1. “The Behavior of Gases”

2. Gases can expand to fill its container, unlike solids or liquids
Compressibility
Gases can expand to fill its container, unlike solids or liquids
The reverse is also true:
They are easily compressed, or squeezed into a smaller volume

3. Compressibility
This is the idea behind placing “air bags” in automobiles
In an accident, the air compresses more than the steering wheel or dash when you strike it
The impact forces the gas particles closer together, because there is a lot of empty space between them.

4. Variables that describe a Gas
The four variables: ( STP)
1. Pressure (P) in kilopascals, 760 mmHg, 760 Torr, and 1 Atm
2. Volume (V) in Liters, ml
3. Temperature (T) in 273 Kelvin
4. Amount (n) in moles

5. Temperature Volume Pressure
Amount of space enclosed by a shape or object
Measure of the average heat or thermal energy of the particles in a substance.
Pressure
Force exerted on a surface per unit area.

6. 1. Amount of Gas
When we inflate a balloon, we are adding gas molecules.
Increasing the number of gas particles increases the number of collisions
thus, the pressure increases
If temperature is constant, then doubling the number of particles doubles the pressure

7. 2. Volume of Gas
In a smaller container, the molecules have less room to move.
The particles hit the sides of the container more often.
As volume decreases, pressure increases. (think of a syringe)
Thus, volume and pressure are inversely proportional to each other

8. 3. Temperature of Gas
Raising the temperature of a gas increases the pressure, if the volume is held constant. (Temp. and Pres. are directly proportional)
The molecules hit the walls harder, and more frequently!

9. Phet Model

10. Kinetic Molecular Theory
KMT

11. 1. The volume of a gas particle is miniscule compared to the distance between themselves and other molecules.
2. Gas particles undergo no intermolecular attractions or repulsions.
3. Gas particles are in continuous, random motion.
4. Collisions between gas particles are perfectly elastic.
5. The average kinetic energy is the same for all gases at a given temperature, regardless of the identity of the gas.

12. The Gas Laws

13. Robert Boyle ( )

14. Boyle’s Law
Gas pressure is inversely proportional to the volume, when temperature is held constant.
Equation: P1V1 = P2V2

15. Graph of Boyle’s Law – page 418
Boyle’s Law says the pressure is inverse to the volume.
Note that when the volume goes up, the pressure goes down

16. A balloon contains 7. 2 L of He. The pressure is reduced to 2
A balloon contains 7.2 L of He. The pressure is reduced to 2.00 atm and the
balloon expands to occupy a volume of 25.1 L. What was the initial pressure
exerted on the balloon?

17. Jacques Charles ( )

18. Charles’s Law
The volume of a fixed mass of gas is directly proportional to the Kelvin temperature, when pressure is held constant.
Temperature Must Be in Kelvin.

19. Converting Celsius to Kelvin
Gas law problems involving temperature will always require that the temperature be in Kelvin. (Remember that no degree sign is shown with the kelvin scale.)
Kelvin = C + 273
°C = Kelvin - 273

20. A balloon is filled with 3. 0 L of helium at 310 K
A balloon is filled with 3.0 L of helium at 310 K. The balloon is placed in an oven where the temperature reaches 340 K. What is the new volume of the balloon?

21. Joseph Louis Gay-Lussac (1778 – 1850)

22. Gay-Lussac’s Law
The pressure and Kelvin temperature of a gas are directly proportional, provided that the volume remains constant.

23. #4. The Combined Gas Law
The combined gas law expresses the relationship between pressure, volume and temperature of a fixed amount of gas.

24. The equation for Avogadro's Law is V/n=k
The equation for Avogadro's Law is V/n=k. V is the volume of the gas, n is the amount of substance of the gas, and k is a proportionality constant.

25. 5. 00 L of a gas is known to contain 0. 965 mol
5.00 L of a gas is known to contain mol. If the amount of gas is increased to 1.80 mol, what new volume will result (at an unchanged temperature and pressure)?

26. Dalton’s Law of Partial Pressures
For a mixture of gases in a container,
PTotal = P1 + P2 + P
P1 represents the “partial pressure”, or the contribution by that gas.
Dalton’s Law is particularly useful in calculating the pressure of gases collected over water.

27. = 6 atm Sample Problem 14.6, page 434
If the first three containers are all put into the fourth, we can find the pressure in that container by adding up the pressure in the first 3:
2 atm
+ 1 atm
+ 3 atm
= 6 atm 1 2 3 4
Sample Problem 14.6, page 434

28. 8. Graham’s Law RateA  MassB RateB  MassA =
The rate of effusion and diffusion is inversely proportional to the square root of the molar mass of the molecules.
Derived from: Kinetic energy = 1/2 mv2
m = the molar mass, and v = the velocity.

29. Ideal Gas Law

30. The Ideal Gas Law Equation: P x V = n x R x T
Pressure times Volume equals the number of moles (n) times the Ideal Gas Constant (R) times the Temperature in Kelvin. R= .0821

31. What volume is occupied by 5
What volume is occupied by 5.03 g of O2 at 28°C and a pressure of 422 mmHg?

32. Density Density is mass divided by volume m V so, m M P V R T D = D =

33. Real Gases
and
Ideal Gases

34. Ideal Gases don’t exist, because:
Molecules do take up space
There are attractive forces between particles
- otherwise there would be no liquids formed

35. Real Gases behave like Ideal Gases...
When the molecules are far apart.
The molecules do not take up as big a percentage of the space
We can ignore the particle volume.
This is at low pressure

36. Real Gases behave like Ideal Gases…
When molecules are moving fast
This is at high temperature
Collisions are harder and faster.
Molecules are not next to each other very long.
Attractive forces can’t play a role.

37. Diffusion is:
Molecules moving from areas of high concentration to low concentration.
Example: perfume molecules spreading across the room.
Effusion: Gas escaping through a tiny hole in a container.
Both of these depend on the molar mass of the particle, which determines the speed.

38. Diffusion: describes the mixing of gases
Diffusion: describes the mixing of gases. The rate of diffusion is the rate of gas mixing.
Molecules move from areas of high concentration to low concentration.
Fig , p. 435

39. Effusion: a gas escapes through a tiny hole in its container
-Think of a nail in your car tire…
Diffusion and effusion are explained by the next gas law: Graham’s

40. Graham’s Law
Sample: compare rates of effusion of Helium with Nitrogen – done on p. 436
With effusion and diffusion, the type of particle is important:
Gases of lower molar mass diffuse and effuse faster than gases of higher molar mass.
Helium effuses and diffuses faster than nitrogen – thus, helium escapes from a balloon quicker than many other gases!