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Chapter 10The Liquid and Solid States

• Changes of State

• Intermolecular Forces

• Properties of Liquids

• Properties of Solids

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Particles are widely

separated with much

free space

Some free space between particles

Little free space between particles

Compressed under moderate pressure

Can be compressed

slightly by moderate pressure

Little or no volume

change under moderate pressure

Particles move through space

Particles move past one another

Particles fixed in place

but vibrate around a fixed position

Fills containerCan be pouredShape remains rigid

Takes shape of container

Takes shape of filled portion of container

Shape not set by container

No fixed shapeNo fixed shapeFixed shape

GasLiquidSolid

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Solids, Liquids, and Gases

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Figure 10.4

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Changes of State• Transitions between these states

• Also called phase changes

• 6 main phase changes:

– Evaporation (also called vaporization)

– Condensation

– Freezing

– Melting (also called fusion)

– Sublimation

– Deposition 10-

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Liquid-Gas Phase Changes

• Evaporation (vaporization)– At the surface of a

liquid, some molecules may have sufficient kinetic energy to escape into the gas state.

– Heat is required to maintain the temperature needed for evaporation.• Evaporation is an

endothermic process.

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Liquid-Gas Phase Changes

• Condensation– Transition from a gas to a liquid

– Occurs when gas particles cannot escape the container and thus, come into contact with a liquid

– An exothermic process (energy is released)

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Figure 10.9 (cropped for only the top portion)

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Equilibrium• A state in which opposing processes occur at

equal rates• An equilibrium is designated by a double arrow,

such as:

liquid gas

• In the above equilibrium, the rates of evaporation and condensation are equal.

• The gas produced by evaporation exerts a pressure on the liquid below it.– At equilibrium, this pressure is called vapor

pressure.• Vapor pressure increases as temperature

increases.10-

vaporization

condensation

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Boiling Point• The temperature at

which boiling occurs

– Boiling occurs when

the vapor pressure equals the external pressure of the

atmosphere

• Normal boiling point

– Occurs when the atmospheric pressure is exactly 1 atm

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Practice – Vapor Pressure Curves

• Consider the vapor pressure curve for propane to the left, which is used as a fuel in barbeque grills.

What is the normal boiling point for propane? What is its boiling point at 0.40 atm?

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Practice Solutions – Vapor Pressure Curves

• What is the normal boiling point for propane?

The normal boiling point is found at a pressure of 1.0 atm. At this pressure, the boiling is about -41°C.

• What is its boiling point at 0.40 atm?

At 0.40 atm, the boiling point is about -60°C.

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Liquid-Solid Phase Changes• Freezing

– The average kinetic energy of a liquid molecules fall when it is cooled

– If the KEav falls low enough, then the molecules become fixed in position into the solid

– Therefore, freezing is the conversion of a

liquid into a solid

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Liquid-Solid Phase Changes• Freezing point

– The temperature at which the freezing of a liquid into a solid state occurs

• As with boiling temperature, the two states are in equilibrium with one another.

solid liquid

– Normal freezing point

• The temperature when the freezing equilibrium is achieved under a pressure of 1 atm.

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melting

freezing

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Freezing

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Figure 10.11

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Melting• Melting

– Phase change from solid to liquid

– Reverse of freezing

– Opposite process of freezing

– Also called fusion

• Melting point– Same temperature

as freezing point

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Figure 10.12

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Solid-Gas Phase Changes• Sublimation

– Evaporation of a solid

– Occurs when a solid has a high vapor pressure

– Solid can change directly from a solid to a gaseous state without going through the liquid state

• Deposition– Can go directly from gas

to solid without passing through the liquid state

– Reverse of sublimation

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Practice – Changes in the State of Matter

• Identify the process shown in the

following diagram:

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Diagram from Practice Problem 10.2

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Practice Solutions – Changes in the State of Matter

• In the starting material, there is a large distance between particles, so we can identify it as a gas. In the ending material, the particles are closely spaced and not random, and thus it is a solid. The gas to solid transition is called deposition.

