is 4.18 J/g-K, and the specific heat of vapor is 1.84 J/g-K. The heat of fusion for water is 6.01 kJ/mol, and the heat of vaporization is 40.67 kJ/mol.

You Try! Ethanol (C2H5OH) melts at -114°C,

and boils at 78°C. The enthalpy of fusion of ethanol is 5.02 kJ/mol, and the enthalpy of vaporization 38.56 kJ/mol. The specific heats of solid and liquid ethanol are respectively .97J/g-K, and 2.3J/g-K. How much heat is required to convert 75 g of ethanol at -120°C to the vapor phase at 78°C?

Critical Temperature and Pressure As temperature rises, gases become

harder to liquefy because their molecules have high KE and are very far apart

Critical Temperature- the highest temperature a substance can exist as a liquid

Critical Pressure- The pressure required to liquefy the gas at the critical temperature.

The stronger the intermolecular forces, the easier a gas will liquefy.

Vapor Pressure Vaporization - change from

liquid to gas at boiling point. Evaporation - change from

liquid to gas below boiling point

Vaporization is an endothermic process - it requires heat.

Energy is required to overcome intermolecular forces pushing the molecules far apart into the gas phase.

Condensation Change from gas to liquid. Achieves a dynamic equilibrium with

vaporization in a closed system. What is a closed system? A closed system means matter

can’t go in or out.

Dynamic Equilibrium The two opposing processes of

condensation and vaporization are occurring at the same rate.

May look as if nothing is happening because there is no net change

The vapor pressure is the pressure exerted by its vapor when the liquid and vapor states of a substance are at dynamic equilibrium

In an open system As the water evaporates, the

molecules spread out and are not recaptured, no equilibrium is reached and the water continues to evaporate until there is nothing left.

Liquids that have high vapor pressure evaporate quickly and are called volatile (gasoline)

Vapor pressure increases with increasing T.

Boiling Point Reached when the vapor pressure

equals the external pressure. Normal boiling point is the boiling

point at 1 atm pressure. Where will water boil faster, at the

top of a hill or the bottom of a valley? Where will pasta cook faster, at the

top of a hill or the bottom of a valley?

Phase Diagrams. A plot of temperature versus

pressure for a closed system, with lines to indicate where there is a phase change.

Shows where equilibrium exists between different states of matter.

Can use it to predict the most stable phase of matter at a given T and P.

Temperature

SolidLiquid

Gas

1 Atm

AA

BB

CCD

D D

Pre

ssur

e

D

SolidLiquid

Gas

Triple Point

Critical Point

Temperature

Pre

ssur

e

SolidLiquid

Gas

This is the phase diagram for water. The density of liquid water is higher

than solid water.

Temperature

Pre

ssur

e

Solid Liquid

Gas

1 Atm

This is the phase diagram for CO2

The solid is more dense than the liquid The solid sublimes at 1 atm.

Temperature

Pre

ssur

e

Solids Two major types. Amorphous- those with much

disorder in their structure. (Rubber and glass)

Crystalline- have a regular arrangement of components in their structure. (Diamond and Quartz)

Melting Points The melting points of crystals are

definite. The melting points of amorphous

solids vary.

Crystals Lattice- a three dimensional grid that

describes the locations of the pieces in a crystalline solid.

Unit Cell-The smallest repeating unit in of the lattice.

Three common types.

Body-Centered Cubic

Face-Centered Cubic

Solids There are many amorphous solids. Like glass. We tend to focus on crystalline solids. two types. Ionic solids have ions at the lattice

points. Molecular solids have molecules. Sugar vs. Salt.

The book drones on about Using diffraction patterns to identify

crystal structures. Talks about metals and the closest

packing model. It is interesting, but trivial. We need to focus on metallic bonding. Why do metal atoms stay together. How there bonding effect their

properties.

Metallic bonding

1s

2s2p

3s

3pFilled Molecular Orbitals

Empty Molecular Orbitals

Magnesium Atoms

Filled Molecular OrbitalsEmpty Molecular Orbitals

The 1s, 2s, and 2p electrons are close to nucleus, so they are not able to move around.

1s

2s2p

3s

3p

Magnesium Atoms


1s

2s2p

3s

3p

Magnesium Atoms

The 3s and 3p orbitals overlap and form molecular orbitals.


1s

2s2p

3s

3p

Magnesium Atoms

Electrons in these energy level can travel freely throughout the crystal.


1s

2s2p

3s

3p

Magnesium Atoms

This makes metals conductors

Malleable because the bonds are flexible.

Carbon- A Special Atomic Solid There are three types of solid carbon. Amorphous- coal uninteresting. Diamond- hardest natural substance

on earth, insulates both heat and electricity.

