10 ppt notes liquids and solids.pptx

LIQUIDS AND SOLIDS

Compared to gases, liquids and solids



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



They are held together by intermolecular

INTERMOLECULAR FORCES

bonded to each other.

 Intermolecular refers to the forces between the

molecules.

 The forces that hold one molecule to another

molecule.

 These forces arise from unequal distribution of

the electrons in the molecule and the electrostatic attraction between oppositely charged portions of molecules.

 These forces cause changes of state by causing

INTERMOLECULAR FORCES

 Physical properties such as melting points, boiling

points, vapor pressures, etc. can be attributed to the strength of the intermolecular attractions present between molecules.

 It works like this: the lower the boiling point (or

vapor pressure or melting point), the weaker the intermolecular attractions; the higher the boiling point, the stronger the intermolecular attractions.

 Hydrogen bonds are not true bonds—they are just

INTERMOLECULAR

FORCES

 Strong – covalent bonding and ionic bonding

DIPOLE - DIPOLE

 the force of attraction that enables two polar molecules to

attract one another. Polar molecules are those which have an uneven charge distribution since their dipole moments do not cancel.

 Compounds exhibiting this type of IMF have higher

melting and boiling points than those exhibiting weaker IMFs.

 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

 Weaker with greater distance.

HYDROGEN BONDING

 the force of attraction between the hydrogen atom of one

molecule and an unshared electron pair on F, O, or N of a neighboring molecule (a special case of dipole-dipole).

 Effects boiling point.

 This is the strongest IMF.

 Never confuse hydrogen bonding with a bonded hydrogen.

 Especially strong dipole-dipole forces when H is attached to

F, O, or N. These three because they have high

electronegativity and they are small enough to get close.

 The unique physical properties of water are due to the fact

HYDROGEN BONDING

(a) The Polar Water Molecule

HYDROGEN BONDING

LONDON DISPERSION FORCES

 Also called induced dipole-induced dipole.

 the force of attraction between two non polar molecules

due to the fact that they can form temporary dipoles.

 Nonpolar molecules have no natural attraction for each

other. Without these forces, we could not liquefy covalent gases or solidify covalent liquids.

 These forces are a function of the number of electrons

in a given molecule and how tightly those electrons are held. Non - polar molecules also exert forces on each other. Electrons are not evenly distributed at every instant in time.

 Have an instantaneous, temporary dipole.

LONDON DISPERSION FORCES

(a) An

instantaneous polarization can occur on atom A

(b) Nonpolar

molecules such as H₂ also can develop

LONDON DISPERSION

FORCES

 Weak, short lived.

 Lasts longer at low temperature.

 Eventually molecules are long enough to make liquids.

 More electrons, more polarizable.

 Bigger molecules, higher melting and boiling points.

 Much, much weaker than other forces.

 Also called van der Waal’s forces.

 Since all molecules have electrons, all molecules have

these forces. These forces range from 5 to 40 kJ/mol.

 The strength of this force increases as the number of

PHASES OF MATTER

The phase of a substance is determined by three factors:

 temperature.  pressure.

LIQUIDS

 Many of the properties due to internal attraction

of atoms.

 Surface tension  Beading

 Capillary action  Viscosity

 Stronger intermolecular forces cause each of

SURFACE TENSION

 The resistance to an increase in its surface area

(polar molecules).

 High surface tension indicates strong IMFs.  Molecules are attracted to each OTHER.

 A molecule in the interior of a liquid is attracted

SURFACE TENSION

 Molecules at the

top are only pulled inside.

 Molecules in the

middle are

BEADING

 If a polar substance is placed on a non-polar

CAPILLARY ACTION

 Spontaneous rising of a liquid in a narrow tube.  Adhesive forces between molecule and glass

overcome cohesive forces between molecules themselves.

 The narrower the tube, the more surface area of

glass, the higher the column of water climbs!

 The weight of the column sets the limit for the

height achieved.

CAPILLARY ACTION

 Hg liquid behaves opposite to water. Water has a

higher attraction for glass than itself so its

meniscus is inverted or concave, while Hg has a higher attraction for other Hg molecules! Its

VISCOSITY

 Resistance to flow (molecules with strong

intermolecular forces).

 Increases with molecular complexity (long C

chains get tangled and larger electron clouds are more polarizable due to the presence of

additional electrons) and increased with increasing IMFs.

 Large forces, more viscous.

