1. Chemistry Chapter 10 notes Physical Characteristics of Gases

2. Kinetic molecular theory of matter All matter is composed of tiny particles which are in constant motion This explains observed properties of matter

3. Kinetic molecular theory of matter (KM) Ideal gas- an imaginary gas which perfectly fits all assumptions of the kinetic molecular theory of matter IE: Ideal gas behaves exactly as a gas should, no deviations Kinetic molecular theory of gases based on 5 assumptions

4. 5 Assumptions 1. Gases consist of large numbers of particles very far apart from one another relative to their size –Most of space occupied by gases is empty space –This explains compressibility of gases and their low density (compared to liquids and solids)

5. 2. Collision between gas particles/particles and gas particles/container are elastic –Elastic collision- no net loss of kinetic energy. –KE is transferred, but total KE of 2 particles remains the same as long as the temperature is constant

6. 3. Gas particles are in continuous, rapid, random motion and therefore have kinetic energy Their KE is high enough to overcome any attractive forces between particles (except near the temperature of condensation)

7. 4. There are no forces of attraction or repulsion between gas particles When gas particles collide, they immediately bounce apart

8. 5. Average kinetic energy of gas particles depends on the temperature of the gas –For any object KE= ½ m v 2 –Where m= mass and v = velocity –All gases at the same temperature have same KE, so lighter particles (H) have higher average speed than heavier particles (O)

9. KM theory and the nature of gases Expansion –Gases fill any container and take it’s shape –Gases have no definite shape or volume Fluidity –Gas particles glide past one another –Behave much like liquids –Gases and liquids are both considered fluids

10. KM theory and the nature of gases Low density –Gases typically have about 1/1000 the density of the same substance in a liquid or solid state Compressibility –Due to their low density gases can be compressed dramatically

11. KM theory and the nature of gases Diffusion –Gases randomly mix with other particles to even distribution –Rates of diffusion depend on the speed of particles, diameter of particles and attractive forces between particles –Lighter gases diffuse more rapidly than heavier ones

12. KM theory and the nature of gases Effusion –Movement of gas particles through a tiny opening –Rates of effusion are directly proportional to the velocity of the particles

13. Real gases Do not behave completely according to kinetic molecular theory 1873 Van der Waals noted that forces between particles of gases caused deviation from ideal gas behavior Deviation is most significant at high pressure and low temperature KM theory holds truest in gases with little attraction between particles (ex. Noble gases)

14. Pressure When describing a gas you must specify characteristics: Volume, temperature, number of molecules and pressure You’ve got the first 3! Pressure is force per unit area on a surface or Pressure = force/area

15. Atmospheric pressure Pressure exerted by gases of the atmosphere At sea level approximately 10.1 N/cm 2 Barometers are used to measure atmospheric pressure Oldest barometer- mercury column measurement expressed in mm of Hg –Normal atmospheric pressure at sea level and 0°C = 760 mm Hg = 1 atmosphere

16. Pressure! SI units for pressure are derived 1 Pascale (Pa) = 1 Newton / meter 2 Pressure often expressed in kilopascals (kPa) 1 atmosphere = 1.01325 x 10 5 Pa (or 101.325 kPa) See table 10-1 on p. 311

17. STP Standard Temperature and Pressure are needed to compare gas volumes STP = 1 atmosphere and 0°C

18. Gas Laws Boyles law Relates pressure and volume of a gas at constant temperature Pressure and volume are inversely proportional PV = k or V=k1/P K is a constant for a given sample of gas

19. Boyles and changing pressure Because k is a constant for a given sample of gas and we know that the product of pressure and volume will always equal k P 1 V 1 = k and P 2 V 2 = k we can set P 1 V 1 equal to P 2 V 2 P 1 V 1 = P 2 V 2 and solve for any one of the 4 values

20. Charles Law Relates temperature and volume of gases at constant pressure 1787 Charles found that volume of a gas changes 1/273 of original volume for each 1°C change in temperature (with a starting point of 0°C and at constant pressure)

21. Charles and absolute zero Kelvin 0 = -273.15°C K= °C + 273.15 This is useful because it is directly proportional to gas volume Charles law: Volume of a fixed sample of gas at constant pressure varies directly with Kelvin temperature

22. Charles… V/T= k or V = kT K is a constant based on quantity of gas and pressure Same thing can be done with Charles for changing volume or temperature as was done with boyles for changing pressures V 1 /T 1 = V 2 /T 2

23. Gay-Lussacs Law Relates pressure and temperature of a gas at constant volume P/T = k or P= kT K is a constant depending on quantity and volume of gas P 1 /T 1 = P 2 /T 2 useful when faced with changing pressures and temperatures

24. Combined gas law Merges three laws just mentioned PV/T = k k is a constant related to the amount of gas P 1 V 1 /T 1 = P 2 V 2 /T 2 if any one quantity is unchanging one of the other gas laws can be derived

25. Daltons combined pressures The total pressure of a mixture of gases is the sum of the individual pressure of each gas alone P T = P 1 + P 2 + P 3 … This can be used no matter how many gases are in combination

26. Law of Combined Pressures Is useful when dealing with gases collected over water Gases collected this way are mixed with water vapor, this exerts water vapor pressure

27. To measure pressure of gas and water vapor in collection bottle, raise bottle until water level in and out are same. At that point pressure inside bottle = atmospheric pressure P atm = P gas + P H 2 O Obtain atmospheric pressure from barometer in lab and subtract water vapor pressure at given temp (from table A8 in book)