Dr. Williamson s notes on Liquids and Solids

CHAPTER 1 and 13 LIQUIDS AND SOLIDS Dr. V. M. Williamson Kinetic-Molecular Theory! All gases are composed of tiny particles (atoms/molecules)! Particles are so small and distance between them so large that the volume of the particles are negligible (zero).! Particles are in constant motion, colliding with each other and with the walls of the container! Particles do not attract or repel each other, and collisions are elastic! The average kinetic energy of the particles is directly proportional to the Kelvin temperature 1 Kinetic-Molecular Description of Phases! Schematic representation of the three common states of matter. Heating Curve at Constant Pressure gas cool liquid cool solid heat heat 3 Solids, Liquids and Gases Properties Solid Liquid Gas Definite shape Yes No No Compressibility Nearly Slight High none Density High Intermediate Low Fluid No Yes Yes Very Diffusing Rate Moderate Rapid slow Vibrational Restricted Slow Particle Motion Rapid Particle Distance Close Medium Far Apart Force of Attraction Strong Medium Negligible A Matter of Forces! Electrostatic in nature! Intramolecular forces: particles (molecules, atoms or ions), relatively strong, responsible for chemical properties. (Ionic, Polar Covalent, and Nonpolar Covalent)! Intermolecular forces: particles (molecules, atoms or ions), relatively weak, responsible for physical properties 1

Dr. Williamson s notes on Liquids and Solids Intramolecular Forces: Primary Bonding Force Model Energy Example (kj/mol) Ionic 400-4000 NaCl + + + + + Covalent + + 150-1100 HCl (polar/nonpolar) O + + + Metallic 75-1000 Cu + + + + + + Intermolecular Forces (IMF): Secondary Bonding! Attractive forces responsible for existence of! Balance between these forces and kinetic energy of particles determine particular physical state or property! Stronger forces lead to boiling points, heats of vaporization, heats of fusion! There are four important intermolecular attractions. Ion-Ion Forces! Predominant in compounds! Very strong, consequently lead to melting points for ionic compounds! The force of attraction between two oppositely charged ions- Governed by Coulomb s Law: ( + )( - q q ) q + F d - and q are the ion charges. d is the distance between the ions. Intermolecular Attractions and Phase Changes! Arrange the following ionic compounds in the expected order of increasing melting and boiling points. IMF = bp or mp NaF, CaO, CaF IMF BP MP You do it! What important points must you consider? 10 Intermolecular Attractions and Phase Changes + - + + - Na F Ca F Ca O Charge First Put in order of increasing melting and boiling points. Melting Points of Some Ionic Compounds Cpd MP( o C) NaF 993 NaCl 801 NaBr 747 KCl 770 BaO 193 Cpd MP( o C) CaF 143 Na S 1180 K S 840 MgO 800 CaO 580 11 NaCl NaF MgO charge, size

Dipole-Dipole Forces! Predominant in molecules Dipole-Dipole Forces! Result from attraction of δ + end of one particle with δ end of other particle! Effective only at short distances ( 1/d 4 )! Weaker since involve only partial charges (approximately 4 kj/mol of bonds); about 1% as strong as primary ionic or covalent bonds.! Molecules tend to align themselves so that the opposite charges are near each other Dipole-Dipole Forces Among Polar Molecules Dipole-Dipole Interactions The polar the molecule, the is its boiling point. Hydrogen Bonding! Special case of interaction! Requirements: Hydrogen covalently bonded to a small, highly electronegative element ( ) Electronegative element with lone pair available for hydrogen-bonding! Typically 15-0 kj/mol (about 5x stronger than other dipole-dipole attractions)! Responsible for anomalously high boiling points of some compounds! (note covalent primary bond is still stronger that any IMF) Hydrogen Bonding Consider H O which is very polar molecule and has hydrogen bonding (H bonded to an,,or ) 18 3

Hydrogen Bonding in C H 5 OH Hydrogen Bonding in NH 3 Hydrogen Bonding in DNA Hydrogen Bonding and Boiling Point London Forces, Dispersion Forces, Instantaneous Dipoles! Named for Fritz London (1900-1954) a German physicist-proposed in 1930! The only attractive force for molecules! London Forces are very forces present in all particles, polar or nonpolar! An uneven distribution of electrons form! Depend on particle size or MM (bigger size, greater number of electrons, greater forces)! Influenced by molecular shape (bigger surface area, greater forces) 3 London Forces Cpd MM BP(K) He 4 4. Ne 0 7 Ar 40 87 Kr 84 11 Xn 131 164 Rn 11! No dipole-dipole forces, yet these can be liquified, so some IM force exists! Imagine cooling these down. Rn will liquefy first 4 4

