States of Matter Grid

1. Kinetic Energy: Kinetic Theory: Energy of Motion ALL MATTER is in constant motion. States of Matter Grid Moderate NONE Very Strong Very tight, close Moderate 99% empty space Vibrate, spin in place Slide past each other Rapid, random straight line definite indefinite indefinite definite definite indefinite incompressible compressible incompressible definite definite definite Does not flow flows flows

2. STATES OF MATTER: Solid, Liquid, Gas, Plasma • SOLIDS: • Most solids are crystalline – • Atoms, ions or molecules arranged in an orderly, repeating 3-D pattern called a… • (CRYSTAL LATTICE). • Unit cell- • Smallest group of particles that keep the shape of the crystal.

3. Allotropes- 2 or more forms of the same element in the same physical state • Example: Solid Carbon • GRAPHITE DIAMOND BUCKEY BALL Each CARBON atom is strongly bonded to 4 other carbons hard PENCIL Widely spaced, linked, hexagon. Weak bonds - soft 60 carbons attached together like a soccer ball. Really flexible

4. Carbon Allotropes So far, it is just of theoretical interest. It is not used in any products or manufacturing processes. One process developed in the Chemistry Division at Argonne National Lab produces smooth thin films of diamond from bucky balls. These films are smoother than those produced by any other method. This some day may be useful in making wear-resistant coatings on things like machine parts.

5. Amorphous solids • Solids that lack an orderly internal structure. Atoms are randomly arranged. • Examples: rubber, asphalt, plastics • Glasses: • Amorphous solids; super cooled liquids • Cooled to a rigid state without crystallizing. Does not melt at a definite temperature but gradually softens. When broken it forms jagged irregular edges.

6. GASES: • KINETIC THEORY OF GASES • 1. There are no attractive forces between gas molecules. • 2. Gases are in constant rapid random straight line motion. • 3. All collisions between gas particles are perfectly elastic. NO ENERGY LOST just transferred from one particle to another.

7. Plasma • Gases heated to very high temperatures. • 1. KE becomes great enough to break molecules into atoms. • 2. As the temp. increases the electrons are stripped off the gaseous atoms. • 3. What results are + ions and free electrons. This is known as plasma. • COLD PLASMA – 50,000K to 1,000,000K • HOT PLASMA- “stars” 10,000,000 to 1,000,000,000K

8. Bose-Einstein • A Bose–Einstein condensate (BEC) is a state of matter of bosons confined in an external potential and cooled to temperatures very near to absolute zero (0 K or -273.15 °C). Under such supercooled conditions, a large fraction of the atoms collapse into the lowest quantum state of the external potential, at which point quantum effects become apparent on a macroscopic scale. • This state of matter was first predicted by Satyendra Nath Bose in 1925. Bose submitted a paper to the Zeitschrift für Physik but was turned down by the peer review. Bose then took his work to Einstein who recognized its merit and had it published under the names Bose and Einstein hence the acronymn. • Seventy years later, the first such condensate was produced by Eric Cornell and Carl Wieman in 1995 at the University of Colorado at BoulderNIST-JILA lab, using a gas of rubidium atoms cooled to 170 nanokelvin (nK)[1] (0.000000170 K or -273.14999983 °C). Eric Cornell, Carl Wieman and Wolfgang Ketterle at MIT were awarded the 2001 Nobel Prize in Physics in Stockholm, Sweden[2]. • When a system of atoms is cooled rather than bosons, the Bose-Einstein condensate is then sometimes called a Super Atom.[3]

9. B. TEMPERATURE • TEMPERATURE AND ENERGY ARE NOT THE SAME THING!!! • Temperature is a measure of • AVERAGE Kinetic Energy • The higher the temperature, the _(Greater)_, avg. kinetic energy of the molecules. • Energy is the ability to do work or supply heat. • Units: calorie or Joule • Oven analogy

10. KELVIN TEMPERATURE SCALE • Based on Absolute Zero. • Used when doing Gas Law calculations because there are no neg. Numbers. • Absolute Zero – • Temperature at which all molecular motion stops! • 0(K) or -273ºC

11. Temperature Scale Comparison Celcius Thirty is hot Twenty is nice Ten is cool Zero is ice

12. Temperature Scale Conversions • Temperature conversions between the three temperature scales: • kelvin / ºCelsius conversions (exact): • kelvin = º Celsius + 273 • ºCelsius = kelvin - 273 • ºFahrenheit / º Celsius conversions (exact): • º F = ºC x 1.8 + 32. • º C = (ºF - 32.) / 1.8

