Lecture Notes Chem 150 K Marr Chapter 12

Lecture Notes Chem 150 - K. Marr Chapter 12 Intermolecular Attractions & the Properties of Liquids & Solids Silberberg 3 ed

Intermolecular Forces: Liquids, Solids, and Phase Changes 12. 1 An Overview of Physical States and Phase Changes 12. 2 Quantitative Aspects of Phase Changes 12. 3 Types of Intermolecular Forces 12. 4 Properties of the Liquid State 12. 5 The Uniqueness of Water 12. 6 The Solid State: Structure, Properties, and Bonding 12. 7 Advanced Materials

Table 12. 1 A Macroscopic Comparison of Gases, Liquids, and Solids State Shape and Volume Compressibility Ability to Flow Gas Conforms to shape and volume of container high Conforms to shape of container; volume limited by surface very low moderate Maintains its own shape and volume almost none Liquid Solid

Chapter 12: Intermolecular Attractions & the Properties of Liquids & Solids 1. Why do the gas laws work with almost any gas? 2. Gases are alike » Mainly empty space Weak intermolecular attractions Solid: Maintains its own shape and volume Liquid: Conforms to shape of container; volume limited by surface Gas: Conforms to shape and volume of container

Why aren’t there liquid laws and solid laws? 1. Little empty space between molecules: Particles close together in S’s and L’s 2. Stronger and quite varied intermolecular attractions in S’s & L’s than in gases. . . Why? » Attractions decrease as distance between molecules increase: I. M. F. a 1/ d 2 3. Attractions dependent on chemical composition – Polar vs Nonpolar molecules e. g. Water vs Carbon Dioxide (sublimes. @ -58. 5 o. C)

Intermolecular Attractions: Bonds between Molecules 1. 2. 3. Much weaker than Chemical Bonding within molecules Chemical Bonds (ionic and covalent) determine chemical properties Intermolecular bonds determine physical properties e. g. density, mp, bp, solubility, vapor pressure, etc.

Kinds of Intermolecular Attractions 1. 2. 3. 4. 5. Dipole-Dipole Attractions Hydrogen Bonds (H-FON Bonds) London Forces (dispersion forces) Ion-Dipole (e. g. spheres of hydration) Chapter 13 Induced Dipole forces Chapter 13 a. Ion induced b. Dipole induced

How do IMF’s affect Heats of Vaporization and Fusion?

Relative magnitudes of forces in Molecular compounds Covalent bonds >>>>> Hydrogen bonding >> Dipole-dipole interactions >>>>> London forces

What kind of IMF? ?

Dipole-Dipole Attractions 1. Dipoles are polar molecules – Molecules w/ polar bonds and asymmetric distribution of charge – What determines if a bond is nonpolar, polar or ionic? – What determines if a molecule with polar bonds is polar or nonpolar 2. 3. Much weaker than covalent bonds Important in maintaining the shape of many biological molecules: e. g. Proteins

Hydrogen Bonds (H-FON Bonds) 1. Special kind of dipole-dipole interaction » Found in HF and molecules containing O-H or N-H bonds 2. 3. 4. 5 x’s the strength of a typical dipole-dipole bond ~ 5% the strength of a covalent bond Important in biological molecules e. g. DNA, proteins

Myoglobin

Hemoglobin

Hydrogen bonding is responsible for the expansion of water when it freezes.

Exercise: Which of the following molecules display hydrogen bonding? 1. 2. 3. 4. Methane, CH 4 methyl ether, CH 3 OCH 3 Hydrogen peroxide, H 2 O 2 methyl alcohol, CH 3 OH

London Forces: Attractions between temporary dipoles Electrostatic Attraction

Random movement of electrons may cause temporary charge imbalances London Force

London (dispersion) forces between nonpolar molecules

Why are London forces the greatest in large molecules?

London Forces (dispersion forces) exist in all molecules 1. 2. Ave. strength <<< Dipole-Dipole interactions Result from temporary charge imbalances a. Due to the random movement of electrons b. Nucleus of one atom attracts electrons from a neighboring atom. c. At the same time, the electrons in one particle repel the electrons in the neighbor and create a short lived charge imbalance.

Relationship between atomic size and the strength of London forces 1. 2. Greatest in large atoms » Electron clouds more easily distorted » Halogens and Noble Gases BP increase w/ molar mass Ion Induced Dipoles » dipoles can be induced by ions » attractions exist between ions and dipoles

London Forces Effect of molecular surface area Cyclopentane, BP = 49. 3 o. C

IMF’s in Nonpolar Organic Molecules What kind of attractive forces are present? What role do molecular size and surface area play? Linear molecules have more surface area than if they are folded into a sphere. 1. 2. 3. • Linear molecules have higher melting and boiling points because of the increased attractions.

