prentice hall © 2003chapter 11 chapter 11 intermolecular forces, liquids and solids chemistry the...

Prentice Hall © 2003 Chapter 11

Chapter 11Chapter 11Intermolecular Forces, Intermolecular Forces,

Liquids and SolidsLiquids and Solids

CHEMISTRY The Central Science

9th Edition

David P. White


• Physical properties of substances understood in terms of kinetic molecular theory:– Gases are highly compressible, assumes shape and volume of

container: • Gas molecules are far apart and do not interact much with each

other.

– Liquids are almost incompressible, assume the shape but not the volume of container:

• Liquids molecules are held closer together than gas molecules, but not so rigidly that the molecules cannot slide past each other.

A Molecular Comparison A Molecular Comparison of Liquids and Solidsof Liquids and Solids


– Solids are incompressible and have a definite shape and volume:

• Solid molecules are packed closely together. The molecules are so rigidly packed that they cannot easily slide past each other.



• Converting a gas into a liquid or solid requires the molecules to get closer to each other:– cool or compress.

• Converting a solid into a liquid or gas requires the molecules to move further apart: – heat or reduce pressure.

• The forces holding solids and liquids together are called intermolecular forces.



• The covalent bond holding a molecule together is an intramolecular forces.

• The attraction between molecules is an intermolecular force.

• Intermolecular forces are much weaker than intramolecular forces (e.g. 16 kJ/mol vs. 431 kJ/mol for HCl).

• When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).

Intermolecular ForcesIntermolecular Forces


Ion-Dipole Forces• Interaction between an ion and a dipole (e.g. water).• Strongest of all intermolecular forces.



Dipole-Dipole Forces• Dipole-dipole forces exist between neutral polar

molecules.• Polar molecules need to be close together.• Weaker than ion-dipole forces.• There is a mix of attractive and repulsive dipole-dipole

forces as the molecules tumble.• If two molecules have about the same mass and size, then

dipole-dipole forces increase with increasing polarity.


Dipole-Dipole Forces



Dipole-Dipole Forces


Note that as the dipole moment increases so does the bp


London Dispersion Forces• Weakest of all intermolecular forces.• It is possible for two adjacent neutral molecules to affect

each other.• The nucleus of one molecule (or atom) attracts the

electrons of the adjacent molecule (or atom).• For an instant, the electron clouds become distorted.• In that instant a dipole is formed (called an instantaneous

dipole).



London Dispersion Forces• One instantaneous dipole can induce another

instantaneous dipole in an adjacent molecule (or atom).• The forces between instantaneous dipoles are called

London dispersion forces.



London Dispersion Forces• Polarizability is the ease with which an electron cloud

can be deformed.• The larger the molecule (the greater the number of

electrons) the more polarizable.• London dispersion forces increase as molecular weight

increases.• London dispersion forces exist between all molecules.• London dispersion forces depend on the shape of the

molecule.



London Dispersion Forces• The greater the surface area available for contact, the

greater the dispersion forces.• London dispersion forces between spherical molecules

are lower than between sausage-like molecules.


London Dispersion Forces



Hydrogen Bonding• Special case of dipole-dipole forces.• By experiments: boiling points of compounds with H-F,

H-O, and H-N bonds are abnormally high.• Intermolecular forces are abnormally strong.



Hydrogen Bonding• H-bonding requires H bonded to an electronegative

element (most important for compounds of F, O, and N).– Electrons in the H-X (X = electronegative element) lie much

closer to X than H.

– H has only one electron, so in the H-X bond, the + H presents an almost bare proton to the - X.

– Therefore, H-bonds are strong.


Hydrogen Bonding


Hydrogen Bonding• Hydrogen bonds are responsible for:

– Ice Floating• Solids are usually more closely packed than liquids;

• Therefore, solids are more dense than liquids.

• Ice is ordered with an open structure to optimize H-bonding.

• Therefore, ice is less dense than water.

• In water the H-O bond length is 1.0 Å.

• The O…H hydrogen bond length is 1.8 Å.

• Ice has waters arranged in an open, regular hexagon.

• Each + H points towards a lone pair on O.


Hydrogen Bonding



Viscosity• Viscosity is the resistance of a liquid to flow.• A liquid flows by sliding molecules over each other.• The stronger the intermolecular forces, the higher the

viscosity.

Surface Tension• Bulk molecules (those in the liquid) are equally attracted

to their neighbors.

Some Properties of LiquidsSome Properties of Liquids


Viscosity


Surface Tension


Surface Tension• Surface molecules are only attracted inwards towards the

bulk molecules.– Therefore, surface molecules are packed more closely than bulk

molecules.

• Surface tension is the amount of energy required to increase the surface area of a liquid.

• Cohesive forces bind molecules to each other.• Adhesive forces bind molecules to a surface.



Surface Tension• Meniscus is the shape of the liquid surface.

