Gases Part 1

1. Gases Part 1

2. Properties of Gases • Gases have very low densities, and may be compressed or expanded easily: in other words, gases expand or compress to completely fill whatever container they are in. • Gases also tend to form homogeneous mixtures, that is they tend to mix completely to form uniform mixtures. Solids and liquids often form heterogeneous mixtures. • Gases exert a pressure or a force where the gas particles collide with a surface. Remember, gas particles are moving randomly with high kinetic energy. Everywhere they collide, the force of the collision is the pressure that the gas exerts.

3. Properties of Gases

4. Properties of Gases

5. Pressure Units of Gases • Scientists typically use atmospheres, atm, where sea level atmospheric pressure (called standard atmospheric pressure) is exactly 1 atm. • Although atm is quite commonly used, the SI unit of pressure is Pascals, Pa, where 1 atm = 101.325kPa. • In the US, our barometers (what we use to measure relative atmospheric pressure), usually read in mm Hg or in torr, where 1 atm = 760 mm Hg = 760 torr.

6. Measuring Pressure of Gases • Scientists use manometers and barometers to measure pressure. • Barometers measure atmospheric pressure. • Manometers measure the pressure of a rxn vessel. • Manometers can be open or closed.

7. Measuring Pressure of Gases

8. Kinetic Molecular Theory of Gases • The special properties of gases have been explained by a series of assumptions, called the Kinetic-Molecular Theory: • Gases consist of large numbers of particles that are in continuous, random, straight-line motion. • The volume of the particles of gases is negligible compared to the volume of “empty” space between the particles. In fact, gas particles typically take up only about 0.1% of the total volume of their containers, the rest being empty space.

9. Kinetic Molecular Theory of Gases • Attractive and repulsive forces between gas particles is negligible. Therefore, there are no appreciable intermolecular forces and all of the gas particles act essentially independently of one another. • For a given temperature, the average kinetic energy of the gas particles is constant, as is the total kinetic energy of all the gas particles. Therefore, gas particle collisions are perfectly elastic.

10. Kinetic Molecular Theory of Gases • The kinetic energy of gas particles depends on temperature. The higher the temperature, the higher the kinetic energy of the particles. At a given temperature, the average kinetic energy of ALL gases is the same. So the average kinetic energy of O2 is the same as He is the same as CO2 so long as they are at the same temperature. Because KE = 0.5mv2, and they have different masses, they must have different velocities, BUT THEY HAVE THE SAME KINETIC ENERGY!

11. Gas Laws • All gases have very similar physical behavior. • What are the gas variables? What can we change about a gas? • These variables and how they relate to one another are described in the Gas Laws: Boyle’s; Charles’; Gay-Lussac’s; Avogadro’s; Combined Law; Ideal Gas Law; and Dalton’s Law of Partial Pressure.

12. Boyle’s Law • Boyle’s Law describes the relationship between the pressure of a gas sample and its volume (n and T constant). • What do you think happens to the volume of a gas when we increase the pressure in an expandable and compressible container? • Why?

13. Boyle’s Law

14. Boyle’s Law: An Inverse Relationship

15. Boyle’s Law • Boyle’s Law is stated as: P1V1 = P2V2 • where 1 are the initial conditions and 2 are the final conditions. • Note that we use Boyle’s Law when we change the volume and pressure. • So if we change the volume, we can calculate the new pressure (or vice versa).

16. Charles’ Law: V and T • Charles’ Law states the relationship between Volume and Temperature (P, n constant). • If we have a balloon filled with gas and we start to heat it, what will happen to the volume of the balloon? • Why? • Charles’ Law is stated as: V1/T1 = V2/T2

17. Charles’ Law: A Linear Relationship • It is important to remember that in all of the Gas Laws, the temperature MUST be in Kelvin! • Why?

18. Gay-Lussac’s Law: P and T • Gay-Lussac’s Law states the relationship between Pressure and Temperature (V, n constant). • If we have a closed steel canister filled with gas and we throw the canister in a fire, what will happen to the pressure of the gases inside the canister and why? • Gay-Lussac’s Law is stated as: P1/T1 = P2/T2 • It is a linear relationship.

19. Avogadro’s Law: n and V (and P as well) • Avogadro’s Law states the obvious relationship between amount and volume. • If we have a balloon and we blow air into it, what happens to the volume? • Avogadro’s Law is stated as: V1/n1 = V2/n2 • It is a linear relationship. • The pressure and amount are also a linear relationship.

20. Combined Gas Law • In real life, think weather balloons, more than 2 gas variable changes. • For example, as a weather balloon ascends, the P, V, and T all change! • So the Combined Gas Law combines all of the prior Laws into one: P1V1/n1T1 = P2V2/n2T2 • or more commonly: P1V1/T1 = P2V2/T2

21. Ideal Gas Law • But what if we just want to know what the pressure of a sample is currently? • If we know the temp, volume, and amount of the gas sample, we can do this using the Ideal Gas Law: PV = nRT • where R is the Ideal Gas Constant • Now you know 3 of 4 variables and solve for the 4th. No change is involved!