13. Temperature Scale Continued • The Kelvin temperature of a substance is a • Direct reading of the average kinetic energy of the molecules! • If you increase the Kelvin temperature of a substance from 100K to 300K the average speed of the molecules will • TRIPLE (3x greater) • Although the numbers are different on the Celsius and Kelvin scales, each 1 increment on the two scales is • The same! 1ºC = 1 K

14. CHANGES OF STATE – PHASE CHANGES(Solid Liquid Gas)(Solid Liquid Gas) • PHASE CHANGES ALWAYS INVOLVE ENERGY CHANGES. • Heat flows from • ___(HOT)____ to ______(COLD)_____! ALWAYS • Average Kinetic Energy (speed) = • KELVIN TEMPERATURE • Is it possible for one water molecule to be at absolute zero in a pot of boiling water? • YES • Explain: Temperature measures AVERAGE KE. Some mc’s are much hotter and some much colder. As they collide their individual KE’s (temperatures) change.

15. Melting and Freezing • Melting • Phase change from solid to liquid. • Melting Point ( MP ) • Temperature at which a solid becomes a liquid. Temp. remains constant during change. • Freezing • Phase change from liquid to solid. • Freezing Point ( FP ) • Temperature at which a liquid becomes a solid. Temp. remains constant during change. energy in melts (endothermic) • MP = FP solid liquid energy out freezes (exothermic) • Ionic Solids: • Have high MP because they have very strong attractive forces. NaCl melts @ >801°C • Molecular Solids: • Low MP because they have 2 forces. The bond that makes the molecule is strong and doesnot break but the weak attractive forces between 2 different molecules does break. Ex. HCl melts @ -112°C • Not all solids Melt! • Many organic solids like wood and animals decompose.

16. - - + + - - + + - - + + Ionic and Molecular Solids Strong covalent bonds between H & O, make up the water molecules. NOT MELTING! Strong ionic bonds holding positive and negative ions together. MELTING!! Weak attractive forces between different water molecules. MELTING!! Ionic Compounds Molecular Compounds

17. Boiling and Condensation • Boiling • Phase change from a liquid to a vapor. (gas) • Boiling Point ( BP ) • Temp. at which a liquid becomes a vapor (gas). Temp. remains constant during the change! • Condensation • Phase change from a vapor to a liquid. • Condensation Point ( CP ) • Temp. at which a vapor becomes a liquid. Temp. remains constant. • energy in boils (endothermic) • BP = CP liquid gas • energy out condensed (exothermic) • Evaporation and Boiling are phase changes from a • Liquid to a vapor.

18. Evaporation and Boiling • Evaporation and Boiling are... • Cooling Processes. • Analyze Evaporation in terms of Kinetic Energy. • When liquid molecules evaporate the • FAST molecules escape. • Therefore, the • SLOW molecules are left behind. • If Average Kinetic Energy = Temperature what can you determine about the temperature of the liquid? • SLOWER MC”S are COOLER! • This is why evaporation is a • COOLING PROCESS.

19. BOILING • Analyze Boiling in terms of Kinetic Energy. • The____________under the pot ____________ • FLAME HEATS • the water, but the boiling takes away the • FAST MOLECULES so the actual boiling process • COOLS the water. • To take energy away means you are • COOLING. • This is why boiling is also a • COOLING PROCESS • EVAPORATION VS BOILING • Evaporation is a • SURFACE phenomenon • whereas boiling occurs • Throughout the whole liquid. • How can you make something evaporate quickly without heating it? • SPREAD IT OUT. Surface mc’s get enough KE to escape.

20. Condensation • Condensation is a phase change from a • VAPOR to a LIQUID. • Condensation is a • WARMING process. • When vapor hits a surface it gives up • ENERGY and makes what it hits • WARMER. • Which molecules are more prone to condense, the fast or slow? • SLOW • Why? (FLY & FLY PAPER) Slow mc’s have little KE and are attracted to other liquid mc’s on the surface. • Explain why it gets warmer outside just before it snows. • Snow is condensation. When you condensate you loose energy. That released energy makes the surrounding air warmer.

21. Sublimation • Phase change from a solid directly to a gas without turning into a liquid. • The solid has a very high VAPOR PRESSURE. • Examples: solid air fresheners, moth balls, dry ice, iodine.

