Liquids and solidsLiquids and solids

10.1 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 Strong

•covalent bonding• ionic bonding

Weak•Dipole dipole•London dispersion forces•Hydrogen bonding

During phase changes the molecules stay intact.

Energy used to overcome forces.

Dipole - Dipole Remember where the polar

definition came from? 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 or ionic bonds.

Weaker with greater distance. Small role in gases.

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. Affects boiling point.

Water

+

-+

Hydrogen Bonding Clip

CH4

SiH4

GeH4SnH4

PH3

NH3 SbH3

AsH3

H2O

H2SH2Se

H2Te

HF

HI

HBrHCl

Boiling Points

0ºC

100

-100

200

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.

Example

H H H HH H H H

+ H H H H

+ +

LD Video

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. Exist in all molecules, much weaker

than other forces. Also called Van der Waal’s forces.

#36

#38

#39

#40 (In Webassign)

10.2 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.

Surface Tension

Molecules in the middle are attracted in all directions.

Molecules at the top are only pulled inside.

Minimizes surface area.

Capillary Action Liquids spontaneously rise in a

narrow tube. Intermolecular forces are cohesive,

connecting like things. Adhesive forces connect to

something else. Glass is polar. It attracts water molecules.

Beading If a polar substance is

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

forces. And Visa Versa

Viscosity How much a liquid resists flowing. Large forces, more viscous. Large molecules can get tangled

up.

Model for Liquids Can’t see molecules so picture

them in motion but attracted to each other.

With regions arranged like solids but•with higher disorder.•with fewer holes than a gas.•Highly dynamic

Phases The phase of a substance is

determined by three things. The temperature. The pressure. The strength of intermolecular

forces. Changes of State Video

10.3 Solids Two major types. Crystalline - have a regular

arrangement of components in their structure. (table salt)

Amorphous - those with much disorder in their structure. Said to be “frozen in place”. (glass, plastic)

10.4 Bonding Models for Metals

Why do metal atoms stay together? How does their bonding effect their properties?

Two Models: Electron Sea Model: A regular array

of metals in a “sea” of electrons. Band (Molecular Orbital) Model:

Electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms.

Electron Sea Model

10.5 C & Si - Atomic Network Solids

Composed of strong directional covalent bonds that are best viewed as a “giant molecule”.

brittle do not conduct heat or electricity carbon, silicon-based graphite, diamond, ceramics, glass

Diamond - hardest natural substance on earth, insulates both heat and electricity.

Graphite - slippery, conducts electricity.

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.

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.

10.6 Molecular solids. S8, P4, CO2, H2O 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. Non-Polar: Large molecules (such as I2 ) can be solids

even without dipoles. Polar: 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.

10.7 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 they

are hard and brittle. Every electron is accounted for so

they are poor conductors - good insulators.

What type of solid will each substance form? #68

Diamond Quartz PH3 NH4NO3

H2 Ar Mg Cu KCl C6H12O6

SF2

10.8 Vapor Pressure Vaporization - change from

liquid to gas at boiling point. Evaporation - change from

liquid to gas below boiling point.

Heat (or Enthalpy) of Vaporization (Hvap) - the

energy required to vaporize 1

mol at 1 atm.

Vaporization is an endothermic process - it requires heat.

Energy is required to overcome intermolecular forces.

Responsible for cool earth. Why we sweat.

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. What the heck is a “dynamic

equilibrium?”

Dynamic equilibrium When first sealed the molecules

gradually escape the surface of the liquid.

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.

Increases with increasing temperature. Decreases with increasing

intermolecular forces. •Bigger molecules (bigger LDF)•More polar molecules (dipole-dipole)

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.

The energy required to accomplish this change is called the Heat (or Enthalpy) of Fusion Hfus

-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

Melting Point Melting point is determined by the

vapor pressure of the solid and the liquid.

At the melting point the vapor pressure of the solid =

vapor pressure of the liquid at 1 atm

Boiling Point Reached when the vapor pressure

equals the pressure of the surrounding atmosphere.

Normal boiling point is the boiling point at 1 atm pressure.

Super heating - Rapid heating above the boiling point allows liquid state to exist above normal boiling point.

Supercooling - Rapid cooling below the freezing point allows liquid state to exist below the normal freezing point.

Practice

10.9 Phase Diagrams A plot of temperature versus pressure for a

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

Critical temperature: temperature above which the vapor can not be liquefied.

Critical pressure: pressure required to liquefy at the critical temperature.

Critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm). Liquid & solid are indistinguishable.

SolidLiquid

Gas

Triple Point

Critical Point

Temperature

Pre

ssur

e

Temperature

SolidLiquid

Gas

1 Atm

AA

BB

CCD

D D

Pre

ssur

e

D

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

Like most substances, bromine exists in one of the three typical

phases. Br2 has a normal melting point of -7.2°C and a normal boiling point of 59°C. The triple point for Br2 is -7.3°C and 40 torr, and the critical point is 320°C and 100 atm. Using this information, sketch a phase diagram for bromine indicating the points described

above. – Based on your phase diagram, order the three phases from

least dense to most dense.

– What is the stable phase of Br2 at room temperature and 1 atm?

– Under what temperature conditions can liquid bromine never exist?

– What phase changes occur as the temperature of a sample of bromine at 0.10 atm is increased from -50°C to 200°C?