 Cyclohexane has a lower viscosity than hexane

TYPES OF SOLIDS

Two major types:

 Crystalline - highly regular arrangement of

components in their structure; example – salts, metals

 Amorphous - considerable disorder in their

structure; example - glass

 Interparticle interactions and the ability to pack

REPRESENTATION OF COMPONENTS

OF A CRYSTALLINE SOLID

 Lattice - a three dimensional grid that describes

the locations of the pieces in a crystalline solid and designates the centers of components (atoms, ions, or molecules) that makes up the substance.

 Unit Cell - The smallest repeating unit in of the

lattice.

REPRESENTATION OF COMPONENTS

OF A CRYSTALLINE SOLID

A. Covalent network solid B. Ionic salt crystal lattice

REPRESENTATION OF COMPONENTS

OF A CRYSTALLINE SOLID

 network covalent—carbon in diamond form—

here each molecule is covalently bonded to each neighboring C with a tetrahedral arrangement. Graphite, on the other hand, exists as sheets that slide and is much softer.

 ionic salt crystal lattice - Coulomb’s Law

dictates the strength of the lattice

 ice - notice the “hole” in the hexagonal structure

X-RAY ANALYSIS OF SOLIDS

 X-ray diffraction — A bending or scattering of

light. The beams of light are scattered from a regular array of points in which the spacing

between the components are comparable with the λ of the light. It is due to constructive

interference when the waves of parallel beams

are in phase and to destructive interference when the waves are out of phase.

 The diffraction pattern can be used to determine

TYPES OF CRYSTALLINE

SOLIDS

Ionic Solid

 contains ions at the points of the lattice that

describe the structure of the solid (think NaCl).

 VERY high MP’s.  Hard.

 Ion-Ion Coulombic forces are the strongest of all

attractive forces.

 “IMF” usually implies covalently bonded substances,

TYPES OF CRYSTALLINE

SOLIDS

 Molecular Solid: discrete covalently bonded

molecules at each of its lattice points (sucrose, ice).

 Atomic Solid: atoms of the substance are

THE BOOK DRONES ON

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?

METALLIC SOLIDS

 Metals are characterized by high thermal and

electrical conductivity, malleability, and ductility.

 These properties are explained by the

METALLIC SOLIDS

 a model that uses hard spheres to represent the

atoms of a metal.

 These atoms are packed together and bonded to each

other equally in all directions.

 It will be easiest for you to understand if you can

imagine taking a cubic box and pouring in marbles. The marbles will layer, perhaps directly on top of one another, but perhaps one layer slides into the

“dimple” made by the first layer so that the two

METALLIC SOLIDS

 In the diagram in each layer, a given sphere is surrounded by six

others.

 a) exhibits aba packing—the second layer is like the first, but it

is displaced so that each sphere in the second layer occupies a dimple in the first layer. The spheres in the third layer occupy dimples in the second layer so that the spheres in the third layer lie directly over those in the first layer hence aba. aba has the hexagonal unit cell shown below and the resulting structure is

METALLIC SOLIDS

BONDING MODELS FOR

METALS

 Remember, metals conduct heat and electricity,

are malleable and ductile, and have high melting points.

 These facts indicate that the bonding in most

metals is both strong and non-directional.

 It is difficult to separate metallic atoms, but easy

BONDING MODELS FOR METALS

 Electron Sea Model: A regular array of metals

in a “sea” of electrons. I A & II A metals pictured at left.  

 Band (Molecular Orbital) Model: Electrons

METALLIC BONDING

1s

2s

2p

3s

3p

Filled Molecular Orbitals

Empty Molecular

Orbitals

Filled Molecular Orbitals

Empty Molecular

Orbitals

The 1s, 2s, and 2p electrons are

close to nucleus, so they are not

able to move around.

1s

2s

2p

3s

3p

Filled Molecular Orbitals

Empty Molecular

Orbitals

1s

2s

2p

3s

3p

Magnesium Atoms

Filled Molecular Orbitals

Empty Molecular

Orbitals

1s

2s

2p

3s

3p

Magnesium Atoms

Electrons in these energy level

can travel freely throughout the

crystal.

l

Filled Molecular Orbitals

Empty Molecular

Orbitals

1s

2s

2p

3s

3p

Magnesium Atoms

This makes metals conductors

Malleable because the bonds are

flexible

.

l

BONDING MODELS FOR

METALS

Metal alloys: a substance that has a mixture of elements and has metallic properties

 substitution alloys - in brass, one third of the atoms

in the host copper metal have been replaced by zinc atoms. Sterling silver - 93% silver and 7% copper. Pewter - 85% tin, 7% copper, 6% bismuth and 2% antimony. Plumber’s solder - 95% tin and 5%

antimony.

 interstitial alloy - formed when some of the

interstices (holes) in the closest packed metal

structure are occupied by small atoms. Steel - carbon is in the holes of an iron crystal. There are many

NETWORK COVALENT SOLID

 Composed of strong directional covalent bonds

that are best viewed as a giant molecule.