London Dispersion Forces: An Illustration Factors Affecting London Forces! The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane).! This is due to the increased surface area in n-pentane. Factors Affecting London Forces! The strength of dispersion forces tends to with molecular weight.! Larger atoms have larger electron clouds, which are easier to polarize. ionic>hydrogen b.>dipole-dipole>london force! Ionic- requires ions. Differentiate between Ionic by charge on ions and size! Hydrogen bonding- requires H bonding to F, O or N (note molecule will be polar due to lone pairs on F,O,N). Differentiate by dipole moment or EN difference.! Dipole-Dipole- requires polar molecules. Differentiate by dipole moment or EN difference.! London Forces- requires nonpolar molecules. Differentiate by molar mass 8 Dipole-Induced Dipole Forces Ion-Dipole Interactions! Ion-dipole interactions are an important force in solutions of ions.! The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents. 5

The Liquid State-Viscosity! Viscosity is the. " For example, compare how water pours out of a glass compared to molasses, syrup or honey.! Oil for your car is bought based on this property. " 10W30 or 5W30 describes the viscosity of the oil at high and low temperatures. The Liquid State! An example of viscosity of two liquids. IMF Vis 31 3 The Liquid State-Surface Tension! Surface tension is a measure of the unequal attractions that occur at the surface of a liquid. Surface Tension Demo IMF ST! The molecules at the surface are attracted unevenly. 33 34 The Liquid State-Capillary Action! Capillary action is the ability of a liquid to rise (or fall) in a glass tube or other container IMF CA 35 Property of the Liquid State: Capillary Action! Drawing up of liquid inside narrow tube when adhesive forces between the liquid and the tube walls exceed cohesive forces within the liquid! Adhesive forces > cohesive forces: concave meniscus; e.g., water in glass! Cohesive forces> adhesive forces: convex meniscus; e.g., mercury in glass! Responsible for intake of water and nutrients by plants from the soil 6

The Liquid State! Capillary action also affects the meniscus of liquids.! The glass has polar S-O bonds that the polar water is attracted to more strongly that to other water molecules Water Hg Evaporation! Evaporation is the process in which molecules escape from the surface of a liquid and become a gas.! Evaporation is temperature dependent. 37 38 Evaporation of Liquids Evaporation of Liquids: Dynamic Equilibrium IMF Evap! Molecules leave liquid surface and enter the gas phase! In closed system, equilibrium is established Liquid Vapor The rate of condensation is equal to the rate of evaporation Vapor Pressure! Pressure exerted by the vapor in equilibrium with its liquid at a particular temperature! Constant with constant temperature as long as both liquid and vapor are present! Independent of container volume! Obeys Dalton s Law of Partial Pressures! Increases with increasing temperature More About Vapor Pressure! Depends on nature of the substance! Compounds with weak intermolecular forces have high vapor pressures and low boiling points (volatile)! Boiling Point: Temperature at which pressure = pressure! Normal Boiling Point: Temperature at which vapor pressure = P ext = 7

Vapor Pressure: A Butane Lighter Vapor Pressure Curves at Different Temperatures Compare at one temperature IMF VP The Liquid State! Vapor Pressure (torr) and boiling point for three liquids at different temperatures. 0 o C 0 o C 30 o C normal boiling point diethyl ether 185 44 647 36 o C ethanol 1 44 74 78 o C water 5 18 3 100 o C! 760 torr = 1 atm, so water has lowest vp VP BP! What are the intermolecular forces in each of these compounds?! Hydrogen bonding > london forces The Liquid State Compound MW(amu) B.P.( o C) CH 4 16-161 C H 6 30-88 C 3 H 8 44-4 n-c 4 H 10 58-0.6 n-c 5 H 1 7 +36 45 46 The Liquid State! Arrange the following substances in order of increasing boiling points. Ne, NH 3, Ar, NaCl, AsH 3 You do it! IMF BP 47 Elevations and Boiling Points! Boiling at different elevations! Boiling is when vapor pressure = external pressure.! What is the atmospheric pressure on the mountain top? It is less than at sea level.! So, less vapor pressure needed, less heat needed to make vp, so lower boiling point! (note: you will have to boil the eggs longer or add salt to get them cooked at this lower bp.) 48 8

