Mr. Hollister HollidayLegacy High School

Regular & Honors Chemistry

Liquids & Solids

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Liquids

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Properties of the States of Matter:Liquids

• High densities compared to gases.• Fluid.

The material exhibits a smooth, continuous flow as it moves.

• Take the shape of their container(s).

• Keep their volume, do not expand to fill their container(s).

• Cannot be compressed into a smaller volume.

Particles in a Liquid

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Liquids• The particles in a liquid are closely

packed, but they have some ability to move around.

• The close packing results in liquids being incompressible.

• But the ability of the particles to move allows liquids to take the shape of their container and to flow. However, they don’t have enough freedom to escape and expand to fill the container(s).

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Properties of Liquids:Surface Tension

• Liquids tend to minimize their surface—a phenomenon we call surface tension.

• This tendency causes liquids to have a surface that resists penetration.

• The stronger the attractive force between the molecules, the larger the surface tension.

Water Strider

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Properties of Liquids:Viscosity

• Some liquids flow more easily than others.

• The resistance of a liquid’s flow is called viscosity.

• The stronger the attractive forces between the molecules, the more viscous the liquid is.

• Also, the less round the molecule’s shape, the larger the liquid’s viscosity.Some liquids are more viscous because

their molecules are long and get tangled in each other, causing them to resist flowing.

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Evaporation• Over time, liquids evaporate—the molecules of

the liquid mix with and dissolve in the air.• The evaporation happens at the surface.• Molecules on the surface experience a smaller

net attractive force than molecules in the interior.

• All the surface molecules do not escape at once, only the ones with sufficient kinetic energy to overcome the attractions will escape.

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Escaping from the Surface• The process of molecules of a

liquid breaking free from the surface is called evaporation.Also known as vaporization.

• Evaporation is a physical change in which a substance is converted from its liquid form to its gaseous form.The gaseous form is called a

vapor.

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Factors Effecting the Rate of Evaporation

• Liquids that evaporate quickly are called volatile liquids, while those that do not are called nonvolatile.

• Increasing the surface area increases the rate of evaporation.More surface molecules.

• Increasing the temperature increases the rate of evaporation.Raises the average kinetic energy,

resulting in more molecules that can escape.

• Weaker attractive forces between the molecules = faster rate of evaporation.

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Reconnecting with the Surface• When a liquid evaporates in a

closed container, the vapor molecules are trapped.

• The vapor molecules may eventually bump into and stick to the surface of the container or get recaptured by the liquid. This process is called condensation.A physical change in which a

gaseous form is converted to a liquid form.

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Dynamic Equilibrium• Evaporation and

condensation are opposite processes.

• Eventually, the rate of evaporation and rate of condensation in the container will be the same.

• Opposite processes that occur at the same rate in the same system are said to be in dynamic equilibrium.

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Evaporation and Condensation

Eventually, the condensation and

evaporation reach the same speed. The air in the flask is

now saturated with water vapor.

Shortly, the waterstarts to evaporate.

Initially the rateof evaporation is

much faster than rate of condensation

When water is justadded to the flask and it is capped, all the water molecules

are in the liquid.

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Vapor Pressure• Once equilibrium is reached,

from that time forward, the amount of vapor in the container will remain the same.As long as you don’t change

the conditions.• The partial pressure exerted

by the vapor is called the vapor pressure.

• The vapor pressure of a liquid depends on the temperature and strength of intermolecular attractions.

Vapor Pressure Curves

• The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure.

• The normal boiling point is the temperature at which its vapor pressure is 760 torr or 1 atm.

Water: A Unique and Important Substance

• Water is found in all three states on Earth.

• As a liquid, it is the most common solvent found in nature.

• Without water, life as we know it could not exist.The search for extraterrestrial

life starts with the search for water.

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Water

• Liquid at room temperature.Most molecular substances that have a molar

mass (18.02 g/mol) similar to water’s are gaseous.

• Relatively high boiling point.• Expands as it freezes.

Most substances contract as they freeze.Causes ice to be less dense than liquid water.

Solids

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Properties of the States of Matter:Solids

• High densities compared to gases.• Nonfluid.

They move as an entire “block” rather than a smooth, continuous flow.

• Keep their own shape, do not take the shape of their container(s).