 Both diamond and graphite are network solids.

 The difference is that diamond bonds with

NETWORK COVALENT SOLID

planet, but when a diamond is “cut” it is actually fractured to make the facets.

CARBON- A SPECIAL ATOMIC

SOLID

There are three types of solid carbon.

1. Amorphous- coal

2. Diamond- hardest natural substance on earth;

colorless; insulates both heat and electricity. hard, colorless and an insulator.

3. Graphite- slippery; black; conducts electricity.

DIAMOND- EACH CARBON IS

SP

HYBRIDIZED, CONNECTED

TO FOUR OTHER CARBONS.

 Carbon atoms are locked

into tetrahedral shape.

 sp3 hybridization and

109.5° bond angles.

 Strong s bonds give the

WHY IS DIAMOND AN

INSULATOR?

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

Empty

MOs

GRAPHITE IS DIFFERENT

 Each carbon is connected

to three other carbons and sp2 hybridized.

 The molecule is flat with

120º angles in fused six member rings.

 The p bonds extend above and below the plane.

 This p bond overlap forms a huge p bonding network

Electrons are free to move throughout these delocalized

orbitals.

SILICON AS A NETWORK

SOLID

 Silicon is to geology what carbon is to biology! The most

significant silicon compounds involve chains with silicon-oxygen bonds.

 silica - empirical formula SiO₂—not at all like its cousin

CO₂. Quartz and some types of sand are silicon dioxide as opposed to a clear colorless gas such as carbon dioxide. Bonding is the difference.

 Silicon cannot use its valence 3p orbitals to form strong π

bonds with oxygen, mainly due to the larger size of the silicon atom and its orbitals—you get inefficient overlap. INSTEAD of forming π bonds, the silicon atom satisfies the octet rule by forming single σ bonds with FOUR

SILICON AS A NETWORK

SOLID

 Each silicon is in the center of a

tetrahedral arrangement of oxygen atoms. This means that although the empirical formula is SiO₂, the

structure is based on a network of SiO₄ tetrahedral with shared oxygen atoms.

 When silica is heated above its MP of

about 1600ºC and cooled rapidly, an amorphous solid forms. We call it glass —it’s really a supercooled, ultra

MOLECULAR SOLIDS

 Molecules occupy the corners of the lattices (not

atoms).

 Different molecules have different forces between

them.

 These forces depend on the size of the molecule

and on the strength and nature of dipole moments.

 Ice and dry ice are examples. Allotropes of sulfur

MOLECULAR SOLIDS

 Characterized by strong covalent bonding within

the molecule yet weak forces between the molecules.

 It takes 6 kJ of energy to melt one mole of solid

water since you only have to overcome hydrogen bonding while it takes 470 kJ of energy to break one mole of O—H bonds.

 Molecules such as CO₂, I₂, P₄, and S₈ have no

dipole moment.

 The melting pt and boiling pt increase since the

MOLECULAR SOLIDS

Those without dipoles:

 Most are gases at 25ºC.

 The only forces are London Dispersion Forces which

are dependent on size of atom.

 Large molecules (such as I₂ ) can be solids even

without dipoles.

 As the size of the molecule increases (radius or molar

MOLECULAR SOLIDS

Those with dipoles:

 Dipole-dipole forces are generally stronger than

London Dispersion forces

 Hydrogen bonding is stronger than Dipole-Dipole

forces.

 No matter how strong the IMF, it is always much,

much weaker than the forces in bonds.

 Stronger forces lead to higher melting and freezing

WATER IS SPECIAL

 Each molecule has two polar O-H bonds.

H

O

H

d

d

-WATER IS SPECIAL

 Each molecule has two polar O-H

bonds.

 Each molecule has two lone pairs of e-

on its oxygen.

 Each oxygen can interact with four

hydrogen atoms.

H

O

H

d

WATER IS SPECIAL

 This gives water an

especially high melting and boiling point.

IONIC SOLIDS

 Stable, high-melting substances held together by

STRONG electrostatic forces that exist between oppositely charged ions.

