21

11.1: A Molecular Comparison of Liquids 11.1: A Molecular Comparison of Liquids and Solidsand Solids

31The forces holding solids and liquids together are called

intermolecular forces.

41

• The covalent bond holding a molecule together is an intramolecular forces• The attraction between molecules is an

intermolecular force• Much weaker than intramolecular forces • Melting or boiling: the intermolecular forces are

broken (not the covalent bonds)

11.2: Intermolecular 11.2: Intermolecular ForcesForces

51

The stronger the attractive forces, the higher the boiling point of the liquid and the melting point of a solid

(low boiling point)

61

Ion-Dipole Forces• Interaction between an ion and a dipole (a polar molecule

such as water)• Strongest of all intermolecular forces (MIXTURES

ONLY!)

71

Dipole-Dipole Forces

• Between neutral polar molecules• Oppositely charged ends of molecules attract• Weaker than ion-dipole forces

• Dipole-dipole forces increase with increasing polarity• Strength of attractive forces is inversely related to

molecular volume

81

London Dispersion Forces• Weakest of all intermolecular forces• Two adjacent neutral, nonpolar molecules

• The nucleus of one attracts the electrons of the other• Electron clouds are distorted• Instantaneous dipole• Strength of forces is directly related to molecular

weight• London dispersion forces exist between all molecules

Examples: 11.11 and 11.13

London dispersion forces depend on the shape of the molecule

• The greater the surface area available for contact, the greater the dispersion forces

101

Hydrogen Bonding

• Special case of dipole-dipole forces

• H-bonding requires H bonded to an electronegative element (most important for compounds of F, O, and N)

Hydrogen Bonding

Boiling point increases with increasing molecular weight. The exception is water (H bonding)

Text, P. 413

Hydrogen Bonding

Solids are usually more closely packed than liquids (solids are more dense than liquids)

Ice is ordered with an open structure to optimize H-bonding (ice is less dense than water)

Text, P. 417; example 11.9

161

Viscosity• Viscosity is the resistance of a liquid to flow

• Molecules slide over each other• The stronger the intermolecular forces, the higher the

viscosity• Viscosity increases with an increase in molecular weight

11.3: Some Properties of 11.3: Some Properties of LiquidsLiquids

171

Surface Tension• Surface molecules are only attracted inwards towards the

bulk molecules • Molecules within the liquid are all equally attracted to

each other

181

• Surface tension is the amount of energy required to increase the surface area of a liquid– Cohesive forces bind molecules to each other (Hg)

– Adhesive forces bind molecules to a surface (H2O)

– If adhesive forces > cohesive forces, the meniscus is U-shaped (water in glass)

– If cohesive forces > adhesive forces, the meniscus is curved downwards (Hg in barometer)

11.4: Phase Changes11.4: Phase Changes

Text, P. 420

(Endothermic)

(Endothermic) (Exothermic)

(Exothermic)(Endothermic)

(Exothermic)

Generally heat of fusion (melting) is less than heat of vaporization (evaporation): it takes more energy to completely

separate molecules, than to partially separate them

Text, P. 420

211

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

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

change (equilibrium)– These points are used to calculate Hfus and Hvap

221

Text, P. 421

Added heat increases the temperature of a consistent state of matterEnergy used for

changing molecular motion, no T change

231

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

temperature• Critical temperature: the minimum temperature for

liquefaction of a gas using pressure• A high C.T. means strong intermolecular forces

• Critical pressure: pressure required for liquefaction

241

• Examples: # 31, 33, WDP # 48

• Other WDP examples: # 44, 46, 50

251

Explaining Vapor Pressure on the Molecular Level

• Some of the molecules on the surface of a liquid have enough energy to escape to the gas phase• After some time the pressure of the gas will be

constant at the vapor pressure (equilibrium)

11.5: Vapor Pressure11.5: Vapor Pressure

261

• 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 (high VP) evaporate rapidly• The higher the T, the higher the average KE, the faster

the liquid evaporates (hot water evaporates faster than cold water)

271

• Vapor pressure increases nonlinearly with increasing temperature

• Clausius-Clapeyron Equation

281

When Temperature changes from T1 to T2, Pressure changes from P1 to P2

• Use the Clausius-Clapeyron Equation to

1. Predict the vapor pressure at a specified temperature

2. Determine the T at which a liquid has a specified VP

3. Calculate enthalpy of vaporization from measurements of VP’s at different temperatures

291

• Sample problems: # 45, WDP # 35

• Other WDP examples: # 36 & 37

301

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

pressure• Normal BP: BP of a liquid at 1atmosphere

• Temperature of boiling point increases as pressure increases

311

• 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

11.6: Phase Diagrams11.6: Phase Diagrams

321

Text, P. 428

Vapor Pressure curve of the liquid (increase

P, increase T)

Stable at low P and high T

Stable at low T and high P

Triple Point: all 3 phases in equilibrium

Beyond this point, liquid and gas phases are indistinguishable

Melting point curve: Increased P favors solid phase; Higher T

needed to melt the solid at higher P

331

The Phase Diagrams

of H2O and CO2

Text, P. 429; Question 49 on P. 444

Line slopes to the left: ice is less dense than water (why?) MP decreases with increased P

341

• Sample Problem: BLBB #51

351

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

molecules, atoms or ions• The smallest repeating unit in a crystal is a unit cell• It has all the symmetry of the entire crystal• Three-dimensional stacking of unit cells is the crystal

lattice• Close-packed structure

11.7: Structures of 11.7: Structures of SolidsSolids

361

Unit Cells

Unit Cells

Primitive cubic: atoms at the corners of a simple cube

• each atom shared by 8 unit cells

391

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 1 unit cell

401

Unit Cells

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

Text, P. 432

2 atoms per cell

4 atoms per cell

1 atom per cell

421

The Crystal Structure of Sodium Chloride

Two equivalent ways of defining unit cell:

Cl- (larger) ions at the corners of the cell, orNa+ (smaller) ions at the corners of the cell

43111.8: Bonding in Solids11.8: Bonding in Solids

Text, P. 435

441

Covalent-Network Solids

Ionic SolidsThe structure adopted

depends on the charges and sizes of the ions

461

Metallic Solids• Various arrangements are possible• The bonding is too strong for London dispersion and

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

and are mobile• Close-packed structure

471

• Amorphous solids (rubber, glass) have no orderly structure– IMFs vary in strength throughout the sample

– No specific melting point

Sample Problems # 53, 69, 71, 73, 75