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Cooling Curve• Shows the phase changes of a substance in a

graph plotting temperature versus heat

removed at constant pressure

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Heating Curve• Shows the phase changes of a substance in a

graph plotting temperature versus heat added

at constant pressure

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Practice – Cooling and Heating Curves

• The diagram to the top right represents the physical states of a substance.

Identify where each state would predominate on the heating curve given to the right.

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Practice Solutions – Cooling and Heating Curves

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Energy Changes• Each change of state takes place at

constant temperature• Any phase change is accompanied by an

energy change– The change in energy for each process is

called the heat of that process (q)

q = mC∆T– Molar heat of fusion

• Energy required to melt 1 mole of a substance

– Molar heat of vaporization• Energy required to evaporate 1 mole of a

substance10-

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Practice – Energy for Phase

Changes• Calculate the heat required to boil

125 g of water at 100.0°C. The molar

heat of vaporization of water is

4.07 x 104 J/mol.

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Practice Solutions – Energy for

Phase ChangesBecause the heat of vaporization is in J/mol,

and since J is the unit for heat, we only need to cancel out moles. To do this, we need to

convert grams of water into moles of water using water’s molar mass, then multiply by the heat of vaporization to find Joules:

125 g H2O x 1 mol H2O x 4.07 x 104 J

18.01 g H2O 1 mol H2O

= 2.82 x 105 J

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Intermolecular Forces• An attractive force that operates between

molecules

• There are many kinds of intermolecular forces:

– London dispersion force

– Dipole-dipole force

– Hydrogen-bonding force

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Polar molecules

containing unpaired

molecules and a hydrogen bonded to

nitrogen, oxygen, or

fluorine

Two dipoles, one containing hydrogen to

an electronegative

element and the other containing an

electronegative element,

attract one another.

Hydrogen-Bonding Force

Polar moleculesPolar molecules

(permanent dipoles)

attract one another

Dipole-Dipole Force

All atoms and molecules

A temporary dipole in one molecule induces

the formation of a

temporary dipole in a nearby molecule and is

attracted to it.

London dispersion

force

OccurrenceType of

Interaction

Type of

Force

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London Dispersion Forces• Instantaneous dipole

– A temporary dipole formed when the electrons in an atom or nonpolar molecule happen to be more on one side in an instant in time,

causing it to be more negative than normal and the opposite side positive

• Induced dipole

– Positive end of the dipole exerts an attractive force on nearby electrons, causing an adjacent atom to develop into another

temporary dipole

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London Dispersion Forces• London dispersion force

– The attraction between temporary dipoles

– Occurs between atoms and molecules

– Only intermolecular force in nonpolar substances

– Tend to be stronger the larger the atom or molecule

– Relatively weak forces

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Practice – London Dispersion

Forces• Which has stronger London

dispersion forces, CH4 or SiH4?

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Practice – London Dispersion

Forces• Which has stronger London

dispersion forces, CH4 or SiH4?

Since silicon is larger than carbon,

SiH4 interactions are stronger than

CH4 interactions. The larger the

electron cloud, the easier it is to

distort it, resulting in larger London

dispersion forces.

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Dipole-Dipole Forces

• Attraction between polar molecules

• Occurs when the partially positive end of one molecule attracts the partially negative end of another molecule

• Generally stronger than London dispersion forces

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Practice – Dipole-Dipole Forces

• Which of the following molecules

experience dipole-dipole forces?

a) SCl2b) CO

c) NH3

d) CCl4

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Practice Solutions –

Dipole-Dipole Forces

• Which of the following molecules experience dipole-dipole forces?a) SCl2

SCl2 has a bent structure similar to water’s structure, and therefore, an overall molecular dipole towards the lone pairs on S. This molecule experiences dipole-dipole forces.

b) COCO has a linear structure with one lone pair on the C and O. Since O is more electronegative than C, this molecule has a molecular dipole towards the O. This molecule experiences dipole-dipole forces.

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Practice Solutions –

Dipole-Dipole Forces• Which of the following molecules

experience dipole-dipole forces?c) NH3

NH3 has a trigonal pyramidal structure, and therefore, an overall molecular dipole towards the lone pair on N. This molecule experiences dipole-dipole forces.

d) CCl4CCl4 has a tetrahedral structure. Since all the atoms around central atom are the same, CCl4 is a nonpolar molecule. This molecule does NOT experience dipole-dipole forces.