The Solid State Normal Melting Point! The melting point is the temperature at which the solid melts (liquid and solid in equilibrium) at exactly 1.00 atm of pressure.! The melting point increases as the strength of the intermolecular attractions increase. IMF MP Heat of Fusion! Heat of fusion is the amount of heat required to melt of a solid at its melting point at constant temperature.! 334!! J o 1.00 g H O! (s) at 0 C 1.00 g HO( l) at 0!+ o C -334 J Heat of crystallization is the reverse of the heat of fusion. IMF H fus 49 50 Molar Heat of Fusion or ΔH fusion! The molar heat of fusion is the amount of heat required to melt of a substance at its melting point.! The molar heat of crystallization is the of molar heat of fusion! 601! J o 1.00 mole H O! (s) at 0 C 1.00 mole HO( l) at 0!+ o C -601 J IMF H fus 51 Heat of Vaporization! Heat of vaporization (ΔH vap ): heat needed to convert one gram of liquid at its boiling point to vapor at the same temperature! Molar heat of vaporization (ΔH vap ): heat needed to convert one mole of liquid at its boiling point to vapor with the same temperature! Specific heat: heat needed to raise temperature of 1 g of substance by one degree Celsius IMF H vap The Solid State! Which requires more energy? NaCl H O NaCl ( s) ( l) or H O ( s) ( l) Phase Diagrams Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases. IMF H fus ionic IMF> hydrogen bonding IMF So H fus of NaCl is than H fus of H O 53 9

Phase Diagrams! The AB line is the liquid-vapor interface.! It starts at the point (A), the point at which all three states are in equilibrium. Phase Diagrams It ends at the critical point (B); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other. Phase Diagrams Each point along this line is the point of the substance at that pressure. Phase Diagrams! The AD line is the interface between liquid and solid.! The at each pressure can be found along this line. Phase Diagrams! Below A the substance cannot exist in the state.! Along the AC line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line. Phase Diagrams (P versus T)! Compare water s phase diagram to carbon dioxide s phase diagram. 60 10

Types of Solids! Classification based on morphology: Crystalline and amorphous Amorphous vs Crystalline Solids SiO crystalline! Classification based on bonding forces: Ionic, molecular, covalent and metallic amorphous Amorphous vs Crystalline Solids Crystalline! Have periodic, orderly arrangement of atoms! Sharp melting points! May be anisotropic; cleavage gives planar faces! display X-ray diffraction patterns which reflect the molecular structure E.g., sugar, salt, ice Amorphous! No well-defined, long-range ordered structure! Broad melting point range! Properties generally isotropic E.g., glass, rubber, plastic, waxes Structure of Crystals! Unit cells are the smallest repeating unit of a crystal. " As an analogy, bricks are repeating units for buildings. 64 Structure of Crystals! There are three basic variations of the cubic crystal system.! Simple cubic unit cells. " The balls represent the positions of atoms, ions, or molecules in a simple cubic unit cell.! Body-centered unit cells! Face-centered unit cells Ionic Solids! Held by bonding forces! Crystalline arrangement of! High melting points (400-3000 o C)! Poor electrical and thermal conductors (become good conductors in molten state)! Examples: CsCl, NaCl, ZnS 65 11

Molecular Solids Held by intermolecular forces Molecular Solids have in each of the positions of the unit cell. Soft Low melting points ( 7 to 400 o C) Poor electrical and heat conductors Examples: CO, C 6 H 6, C 60 Molecular Solids Covalent Solids Also referred to as Atoms held by covalent bonds to each other Very hard; high-melting (100 to 4000 o C) Poor heat and electrical conductors Covalent Bonding! Some examples of covalent solids are:! Diamond, graphite, SiO (sand) 70 Metallic Solids! Metallic Solids may be thought of as positively charged nuclei surrounded by a sea of electrons.! The positive ions occupy the crystal lattice positions.! Soft to very hard! Wide range of melting points (-39 to 4000 o C)! High heat and electrical conductivity! Examples: Li, Cu, Ni, Ag.. (metals) 71 Crystalline Solids 7 1

ID Type of Bonding in Solids! Fe(s)! Ar(s)! Diamond, C(s)! Ice(s)! CO (s)! NaCl(s)! NH 4 NO 3 (s) 73 13