– If adhesive forces are greater than cohesive forces, the liquid surface is attracted to its container more than the bulk molecules. Therefore, the meniscus is U-shaped (e.g. water in glass).

– If cohesive forces are greater than adhesive forces, the meniscus is curved upwards.

• Capillary Action: When a narrow glass tube is placed in water, the meniscus pulls the water up the tube.



• Surface molecules are only attracted inwards towards the bulk molecules.

• Sublimation: solid gas.• Vaporization: liquid gas.• Melting or fusion: solid liquid.• Deposition: gas solid.• Condensation: gas liquid.• Freezing: liquid solid.

Phase ChangesPhase Changes


Energy Changes Accompanying Phase Changes

• Sublimation: Hsub > 0 (endothermic).

• Vaporization: Hvap > 0 (endothermic).

• Melting or Fusion: Hfus > 0 (endothermic).

• Deposition: Hdep < 0 (exothermic).

• Condensation: Hcon < 0 (exothermic).

• Freezing: Hfre < 0 (exothermic).




• Generally heat of fusion (enthalpy of fusion) is less than heat of vaporization:– it takes more energy to completely separate molecules, than

partially separate them.




• All phase changes are possible under the right conditions.• The sequence

heat solid melt heat liquid boil heat gas

is endothermic.• The sequence

cool gas condense cool liquid freeze cool solid

is exothermic.



Heating Curves• Plot of temperature change versus heat added is a heating

curve.• During a phase change, adding heat causes no

temperature change.– These points are used to calculate Hfus and Hvap.

• Supercooling: When a liquid is cooled below its melting point and it still remains a liquid.

• Achieved by keeping the temperature low and increasing kinetic energy to break intermolecular forces.



Critical Temperature and Pressure• Gases liquefied by increasing pressure at some

temperature.• Critical temperature: the highest temperature for

liquefaction of a gas using pressure.• Critical pressure: pressure required for liquefaction.



Critical Temperature and Pressure



Explaining Vapor Pressure on the Molecular Level

• Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the bulk liquid.

• These molecules move into the gas phase.• As the number of molecules in the gas phase increases,

some of the gas phase molecules strike the surface and return to the liquid.

• After some time the pressure of the gas will be constant at the vapor pressure.

Vapor PressureVapor Pressure



• Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface.

• Vapor pressure is the pressure exerted when the liquid and vapor are in dynamic equilibrium.

Volatility, Vapor Pressure, and Temperature• If equilibrium is never established then the liquid

evaporates.• Volatile substances evaporate rapidly.



Volatility, Vapor Pressure, and Temperature

• The higher the temperature, the higher the average kinetic energy, the faster the liquid evaporates.


Volatility, Vapor Pressure, and Temperature



Vapor Pressure and Boiling Point• Liquids boil when the external pressure equals the vapor

pressure.• Temperature of boiling point increases as pressure

increases.



Vapor Pressure and Boiling Point• Two ways to get a liquid to boil: increase temperature or

decrease pressure.– Pressure cookers operate at high pressure. At high pressure the

boiling point of water is higher than at 1 atm. Therefore, there is a higher temperature at which the food is cooked, reducing the cooking time required.

• Normal boiling point is the boiling point at 760 mmHg (1 atm).



• Phase diagram: plot of pressure vs. Temperature summarizing all equilibria between phases.

• Given a temperature and pressure, phase diagrams tell us which phase will exist.

• Any temperature and pressure combination not on a curve represents a single phase.

Phase DiagramsPhase Diagrams


• Features of a phase diagram:– Triple point: temperature and pressure at which all three phases

are in equilibrium.

– Vapor-pressure curve: generally as pressure increases, temperature increases.

– Critical point: critical temperature and pressure for the gas.

– Melting point curve: as pressure increases, the solid phase is favored if the solid is more dense than the liquid.

– Normal melting point: melting point at 1 atm.



The Phase Diagrams of H2O and CO2




• Water:– The melting point curve slopes to the left because ice is less

dense than water.

– Triple point occurs at 0.0098C and 4.58 mmHg.

– Normal melting (freezing) point is 0C.

– Normal boiling point is 100C.

– Critical point is 374C and 218 atm.




• Carbon Dioxide:– Triple point occurs at -56.4C and 5.11 atm.

– Normal sublimation point is -78.5C. (At 1 atm CO2 sublimes it does not melt.)

– Critical point occurs at 31.1C and 73 atm.



Unit Cells• Crystalline solid: well-ordered, definite arrangements of

molecules, atoms or ions. • Crystals have an ordered, repeated structure.• The smallest repeating unit in a crystal is a unit cell.• Unit cell is the smallest unit with all the symmetry of the

entire crystal.• Three-dimensional stacking of unit cells is the crystal

lattice.

Structures of SolidsStructures of Solids

Unit Cells



Unit Cells• Three common types of unit cell.

