Physical Characteristics of Gases

1. Physical Characteristics of Gases Chapter Ten

2. Kinetic-Molecular Theory of Matter • based on the idea that particles of matter are always in motion • explains the properties of solids, liquids, and gases in terms of the energy of particles and the forces that act between them

3. Kinetic-Molecular Theory of Gases • helps understand the behavior of gas molecules and the physical properties of gases • provides a model of an ideal gas Ideal Gas: an imaginary gas that perfectly fits all the assumptions of the kinetic-molecular theory

4. Five Assumptions 1. Gases consist of large #s of tiny particles that are far apart relative to their size. -volume 1000 times greater than that of liquid or solid; much of volume is empty space; lower densities; easily compressed 2. Collisions between gas particles and between particles and container walls are elastic. Elastic collision – one in which there is no net loss of kinetic energy. Kinetic energy is transferred between particles during collisions.

5. 3. Gas particles are in continuous, rapid, random motion. -they therefore possess kinetic energy 4. There are no forces of attraction or repulsion between gas particles 5. The average KE of gas particles depends on the temp of the gas. Speeds and energies increase with temp increase, and decrease with temp decrease. ALL gases at the same temp have the same avg KE. *Many gases behave nearly ideally if pressure is low and temp is high.

6. Properties of Gases • Expansion • No definite shape or volume • Completely fill any container and take its shape • No significant attraction or repulsion • Fluidity • Because attractive forces are insignificant, gas particles glide easily past one another. This makes them like liquids. Because gases and liquids flow, they are both referred to as fluids. • Fluids – gases and liquids that flow

7. Properties cont. • Low Density • Particles are very far apart in gases, thus they have very low densities (~1/1000 that of same substance as liquid or solid) • Compressibility • During compression, gas particles that are very far apart are crowded closer together. The volume of a given sample of a gas can be greatly decreased.

8. Properties • Diffusion • Spontaneous mixing of the particles of 2 substances caused by their random motion • Rate of diffusion depends on 3 properties of the gas particles: their speeds, their diameters, and the attractive forces between them • Effusion • A process by which gas particles pass through a tiny opening • Rates of effusion directly proportional to the velocities of the particles (molecules of low mass effuse faster than molecules of high mass)

9. Deviations of Real Gases • Real gas – a gas that does not behave completely according to the assumptions of the kinetic-molecular theory • Johannes van der Waals – pointed out that particles of real gases occupy space and exert attractive forces. At high pressures and low temps the deviation from ideal behavior may be considerable. • KM theory more likely to hold true for gases w/little attraction. He and Ne show ideal behavior over wide range of temps and pressures. The more polar the gas molecule, the greater the forces and the less ideal its behavior.

10. To describe a gas fully, you need to state four measurable quantities: volume, temperature, number of molecules, and pressure. • Star this – you have to know it

11. Pressure -the force per unit area on a surface -the smaller the area of contact, the greater the pressure Newton – (SI) N – the force that will increase the speed of a one-kg mass by one meter per second each second it is applied (at Earth’s surface, each kg exerts 9.8 N of force due to gravity) Atmospheric pressure – sum of individual pressures of various gases in atmosphere (mostly N and O); at sea level is about 1.0 N/cm2

12. Measuring Pressure • Barometer • A device used to measure atmospheric pressure • First introduced by Evangelista Torricelli in the 1600s

13. The Barometer • mercury in the tube falls only until the pressure exerted by its weight equals the pressure exerted by the atmosphere • exact height of mercury in tube depends on atmospheric pressure • **at sea level and zero degrees, average atmospheric pressure can support a 760 mm column of mercury

14. Units of Pressure 1 atm = 760 mm Hg = 760 torr = 101.325 kPa = 1.01324 x 105 Pa

15. The Gas Laws Simple mathematical relationships between the volume, temperature, pressure, and amount of a gas.

16. For gases at constant temperature: • Pressure is caused by moving particles hitting the container walls. Decreasing volume w/same # of particles increases collisions therefore increasing pressure. Doubling pressure – volume cut by one-half • Reducing pressure by one-half, volume doubles • As one variable increases, the other decreases • Inverse relationship – graphs a curve (hyperbola)

17. Boyle’s Law -states that the volume of a fixed mass of gas varies inversely w/the pressure at constant temperature PV = k (k is constant for a given sample of gas) If P changes, V will change, but k will remain constant. P1V1 = k = P2V2 -most useful for changing conditions of a given sample of a gas Hint: We Boil Peas and Vegetables

18. Boyle’s Law Relationship

19. Charles’s Law • If pressure is constant, gases expand when heated • Relationship discovered by Jacques Charles in 1787 • All gases expand to the same extent when heated thru the same temp interval • Real gases can’t be cooled to -273C; before reaching that temp, intermolecular forces exceed KE of molecules and the gases condense to form liquids and solids • Avg KE of gas molecules is most closely related to Kelvin temp

20. Charles’s Law -states that the volume of a fixed mass of gas at constant pressure varies directly with the Kelvin temperature Gas volume and Kelvin temperature are directly proportional to each other. V/T = k or V = kT For changing conditions: V1/V2 = T1/T2 Hint: Charlie’s angels are on TV

21. Charles’s Law Relationship Directly proportional – graphs a straight line

22. Gay-Lussac’s Law Joseph Gay-Lussac recognized this relationship in 1802: -for a fixed quantity of gas at constant volume, pressure should be directly proportional to the Kelvin temp, which depends directly on avg KE Gay-Lussac’s law: the pressure of a gas at constant volume varies directly with the Kelvin temp P = kT or P/T = k

23. Gay-Lussac’s Law Relationship

24. A gas sample often undergoes changes in T, P, and V all at the same time. Boyle’s, Charles’s, and Gay-Lussac’s laws can be combined into a single expression that is useful in such situations. Combined Gas Law -expresses the relationship between P, V, and T of a fixed amount of gas PV = K Boyle’s V/T = k Charles’s P/T = k Gay-Lussac’s therefore Hint: Peas and Vegetables on the Table PV/T = k Combined

25. Dalton’s Law of Partial Pressures Partial pressure – the pressure of each gas in a mixture Dalton’s law of partial pressures – states that the total pressure of a mixture of gases is equal to the sum of the partial pressures of the component gases. True for any # of gases present.

26. Gases Collected by Water Displacement