Chapter 12

1. Chapter 12 States of Matter

2. Kinetic Molecular Theory • Composition and structure (types of atoms & arrangments) determine chem properties of matter • Also affect physical properties • Kinetic Molecular Theory – describes behavior of matter in terms of particles in motion • Kinetic is Greek for “to move” • Ludwig Boltman & James Maxwell proposed model to explain properties of gases

3. Kinetic Molecular Theory • Several assumptions are made by the model: • Particle Size • Gases consist of small particles separated by empty space. • Volume of particles are small compared to volume of empty space • No significant attractive/repulsive forces since particles are so far apart • Particle Motion • Particle Energy

4. Kinetic Molecular Theory • Several assumptions are made by the model: • Particle Size • Particle Motion • Particles are in constant, random motion that move in straight lines until they collide • Collisions between gas particles are elastic. No kinetic energy is lost (may be transferred). • Particle Energy

5. Kinetic Molecular Theory • Several assumptions are made by the model: • Particle Size • Particle Motion • Particle Energy • Mass and velocity determine kinetic energy • m=mass • v=velocity (speed and direction) • Temperature measure of average KE in a sample of matter

6. Behavior of Gases • KMT explains behavior of gases • Constant motion of particles allows gas to expand and fill container • Low Density () • Gas with lower density => fewer molecules than another element in the same volume • Compression/Expansion • Compression – reduce volume • Air is compressible • Expansion – larger volume

7. Behavior of Gases • Diffusion & Effusion: • Since gas particles are not attracted to one another, they easily pass by each other • When gas flows from one space to another space already occupied by a gas, the gases will mix until evenly distributed (diffusion) • Diffusion – movement of one material through another • Particles diffuse from area of high concentration to low concentration • Ex: food coloring in water, smell when cooking, etc.

8. Behavior of Gases • Diffusion & Effusion ~ cont.: • Effusion – gas escapes through tiny opening • Graham’s Law of Effusion – • means “inversely proportional to”

9. Behavior of Gases • Heavier particles diffuse slower than lighter particles • Therefore, Graham’s Law can also set up a proportion to compare diffusion rates of 2 gases:

10. Example 1 Ammonia has molar mass = 17 g/mol, hydrogen chloride has molar mass of 36.5 g/mol. What is ratio of diffusion rates? **On the right track as HCl is about twice as massive as ammonia!

11. Example 2 Calculate ratio of effusion rates for nitrogen (N2) and neon (Ne). Soln: 0.849

12. Gas Pressure • Pressure – force per unit area • Air Pressure (atmospheric) – pressure is exerted in all directions and varies in different locations due to gravity changes (at different elevations). • P is lower at higher altitudes • At sea level,

13. Gas Pressure • Measuring Air Pressure: • Barometer – instrument used to measure atmospheric pressure • At sea level, ~760 mmHg • Changes in air temp or humidity changes pressure! • Manometer – instrument used to measure gas pressure in closed container

14. Units of Pressure • Units of Pressure: • SI unit is a pascal (Pa) • ** • At sea level, average air pressure is 101.3 kPa when

15. Comparison of Pressure Units

16. Dalton’s Law of Partial Pressure • Dalton’s Law of Partial Pressures – Total pressure of mixture of gases = sum of the pressures of all gases in mixture • Each gas in a mixture exerts pressure independently of other gasses

17. Example 3 A mixture of O2 gas, CO2 and N2 has a total pressure of 0.97 atm. What is the partial pressure of O2 if the partial pressure of CO2=0.70 atm and N2=0.12 atm?

18. Chapter 12.2 Forces of Attraction

19. Intermolecular Forces • Intramolecular Forces: • Attractive forces that hold particles together in ionic, metallic, and covalent bonds • “intra-” means forces within • Intermolecular Forces • The forces occur between or among particles • 3 types: • Dispersion Forces: weak forces that result in temporary shifts in density of electron clouds • Dipole-Dipole – attraction between oppositely charged regions of polar molecules • Hyrdrogen Bonding – (type of dipole-dipole) bond that exists between H and one of the following: F, O, N

20. Dispersion Forces • Weak forces • Sometimes called London forces after physicist Fritz London • Electrons in electron cloud are in constant motion • When 2 electron clouds come in contact, they repel • For q quick moment, the electron density is greater in one region forming a temporary dipole • All particles consist of dispersion forces • As size of particle increase, dispersion forces increase too (stronger)

21. Dipole-Dipole Forces • Polar molecules contain permanent dipoles (some regions are always partially pos/neg) • Neighboring molecules orient themselves so oppositely charged particles will align

22. Hydrogen Bonds • Typically stronger than dipole-dipole and dispersion • H bonding only occurs with F, O, N (FON or “phone”)

23. Comparison

24. Chapter 12.3 Liquids and Solids

25. Liquids • Recall, (Chap 3) • Take shape of container! • Forces of attraction keep molecules closely packed in fixed volume, but not in fixed position. • Density and Compression: • Liquids are much denser than gases - stronger intermolecular forces holds particles together. • Large amounts of pressure must be applied to compress liquids to very small amounts. • Liquids considered incompressible

26. Liquids • Fluidity - ability to flow and diffuse; liquids and gases are fluids. • Liquids diffuse slower than gases at the same temp due to intermolecular forces interfering with flow • Viscosity - measure of resistance of a liquid to flow • Determined by type of intermolecular forces, size and shape of particles, and temp • Particles in liquid are close enough for attractive forces to slow their movement as they flow past one another • The stronger the intermolecular attractive forces, the higher the viscosity

27. Liquids • Viscosity~Cont. • Larger molecules create greater viscosity. • Long chains of molecules result in higher viscosity. • Increasing the temp decreases viscosity - adds energy that allows molecules to overcome intermolecular forces and flow more freely.

