Daniel L. RegerScott R. GoodeDavid W. Ball

http://academic.cengage.com/chemistry/reger

Chapter 11Liquids and Solids

• Intermolecular forces are the attractions that hold molecules together in the liquid and solid states.

State Volume Shape of State Density Compressibility

gas assumes shape and volume of container

low easily compressed

liquid definite volume, assumes shape of container

high nearly incompressible

solid both definite shape and volume

high nearly incompressible

Characteristic Properties of Gases, Liquids, and Solids

Kinetic Molecular Theory• What was kinetic molecular theory?• Intermolecular forces:• Do not change with temperature• However, kinetic energy does

• So…

Physical State Relation Between Energy of Attraction and Kinetic Energy of Molecules

solid kinetic energy << energy of attraction

liquid kinetic energy ≈ energy of attraction

gas kinetic energy >> energy of attraction

Example• At room temperature, chlorine is a gas, bromine is a liquid,

and iodine is a solid. Arrange the molecules in order of increasing intermolecular forces.

• Intermolecular forces determine phase, but temperature and pressure can influence phase too.

• Definitions:• Evaporation is the process by which molecules

escape from the liquid to the gas phase.• Condensation is the process by which

molecules go from the gas phase to the liquid phase.

Phase Changes

• The vapor pressure is the partial pressure of the gas when the rate of evaporation equals the rate of condensation.

Vapor Pressure

• A state of dynamic equilibrium is one in which the two opposing changes occur at equal rates, so no net change is apparent.

Dynamic Equilibrium

At constanttemperature

• As temperature increases, the vapor pressure of a liquid increases.• The stronger the intermolecular

forces, the lower the vapor pressure of the liquid at any temperature.

Factors that Affect Vapor Pressure

Vapor Pressure Curves

(a) diethyl ether (b) ethanol (c) water

Which substance has the weakest intermolecular forces?

Boiling Point• The boiling point of a liquid is the

temperature at which the vapor pressure is equal to the external pressure.• The normal boiling point of a liquid is the

temperature at which its equilibrium vapor pressure is equal to 1 atmosphere.• At the boiling point, bubbles filled with

vapor form below the surface of the liquid.

Other Vaporization Properties• Enthalpy of vaporization (Hvap) is the

enthalpy change that accompanies the conversion of one mole of a substance from a liquid to a gas at constant temperature.

• The critical temperature is the maximum temperature at which a substance can exist in the liquid state.

• The critical pressure is the minimum pressure needed to maintain the liquid state up to the critical temperature.

• As the strength of intermolecular forces increase:• vapor pressure of the liquid decreases;• boiling point increases;• enthalpy of vaporization increases;• critical temperature increases.

Vaporization and Intermolecular Forces

Example Calculation• Butane boils at a temperature of -0.6°C and has a

ΔHvap = 22.3 kJ/mol. How much energy is necessary to boil 150 g of butane?

• The changes of a substance from liquid to solid (freezing) and from solid to liquid (melting or fusion) are also opposing changes that involve a dynamic equilibrium.

Liquid-Solid Equilibrium

• The melting point of a substance is the temperature at which the solid and liquid phases are in equilibrium when the pressure is one atmosphere.• There is very little effect of pressure on the

melting point of a solid.

• The enthalpy of fusion (Hfus) is the enthalpy change that accompanies the change of one mole of solid into liquid at constant temperature.

Definitions

Heating Curves• A heating curve is a graph of temperature

of a sample versus heat added.

Only One Phase is Present• When only one phase is present (up to A,

B to C, D and after), then q = mCsT.

• During a phase transition (A to B, C to D), the temperature remains constant and q = H of the transition.

Phase Transitions

• In the heating curve below, identify the phase transition between A and B.

Example: Heating Curve

Example: Heating Curve• From the heating curve below, determine

which phase (solid, liquid, or gas) has the largest specific heat.

Solid-Gas Equilibrium• Sublimation is the direct conversion of a

substance from the solid to the gas phase.• Deposition is the reverse of the sublimation

process.

