The Kinetic Theory of Matter explains the properties of solids, liquids, & gases

1. The Kinetic Theory of Matter explains the properties of solids, liquids, & gases

2. The Kinetic Theory of Matter Based on idea that particles of matter are in constant motion. Describes properties of solids, liquids, & gases in terms of the FORCE of the particles The constant random motion of tiny particles called Brownian motion

4. Physical Behavior of Matter Section 10.1

5. States of Matter • Four states of matter

6. Solids • Particles-closely packed; can’t be compressed • Voids - extremely small • Particle motion is vibratorymotion; definite shape & volume • Apply heat-particles vibrate more & move SLIGHTLY farther apart; causes solid to expand

7. Kinetic Model of Solids STRONG intermolecular forces result in rigid structure of solids Particles move (vibrate) but not past each other Particles occupy fixed 3-D positions that repeat throughput the solid 3-D arrangement = crystal lattice

8. Liquids • Flowing matter w/ definite volume & indefinite shape • Particles have weak bonds that keep them close; more space to move; particles able to move relative to each other • When heat applied, liquids expand a little

9. Kinetic Model of Liquids Particles of liquid slide past each other; consider magnetized spheres Intermolecular forces maintain their volume NOT shape

10. Gases • Flow, too (consider wind) • Particles far apart; complete freedom of movement • Motion is random • No definite shape/volume • Easily compressed into smaller volume • Expand & contract in response to temperature changes more so than liquids & solids

11. Kinetic Model of Gases Particles in gas in constant, random motion Change direction ONLY when they strike wall of container OR another gas particle (Air-hockey puck) Density (M/V) in gas lower than solid Few particles in gas vs solid of same volume (due to space between particles)

12. 5 Assumptions of the Kinetic Theory Gases are made of molecules in constant, random movement. LARGE portion of the volume of a gas = empty space. The volume of all gas molecules, in comparison, is negligible.

13. 5 Assumptions of the Kinetic Theory The molecules show no forces of attraction or repulsion (UNLIKE solids & liquids). No energy is lost in collision of molecules; the impacts are completely elastic. The temperature of a gas is the average KE of all of the molecules.

14. Ideal Gases • Ideal gases = gases that obey the assumptions of the kinetic theory • Except for temperatures extremes, most real gases behave like ideal gases • At temperature extremes, forces between particles & particle size begin to matter • At temperature extremes gases no longer follow the assumptions of the kinetic theory

15. Gases & Pressure, Temperature, & Volume • KToM explains gas pressure = the total force exerted by gas molecules colliding against the walls of a container. • IF the container can expand, like a balloon/tire, in pressure can the volume; THUS the balloon/tire will get BIGGER . • If you the temperature of the gas, the KE of its molecules &, the pressure/volume

16. Gases & Pressure, Temperature, & Volume • There is a relationship between pressure, volume and temperature in an ideal gas • If you the pressure & hold the volume constant, the temperature (principle of a refrigerator) • If you the temperature & hold the pressure constant, the volume (heating a balloon)

17. Earth’s Atmosphere and Pressure • We’re at the bottom of an ocean of air • Atmospheric pressure = force exerted on us by molecules of air (14.7 lbs/square inch) • Atmospheric pressure related to column of air; • As elevation ↑ pressure↓ • As elevation ↓, pressure ↑ • How does atm pressure affect YOU?

18. Other Forms of Matter • Amorphous Solids = arrangement of molecules is fairly random; so, crystal lattice is loosely packed ; haphazard, disjointed • Examples = GLASS, cotton candy Wax/Candles

19. Liquid Crystals • When solids melt, the crystal lattice disintegrates; particles lose their 3-D pattern • Liquid crystals-NOT liquid OR solid • When melted LCs lose their rigid organization in 1 or 2 dimensions NOT all 3 dimensions • Interparticle forces in liquid crystals are relatively weak; when forces in lattice are broken, crystals can flow like liquids • Liquid crystal displays (LCDs) used in TVs, watches, calculators, thermometers, etc.

20. Plasmas • Form at very temperatures • Plasma = gas that has been energized; some e- break free from, but travel w/their nucleus • Plasma = free e- & ions of that element. • Gases can become plasmas in several ways, ALL include pumping the gas w/ energy. • Examples = stars, fluorescent tubes, neon lights, etc.

