States of Matter Chapter 12

1. States of MatterChapter 12

2. 12.1 - Kinetic Molecular Theory • Particles in an ideal gas… • have no volume. • have elastic collisions. • are in constant, random, straight-line motion. • don’t attract or repel each other. • have an avg. KE directly related to Kelvin temperature.

3. Real Gases • Particles in a REAL gas… • have their own volume • attract each other • Gas behavior is most ideal… • at low pressures • at high temperatures • in nonpolar atoms/molecules

4. Characteristics of Gases • Gases expand to fill any container. • random motion, no attraction • Gases are fluids (like liquids). • no attraction • Gases have very low densities. • no volume = lots of empty space

5. Characteristics of Gases • Gases can be compressed. • no volume = lots of empty space • Gases undergo diffusion & effusion. • What’s the difference?

6. K = ºC + 273 ºF -459 32 212 ºC -273 0 100 K 0 273 373 Temperature • Always use absolute temperature (Kelvin) when working with gases.

7. Pressure Which shoes create the most pressure?

8. Aneroid Barometer Mercury Barometer Pressure • Barometer • measures atmospheric pressure

9. U-tube Manometer Bourdon-tube gauge Pressure • Manometer • measures contained gas pressure

10. Pressure • KEY UNITS AT SEA LEVEL 101.325 kPa (kilopascal) 1 atm 760 mm Hg 760 torr 14.7 psi All of these amounts are equal!

11. Convert the following pressures and temperatures • Convert temps: • 36°C  K • 237 K  °C • 98.6°F  °C • Convert pressures: • 765 mmHg  torr • 2.3 atm  mmHg • 560 kPa  atm

12. Dalton’s Law of Partial Pressures • The total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture. Ptotal = P1 + P2 + P3 + … Pn

13. Dalton’s Law Practice • A mixture of oxygen (O2), carbon dioxide (CO2) and nitrogen (N2) has a total pressure of 0.97 atm. What is the partial pressure of O2 if the partial pressure of CO2 is 0.70 atm and the partial pressure of N2 is 0.12 atm? Ptotal = PN2 + PCO2 + PO2

14. Intermolecular Forces

15. 12.2 - Definition of IMF Intermolecular forces = Attractive forces between molecules • Much weaker than chemical bonds within molecules

16. Some quick notes on Intermolecular Forces • Intermolecular Forces (IMF) – the forces between molecules • Strongest in solids – hold molecules tightly together and in place • Weakest in gases – gas particles are far apart

17. Types of IMF

18. Types of IMF • London Dispersion Forces – occur between ALL molecules View animation online.

19. + - Types of IMF • Dipole-Dipole Forces – occur in polar molecules View animation online.

20. Types of IMF • Hydrogen Bonding – occur between molecules containing a hydrogen atom bonded to a small highly electronegative atom with at least one lone electron pair

21. 12.3 - Liquids & Solids

22. Liquids • Essentially incompressible, density in between solids and gases • Fluidity – like gases, are considered fluids because they flow and diffuse • Viscosity – a measure of the resistance to flow (i.e., honey has a high viscosity)

23. Liquids • Surface tension – a measure of the inward pull by particles in the interior of the liquid on the particles on the surface of the liquid • Surfactants – lower the surface tension of water (ex: soaps and detergents)

24. Liquids • Cohesion – forces of attraction between identical molecules • Adhesion – forces of attraction between molecules that are different • Capillary action – movement of a liquid upward into a narrow tube (capillaries)

25. The Meniscus, Adhesion & Cohesion

26. Solids • More dense than liquids and gases • Incompressible • Crystalline solid – atoms, ions, or molecules are arranged in an orderly, geometric structure (ex: salt) • Amorphous solids – particles are not arranged in a regular, repeating pattern (ex: glass)

27. Amorphous Solids

28. 12.4 - Phase Changes

29. Phase Changes • Evaporation • molecules at the surface gain enough energy to overcome intermolecular forces (IMF) • Volatility • measure of evaporation rate (volatile substances evaporate easily) • depends on temp & intermolecular forces (IMF)

30. Phase Changes • Equilibrium • trapped molecules reach a balance between evaporation & condensation

31. temp v.p. IMF v.p. Phase Changes • Vapor Pressure • pressure of vapor above a liquid at equilibrium • depends on temp & IMF • directly related to volatility v.p. temp

32. Patm b.p. IMF b.p. Phase Changes • Boiling Point • temp at which v.p. of liquid equals external pressure • depends on Patm & IMF • Normal B.P. - b.p. at 1 atm

33. IMF m.p. Phase Changes • Melting Point • equal to freezing point • Which has a higher m.p.? • polar or nonpolar? • covalent or ionic? polar ionic

34. Phase Changes • Sublimation • solid  gas • Does NOT go thru the liquid state • EX: dry ice, mothballs, solid air fresheners

35. Phase Diagrams • Show the phases of a substance at different temps and pressures.

36. Gas - KE  Boiling - PE  Liquid - KE  Melting - PE  Solid - KE  Heating Curves

37. Heating Curves • Temperature Change • change in KE (molecular motion) • depends on heat capacity • Heat Capacity • energy required to raise the temp of 1 gram of a substance by 1°C • Water has a very high heat capacity

38. Heating Curves • Phase Change • change in PE (molecular arrangement) • temp remains constant • Heat of Fusion (Hfus) • energy required to melt 1 gram of a substance at its m.p.

39. Heating Curves • Heat of Vaporization (Hvap) • energy required to boil 1 gram of a substance at its b.p. • usually larger than Hfus…why?

40. Ideal Gas Law PV=nRT UNIVERSAL GAS CONSTANT R=0.0821 Latm/molK You don’t need to memorize this value!

41. Ideal Gas Law Problems • Calculate the pressure in atmospheres of 0.412 mol of helium (He) at 16°C & occupying 3.25 L. GIVEN: P = ? atm n = 0.412 mol T = 16°C = 289 K V = 3.25 L R = 0.0821Latm/molK WORK: PV = nRT P(3.25)=(0.412)(0.0821)(289) L mol Latm/molK K P = 3.01 atm