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Boiling Points

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Hydrogen Bonding• Special type of dipole-

dipole force

• Only occurs in molecules that contain hydrogen bonded to a small, highly electronegative element

• Stronger than a regular dipole-dipole force

• Important force in living systems by stabilizing

molecular shapes

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Hydrogen Bonding

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Figure 10.21

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Practice – Hydrogen Bonding

• Identify the molecules from the

following list that experience

hydrogen bonding in the pure liquid

state: N(CH3)3, CH3CO2H, and HOCl.

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Practice Solutions – Hydrogen Bonding

• Identify the molecules from the following list that experience hydrogen bonding in

the pure liquid state: N(CH3)3, CH3CO2H, and HOCl.

CH3CO2H and HOCl experience hydrogen bonding because the hydrogen atoms in

these molecules are bonded to oxygen, a small, highly electronegative element.

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Practice Solutions – Hydrogen Bonding

• Identify the molecules from the following list that experience hydrogen bonding in the pure liquid state: N(CH3)3, CH3CO2H, and HOCl.

In N(CH3)3 the hydrogen atoms are bonded to carbon, which has a relatively similar electronegativity. Therefore, N(CH3)3 does not experience hydrogen bonding.

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Trends in Intermolecular Forces• Remember:

– London forces exist between all atoms and molecules.

– Dipole-dipole moments exist in only polar compounds.

– Hydrogen bonds only exist in polar compounds that contain hydrogen.

• In terms of strength (magnitude), intermolecular forces compare as shown below:

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Figure 10.23

Molecules

experiencing

London dispersion forces

only

Molecules also

experiencing dipole-dipole

forces

Molecules also

experiencing hydrogen

bonding

< <

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Practice – Trends in Intermolecular Forces

• Consider the substance CO and HF.

Which has the stronger

intermolecular forces? Which has

the higher melting point? Boiling

point? Vapor pressure?

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Practice Solutions – Trends in Intermolecular Forces

• Consider the substance CO and HF. Which has the stronger intermolecular forces? Which has the higher melting point? Boiling point?

Vapor pressure?

HF has the stronger intermolecular forces because it has hydrogen bonding, whereas CO

only experiences dipole-dipole forces. Higher temperatures are required to overcome the stronger forces in HF, therefore it will have the

higher melting and boiling points. 10-

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Practice Solutions – Trends in Intermolecular Forces

• Consider the substance CO and HF. Which has the stronger intermolecular forces? Which has the higher melting point?

Boiling point? Vapor pressure?

Because of the lower energy required for evaporation, at any temperature CO will

have a higher vapor pressure than HF.

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Properties of Liquids• Are related to the distance

between particles and to intermolecular forces

• Particles in a liquid are much closer together than the particles in a gas

• Liquid particles are not fixed, as they are in a solid

• Three common properties of liquids: – Density– Viscosity– Surface tension

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Density• Remember:

• Densities of the states of matter are related to the distance between particles

• Most substances are denser as solids

than as liquids because their molecules or atoms are closer together

– Water is an exception in that ice is less dense than liquid water

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volume

mass d =

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Viscosity• Liquids and gases are fluids.

– A fluid is any substance that can flow

• Viscosity is the resistance of a

substance to flow.

– Generally, the viscosity of liquids is low

• Viscosities generally vary by

increasing with the magnitude of

their intermolecular forces and with

molecular size.10-

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Surface Tension• The amount of work

required to increase the surface area of a liquid by a unit amount

• Causes a liquid surface to behave like a stretched membrane

• The greater the intermolecular forces in a liquid, the greater the surface tension

• Surface tension decreases as temperature increases

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Figure 10.27

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Meniscus• Either a concave or convex curved surface of a liquid

produced by intermolecular forces• A concave surface occurs when the intermolecular

forces between the liquid and the glass are greater than the intermolecular forces among the liquid molecules

• A convex surface occurs when the intermolecular forces between the liquid and the glass are less than the intermolecular forces among the liquid molecules

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Amorphous and Crystalline Solids• Amorphous solid

– Occurs when the temperature of a liquid drops rapidly, resulting in the particles solidifying in a partially disordered state

– Particles are somewhat randomly arranged– Lacks regular form

• Crystalline solid– Occurs when a liquid solidifies slowly,

allowing the array of particles to become well ordered

– Most solids are of this type– Results in the symmetrical arrangement of

planar faces

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Amorphous and Crystalline Solids

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Figure 10.31

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Crystals and Crystal Lattices• Crystal

– An orderly, repeating, three-dimensional assembly of fundamental particles (atoms,

molecules, or ions)

• Crystal lattice

– Pattern forms by repeating arrays of crystal

– Also called crystal structure

– Atoms, molecules, or ions are arranged in close-packed structures

• Orderly arrangement that is an efficient way to use space

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Other Close-Packed Arrangements

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Types of Crystalline Solids• Solids show a wide range of properties such

as:– Melting point

– Hardness

– Malleability

• The types of forces holding the fundamental particles together help explain the properties

• 4 Types of crystalline solids– Metallic solids

– Ionic solids

– Molecular solids

– Network solids

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Low melting point; soft; poor heat conductors; electrical

insulators

London dispersion

forces

Nonpolar

molecules

Very high melting point; very hard; somewhat brittle; non- or

semiconductors

Covalent

bondsAtomsNetwork

Low to moderate melting point; variable hardness; may

be brittle; nonconductors

Dipole-dipole

forces

Polar

molecules

Molecular

High melting point; hard,

brittle; nonconductors when

solid; electrical conductors when melted

Ionic bondsCations and

anionsIonic

Low melting point & soft; or high melting point & hard;

good heat & electrical conductors; malleable &

ductile

Attractions

between nuclei and delocalized

electrons

AtomsMetallic

Fundamental

Particles

Attractive

ForcesProperties

Type of

Solid

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Practice – Types of Crystalline Solids

• A crystalline solid is very hard, does

not conduct electricity when solid or

when melted, and has a high melting

point. What type of solid is it?

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Practice Solutions – Types of Crystalline Solids

• A crystalline solid is very hard, does not conduct electricity when solid or when

melted, and has a high melting point. What type of solid is it?

This solid is a network solid. Network

solids are the only crystalline solids that have high melting points, are very hard,

and have the possibility of

nonconductors.

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Metallic Solids• Valence electrons

move freely through all parts of a metal

• Attractions between atoms of a metal are delocalized, and therefore it is easy to move atoms by applying force

• Ductile and malleable

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Alloys• Forms when a metal is mixed with one or

more additional metallic or nonmetallic elements

• Have properties different than those of their parent elements

• Table 10.4

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Ionic Solids• Contain cations and

anions arranged in crystalline solids

• Electrostatic forces (ionic bonds) hold together ionic crystals

• High melting points• Hard and brittle• Solids are not electrical

conductors, but melted or dissolved, they become good conductors

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Superconductors• Offer no resistance to the conduction of electrical

current

• Repel magnetic fields

• Although the nature of superconductors is not entirely understood, the crystal structure of the solid material does play a role

• Often formed from silicon and/or other metalloids

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Molecular Solids• Intermolecular forces between molecules hold

a molecular solid together

– Nonpolar solids are held together by London dispersion forces

• Form soft crystals with low melting points

• Electrical insulators

• Poor conductors of heat

– Polar solids are held together by London dispersion forces and dipole-dipole forces

• Typically harder than nonpolar solids

• Low to moderate melting points

• Electrical insulators (no ions)

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Molecular Solids

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Network Solids• Consists of a giant

molecule that forms the entire crystal

• Formed by metalloids or carbon

• Strong covalent bonds connect the atoms in a network solid

• Poor electrical conductors• High melting points• Very hard• Have very stable, three–

dimensional structures• Many have a diamond

structure, or some derivative of it

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Network Solids

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Practice – Structures of Solids• Identify the type of solid shown by

the following molecular-level image.

What types of forces hold the

particles together in this solid?

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Practice Solutions – Structures of Solids

• This diagram shows a molecular solid.

Since it is made of nitrogen atoms exclusively, it is nonpolar and is therefore

held together by weak London forces.