• Enthalpy of sublimation (Hsub) is the enthalpy change for the conversion of one mole of substance from solid to gas.

• Hsub = Hfus + Hvap

Enthalpy Diagram for Phase Changes

• A phase diagram is a graph of pressure versus temperature that shows the region of stability for each phase.

A Phase Diagram

• There is a unique combination of pressure and temperature, called the triple point (T), at which all three phases (solid, liquid, gas) are at equilibrium.

Triple Point

• The melting point of a substance changes very little with pressure.• The effect of pressure on the melting

point of a substance depends on the relative density of the two phases.

Melting Point and Pressure

• If the solid is denser than the liquid (which is the more common case), the melting point increases with increasing pressure.

• If the liquid is denser than the solid (as in H2O), the melting point decreases with pressure.

Melting Point and Pressure

• Electrostatic forces account for all types of intermolecular attractions. There are three types of attractions:• Dipole-dipole attractions• London dispersion forces• Hydrogen bonding

Intermolecular Attractions

• Dipole-dipole attractions result from electronic forces between molecular dipoles:

Dipole-Dipole Attractions

• London dispersion forces arise from the attractions between instantaneous dipoles and induced dipoles.

London Dispersion Forces

Dispersion Forces and Periodic Trends

• Polarizability is the ease with which a charge distorts the electron cloud in a molecule.• Polarizability generally increases with the

number of electrons in the molecule.

• For related series of molecules, London dispersion forces increase going down any group in the periodic table.

Substance Molar Mass Boiling Point (C)CH4 16 -184

SiH4 32 -112

GeH4 77 -90

SnH4 123 -52

F2 38 -188

Cl2 71 -35

Br2 160 59

I2 254 184

Boiling Points of Some Nonpolar Substances

• The unexpectedly high boiling points of water, ammonia, and hydrogen fluoride requires another kind of intermolecular force.

Hydrogen Bonding

• Hydrogen bonding occurs between a hydrogen atom bonded to N, O, or F, and a lone pair of electrons on a second N, O, or F.• Hydrogen bonds are

sometimes shown as dotted lines.

Hydrogen Bonding

• Hydrogen bonding causes ice to have a lower density than liquid water.

Structure of Solid Water

• Identify the kind of intermolecular forces, and predict which substance in each pair has the stronger forces of attraction.

(a) BF3, BBr3 (b) C2H5OH, C2H5Cl

Example: Intermolecular Forces

• Surface tension is the energy needed to increase the surface area of a liquid.• Surface tension results from intermolecular

interactions.

Liquids: Surface Tension

Liquids: Capillary Action• Capillary action causes water to rise in

a small diameter glass tube.• Capillary action is the result of a

competition between:• cohesion: the attraction of molecules for

other molecules of the same substance.• adhesion: the attraction of molecules for

other molecules of a different substance.

Capillary Action• Water rises because adhesion is stronger

than cohesion.• Mercury is lowered because cohesion is

stronger than adhesion.

• Viscosity is the resistance of a fluid to flow.• The stronger the intermolecular forces of

attraction, the greater the viscosity.• Other factors contribute to viscosity as well,

like structure, size, and shape of molecules.

Liquids: Viscosity

Solids• A crystalline solid: the units that

make up the solid are arranged in a very regular, repeating pattern.• Ionic compounds, metals, and solids of

small molecules are usually crystalline.

• An amorphous solid lacks the long range order of a crystalline solid.• Most plastics are amorphous solids.

(they are polymers)

• Crystalline solids can be classified by the nature of the forces that hold the units together in a regular arrangement.• These forces are usually referred to

as crystal forces.

Crystalline Solids

Molecular Solids• Molecular solids consist of atoms or

small molecules held together by van der Waals forces and/or hydrogen bonding.• Because these crystal forces are fairly

weak, molecular solids are generally soft and low-melting.• Examples are CO, Ar, I2, and most

organic molecules.

