1. Phases of Matter Chp 13 and 14

2. Phases of Matter  Solid – molecules are held tightly together by intermolecular forces, molecules move slowly  Liquid – some intermolecular forces still exist, but they are becoming weaker as molecules speed up  Gas – all intermolecular forces have been broken, molecules move very quickly and rarely interact  It take ENERGY to break those intermolecular forces and cause a phase change!!!

3. Phase Changes  Evaporation – liquid to gas, is a cooling process since hottest molecules escape (why we sweat) –Boiling – evaporation throughout a substance, it can be decreased as air pressure decreases  Condensation – gas to liquid, is a warming process since hot gas molecules bring their energy with them  Melting – solid to liquid  Freezing – liquid to solid  Sublimation – solid to gas, skipping liquid phase –This occurs especially at low pressures since the boiling point becomes lowered so much

4. Intermolecular forces  Exist between molecules  3 types: –Dipole-dipole: occurs between polar molecules –Hydrogen bond: special dipole that occurs when hydrogen is involved (stronger than regular dipole- dipole interactions) –London dispersion force: occurs for noble gases and nonpolar molecules, very weak and doesn’t last long  The stronger the intermolecular forces involved, the higher the boiling point since more energy is needed to break those interactions.

5. Heating/Cooling Curves  Graph that represents at what temperature phase changes occur and how heat and temperature are related  Note, the temperature doesn’t change during a phase change (boiling water stays at 100 o C until is becomes a gas) –Heat of fusion – energy needed to turn a solid into a liquid (or released when a liquid turns into a solid)  6.02 kJ/mol for water –Heat of vaporization – energy needed to turn a liquid into a gas (or released when a gas turns into a liquid)  40.6 kJ/mol for water –HOV is always higher than HOF because it takes more energy to completely break the intermolecular forces

6. Using HOV and HOF  Ex. How much energy is required to melt 8.5 g of ice at 0 o C? 8.5 g1 mol6.02 kJ =2.8 kJ 18 g1 mol

7. Combining temperature and phase changes  Ex. How much energy is required to heat 25 g of water from 25 o C to 100 o C as steam? –2 parts: heat 75 o C, then hov to get steam –Part 1 Q = sm  T Q = 4.184 (25)(75) = 7800 J or 7.8 kJ –Part 2 25 g1 mol40.6 kJ= 57 kJ 18 g1 mol –Now, we add since both must occur 7.8 kJ + 57 kJ = 64.8 kJ

8. Pressure  All gases, including the atmosphere, exert a pressure –Result of gravity pulling on the gas’s mass –760 mm Hg at sea level on Earth  Has a variety of units –1 standard atmosphere = 1 atm = 760 mm Hg = 760 torr = 101,325 Pa = 14.69 psi  We can convert between units using a T- chart –We typically use pascals in science

9. Unit conversion  Ex. What is 7.3 atm in mm Hg? 7.3 atm760 mm Hg= 5548 mm Hg 1 atm  Ex. What is 7.3 mm Hg in Pascals? 7.3 mm Hg101325 Pa=973.2 Pa 760 mm Hg

10. Boyle’s Law  As the volume of a gas is decreased, it’s pressure increases  P 1 V 1 = P 2 V 2  Ex. If a 1.5 L sample of freon gas at 56 torr is increased to a pressure of 150 torr, what is the new volume of the gas? 56 (1.5) = 150 (V) V = 0.56 L *Any units are fine, as long as they are the same on both sides of the equation.

11. Charles’ Law  As the temperature of a gas increases, it’s volume increases (temp must be in K)  At -273 o C, all gases occupy a volume of 0, which is impossible, so that is coldest temp possible (absolute zero)  V 1 = V 2 T 1 T 2  Ex. If a 2 L gas sample at 298 K is cooled to 278 K, what is it’s new volume? 2 = V 2 298 278 298 278 V 2 = 1.9 L

12. Combined Gas Law  Stick charles’ law and boyle’s law together  Pressure can be in any unit, as long as it’s the same on both sides, but temperature must be in kelvin  P 1 V 1 = P 2 V 2 T 1 T 2 T 1 T 2

13. Avogadro’s law  As the number of moles of gas increases, so does the volume  V 1 = V 2 n 1 n 2  If 0.5 mol of oxygen occupy 12.2 L, what volume would 0.33 mol of oxygen occupy? 12.2 = V 2 0.5 0.33 0.5 0.33 V 2 = 8.1 L

14. Ideal Gas Law  Combines all gas laws into one  PV = nRT R = 0.08206 Latm/molK *so, temp must be in kelvin, volume must be in liters and pressure must be in atm  Ex. A 8.56 L sample of hydrogen gas at 0 o C has a pressure of 1.5 atm. How many moles are present in the sample? 1.5 (8.56) = n (0.08206) (273) n = 0.57 mols

15. Some Terms  STP – standard temperature at pressure –1 atm and 0 o C  Molar volume – the volume that one mole of any gas takes up at STP –22.4 L at STP  Kinetic molecular theory – predicts why gases have the properties that they do –As temp increases, particles move faster and collide with the container more often and interact with each other less

16. Dalton’s Law of Partial Pressure  When gases are mixed, their pressures add together to create a new pressure in the container –Occurs because gases move so fast, what the particle is doesn’t matter as much as how many there are  P total = P 1 + P 2 …  P total = n total (RT) V

17. An example  If 12 g of oxygen and 46 g of He are pumped into a 5.0 L container at 25 o C, what will the total pressure in the container be? –First g convert to moles using molar mass 12 g 1 mol = 0.375 mol 46 g 1 mol = 11.5 mol 32 g 4 g 32 g 4 g –Next o C becomes K by adding 273 25 + 273 = 298 K –Then plug into formula P total = (.375 + 11.5) (0.08206)(298) 5 P total = 58.1 atm