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2. Gases & Atmospheric Chemistry Gases; a unique state of matter following their own laws and displaying interesting chemical behaviour http://www.youtube.com/watch?v=Zz95_VvTxZM

3. Gases Are Special  State of Matter: Gases can be compressed, solids & liquids cannot  Kinetic Molecular Theory: All particles of solids, liquids & gases display constant random motion 3 SolidsLiquidsGases Types of Motion VibrationalVibrational, Rotational & Translational Strength of Attraction StrongestIntermediateWeakest Organization of Particles Highly OrganizedIntermediateVery Low

4. Temperature  Measures the average kinetic energy of particles in a substance  Kinetic energy is the energy of movement  The temp. of a gas, greatly affects its behaviour  Measured in Kelvins (K) 0 o C = 273 K 4

5. Pressure  The force exerted on a surface, per unit of area  The standard unit (SI) of pressure is the Pascal (Pa)  1 Pa = 1 N/m 2, (1 Newton of force exerted over a 1 m 2 surface) 5

6. Atmospheric Pressure  The pressure of the large mass of air pressing down on the surface of the Earth is called Atmospheric pressure  Standard atmospheric pressure at sea level is 101,325 Pa, because this is such a large number, we usually express it as 101.325 kPa (kilopascals)  STP: standard temp & press;  0 o C & 101.325kPa  SATP: standard ambient temp & press;  25 o C & 100 kPa 6

7. Units of Pressure  Because 101.3 kPa is standard, we can say 101.3 kPa = 1 atm (atmosphere)  Other units of pressure include:  Millimetres Mercury; 760 mmHg = 1 atm (used in biology)  Torr; 760 torr = 1 atm (used in physics)  Pounds per Square Inch; 1 atm = 14.696 psi (used in industry) 7

8. Boyle’s Law  At any constant temperature, the multiplication product of the pressure and the volume of any size sample of any gas is a constant.  To express it mathematically, we use the equation: P 1 V 1 = P 2 V 2  The pressure and the volume are inversely proportional; as the pressure increases the volume of the sample of gas must decrease 8

9. Boyle’s Law  P vs. V  P vs. 1/V 9

10. Charles’ Law  At constant pressure, the mathematical product of the temperature and the inverse volume of any size sample of any gas is a constant  To express it mathematically, we use the equation: V 1 T 2 = V 2 T 1  The pressure and the volume are directly proportional; as the temperature increases the volume of the sample of gas must also increase 10

11. Charles’ Law  V vs. T 11

12. Gay-Lussac’s Law  At constant volume, the mathematical product of the temperature and the inverse pressure of any size sample of any gas is a constant  To express it mathematically, we use the equation: P 1 T 2 = P 2 T 1  The pressure and temp are directly proportional; as the temp increases the pressure of the sample of gas must also increase 12

13. Gay-Lussac’s Law  P vs. T 13

14. The Combined Gas Law  Encompasses Charles + Boyle + Gay-Lussac together for a constant amount of gas  To express it mathematically, we use the equation: P 1 V 1 T 2 = P 2 V 2 T 1  Keeping P, V or T constant is difficult to do in the lab. The Combined Law allows us to bypass this 14

15. Avogadro’s Law  The volume of gas is directly proportional to the amount of gas present  Example: 1 mole of O 2 will occupy the SAME volume as 1 mole of CO 2, under the same conditons of pressure and temperature  To express it mathematically, we use the equation: V 1 n 2 = V 2 n 1  Any 1 mol sample of gas occupies 22.4 L at 0 o C and 1 atm pressure (STP)  Any 1 mol sample of gas at occupies 24.8 L at 25 o C and 100 kPa pressure (SATP) 15

16. Ideal Gas Law  All gases, no matter the chemical, show remarkably similar properties  Pressure, volume, temperature and molar amounts of gas yield the following: PV/nT = R  R is the universal gas constant: when using kPa, R = 8.3143510 kPa L/mol K when using atm, R = 0.08206 L atm/mol K  This equation is the most important, because it allows us to use easily measureable values (P,V & T) to determine molar amounts (n) 16

17. Dalton’s Law of Partial Pressure  When Dalton was conducting his atomic theory studies, he also included studies of the behavior of gases in 1803.  For a mixture of gases in a container, the total pressure exerted is the sum of the pressures that each gas would exert if it were alone. This law can be expressed in equation form as: p = p1 + p2 + p3 +... where p is the total or measured pressure and p1, p2,... are the partial pressures of the individual gases  For air, an appropriate form of Dalton's law would be: p(air) = p(N 2 ) + p(O 2 ) + p(CO 2 ) +... 17

18. Composition of Dry Air at Sea Level ComponentMole PercentMolar Mass N2N2 78.08 28.013 O2 O2 20.94831.998 Ar0.93429.948 CO 2 0.0314 44.010 Ne 0.001818 20.183 He 0.0005244.003 CH 4 0.00216.043 Kr0.00011483.80 H2H2 0.000052.016 N2ON2O0.0000544.013 Xe0.0000087131.30

19. Gas Reactions  Because of Avogadro’s Law, reactions with gases are easy to work with in terms of stoichiometry 2 CO(g) + 1 O 2 (g)  2 CO 2 (g) Ex: If we start with 65.0L of CO; because of the 2 CO: 1 O 2 ratio, we can easily predict that 32.5L of O 2 will be required to fully react 19