• Keep their own volume, do not expand to fill their container(s).

• Cannot be compressed into a smaller volume.

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Solids• Some solids have their particles

arranged in an orderly geometric pattern. We call these crystalline solids.Salt and diamonds.

• Other solids have particles that do not show a regular geometric pattern over a long range. We call these amorphous solids.Plastic and glass.

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Crystalline Solids

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Types of Crystalline Solids

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Molecular Crystalline Solids

• Molecular solids are solids whose composite units are molecules.

• Solid held together by intermolecular attractive forces.Dispersion, dipole-dipole,

or H-bonding.• Generally low melting

points

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Ionic Crystalline Solids• Ionic solids are solids whose

composite units are formula units.• Solid held together by

electrostatic attractive forces between cations and anions.Cations and anions arranged in a

geometric pattern called a crystal lattice to maximize attractions.

• Generally higher melting points than molecular solids.Because ionic bonds are stronger

than intermolecular forces.

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Atomic Crystalline Solids

• Atomic solids are solids whose composite units are individual atoms.

• Solids held together by either covalent bonds, dispersion forces, or metallic bonds.

• Melting points vary depending on the attractive forces between the atoms.

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Practice—Classify Each of the Following Crystalline Solids as Molecular, Ionic, or Atomic.

• H2O(s)

• Si(s)

• C12H22O11(s)

• CaF2(s)

• Sc(NO3)3(s)

• H2O(s)—molecular.

• Si(s)—atomic.

• C12H22O11(s)—molecular.

• CaF2(s)—ionic.

• Sc(NO3)3(s)—ionic.

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Types of Atomic Solids

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Types of Atomic Solids:Covalent Network

• Covalent atomic solids have their atoms attached by covalent bonds. Effectively, the entire solid is one giant molecule.

• Because covalent bonds are strong, these solids have very high melting points.

• Because covalent bonds are directional, these substances tend to be very hard.

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Types of Atomic Solids:Nonbonding

• Nonbonding atomic solids are held together by dispersion forces.

• Because dispersion forces are relatively weak, these solids have very low melting points

• All the noble gases form nonbonding atomic solids.

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Types of Atomic Solids:Metallic

• Metallic solids are held together by metallic bonds.

• Metal atoms release some of their electrons to be shared by all the other atoms in the crystal.

• The metallic bond is the attraction of the metal cations for the mobile electrons.Often described as islands of

cations in a sea of electrons.

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Metallic Bonding• The model of metallic bonding

can be used to explain the properties of metals.

• The luster, malleability, ductility, and electrical and thermal conductivity are all related to the mobility of the electrons in the solid.

• The strength of the metallic bond varies, depending on the charge and size of the cations, so the melting points of metals vary as well.

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Practice—Decide if Each of the Following Atomic Solids Is Covalent, Metallic, or

Nonbonding.

• diamond

• neon

• iron

• diamond covalent.

• neon nonbonding.

• iron metallic.

Phase Changes

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Heating Curve Diagrams

• The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other.

• The temperature of the substance does not rise during the phase change.

Phase DiagramsPhase diagrams display the state of a

substance at various pressures and temperatures.

Phase Diagrams

• The AB line is the liquid-vapor interface.• It starts at the triple point (A), the point at

which all three states are in equilibrium.

Phase DiagramsIt ends at the critical point (B); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

Phase DiagramsEach point along this line is the boiling point of the substance at that pressure.

Phase Diagrams• The AD line is the interface between liquid and

solid.• The melting point at each pressure can be

found along this line.

Phase Diagrams• Below A the substance cannot exist in the

liquid state.• Along the AC line the solid and gas phases are

in equilibrium; the sublimation point at each pressure is along this line.

Phase Diagram of CO2

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Sublimation

• Sublimation is a physical change in which the solid form changes directly to the gaseous form.Without going through the

liquid form.

• Like melting, sublimation is endothermic.

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Intermolecular Forces

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Why Are Molecules Attracted to Each Other?

• Intermolecular attractions are a result of attractive forces between opposite charges.

• + ion to – ion.• + end of one polar molecule to − end of another

polar molecule.H-bonding is especially strong.Even nonpolar molecules will have temporary induced

dipoles.• Larger charge = stronger attraction.

