1. Gases Liquids and Solids

2. Kinetic Molecular Theory of Matter 1) All matter is composed of small particles 2) The particles are in constant motion and therefore possess kinetic energy 3) The particles can have attractions and repulsions between them 4) The kinetic energy of the particles increases as temperature increases 5) The particles collide elastically and can transfer energy to each other during these collisions

3. Consider the phases of matter in light of the kinetic molecular theory (kinetic energy verses attractions)

4. Pressure: a measurement in the gas phase Pressure is the force applied to a given area Pressure is the force applied to a given area P = F / A P = F / A Pressure can increase by increasing the force or decreasing the area Pressure can increase by increasing the force or decreasing the area Units: Units: mm Hg atmospheres (atm) torr

5. Units of Pressure 1 mm Hg / 1 torr 1 atm / 760 mm Hg 1 atm / 760 torr 1 atm is a standard unit of pressure and corresponds to pressure at sea level A Barometer

6. Boyle’s Law At constant temperature As volume goes down, pressure goes up (Think about squeezing a balloon) P 1 x V 1 = P 2 x V 2 Pressure and volume can have any units as long as they are the same on both sides

7. Boyle’s Law in Practice Filling a syringe Increasing the volume of the syringe decreases its pressure so the liquid flows in This is how inhaling works too. Expanding your lungs creates more volume and less pressure so air from the room rushes in

8. Charles’s Law At constant pressure As temperature goes down, volume goes down (Think about tires in the winter and summer) Volume can have any units as long as they are the same on both sides. Temperature must be Kelvin (K = C + 273) V1V1 V2V2 T1T1 T2T2 =

9. Combined Gas Law If you know this… you know both Boyle’s and Charles’s These types of problems give one set of conditions… (V 1, T 1 etc….) then a new set of conditions…. (V 2, T 2 etc….) And you have one single variable to solve for P 1 V 1 P 2 V 2 T 1 T 2 =

10. Try some A gas has a volume of 250.0 mL at 1.75 atm of pressure. If the pressure is changed to 765 mm Hg, what is the new volume in mL? A gas occupies 500.0 L at 1.00 atm when the temperature is 61 C. In August the temperature climbs to 99 C and the gas is in the same container with the same volume. What is it’s new pressure?

11. Ideal Gases Ideal gases: Have perfectly elastic collisions with no attractions at all when colliding Occupy no volume compared to the volume of the container Real gases however do have volumes and do have attractions When are real gases like ideal gases? Under conditions of large container volume and/or high temperature.

12. Ideal Gas Law PV = nRT P - pressure in atm V – volume in L n – number of moles of gas T – temperature in Kelvin R – ideal gas constant 0.0821 L ∙ atm / mol ∙ K These problems have only one set of conditions and you solve for one variable

13. Keep in mind….. Pressure – may need to convert mm Hg or torr into atm Volume – may need to convert into liters n – may need to convert grams into moles Temperature – may need to convert celsius into Kelvin

14. Try this…. How many liters would be occupied by 12.05 moles of H 2 gas that is at 125° C and exerts 775 mmHg of pressure?

15. Daltons Law of Partial Pressures In a mixture of gases, the total pressure is the sum of the pressures of the individual gases P T = P 1 + P 2 + P 3 …..

16. Phase Changes

17. The boiling process Volatility – how easily something evaporates Boiling point – the temperature at which a liquid boils Vapor pressure – the pressure exerted by a vapor as it evaporates from a liquid Boiling occurs when: vapor pressure = surrounding pressure

18. Boiling point of water at different locations (different pressures)

19. Intermolecular forces Ion – IonStrongest Ion - Dipole Dipole - Dipole London dispersionWeakest What is present in each of these?