Predicting the Relative Boiling Points of Substances 1. 2. The substance with the strongest intermolecular attractions will have the higher BP. Why? More energy needed to separate molecules higher boiling temperature e. g. Halogens and noble gases

Cooling Curve: H 2 O (g) H 2 O (s) What is happening in to KE and PE at each part of the curve? DHvap = + 40. 7 k. J/mol DHfus = + 6. 01 k. J/mol

Molar Heat of Vaporization, DHvap Heat absorbed when one mole liquid is changed to one mole of vapor at constant T and P » Depends on strength of IMF’s » Endothermic (results in PE elevation) » For water: DHvap = + 40. 7 k. J/mol @ 100 o. C

Molar Heat of fusion, DHvap Heat absorbed when 1 mole solid is changed to 1 mole of liquid at constant T and P. » Depends on strength of IMF’s » Endothermic (results in PE elevation) » For water: DHfus = + 6. 01 k. J/mol

Quantitative Aspects of Phase Changes Within a phase A change in heat is associated with a change in average KE and, therefore, a change in temperature. q = (mass)(Specific Heat)(Dt) During a phase change A change in heat occurs at constant temperature, which is associated with a change in PE, as the average distance between molecules changes--Bond IMF formation is exothermic, IMF breaking is endothermic q = (moles of substance)(enthalpy of phase change)

Calculation of the Heat of Fusion of Ice 1. Use the data below to calculate the heat of fusion of water. – – 2. A piece of ice at zero Celsius melts in 100. 0 g water until the water’s temperature also becomes zero. o Initial water temp. = 44. 0 C Mass of ice that melted = 56. 0 g. o Specific heat of water, CH 2 O = 4. 184 J/g C Calculate the % error and explain the source of the error. DHfus H 2 O = + 6. 01 k. J/mol

Application Questions 1. 2. o The molar heat of vaporization of water at 25 C is 43. 99 k. J/mol. How many kilojoules of heat would be o required to vaporize 125 m. L (125 g) of water at 25 C? » Answer: 305 k. J How much heat would be needed to convert 125 m. L o o water at 25. C to steam at 100. 0 C? The heat of vaporization of water at 100. 0 o. C is 40. 657 k. J/mol and o the specific heat of water is 4. 184 J/g C. » Answer: 321 k. J

General Properties of Liquids and Solids 1. 2. Macroscopic properties depend on Microscopic properties of L’s and S’s » Molecules tightly packed » Strong intermolecular attractions

Macroscopic properties of S’s and L’s 1. 2. Compressibility » Little to none. Why? Diffusion (ability to flow) » Slow in liquids (T 6) – Like moving in a crowded room » Nonexistent in Solids – Particles not free to move

Macroproperties Liquids Why are raindrops spherical? » Increases stability by maximizing the number of IMF’s and decreasing surface tension » Surface Tension = • E needed to increase the surface area of a liquid by a given amount (J/m 2) • Depends on nature of intermolecular forces

The Molecular Basis of Surface Tension Surface molecules experience fewer intermolecular forces than interior molecules Liquids minimize surface area by forming spherical surfaces Lowers PE, thus increases stability ( e. g. Raindrops, overfilled glass) • Surface molecules are at higher P. E. than interior molecules Recall: Bond formation results in P. E. lowering

Macroproperties Liquids 1. 2. Wetting of a Surface by a Liquid » Spreading of a Liquid across a surface » Caused by attraction of liq. molecules to surface molecules Why are Liquids with a Low Surface Tension Good Wetters? . . . . – Low Surface Tension means weak IM Forces 3. Which wets solids surfaces better, hydrocarbons (gasoline/oil ) or water? . . . .

Evaporation and Sublimation Where molecules leave surface and enter vapor space around them » Evaporation: L Vapor » Sublimation: S Vapor

Evaporation: 1. 2. 3. 4. L Vapor (T 9) Factors that affect the rate of evaporation » Surface area, Temp. , Strength of IMF’s Why does evaporation occur at temp. ’s below the BP? Why does an increase in temp. increase the rate of evaporation? Why does sweating cool you? » Why does a fan cool you?

Effect of Temperature on the Distribution of Molecular Speeds in a Liquid

Evaporation: L Vapor : Application Questions (T 10) 1. 2. If the liquids are water and ethanol, which one is liquid A? Liquid B? Why does the evaporation of ethanol from the skin feel cooler than the evap. of water? » Which removes more K. E. ? 3. 4. Which liquid will evaporate more quickly, acetone or ethanol? Which is more difficult to clean up, a marine spill of low or high MW crude oil?

Application Questions (T 12) 1. 2. How can snow or ice cubes in the freezer disappear w/o melting? Why do moth balls (naphthalene) need to be replaced periodically?

Changes of State & Dynamic Equilibrium 1. Change of State (T 11) » Change from one physical state to another S L G or S G – Occur under conditions of Dynamic Equilibriium – Two opposing events occurring @ equal rates 2. Application Questions » Why does fanning cause you to feel cooler? » Why does sweating in the tropics cool you less than sweating in the desert?