22. PHASE CHANGES ALWAYS INVOLVE ENERGY CHANGES. • Energy In = Endothermic process = Cooling Process • SOLID LIQUID GAS • Energy out = Exothermic Process = Warming process • The following phase changes are cooling processes • MELTING, BOILING, EVAPORATING • Because they absorb energy leaving the surroundings cooler. (High KE mc’s leave and the slower cooler mc’s are left behind. • The following phase changes are warming processes • Freezing and Condensing • Because they release energy making the surroundings warmer (KE is absorbed which adds energy and makes the surroundings warmer.

23. Phase Diagram • Shows the relationship between solid, liquid and vapor phases in a sealed container. Each sections shows a pure phase. Equilibrium 2 phases existing at the same time at a certain temp & pressure. (line separating 2 regions.) Triple Point Only condition that allows all 3 phases to exist at the same time. (where lines intersect)

24. Gas Pressure • Atmospheric pressure – • Due to the collisions of air mc’s on an object. • Collisions of gas molecules are • Perfectly elastic which means • No energy is lost (just transferred)

25. Measuring air pressure –Barometers 760mmHg Eudiometer Below sea level Cave (2atm) Higher than sea level on Moon (0atm) At sea level (1atm)

26. Pressure Unit Equalities • 1 atm (atmosphere) • = 760 mm Hg • = 30 “ Hg • = 14.7 psi (pounds per square inch) • = 101,3 kPA (kilo Pascal) metric unit • Using these equalities, it is possible to convert from one unit of pressure to another.

27. Examples of Unit Conversions = 2343 = 2340mmHg • Convert 92.5 “ Hg to mm Hg • Equality: 760 mm Hg = 30 ” Hg • 92.5 “ Hg760 mmHg 1 30”Hg • Convert 2.25 atm to psi • Equality: 1 atm = 14.7psi • 2.25atm14.7psi 1 1atm • Convert 5.25 atm to kPA • Equality: 1atm = 101.3 kPA • 5.25 atm101.3 kPA 1 1 atm • Convert 222.2 kPA to mmHg • Equality: 101.3 kPA = 760 mmHg • 222.2 kPA760 mmHg 1 101.3 kPA = 33.075 = 33.1psi = 531.825 = 532 kPA = 1667.05 = 1667mmHg

28. Factors involved in Boiling • Boiling is affected by 2 factors: • Temperature • Pressure • vapor pressure- • Is the pressure above a liquid in a SEALED CONTAINER! • boiling point – • Temp. at which the vapor pressure of a liquid = external atmospheric pressure.

29. Diagram/Explanationof a pot of water trying to boil. 1 atm Air pressure pushing down 1atm Air pressure pushing down 1 atm Air pressure pushing down • 1. @ 70ºC air pressure greater than vapor pressure of liquid. No Boiling = No Cooling Water continues to heat up. • 2. @ 98ºC air pressure still greater than the vapor pressure of liquid. No Boiling = No Cooling Water continues to heat up. • 3. @ 100ºC air pressure = vapor pressure. Boiling occurs. Water Cools. Temperature remains constant!

30. Why does a pressure cooker cook your food faster? • Sealed lid! • As water heats up and KE increases the air pressure inside pushing down on the liquid increases. • More air pressure means liquid can not boil at 100ºC. No Boiling = No cooling • Water must heat up to a temp. > 100ºC in order for the air pressure = vapor pressure • Higher temp. boiling cooks food faster!

31. How Can You Make Water Boil at Room Temperature? • Boiling at room temp. demo. • The vacuum pump removes air mc’s. • No air mc’s = no collisions. • No collisions = no air pressure. • Therefore the temperature at which the temp. of the water = external pressure is much lower than 100ºC. • Because boiling is a cooling process the temp. of the water decreases as the water boils in a vacuum.

32. Boiling Water on a Mountain • Less air mc’s = • Less collisions = • Less air pressure = • Lower temp. for vapor pressure = air pressure • Boil water at lower temperature! • Will your food cook faster or slower? • Slower • Boiling is a cooling process it’s the heat (temp.) that’s cooking the food. • Lower Temperature = Longer Cooking Time!

33. How does Pressure and Temperature change in a Closed Container as Ice is heated, Melts and turns to Steam. 5. • 1. Ice @ 0ºC. Temp. remains constant as ice melts. Pressure inside container ↑. • 2. Ice and Water at 0ºC. Remains constant until completely liquid. Pressure inside container ↑. • 3. Water only. Temperature ↑ until vapor pressure of the liquid = air pressure that has built up in the sealed container. Pressure inside container ↑. • 4. Water and Steam at 100ºC. Remains Constant until completely steam. Pressure inside container ↑. • 5. Steam only. Temperature ↑ and Pressure inside container ↑until pressure builds up so much container explodes! 3. 2. 4. 1.

34. The End!