Graphite- slippery, conducts electricity.

How the atoms in these network solids are connected explains why.

Diamond- each Carbon is sp3hybridized, connected to four

other carbons. Carbon atoms are

locked into tetrahedral shape.

Strong bonds give the huge molecule its hardness.

Why is it an insulator?

Empty MOs

Filled MOs

EThe space between orbitals make it impossible for electrons to move around

Each carbon is connected to three other carbons and sp2 hybridized.

The molecule is flat with 120º angles in fused 6 member rings.

The bonds extend above and below the plane.

Graphite is different.

This bond overlap forms a huge bonding network.

Electrons are free to move through out these delocalized orbitals.

The layers slide by each other.

Molecular solids. Molecules occupy the corners of the

lattices. Different molecules have different

forces between them. These forces depend on the size of

the molecule. They also depend on the strength

and nature of dipole moments.

Those without dipoles.

Most are gases at 25ºC. The only forces are London Dispersion

Forces. These depend on size of atom. Large molecules (such as I2 ) can be

solids even without dipoles.

Those with dipoles. Dipole-dipole forces are generally

stronger than L.D.F. Hydrogen bonding is stronger than

Dipole-dipole forces. No matter how strong the

intermolecular force, it is always much, much weaker than the forces in bonds.

Stronger forces lead to higher melting and freezing points.

Water is special

HO

H

-

Each molecule has two polar O-H bonds.

Water is special

HO

H

Each molecule has two polar O-H bonds.

Each molecule has two lone pair on its oxygen.

Water is special Each molecule has two polar

O-H bonds. Each molecule has two lone

pair on its oxygen. Each oxygen can interact with

4 hydrogen atoms.

HO

H

Water is special

HO

H

HO

H

HO

H

This gives water an especially high melting and boiling point.

Ionic Solids The extremes in dipole dipole forces-

atoms are actually held together by opposite charges.

Huge melting and boiling points. Atoms are locked in lattice so hard and

brittle. Every electron is accounted for so they

are poor conductors-good insulators.

Dynamic equilibrium When first sealed the molecules gradually

escape the surface of the liquid.


escape the surface of the liquid. As the molecules build up above the

liquid some condense back to a liquid.


escape the surface of the liquid. As the molecules build up above the

liquid some condense back to a liquid. As time goes by the rate of

vaporization remains constant but the rate of condensation increases because there are more molecules to condense.


escape the surface of the liquid As the molecules build up above the

liquid some condense back to a liquid. As time goes by the rate of

vaporization remains constant but the rate of condensation increases because there are more molecules to condense.

Equilibrium is reached when

Rate of Vaporization = Rate of

Condensation Molecules are constantly changing

phase “Dynamic” The total amount of liquid and vapor

remains constant “Equilibrium”

Dynamic equilibrium

Vapor pressure The pressure above the liquid at

equilibrium. Liquids with high vapor pressures

evaporate easily. They are called volatile.

Decreases with increasing intermolecular forces.

• Bigger molecules (bigger LDF)

• More polar molecules (dipole-dipole)

Vapor pressure Increases with increasing

temperature. Easily measured in a barometer.

Dish of Hg

Vacuum

Patm=

760 torr

A barometer will hold a column of mercury 760 mm high at one atm

Dish of Hg

Vacuum

Patm=

760 torr

A barometer will hold a column of mercury 760 mm high at one atm.

If we inject a volatile liquid in the barometer it will rise to the top of the mercury.

Dish of Hg

Patm=

760 torr

A barometer will hold a column of mercury 760 mm high at one atm.

If we inject a volatile liquid in the barometer it will rise to the top of the mercury.

There it will vaporize and push the column of mercury down.

Water

Dish of Hg

736 mm Hg

Water Vapor

The mercury is pushed down by the vapor pressure.

Patm = PHg + Pvap

Patm - PHg = Pvap

760 - 736 = 24 torr

Temperature Effect

Kinetic energy

# of

mol

ecu

les

T1

Energy needed to overcome intermolecular forces

Kinetic energy

# of

mol

ecu

les

T1


T1

T2

At higher temperature more molecules have enough energy - higher vapor pressure.


Mathematical relationship

ln is the natural logarithm

• ln = Log base e

• e = Euler’s number an irrational number like

Hvap is the heat of vaporization in J/mol

lnP

P =

H

R T-T

T1vap

T2vap

vap

2 1

1 1

R = 8.3145 J/K mol. Surprisingly this is the graph of a

straight line. (actually the proof is in the book)

Mathematical relationship

lnP

P =

H

R T-T

T1vap

T2vap

vap

2 1

1 1

Chapter 11

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