1. Gases, Liquids and Solids Bettelheim, Brown, Campbell and Farrell Chapter 5

2. Overview Phase Changes Gases Intermolecular Forces –London Dispersion, Dipole-dipole, Hydrogen-bonding Liquids Solids

3. State characteristics Gases –No definite shape or volume; flows –Particles far apart; no interactions Liquids –Definite volume, no fixed shape, flows –Particle in contact; moderate interactions Solids –Definite shape, volume; does not flow –Particles in contact; strong interactions

4. Solid LiquidGas All particles Particles Particles do touching touchingnot touch No empty Some openLots of empty spaces spacesspace between between betweenthem them Strong ModerateNo Interactions interactionsinteractions

5. Changes of State Melting: Change from solid to liquid state + heat Freezing: Change from liquid to solid state - heat Evaporation: Change from liquid to gas state + heat Condensation: Change from gas to liquid state - heat Sublimation: Change from solid to gas state + heat (Skips liquid state) Deposition: Change from gas to solid state - heat (Skips liquid state)

6. Changes of State Heat of Fusion: The amount of heat needed to change 1 g of a substance from the solid to the liquid state (at melting point) Heat of Vaporization: The amount of heat needed to change 1 g of a substance from the liquid to the gaseous state (at boiling point)

7. Heat and Phase Changes Example: Naphthalene, an organic substance often used in mothballs, has a heat of fusion of 35.7 cal/g. How much heat, in kilocalories, is required to melt 38.4 g of naphthalene?

9. Heat and Phase Changes Example: Naphthalene, an organic substance often used in mothballs, has a heat of fusion of 35.7 cal/g and a molar mass of 128.0 g/mol. How much heat, in kilocalories, is required to melt 0.300 mol of naphthalene?

10. Phase Changes The heating curve of ice Heat of fusion Heat of vaporization

11. Kinetic Molecular Theory of Gases Gas particles have negligible volume compared to volume gas occupies No attraction between particles Particles move through space in straight lines Kinetic energy (KE) proportional to temperature (K) Can collide with container or each other Total KE before collision = KE after collision Collisions with walls of container exert pressure

12. Gases Gas pressure:Gas pressure: the force per unit area exerted against a surface –most commonly measured in millimeters of mercury (mm Hg), atmospheres (atm), and torr

13. Gas Pressure A mercury barometer

14. Gas Laws Boyle’s law:Boyle’s law: Volume and pressure are inversely proportional (fixed mass and gas at a constant temperature

15. Gas Laws Charles’s Law: Charles’s Law: Temperature and volume are directly proportional (fixed mass and pressure); Temperature is in kelvins (K)

16. Gas Laws Gay-Lussac’s Law:Gay-Lussac’s Law: Pressure and temperature are directly proportional (fixed mass and volume); Temperature in kelvins (K)

17. Combined Gas Law combined gas lawBoyle’s law, Charles’s law and Gay- Lussac’s law can be combined into one law called the combined gas law

18. Gas Laws Problem:Problem: a gas occupies 2.00 L at 5.00 atm. Calculate its volume when the pressure is 10.0 atm. (Assume no change in temperature.)

19. Gas Laws Problem:Problem: a gas occupies 2.00 L at 5.00 atm. Calculate its volume when the pressure is 10.0 atm. (Assume no change in temperature.)

20. Gas Laws Avogadro’s law:Avogadro’s law: volume of gas is directly proportional to its molar amount at a constant pressure and temperature –Volume  number of moles (n) V/n = constant –One mole of ANY gas at STP occupies 22.4 L –Standard temperature and pressure (STP) are 0°C (273 K) and 1 atm pressure

21. Ideal Gas Law Ideal gas law:Ideal gas law: Can be used for a single sample that does not change PV = nRT P = pressure of the gas in atmospheres (atm) V = volume of the gas in liters (L) n = moles of the gas (mol) T = temperature in kelvins (K) ideal gas constant R = ideal gas constant (a constant for all gases)

22. Ideal Gas Law The value of R is determined using the fact that one mol of any gas at STP occupies 22.4 L –Problem: –Problem: 1.00 mol of CH 4 gas occupies 20.0 L at 1.00 atm. What is the temperature of the gas in kelvin?

23. Ideal Gas Law –Problem: –Problem: 1.00 mol of CH 4 gas occupies 20.0 L at 1.00 atm. What is the temperature of the gas in kelvin?

24. Ideal Gas Law –Problem: –Problem: 1.00 mol of CH 4 gas occupies 20.0 L at 1.00 atm. What is the temperature of the gas in kelvin?

25. Gas Laws Dalton’s law of partial pressures:Dalton’s law of partial pressures: the total pressure, P T, of a mixture of gases is the sum of the partial pressures of each individual gas

26. A tank contains N 2 at 2 atm and O 2 at 1 atm. If we add CO 2 until the total pressure is 4.6 atm, what is the partial pressure of CO 2 ?

27. Intermolecular Forces Forces between molecules affect physical properties such as: –Boiling point (bp) –Melting point (mp) –Solubility

28. Intermolecular Forces –Occurs because of electrostatic attractions between between positive and negative “poles” of molecules

29. London Dispersion Forces London dispersion forces are the attraction between very temporary induced dipoles

30. London Dispersion Forces –Exist between all atoms and molecules –In general, their strength increases as the mass and number of electrons in a molecule increases –Even though these forces are very weak, they contribute significantly to the attractive forces between large molecules because they act over large surface areas

31. Dipole-Dipole Interactions The electrostatic attraction between positive and negative dipoles –consider butane and acetone, compounds of similar molecular weight

32. Hydrogen Bonds Hydrogen bond: a hydrogen covalently bonded to an atom of high electronegativity (O, N, F) is attracted to another O, N or F Affects bp, mp and solubility

33. Liquids –When distances decrease so that almost all molecules touch or almost touch, the gas condenses to a liquid –The position of molecules in a liquid is random and there is irregular space between them into which other molecules can slide; this causes liquids to be fluid –Surface tension can result

34. Evaporation/Condensation –If a molecule is moving rapidly (has a high KE) and moving upward, it can escape the liquid and enter the gaseous space above it vapor pressure –Eventually, the number of gaseous molecules will reach an equilibrium with the number of liquid molecules- the partial pressure of the vapor in equilibrium with the liquid: vapor pressure –Boiling point: –Boiling point: the temperature at which the vapor pressure of a liquid equals the atmospheric pressure

35. Evaporation/Condensation

36. Solids crystallization –Upon cooling, molecules come so close together and attractive forces between them become so strong that random motion stops and a solid is formed : crystallization –Can be crystalline or amorphous

37. Types of Solids