1. Gas Laws Chapter 14

2. To talk about how gases act in different situations, we must understand: The higher the temp, the faster molecules move The faster the particles move, the farther away each other they become Pressure is defined as the number of times the particles of the gas hit the walls of the container.

3. Boyles Law Robert Boyle-Irish Scientist Volume-measured in L Pressure measured in atm, kPa, or mmHg 1 atm=101.3 kPa= 760 mmHg Volume is inversely proportional to pressure P 1 V 1 =P 2 V 2

4. Why? If the temperature is constant, then the particles are moving at the same speed. However, if the volume is decreasing, then the particles have LESS room, so they hit the walls of the container wall more often, therefore the pressure goes up. EX: If 40 mL of a gas is at 760 mmHg, what will the volume be at a pressure of 800 mmHg?

5. What would a graph of Boyle’s Law look like?

6. Charle’s Law Effects of temperature change on volume at constant pressure. At constant pressure, the volume of a gas will increase in direct proportion to the temperature (in K).

7. Why? As the temperature goes up, particles move faster, so they hit the walls of the container more often, so the pressure going outward increases. Since pressure going inward stays the same (atmospheric), the volume increases. EX: At constant pressure, what will the new volume of 50 mL of a gas be if the temp is increased from 27 degrees C to 127 degrees C?

8. What would a graph of Charles Law look like?

9. Gay-Lussac’s Law At constant volume, pressure and temperature (K) of a gas is in direct proportion. As temperature increases, pressure increases.

10. What would a graph of Gay-Lussac’s Law look like?

11. Example Gas at a constant volume has a pressure of 380 mmHg. The temperature is increased from 50 degrees C to 100 degrees C. What is the new pressure?

12. Combined Gas Law

13. STP=Standard Temperature and Pressure – Standard Temperature =273K – Standard Pressure = 1 atm EX: If we measure a gas at a volume of 100 mL, a pressure of 788 mmHg, and a temperature of 21 degrees C, what would the volume be at STP?

14. A nice trick PTV

15. Ideal Gas Law P=Pressure (must be in atm) V=Volume (must be in L) n=moles of a gas R=ideal gas constant R=.0821 Latm/molK T=Temperature (must be in K)

16. Kinetic Theory of Gases All gases are composed of separate, tiny (invisible) particles called molecules. Gas molecules are in constant, rapid, straight- line motion, and therefore posess kinetic energy. These molecules collide with each other and the walls of the container. Kinetic energy does NOT change because of a collision…..this is called an elastic collison.

17. Kinetic Theory Cont. The molecules of a gas display no attraction or repulsion for one another. The average kinetic energy is directly proportional to the temperature of the gas.

18. Ideal Gas vs. Real Gas A gas that exactly conforms to KMT is called a perfect or ideal gas. No gas exists for which ALL of these assumptions are true. Why? – Gas particles do attract and repulse on another – Gas particles do have mass and volume. REAL GASES DIFFER MOST FROM AN IDEAL GAS UNDER LOW TEMPS AND HIGH PRESSURES.

19. Dalton’s Law of Partial Pressures The total pressure of a system is the result of all the INDIVIDUAL pressures of EACH gas in that system. EX: In the air around us, the gases in most abundance are Ar (.15 atm), N 2 (.7 atm), O 2 (.10 atm), and CO 2 (.05 atm). The total pressure, at sea level, is 1 atm.

20. Dalton’s Law Cont. Equation:

21. Grahams Law Diffusion vs. Effusion – Diffusion-movement of a gas from higher concentration to low concentration – Effusion-escape of a gas from a container through a hole

22. Graham’s Law of Effusion In general, gases of lower molar mass diffuse and effuse faster than gases of higher molar mass. Grahams Law of Effusion states that the rate of effusion of a gas is inversely proportional to the square root of the gas’s molar mass. Grahams Law can be written as follows for two gases, A and B: