1. Chapter 10 Physical Characteristics of Gases

2. 10.1 The Kinetic-Molecular Theory of Matter In the late 19 th century the Kinetic-Molecular Theory was developed to describe the behavior of matter (solid, liquid, gas). Theory is based on the assumption that particles of matter are ALWAYS in motion.

3. The Kinetic-Molecular Theory of Gases This model describes a “ideal gas”: an imaginary gas that perfectly fits the assumptions of the kinetic-molecular theory. 5 assumptions: 1.Gases consist of large #s of tiny particles that are FAR apart relative to their size.

4. 2. Collisions between gas particles & the container wall are ELASTIC collisions. No net energy is lost. 3. Gas particles are in constant, rapid, random motion. 4. There are NO forces of attraction/repulsion between the gas particles. 5. The kinetic energy of gas particles depends on the temperature. Higher temp = more KEKE=1/2 mv 2 Lower temp = less KE http://www.chm.davidson.edu/vce/kineticmoleculartheory/PT.html

5. The Kinetic-Molecular Theory & the Nature of Gases Expansion: Because gases have no definite shape, it will spread and take the shape of the container. Fluidity: Because the attractive forces are negligible between gas particles, they can flow over one another. Therefore, gases are considered a fluid.

6. Low Density: Because the gas molecules are “far” apart, they have extremely low densities. Compressibility: Because of the distance between gas molecules, we can force the molecules closer together- compress them.

7. Diffusion: Gases can spread out and mix. The mixture will be spontaneous & uniform. Effusion: When a gas escapes a container though small holes.

8. Deviations of REAL Gases from Ideal Behavior At high pressures and low temperatures… real gases do not behave like an Ideal gas. – At high pressures, the distance between the molecules becomes much smaller thus, the volume of the gas molecules becomes important. – At low temperatures, the gas molecules slow down and now become attracted to each other. We no longer have perfectly elastic collisions.

9. 10.2 PRESSURE To describe gases fully, you need 4 quantities: – Volume (L, mL) – Temperature (K) – # of molecules (moles/Avogadro’s #) – Pressure Pressure- the force per unit area on a surface Pressure = Force/ area SI unit- pascal (Pa)

10. Atmospheric pressure- pressure exerted by the column of air in the atmosphere.

11. Measuring Pressure Barometer- instrument used to measure atmospheric pressure. You want to measure the height of the mercury in the tube. Usually as mmHg.

12. Units of Pressure There are several different units of pressure that are used in labs around the world today. 1 PaPascal = 1 N/m 2 (The SI unit) 1 psi= 1 lb/in 2 (English unit) = 6,891 Pa 1 Bar= 100,000 N/m 2 = 100 kPa ≈ 1 atm 1 Torr= 1 mm Hg = 133.3 Pa 1 atm= 101.3 kPa = 760 mm Hg = 29.92 in Hg = 14.70 psi = 760 torr = 101,325 Pa

13. Standard Temperature and Pressure STP To make it easier to study and describe gas behavior we assume the gas is at Standard Temperature and Pressure. – 273 K or 0°C – 1atm or 760 mmHg or 101,325 Pa

14. 10.3 The Gas Laws Gas Laws: simple mathematical relationships btwn the volume, temperature, pressure and amount of gas. There are 4 scientists that made great strides in understanding gas behavior. 1. Robert Boyle 2. Jacques Charles 3. Joseph Gay-Lussac 4. John Dalton

15. Boyle’s Law: Pressure-Volume Relationship English Chemist/Physicist investigated the relationship between pressure and volume. inversely Boyle’s Law- the volume of a fixed mass of gas varies inversely with the pressure at constant temperature. P 1 V 1 = P 2 V 2

16. Example Problem A gas at a pressure of 608 mmHg is in a 545 cm 3 volume container. The volume is increased to 1065 cm 3 with no change in temperature. What is the new pressure? P 1 V 1 = P 2 V 2 (608 mmHg) × (545 cm 3 ) = (P 2 ) × (1065 cm 3 ) P 2 = (608 mmHg) × (545 cm 3 ) (1065 cm 3 )

17. Charles’s Law: The Volume - Temperature Relationship A French physicist who used a cylinder with a moveable piston to test the relationship between Temperature and Volume. directly Charles’s Law: the volume of a fixed mass of gas at constant pressure varies directly with the Kelvin temperature. V 1 T 2 = V 2 T 1

18. Practice Problems What will the volume of a gas be at 355K if its volume at 273K is 8.57L? Assume pressure remains unchanged. V 1 T 2 = V 2 T 1 (8.57L) × (355K) = V 2 × (273K) V 2 = (8.57L) × (355K) (273K)

19. Gay-Lussac’s Law: Pressure – Temperature Relationship French Chemist & Physicist. directly Gay-Lussac’s Law: The pressure of a fixed mass of gas at constant volume varies directly with the Kelvin temperature. P 1 = P 2 T 1 T 2

20. Practice Problem At 120. °C, the pressure of a simple of nitrogen is 1.07 atm. What will be the pressure at 205 °C, assuming constant volume? P 1 T 2 = P 2 T 1

21. The Combined Gas Law The Combined Gas law: the relationship between pressure, volume and temperature of a fixed amount of gas. P 1 V 1 = P 2 V 2 T 1 T 2

22. Practice Problems The volume of a gas is 27.5 mL at 22.0°C and 0.974 atm. What will the volume be at 15.0 °C and 0.993 atm? P 1 V 1 T 2 = P 2 V 2 T 1

23. Dalton’s Law of Partial Pressure The sum of the partial pressures of all the components in a gas mixture is equal to the total pressure of the gas mixture. P T = P a + P b + P c + …

24. Practice Problem What is the atmospheric pressure if the partial pressures of N 2, O 2 and Ar are 604.5 mmHg, 162.8 mmHg and 0.5 mmHg respectively? P T = P a + P b + P c + … P T = 604.5 + 162.8 + 0.5