28. Liquids • Surface Tension - energy required to increase the surface area of a liquid by a given amount. • Particles in middle can be attracted to particles above, below and on either side • No attractive forces above for particles on surface to balance forces from below • Net forces is pulled down => surface pulled tight • Stronger attractive forces means stronger surface tension **Soaps/detergents decrease surface tension by disrupting H-Bonds. When bond is broken, water spreads out allowing dirt to be carried away! • Surfactants- compounds that lower the surface tension of water.

29. Liquids **When water placed in glass tube, the surface of the water is not straight (meniscus) • Cohesion - force of attraction between identical molecules • Adhesion - force of attraction between molecules that are different • Capillary action- upward movement of liquid into a narrow cylinder, or capillary tube (how paper towels work!)

30. Solids • Recall: • Definite volume • Definite shape • Most solids are more dense than liquids. • Ice is not more dense than water. • Particles in solid ARE in constant motion! • Particles in a solid vibrate in a fixed position • To be solid vs. liquid at room temp, there must be strong attractive intermolecular forces • These forces limit motion of particles in vibration creating order

31. Solids • Crystalline Solids - solids with atoms, ions, or molecules arranged in orderly, geometric shape • Location of particles is are represented by points on crystal lattice frame work

32. Solids • Unit Cell – smallest arrangement of atoms in crystal lattice that has same symmetry as whole crystal (building block) • Crystal shapes differ since surfaces/faces of unit cells do not always meet at right angles

33. Solids • 5 Categories of Crystalline Solids: • Molecular Solids • Solids held together by dispersion, dipole-dipole or H-Bonding • Most not solid at room temp • Poor conductors of heat and electricity • Covalent Network Solids • Form multiple covalent bonds (C, Si, etc) • Ex: carbon forms diamonds • Ionic Solids • Metallic Solids • Amorphous Solids

34. Solids • 5 Categories of Crystalline Solids: • Ionic Solids • Each ion is surrounded by ions of opposite charge • High MP and Hardness • Strong but brittle • Metallic Solids • Positive metal surrounded by sea of electrons • Malleable and ductile • Good conductors of heat and electricity • Amorphous Solids • Particles not arranged in regular, repeating pattern (no crystals) • Generally form when molten material cools too quickly to allow crystals to form • Ex: glass, rubber, plastics

35. Chapter 12.4 Phase Changes

36. Phase Changes • Phase changes require energy! • When Energy is added/removed from system, one phase can change into another

37. Phase Changes • Melting: • What happens when you put ice cubes in water? • Water is @ high temp & heat flows from water to ice! • Heat is transfer of energy from object with higher temp to object with lower temp • At ice’s MP, the E ice absorbs disrupts the H-bonds that hold water molecule together in ice crystal • When enough E is absorbed, the H-bonds break and move apart to enter liquid phase • Melting Point – temp in which forces holding crystal lattice together are broken & become liquid

38. Phase Changes • Vaporization: • Process by which liquid changes from liquid to a gas/vapor • When ice melts, temp of ice and water remain constant until ice is melted completely • If heat still being added, THEN temp will increase • **Vapor – gas phase of substance that is usually a liquid at room temp • If input of E is gradual, molecules tend to escape from surface of liquid. When vaporization occurs only at surface of a liquid, you get evaporation.

39. Phase Changes

40. Phase Changes • Vapor Pressure – pressure exerted by a vapor over a liquid • Boiling Point – temp where VP of liquid = the atmospheric pressure • At BP, bubbles begin to form indicating an increase in E • Sublimation – process of a solid changing directly to a gas • Ex: dry ice, moth balls, air fresheners, etc.

41. Phase Changes • Phase changes that Release E: • Freezing Point – temp where liquid is converted into crystalline solid • Heat is removed from system, KE decreases & velocity decreases • H-bonding increases creating a solid • Condensation – process where gas/vapor becomes a liquid • Vapor loses E, velocity decreases, H-bonding increase • Ex: condensation on glass, dew on grass, etc.

42. Phase Changes • Deposition – process of gas/vapor changing to solids without becoming a liquid • Ex: frost

43. Phase Changes • Phase Change Diagrams – graphs of P vs. T that show phase of different substances under different conditions • temp & pressure control phase of substance • Pts on curve is where 2 phases coexist • Triple Point – point that represents temp and pressure where all 3 phases exist • Critical Point – point of critical temp and pressure water can NOT exist as liquid

44. Phase Changes