Hello! Welcome to your 2 nd semester of Chemistry! You will know your grade when you get your report card 1. Sit in the same seat. I will make changes as needed --- maybe a new seating chart? 2. All the rules are the same as last semester. All class documents are on Sharepoint. 3. Did you get your new packet? You’ll need it out today…. 4.“How can I get a better grade this semester?” a. Do practice problems WITH ME, don’t WATCH ME. b. Do homework IN CLASS if there’s time. LOOK BACK ON YOUR NOTES and ASK YOUR NEIGHBORS and don’t waste class time chatting. c. Use your brain. If you can simply focus, you’ll realize that chemistry isn’t really that hard. But you DO need to use your brain.

1) Which of the choices shows the 3 phases of matter correctly ranked by amount of kinetic energy? a. liquid < gas < solid b. solid < liquid < gas c. solid > liquid > gas d. gas > solid > liquid 2) Which of the choices shows the 3 phases of matter correctly ranked by density? a. liquid < gas < solid b. solid < liquid < gas c. solid > liquid > gas d. gas > solid > liquid Warmup (2 minutes)

Properties of Gases

IM Forces keep molecules “stuck together” Kinetic Energy: energy due to the motion of an object

K.E. and IMFs: 3 States of Matter Solids have: - molecules which are tightly packed; strong IM forces - low amount of kinetic energy; can only ‘vibrate’ Liquids have: - molecules which are tightly packed but IM forces stretched/broken - medium amount of kinetic energy; molecules “flow” Gases have: - no IM forces between molecules, which are free to move independently of one another - TONS of kinetic energy; molecules move wherever

The Kinetic Molecular Theory of Gases makes 5 assumptions about ideal gas behavior: 1. A gas is considered to be composed of tiny hard spheres 2. Molecules are far enough apart that we can ignore their volume. 3. Gas molecules have a lot of kinetic energy and are constantly in motion

5. There are NO forces of attraction or repulsion between gas particles because they move quickly in straight lines 4. No energy is lost when particles collide with container walls or each other

6 Physical Properties of Gases Low Density –molecules have mass in a larger amount of volume Effusion – movement of molecules through a tiny opening. Expansion –molecules expand in volume to fill a larger space. Compression –Volume can be decreased to fill a smaller space. Fluidity –Molecules flow past each other without getting stuck together.

Diffusion: the tendency for a molecule to move from an area of high to low concentration Molar Mass (g/mole) Diffusion Rate (m/s) Make a rough graph of this data to see the relationship between molar mass and diffusion rate

H2H2 Ne H2OH2O N2N2 O2O2 CO 2 Grahm’s Law: molecules of low molar mass diffuse more rapidly than molecules of greater molar mass

Heat Gas Temp: Pressure: Temp: Pressure: Low High Low Pressure is caused by collisions of the molecules with the sides of a container. The more often molecules of air strike a single spot, the more pressure is applied there!

How exactly do we use a barometer to measure atmospheric pressure? 1 atm = 760 mm Hg 760 mm 1 atm Pressure Dish of Mercury Column of Mercury

Units of Pressure  atm (atmospheres)  mm Hg (millimeters of Mercury)  kPa (kiloPascals) 1 atm = 760 mm Hg = 101.3kPa Silly Suzy and Bozo Joe are arm wrestling! Suzy exerts a pressure of 1890 mmHg. Joe exerts a pressure of 140 kPa. Who will probably win? 1890 mmHg (101.3 kPa) = 252 kPa (760 mmHg)

Temperature measures the average KE Faster molecules, higher temperature. 1. If you change temperature from 300 K to 600 K, what will happen to the KE of the sample? The kinetic energy doubles. Average KE of a sample is directly proportional to the temperature in Kelvin 2. If you change temperature from 300ºC to 600ºC, what will happen to the KE? KE doesn’t double: 873 K is NOT twice 573 K 3. At what temperature would molecular motion stop? At 0 K (or -273 ⁰C), the KE = 0 Joules Kinetic Energy and Temperature

STP Standard Temperature (273 K or 0˚C) and Pressure (101.3 kPa or 1 atm or 760 mmHg Temperature and air pressure can vary from one place to another on the Earth, and can also vary in the same place with time. It is necessary to define standard conditions for temperature and pressure: