Chapter 12

1. Chapter 12 Intermolecular Attractions and the Properties of Liquids and Solids

2. Gases, Liquids and Solids

3. Intermolecular Attractions • Intermolecular forces depend on distance • Gases have very small attractive forces • Solids/liquids have larger attractive forces since molecules are closer to each other • Intermolecular forces - attractions between two molecules • Intramolecular forces - chemical bonds that hold molecule together • Intermolecular forces weaker than Intramolecular forces

4. van der Waals’ Forces • HCl molecules • H and Cl atoms held tightly by covalent bond • Strength of chemical bond keeps molecule intact • Attractions between HCl molecules are weaker (4% as strong) • Attraction between molecules determine physical properties • Notice disorientation!

5. Dipole-dipole attractions • HCl(g) - polar molecule with partial charges • Polar molecules tend to line up so partial negative and near partial positive • Still net attraction!! (Dipole-dipole!) • Why weak? • Charges associated are only partial charges • Ordinary temperatures (Thermal energy) causes the dipoles to be somewhat misaligned reducing effectiveness of attractions

6. Hydrogen Bonds • Important Dipole-dipole attraction when hydrogen bonds to very small, highly electronegative atom • Think FON (HF, OH, and NH) • Why Hydrogen bonding? • Ends of bond contain substantial positive and negative charges • Charges highly concentrated due to small size • Positive ends can get very close to negative of another molecule due to small size

7. Hydrogen Bonds in Water • In Liquid water - molecules experience hydrogen bonds that continually break and re-form • As water freezes, molecules become locked and participate in 4 hydrogen bond • Resulting structure has larger volume than liquid water • Ice cubes float in more dense liquid

8. London Dispersion forces • Nonpolar molecules still have attraction (although weak) to hold substance together • 1930 - Fritz London, German Scientist • Nonpolar substances can still have attraction • Atoms constantly moving • Motion in one particle affects neighboring particles • Electrons repel and push away • At any given moment, the electron density of molecule can be unsymmetrical • At particular instant, instantaneous dipole!

9. Induced Dipoles • As instantaneous dipole forms, causes electron density of neighbor to be unsymmetrical • Also forms a dipole (called INDUCED DIPOLE) • Always causes positive of one to be near negative of another

10. Very short lived attraction • Dipoles vanish as they are formed but will form in other location • Over period of time, there is a net, overall attraction but relatively weak • Called London dispersion forces or instantaneous dipole-induced dipole attractions • Distinguished from permanent dipole-dipole

11. Strengths of London forces • Measure using boiling point • Polarizability • Measure of the ease the electron cloud is distorted • As volume of electron cloud increases, polarizability increases • As atom size increase, higher London forces • Number • For molecules containing same elecments, London forces increase with number of atoms • BPhexane > BPpropane • Molecular Shape

13. Intermolecular forces and tightness of packing affect the properties of liquids and solids • Compressibility and diffusion depend primarily on tightness of packing • Most physical properties depend primarily on strengths of intermolecular attractions • Rate of evaporating depends on surface area, temperature, and strengths of intermolecular attractions

14. Compressibility and Diffusion • Compressibility - measure of the ability of substance to be forced into smaller volume • Incompressible • Solids and liquids have no empty volume • Diffusion • Quick in gases • Slow in liquids • Almost nonexistent in solids

15. Surface tension • A property related to the tendency of a liquid to seek a shape that yields the minimum surface area • The shape with minimum surface area is sphere • Why? • Molecules within liquid surrounded by densely packed molecules • Whereas surface molecules have neighbors beside and below it • Surface molecules are attracted to fewer neighbors

16. A molecule at the surface has a higher potential energy than a molecule in the bulk of the liquid • Remember a system becomes more stable when its potential energy decreases • For a liquid, reducing surface area (reducing the number molecules at surface area) lowers potential energy • Lowest energy achieved when liquid has smallest surface area • Surface tension of a liquid is proportional to energy needed to expand surface area

17. Surface tension every day • Water in rim • “Invisible Skin” • Soap, Pepper and Water

18. Wetting of surface by a liquid • Wetting - spreading of liquid across a surface to form a thin film • To occur, the intermolecular attractive forces between the liquid and the surface must be of about the same strength as forces within liquid itself • Think glass coated • SURFACTANTS - drastically lower the surface tension of water • Water is “wetter”

20. Viscosity • Viscosity - resistance to a change in form of a liquid • “Internal friction” of material • Factors • Temperature (Temp decreases; viscosity increases) • Molecular size • Tangling • Attractions • Acetone vs. Ethlyene glycol

21. Evaporation and sublimation • Important factor: Change of State! • Evaporation - for liquid, tendency to undergo change to gas • Sublimation - solid to gas change of state

22. Evaporation and cooling • Evaporation causes cooling effect • Rate of evaporation per unit of surface area of a given liquid is greater at a higher temperature • The weaker the intermolecular attractive forces, the faster the rate of evaporation at a given temperature

23. Change of state - substance is transformed from one physical state to another • Physical equilibrium similar to chemical equilibrium • Evaporation and condensation (change of vapor to liquid) cause equilibrium • Evaporation increases number of molecules in vapor • Condensation decreases number of molecules in vapor • Melting Point - solid to liquid

24. Vapor Pressure • Vapor pressure - the pressure that molecules exert when a liquid evaporates • Equilibrium vapor pressure • Occurs in closed container • Rates of evaporation and condensation are equal • Concentrations of molecules in vapor remains constant and the vapor exerts constant pressure

25. Factors that affect Vapor Pressure • VP is solely function of evaporation per unit area of liquid’s surface • If rate large • Large conc of molecules in vapor state necessary for eq. • Also, VP is high, and evaporation rate high • As temp increases, rate and VP increase • Can use VP as indication of relative strengths of attractive forces in liquids

26. Boiling Point • How would you check for boiling? • Increasing heat just increases amt of bubbles • Any pure liquid remains at constant temperature at boiling point • Why do liquids boil? • Bubbles contain vapor!! • As liquid evaporates, pressure pushes • Opposing force is pressure of atmosphere • The temp at which vp of liquid is equal to prevailing atmospheric pressure

27. Why does water boil at lower temp in Denver than NY? • Normal boiling point - boiling point of liquid at 1 atm • Relates to intermolecular attractions • When attractive forces are strong, the liquid has low vp and therefore, must be heated to higher temp • High boiling points result from strong intermolecular attractions

28. Heating and cooling curves • Heating Curve Application • Heat at constant rate • Diagram

29. Le Châtelier’s Principle • When a dynamic equilibrium in a system is upset by a disturbance, the system responds in a direction that tends to counteract the disturbance and, if possible, restore equilibrium • Heat + Liquid  vapor • Position of equilibrium

30. Phase Diagrams • Graphical representation of phase equilibria • Triple point - all three phases exist • Critical point

31. Phase Diagrams for Water