 Coulomb’s Law dictates the strength of the

electrostatic force.

 The extremes in dipole dipole forces - atoms are

actually held together by opposite charges.

 Huge melting and boiling points.

 Atoms are locked in lattice → hard and brittle.  Every electron is accounted for so they are poor

PRACTICE ONE

Using Table 10.7, classify each of the following substances according to the type of solid it forms.

a. Gold

b. Carbon dioxide

c. Lithium fluoride

VAPOR PRESSURE

 Vaporization - change from liquid to gas at

boiling point

 Evaporation - change from liquid to gas below

boiling point.

 Boiling – change from liquid to gas at the

boiling point - – energy is added.

 ENDOTHERMIC – energy must be

absorbed.

 Heat (or Enthalpy) of Vaporization ( H∆ _vap) –

WATER’S H

∆

 Water’s heat of vaporization is 40.7 kJ/mol. This

is a large value.

 Water makes life on this planet possible since it

acts as a coolant.

 The reason its H_vap is so large has everything to

do with hydrogen bonding. The IMFs in water are huge, thus a great deal of the sun’s energy is

needed to evaporate the rivers, lakes, oceans, etc. of Earth.

 Perspiration is a coolant for animals possessing

CONDENSATION

 EXOTHERMIC

 Achieves a dynamic equilibrium with vaporization in a

DYNAMIC EQUILIBRIUM

a.

When first sealed, the molecules gradually

escape the surface of the liquid.

b.

As the gas molecules build up above the

DYNAMIC EQUILIBRIUM



As time goes by the rate of

DYNAMIC EQUILIBRIUM

Equilibrium is reached when:

Rate of Vaporization = Rate of

Condensation

 Molecules are constantly changing phase:

Dynamic

 The total amount of liquid and vapor remains

VAPOR PRESSURE

 The pressure caused by the gas above the liquid at

equilibrium.

 Liquids with high vapor pressures evaporate easily.

They are called volatile. They have weak IMFs.

 Vapor pressure decreases with increasing IMFs.

VAPOR PRESSURE

 Increases with increasing temperature.  Easily measured in a barometer.

 Heat the particles up, speed the up, move them

out.

 Increasing the temperature increases the KE

which facilitates escape and the speed of the escapees.

 They bang into the sides of the container with

more frequency and more energy.

 More molecules can attain the energy needed to

Dish of Hg

Vacuum

P

=

760 torr

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

Dish of Hg

P

=

760 torr

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

Dish of Hg

736

mm Hg

Water Vapor

 The mercury is pushed down

by the vapor pressure.

 P_atm = P_Hg + P_vap

 P_atm - P_Hg = P_vap

TEMPERATURE EFFECT

Kinetic energy

#

of

m

ol

ec

u

le

s

T

Energy needed to

Kinetic energy

#

of

m

ol

ec

u

le

s

T

Energy needed to

overcome intermolecular

forces

T

T



At a higher temperature more molecules have

enough energy - higher vapor pressure.

Energy needed to

MOLAR MASS AND VAPOR

PRESSURE

 As Molar Mass increases, Vapor Pressure

decreases.

 Because as molecules increase in molar mass,

they also increase in the number of electrons.

 As the number of electrons increase, the

polarizability of the molecule increases so more induced dipole-induced dipole or dispersion forces exist, causing stronger attractions to form

between molecules.

 This decreases the number of molecules that

BIG EXCEPTION

 Hydrogen bonding causes a major exception.  Its presence greatly increases the IMFs of the

liquid.

 Water has an incredibly low VP for such a light

MATHEMATICAL RELATIONSHIP

ln is the natural logarithm

 ln = Log base e

 e = Euler’s number - an irrational number

like p

R - universal gas constant, but use the “energy” R

= 8.3145 J/K mol.

C – a constant characteristic of the liquid. D_{Hvap is the heat of vaporization in J/mol.}

MATHEMATICAL

RELATIONSHIP

PRACTICE TWO

 Using the plots in graph in notes, determine

MATHEMATICAL

RELATIONSHIP

If we know the values of ΔH_vap and VP at one

temperature, we can solve the expression for the

constant, C and set a second expression for T₂ equal to the first since the value of C is NOT dependent upon temperature:

This form is called the Clausius- Clapeyron equation.

MATHEMATICAL

RELATIONSHIP

 The Clausius-Clapeyron equation allows us to

estimate the vapor pressure at another

temperature, if the vapor pressure is known at some temperature, and if the enthalpy of

vaporization is known.