– Primitive cubic, atoms at the corners of a simple cube,• each atom shared by 8 unit cells;

– Body-centered cubic (bcc), atoms at the corners of a cube plus one in the center of the body of the cube,

• corner atoms shared by 8 unit cells, center atom completely enclosed in one unit cell;

– Face-centered cubic (fcc), atoms at the corners of a cube plus one atom in the center of each face of the cube,

• corner atoms shared by 8 unit cells, face atoms shared by 2 unit cells.


Unit Cells


The Crystal Structure of Sodium Chloride

• Two equivalent ways of defining unit cell:– Cl- (larger) ions at the corners of the cell, or

– Na+ (smaller) ions at the corners of the cell.

• The cation to anion ratio in a unit cell is the same for the crystal. In NaCl each unit cell contains same number of Na+ and Cl- ions.

• Note the unit cell for CaCl2 needs twice as many Cl- ions as Ca2+ ions.



The Crystal Structure of Sodium Chloride



Close Packing of Spheres• Solids have maximum intermolecular forces.• Molecules can be modeled by spheres.• Atoms and ions are spheres.• Molecular crystals are formed by close packing of the

molecules.• We rationalize maximum intermolecular force in a crystal

by the close packing of spheres.



Close Packing of Spheres• When spheres are packed as closely as possible, there are

small spaces between adjacent spheres.• The spaces are called interstitial holes.• A crystal is built up by placing close packed layers of

spheres on top of each other.• There is only one place for the second layer of spheres.



Close Packing of Spheres• There are two choices for the third layer of spheres:

– Third layer eclipses the first (ABAB arrangement). This is called hexagonal close packing (hcp);

– Third layer is in a different position relative to the first (ABCABC arrangement). This is called cubic close packing (ccp).



Close Packing of Spheres



Close Packing of Spheres• Each sphere is surrounded by 12 other spheres (6 in one

plane, 3 above and 3 below).• Coordination number: the number of spheres directly

surrounding a central sphere.• Hexagonal and cubic close packing are different from the

cubic unit cells.• If unequally sized spheres are used, the smaller spheres

are placed in the interstitial holes.



• There are four types of solid:– Molecular (formed from molecules) - usually soft with low

melting points and poor conductivity.

– Covalent network (formed from atoms) - very hard with very high melting points and poor conductivity.

– Ions (formed form ions) - hard, brittle, high melting points and poor conductivity.

– Metallic (formed from metal atoms) - soft or hard, high melting points, good conductivity, malleable and ductile.

Bonding in SolidsBonding in Solids


Molecular Solids• Intermolecular forces: dipole-dipole, London dispersion

and H-bonds.• Weak intermolecular forces give rise to low melting

points.• Room temperature gases and liquids usually form

molecular solids and low temperature.• Efficient packing of molecules is important (since they

are not regular spheres).



Covalent-Network Solids

forces between particles: covalent bonds.• Atoms held together in large networks.

• Examples: diamond, graphite, quartz (SiO2), silicon carbide (SiC), and boron nitride (BN).

• In diamond: – each C atom has a coordination number of 4; each C atom is

tetrahedral; there is a three-dimensional array of atoms.

– Diamond is hard, and has a high melting point (3550 C).



Covalent-Network Solids



Covalent-Network Solids• In graphite

– each C atom is arranged in a planar hexagonal ring;

– layers of interconnected rings are placed on top of each other;

– the distance between C atoms is close to benzene (1.42 Å vs. 1.395 Å in benzene);

– the distance between layers is large (3.41 Å);

– electrons move in delocalized orbitals (good conductor).



Ionic Solids• Ions (spherical) held together by electrostatic forces of

attraction.• There are some simple classifications for ionic lattice

types.


Ionic Solids


Ionic Solids• NaCl Structure

• Each ion has a coordination number of 6.

• Face-centered cubic lattice.

• Cation to anion ratio is 1:1.

• Examples: LiF, KCl, AgCl and CaO.• CsCl Structure

• Cs+ has a coordination number of 8.

• Different from the NaCl structure (Cs+ is larger than Na+).

• Cation to anion ratio is 1:1.



Ionic Solids• Zinc Blende Structure

• Typical example ZnS.

• S2- ions adopt a fcc arrangement.

• Zn2+ ions have a coordination number of 4.

• The S2- ions are placed in a tetrahedron around the Zn2+ ions.

• Example: CuCl.



Ionic Solids• Fluorite Structure

• Typical example CaF2.

• Ca2+ ions in a fcc arrangement.

• There are twice as many F- per Ca2+ ions in each unit cell.

• Examples: BaCl2, PbF2.



Metallic Solids• Metallic solids have metal atoms in hcp, fcc or bcc

arrangements.• Coordination number for each atom is either 8 or 12.• Problem: the bonding is too strong for London dispersion

and there are not enough electrons for covalent bonds.• Resolution: the metal nuclei float in a sea of electrons.• Metals conduct because the electrons are delocalized and

are mobile.


prentice hall © 2003chapter 11 chapter 11 intermolecular forces, liquids and solids chemistry the...

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