Vapor Pressure of Solids & Liquids 1. Vapor Pressure exerted by a vapor in a closed flask in equilibrium w/ its liquid (T 11) 2. Measurement of vapor pressure (T 13 a)

Rateevap > Ratecond Rateevap = Ratecond Liquid – Gas Equilibrium

• Only Temperature and Intermolecular Forces affect Vapor Pressure • The % of molecules with sufficient K. E. to escape the liquid surface increases with Temperature

Factors affecting Vapor Pressure 1. 2. Strength IMF’s (T 13 b) Temperature » At higher temps a higher % of the molecules have suffiecient K. E. to escape the liquid surface 3. Factors not Affecting VP » Surface Area – Increases the rate of evaporation and condensation equally » Amount of Liquid – Evaporation occurs at surface

VP in Solids 1. Due to vibrations of surface molecules Sublimation (T 12) Examples • Snow and ice cubes • Dry ice (solid carbon dioxide)

Boiling Points of Liquids Boiling Point 1. • • Temp. at which VP of a liquid equals atmospheric pressure (T 14) Depends on o o 2. 3. Strength of IMF’s Atmospheric pressure Why do liquids Boil? Why do bubbles form on sides first?

Application Questions 1. 2. 3. Explain why water can exist at temps above its normal BP in a car’s radiator. Explain why a pressure cooker can cook a beef stew faster than in a normal pot. Explain why water, hydrogen fluoride and ammonia have much higher BP’s than one might predict by size alone (T 15)

Clausius-Clapeyron Equation 1. Relates the Vapor Pressure of a liquid to the heat of vaporation, -DHvap Ln P = -DHvap/RT + constant or Ln (P 1/P 2) = -DHvap /R(1/T 2 - 1/T 1) Note: P = Vap Pressure; R = 8. 314 J/mol K 2. Our next lab experiment will involve the use of the Clausius-Clapeyron Equation to calculate the DHvap for methanol and ethanol

Clausius-Clapeyron Eqn: ln P = -DHvap/RT + c P = Vapor Pressure R = 8. 314 J/mol K

Dynamic Equilibrium and Le Chatelier’s Principle: When a stress is applied to a system at equilibrium, the equilibrium will shift to relieve that stress and, if possible, restore equilibrium. » Position of Equilibrium: Refers to relative amount of reactants and products Application Questions. . . . 1.

Le Chatelier’s Principle : Application Questions 1. 2. 3. Use L. C. P. to predict how an increase in temperature will affect the vapor pressure of a solid. Use L. C. P. to predict how a decrease in temperature will affect the vapor pressure of a liquid. Use L. C. P. to predict how an increase in atmospheric pressure will affect the vapor pressure of a liquid.

Phase Diagrams Objective: Interpret phase diagrams and show a phase diagram can be used to represent thermodynamic relationship between the three states of matter for a particular substance.

Phase Diagrams (T 16) 1. 2. 3. 4. Lines represent equilibrium between phases Triple point » T and P at which all three phases present @ equilibrium Critical Temperature » Temp at which liquid phase can not be distinguished from its vapor Supercritical Fluid » Fluid at a Temp > Critical Temp. – e. g. Supercritical carbon dioxide Used to decaffeinate coffee and tea

Application Questions 1. 2. 3. What phase will occur if water at – 20. C and 2. 15 torr is heated to 50. o. C under constant pressure? What phase will water be in if it is at a pressure of o 330 torr and a temperature of 50 C? Use L. C. P. to explain why the MP of ice decreases as pressure increases. o

Application Questions (T 17) 1. 2. At what temperature does Dry ice sublime? What effect does an increase in pressure have on the melting point of carbon dioxide. Use L. C. P. to explain why.

Crystalline Solids Objective: • • Describe how atoms, molecules, or ions are arranged in crystalline solids Unit Cells to Know – Simple cubic (T 18) – Face-centered cubic (T 19, 20, 21) – Body-centered cubic (T 21)

X-Ray Diffraction Objective: Describe the use of X-Ray diffraction to determine the structure of crystals (T 23 a)

Physical Properties and Crystal Types Objective: Relate the properties of solids to crystal type » Understand Table 12. 5, (T 24 a) Crystal Types • Ionic • Molecular • covalent (Network) • Metallic (T 24 c)

Application Questions 1. 2. Boron nitride, which has the empirical formula BN, o melts under pressure at 3000 C and is as hard as diamond. What is the probable crystal type of BN? Crystals of elemental sulfur are easily crushed and o melt at 113 C to give a clear yellow liquid that does not conduct electricity. What is the probable crystal type for solid sulfur?

Physical Properties of Graphite vs. Diamond Property Density (g/m. L) Hardness Color Electrical Conductivity DHcomb (k. J/mol) Graphite Diamond 2. 27 Very soft 3. 51 Very hard Shiny black Colorless/transparent High None -393. 5 -395. 4

Which Structure is Diamond? Graphite?

Uses of Unit Cells • Unit cells may be used to determine the a. Empirical formula of an ionic compound E. g. Problems 12. 90 and 12. 91, page 480 Silberberg 3 rd ed. b. Molar Mass of a substance E. g. Problem 12. 93, page 480 Silberberg 3 rd ed. c. Density of a substance (divide the mass of the atoms per unit cell by the volume of the unit cell)