PRACTICE THREE

The vapor pressure of water at 25°C is 23.8 torr, and the heat of vaporization of water at 25°C is

SUBLIMATION

 solids also have vapor pressures.

 Some solids go directly to the vapor phase at

1atm, skipping the liquid phase all together!

 ENDOTHERMIC

 Examples: Iodine and dry ice (solid CO₂)

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

HEATING CURVE FOR WATER

Ice

Water

Vapor

HEAT OF FUSION

D

H

 the enthalpy change that occurs at the melting

point (or freezing point).

 This energy is clearly going into increasing the

PE of the molecules since the temperature or average KE of the molecules has plateaued or is staying the same.

 vapor pressure of solid = vapor pressure of liquid

(EQ is established).

 On the plateaus, calculate the energy change

using q = ΔH_{[vap or fus]}m

 On the slants, calculate the E change using q =

HEATING CURVE FOR WATER

Heat of

Fusion

Heat of

Vaporization

Slope is Heat Capacity

-40 -20 0 20 40 60 80 100 120 140

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.

 Molecules break loose from lattice points and

solid changes to liquid.

 Temperature remains constant during ANY

MELTING POINT

 The melting and boiling

points of water are

determined by the vapor pressures of the solid and liquid states.

 Below zero - VP of ice has a

larger temperature

dependence. This means the VP of ice increases more

rapidly than the liquid’s VP for each increase in

temperature.

 A point is eventually reached

where the VP solid = VP liquid.

 We call this temperature the

melting pt.

“NORMAL”



normal melting point

—the

temperature at which the VP solid = VP

liquid

AND

P total = 1atm



normal boiling point

—the

BOILING POINT

 Reached when the vapor pressure equals the

external pressure.

 Normal boiling point is the boiling point at 1 atm

pressure.

SUPERCOOLING

 the substance is at a temperature below its

freezing point, yet it remains a liquid.

 Usually happens when the cooling has been

gradual and the degree of organization needed to form crystals has not happened.

 At some point, organization happens and the

SUPERHEATED

 the substance is at a temperature above its

boiling pt, yet it remains a liquid.

 Usually happens when heated very rapidly (like

in a microwave oven) and bubbles form in the interior with high internal pressures.

 They often burst before reaching the surface

making quite a mess.

 This is called bumping in the lab. Prevent it by

adding boiling chips to the flask. These chips are porous and have air trapped in them, upon

Solid

Water

Liquid

Water

Water Vapor

Vapor

Solid

Water

Liquid

Water

Water Vapor

Vapor



If the vapor pressure of the solid is higher than

that of the liquid, the solid will release

Solid

Water

Liquid

Water

Water Vapor

Vapor



While the molecules will condense



This can only happen if the temperature is

above the freezing point since the solid is

turning to liquid.

Solid

Water

Liquid

Water



If the vapor pressure of the liquid is higher

than that of the solid, the liquid will release

molecules to achieve equilibrium.

Solid

Water

Liquid

Water

Solid

Water

Liquid

Water

Water Vapor

Vapor



While the molecules condense to a



The temperature must be below the

freezing point since the liquid is turning

to a solid.

Solid

Water

Liquid

Water



If the vapor pressure of the solid and liquid are

equal, the solid and liquid are vaporizing and

condensing at the same rate = melting point.

Solid

Water

Liquid

Water

PHASE DIAGRAMS

 Represent phases of matter as a function of

temperature and pressure.

 Critical temperature: temperature above

which the vapor cannot be liquefied.

 Critical pressure: pressure required to liquefy

AT the critical temperature.

 Critical point: critical temperature and

pressure coordinates (for water, Tc = 374°C and 218 atm).

 Triple point: the point where all three phases

Solid

Liquid

Gas

Triple

Point

Critical

Point

Temperature

Pr

es

su

Solid

Liquid

Gas

 This is the phase diagram for water.

 The density of liquid water is higher than solid

water.

Temperature

Pr

es

su

re

Solid

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

P

re

ss

ur

EFFECT OF PRESSURE

The wire will exert pressure on the block, melt it and begin a journey downward through the block due to the force of gravity acting on the weights. After the wire has left the

REMEMBER!!

 EACH PHASE BOUNDARY REPRESENTS

AN EQUILIBRIUM SET OF PRESSURE ANDTEMPERATURE CONDITIONS!!

 Be sure and use the word equilibrium in your