21. Examples of Plasma

22. Energy and Changes of State Section 10.2

23. Temperature & Kinetic Energy & Particle Motion • Temperature = measure of the average kinetic energy of particles in a material • When heated liquid & gas particles have more kinetic energy BUT not all particles have the same kinetic energy; particles are moving at different speeds • Generally, as temperature matter moves to a more active state; as temperature matter moves to a less active state

24. The Kelvin Scale • Absolute zero = temperature at which a substance would have zero (or very little) kinetic energy • Kelvin Scale = used for temperature; it is defined so temperature of a substance is directly proportional to the average kinetic energy of the particles • 0 on Kelvin scale = absolute zero & measured as Kelvins; divisions on Fahrenheit & Celsius scale are measured in degrees • Celsius degree & Kelvins = the same size; absolute zero = -273.150C • Kelvin scale measures everything ABOVE absolute zero; all numbers are positive

25. Temperature Conversions • When converting from kelvin (K or TK) to Celsius (C or TC), and vice versa, the magic number is 273!! • K= (0C+ 273); K= (150C+ 273) = 288 K • 0C= (K- 273); 0C= (320 K - 273) = 470C

26. Mass & Speed of Particles KE of gas depends on mass & speed of particles 1. Gases at SAME temp have SAME average KE 2. LARGER gas molecule simply moves SLOWER than SMALL gas molecule Ex: O2 = 16x more massive than H2; at SAME temp, H2 moves FASTER than O2

27. Mass & Speed of Matter • Random motion causes particles to spread out to fill a container • DIFFUSION = the process in which these particles fill a space • Particles move from areas of high concentration to areas of low concentration • Rate of diffusion of a gas dependent upon the KE of that gas/substance

28. Changes of State TRIPLE POINT = The single specific temperature & pressure at which all 3 phases can co-exist CRITICAL POINT = The conditions where gas & liquid become indistinguishable Different phases of a system may be represented using a phase diagram. Axes of the diagrams are typically pressure & temperature

29. Phase Diagram for Water

30. Changes of State • EVAPORATION • particles of a liquid form a gas by escaping from the surface • 3 things affect evaporation rate? Area of the surface, temperature, humidity • Volatile liquids evaporate quickly (perfumes, paint) • As liquids evaporate, they cool Heating Curve-based on standard temp & pressure

31. Changes of State MELTING The process of heating a SOLID substance to a point where it turns LIQUID. FREEZING is the opposite of melting. It is the process of REMOVING heat from a liquid & turning the liquid into a solid. The freezing point is the SAMETEMPERATURE as the melting point. Heating Curve-based on standard temp & pressure

32. Changes of State • Sublimation-process by which particles in a solid change to gaseous state w/o melting • Condensation-reverse of evaporation; gaseous particles become close (condense) & form a liquid

33. Specific Heat • To change the temperature of a SOLID 2.1 Joules/g0C • To change the temperature of a LIQUID 4.2 Joules/g0C • To change the temperature of a GAS 2.02 Joules/G0C Heat = mass x specific heat x temperature change Q = m x c x (Tf –Ti)

34. Heat of Vaporization • The amount of heat required (absorbed by the liquid) to convert unit mass of a liquid into its vapor w/o a change in temperature. • 2260 Joules = ENREGY needed to move the molecules in 1 g of water FAR enough apart that they form water vapor (JOULE, J,= SI unit of energy required to lift a 1-g mass 1m against the force of gravity) • Heat of Vaporization (Hv) of H2O = 2260 J/g

35. Heat of Vaporization of Water Hv = 2260 J /g • The diagram right shows the uptake of heat by 1 kg of H2O, from ice at -50 ºC  to steam above 100 ºC. A: Rise in temp. as ice absorbs heat.B: Absorption of latent heat of fusion.C: Rise in temp. as liquid H2O absorbs heat.D: Water boils & absorbs latent heat of vaporization.E: Steam absorbs heat & thus increases its temperature.

36. Heat of Fusion • The heat nrg which must be removed to solidify a liquid or added to melt a solid • Melting point=temperature of the solid when its crystal lattice begins to break apart (intermolecular forces are overcome & solid becomes a liquid) • Freezing Point= temperature of liquid when it begins to form a crystal lattice & becomes a solid

37. Heating Curve

38. Phase Diagram