Intermolecular Forces

The attractions between molecules are not nearly as strong as the intramolecular attractions that

hold compounds together.

Intermolecular Forces

They are, however, strong enough to control physical properties such as boiling and melting

points, vapor pressures, and viscosities.

Intermolecular Forces

These intermolecular forces as a group are referred to as van der Waals forces.

van der Waals Forces• London dispersion forces• Dipole-dipole interactions• Hydrogen bonding

Johannes van der Waals

Types of Intermolecular ForcesType of

force

Relative

strength

Present

in Example

London

Dispersion

force

Weak, but

increases

with molar

mass

All atoms

and

molecules

H2

Dipole–

Dipole

force

Moderate

Only

polar

molecules

HCl

Hydrogen

BondStrong

Molecules

having H

bonded to

F, O, or N

HF

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Dispersion Forces• Also known as London forces or

instantaneous dipoles.• Caused by distortions in the electron cloud of

one molecule inducing distortion in the electron cloud on another.

• Distortions in the electron cloud lead to a temporary dipole.

• The temporary dipoles lead to attractions between molecules—dispersion forces.

• All molecules have attractions caused by dispersion forces.

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Instantaneous Dipoles

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Strength of the Dispersion Force

• Depends on how easily the electrons can move, or be polarized.

• The more electrons and the farther they are from the nuclei, the larger the dipole that can be induced.

• Strength of the dispersion force gets larger with larger molecules.

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Noble Gas Molar Mass

(g/mol)

Boiling Point

(K)

He 4.00 4.2

Ne 20.18 27

Ar 39.95 87

Kr 83.80 120

Xe 131.29 165

Dispersion Force and Molar Mass

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-300

-250

-200

-150

-100

-50

0

50

100

150

200

250

1 2 3 4 5 6

Bo

ilin

g P

oin

t, °

C

Period

Relationship Between Dispersion Force and

Molecular Size

BP, Noble Gas

BP, Halogens

BP, XH4

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Practice—The Following Are All Made of Non–Polar Molecules. Pick the Substance in Each

Pair with the Highest Boiling Point.

• CH4 or C3H8

• BF3 or BCl3

• CO2 or CS2

• CH4 or C3H8.

• BF3 or BCl3.

• CO2 or CS2.

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Permanent Dipoles• Because of the kinds of

atoms that are bonded together and their relative positions in the molecule, some molecules have a permanent dipole.Polar molecules.

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Dipole-to-Dipole Attraction

• Polar molecules have a permanent dipole.A + end and a – end.

• The + end of one molecule will be attracted to the – end of another.

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Molar Mass

(g/mol)

Boiling

Point, °C

Dipole

size, D

CH3CH2CH3 44 -42 0

CH3-O-CH3 46 -24 1.3

CH3 - CH=O 44 20.2 2.7

CH3-CN 41 81.6 3.9

Polarity and Dipole-to-Dipole Attraction

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Attractive Forces

+ - + - + - + -

++

++

__

__

Dispersion forces—All molecules.

Dipole-to-dipole forces—Polar molecules.

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Hydrogen Bonding

• Hydrogen atoms bound to a N, O or F atom have strong intermolecular attractions.Unusually high melting and boiling

points.Unusually high solubility in water.

• This kind of attraction is called a hydrogen bond.

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Properties and H-Bonding

Name Formula

Molar

mass

(g/mol)

Structure

Boiling

point,

°C

Melting

point,

°C

Solubility

in water

Ethane C2H6 30.0 -88 -172 Immiscible

Ethanol CH4O 32.0 64.7 -97.8 Miscible

H C

H

H

C H

H

H

H C

H

H

O H

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Intermolecular H-Bonding

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H-Bonds vs. Chemical Bonds

• Hydrogen bonds are not chemical bonds.• Hydrogen bonds are attractive forces

between molecules.• Chemical bonds are attractive forces that

make molecules.

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-200

-150

-100

-50

0

50

100

150

1 2 3 4 5

Bo

ilin

g P

oin

t, °

C

Period

Relationship Between H-Bonding and

Intermolecular Attraction

BP, HX

BP, H2X

BP, H3X

BP, XH4

CH4

NH3

HF

H2O

SiH4

GeH4

SnH4

H2S H2Se

H2Te