1. STATES OF MATTER:  Kinetic Theory- the tiny particles in all forms of matter are in constant motion.  Kinetic Energy: energy of motion  SOLIDS:   The particles in solids:  Are packed together (often in an organized pattern called a crystal lattice)  Are held together by strong forces (therefore have high melting points)  Vibrate about fixed points

2. Crystal  Solid with atoms, ions or molecules arranged in an orderly, repeating 3-D pattern.  Unit Cell- smallest group of particles within a crystal that retains the shape of the crystal  Crystals are classified in to 7 systems- most of which are square cubic or rectangular cubic.

3. Allotropes-Allotropes-  Example: Solid Carbon  GRAPHITE DIAMOND BUCKEY BALL  Oxygen; O 2 gas or O 3 gas (ozone) A molecular form of an element that exists in 2 or more different forms in the same physical state PENCIL Widely spaced, linked, hexagon. Weak bonds - soft Each CARBON atom is strongly bonded to 4 other carbons hard 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  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. Liquids  Most liquids are polar molecular compounds. (polar: to have + and – areas)  The particles in a liquid:  Are packed together (almost as closely as a solid)  Are held together by weak attractive forces*(therefore have low melting points)  Slide past each other (vibrate and spin in fixed positions.

7. Intermolecular Forces  Attractive forces between molecules  The positve end of one molecule is attracted to the negative end of an adjacent molecule.  Weakest: London dispersion forces  Stronger; dipole – dipole interactions  Strongest; Hydrogen- bonding

8. GASES: Most gases are non-polar molecular compounds  The particles in gases:  Are very far apart (gases are about 99.99% empty space  Have no attractive forces (therefore have extremely low melting points.)  Move very rapidly and in constant random straight line motion.  Diffuse from areas of high concentration to areas of low concentration (smaller lighter molecules diffuse at a faster rate than larger heavier molecules)

9. Perfectly Elastic Collisions  Collisions between gas molecules in which no energy is lost  (energy is transferred from one particle to another)

10. Summation of States of Matter SolidsLiquidsGases Attractive Forces Forces b/t particles Molecular Packing Molecular Motion characteristicsIncompressible Definite shape Definite volume Nearly incompressible Indefinite shape Definite volume Compressible Indefinite shape Indefinite volume Very Strong NONE Bounce off each other weak Very tight, close 99% empty space Very far apart Closely packed Vibrate in place Slide past each other Rapid, random straight line motion Ave. speed = 1000mph

11. Plasma  1. occurs at extremely high temperatures. (millions of degrees Celsius) KE becomes great enough to break molecules into atoms.  2. At these temperatures electrons have been removed from the gaseous atoms.  3. The resulting fluid of bare nuclei (+ ions) and free electrons is called plasma.  COLD PLASMA – 50,000K to 1,000,000K  HOT PLASMA- “stars” 10,000,000 to 1,000,000,000K

12. 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. state of matterbosonspotentialtemperaturesabsolute zeroKsupercooled quantum state of matterbosonspotentialtemperaturesabsolute zeroKsupercooled quantum state  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. Satyendra Nath BoseSatyendra Nath Bose  Seventy years later, the first such condensate was produced by Eric Cornell and Carl Wieman in 1995 at the University of Colorado at Boulder NIST-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]. Eric CornellCarl WiemanUniversity of Colorado at BoulderNISTJILArubidiumnanokelvin[1]KCEric CornellCarl WiemanWolfgang KetterleMITNobel Prize in Physics[2] Eric CornellCarl WiemanUniversity of Colorado at BoulderNISTJILArubidiumnanokelvin[1]KCEric CornellCarl WiemanWolfgang KetterleMITNobel Prize in Physics[2]  When a system of atoms is cooled rather than bosons, the Bose- Einstein condensate is then sometimes called a Super Atom.[3] atomsbosonsSuper Atom[3]atomsbosonsSuper Atom[3]

13. KINETIC ENERGY AND TEMPERATURE  TEMPERATURE AND ENERGY ARE NOT THE SAME THING!!!  Temperature is a measure of AVE. Kinetic Energy  The higher the Temp, the greater the Kinetic Energy  Kelvin temperature is directly proportional to Average Kinetic Energy.  Celcius vs Kelvin Scale; Kelvin = °C + 273  A 1° increment on the Kelvin scale = 1° on Celcius  3x Kelvin Temperature = 3x Kinetic Energy  Absolute Zero = temp. all molecular motion stops

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

15. Pressure  Gas pressure – collisions of gas particles with the surface of an object.  Atmospheric Pressure – collision of “air molecules with the surface of an object.  Barometer – instrument  Barometer – instrument measures the height of a column of mercury supported by air pressure

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

17. Boiling point – temperature at which the vapor pressure of a liquid = external pressure Can boil by: increasing temperature  increases KE of mc’s or decreasing external pressure  so mc’s already have enough KE As elevation increases, atmospheric pressure decreases so boiling point decreases.

18. CHANGES OF STATE – PHASE CHANGES are phase changes that ALWAYS involve energy changes. CHANGES OF STATE – PHASE CHANGES are phase changes that ALWAYS involve energy changes. Energy in = Endothermic (cools surroundings) (melting) evaporation or boiling SOLID LIQUID GAS (freezing) condensation Energy out = Exothermic (warms surroundings) Melting, evap. and boiling are cooling processes Freezing and condensation are warming process

19. Evaporation VS Boiling sublimation – the conversion of a solid to a gas without passing through the liquid state Substances with very weak intermolecular forces are unable to hold molecules together so they are able to spread apart, preventing them from having a liquid phase, and forcing them to sublime.

20. Phase Diagram During a phase change there is no temperature change because all of the heat energy is being converted into kinetic energy as the motion of the molecules increases.

21. 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)