• In a covalent network solid, all of the atoms in a crystal are held together by covalent bonds.• Solids of this kind are high melting and

often very hard because strong covalent bonds hold the atoms together.• Some examples of covalent network

solids are diamond (C), boron nitride (BN), and silicon dioxide (SiO2).

Covalent Network Solids

• Allotropes are two or more molecular or crystalline forms of an element in the same physical state.• O2 (oxygen) and O3 (ozone) are

examples of gas-phase allotropes.• Many elements have two or more

allotropes in the solid phase: C, S, P, Sn, among others.

Allotropes

• Graphite and diamond are allotropes of carbon that have different covalent network structures.

Allotropes of Carbon

• An ionic solid consists of oppositely charged ions, held together by strong electrostatic interactions.• Ionic solids are high melting and usually

brittle – they tend to shatter under impact.• Binary compounds made up of a metal

and a nonmetal are in this category.

Ionic Solids

• Metallic solids are formed from metal atoms, and are characterized by high thermal and electrical conductivity, metallic luster, and malleability.• A special kind of bonding, metallic

bonding, is needed to account for these unique properties.

Metallic Solids

• The electron sea model for metallic bonding views the solid as metal ions in a “sea” of electrons formed from the valence shell electrons.• The electrons are very mobile and

adequately account for the conductivity and malleability of metals.• Another model for metallic bonding will

be discussed in Chapter 20.

Metallic Bonding

Properties of Solids - SummaryType of Solid Molecular Covalent

NetworkIonic Metallic

Structural unit

Atoms or molecules

Atoms Ions Atoms

Attractive forces

Intermolecular forces

Localized covalent bonds

Ionic bonds Metallic bonds (delocalized)

Melting points

Low melting, often gases or liquids at room temperature

High melting High melting Variable, from low to very high

Character Soft Hard and brittle Brittle Malleable

Electrical conductivity

Poor Variable, depending on structure

Poor in solid, but good when molten

Very high

The Bragg Equation• The distances between layers of atoms in a

crystal, as measured by x ray diffraction, are given by the Bragg equation:

where = wavelength of x rays, d = distance between layers of atoms, = angle of x ray diffraction, and n is a whole number called the order.

sin2dn

Example: The Bragg Equation• X rays ( = 154 pm) are diffracted by a

crystal at an angle of 18.5. Assuming n = 1, calculate the distance between the layers of atoms that cause this diffraction.

Crystal Structure• The arrangement of the units (atoms or

molecules) is described by the unit cell – a small regular geometric figure that defines the repeating pattern in the crystal.• The location of every particle in the crystal

can be determined from the size and shape of the unit cell.

The Unit Cell

Single unit cell Crystal lattice

The Unit Cell• Each unit cell is defined by the length of

the edges (a, b, and c) and the angles between them (, , and ).

The Cubic Unit Cells• For simplicity, only the three cubit unit

cells are considered (a = b = c, = = = 90).

Crystalline Solids

There are several types of basic arrangements in crystals, like the ones depicted above.

• Calculate the density of nickel, which crystallizes in a face-centered cubic cell with an edge length of 351 pm.

Example: Density from Crystal Data

Closest Packing Structures• Closest packing is the arrangement of

spheres in the most efficient manner, and results in the smallest empty space.• There are two closest packing

arrangements, called hexagonal close packing (HCP) and cubic close packing (CCP).• In both of these arrangements, each atom

has twelve nearest neighbors.

Close Packing Structures

Ionic Crystal Structures• Cations and anions alternate in ionic

crystals, to maximize the attractive interactions and minimize the repulsions.• In an ionic crystal lattice, the composition of

the unit cell must correspond to the formula of the compound. For example, in NaCl the ratio of cations to anions is 1:1; in CaBr2, the ratio of cations to anions is 1:2.

Ionic Unit Cell of NaCl• NaCl has an FCC arrangement of Cl-

ions. What about the Na+?

Example: Ionic Unit Cell

• Calcium fluoride has a unit cell with a face-centered cubic arrangement of the Ca2+ ions. How many F- ions are present in the unit cell?