1. Warm-up #6
2. The balloon in the previous problem will burst if its volume reaches L. Given the initial conditions specified in that problem, determine at what temperature, in degrees Celsius, the balloon will burst if its pressure at that bursting point is atm.
Identify what you know
V1=250.0L
T1= 22°C = = 295K
P1= mm Hg
V2=400.0L
T2= ?
P2= .475 atm = .475 atm x 760 mmHg P2 =361 mmHg
1 atm
Solve for what you don’t know T2= 230 K or -43 C
T2= T1 x P2V2
P1V1
T2=230.3 K

2. Review Homework
See overhead

3. Ideal Gases & Ideal Gas Law

4. Ideal Gas Law- not in notes
When dealing with gas, an equation was used to relate all of the factors needed in order to solve a gas problem. This equation is known as the Ideal Gas Equation.
However, anything ideal does not exist and 2 assumption were made:
1) the gas particles have no forces acting among them
2) the particles do not take up any space (volume is ignored)
An ideal gas is a hypothetical gas dreamed by chemists because it is be easier if things like intermolecular forces do not exist to complicate the simple Ideal Gas Law.
This definition of an ideal gas contrasts with the Non-Ideal Gas definition, which describes how gases actually behaves in reality.
For now, let us focus on the Ideal Gas laws.

5. No kinetic energy is lost in elastic collisions
Ideal Gases
Ideal gases are imaginary gases that perfectly fit all of the assumptions of the kinetic molecular theory.
Gases consist of tiny particles that are far apart relative to their size.
Collisions between gas particles and between particles and the walls of the container are elastic collisions
No kinetic energy is lost in elastic collisions

6. Ideal Gases (continued)
Gas particles are in constant, rapid motion. They therefore possess kinetic energy, the energy of motion
There are no forces of attraction between gas particles
The average kinetic energy of gas particles depends on temperature, not on the identity of the particle.

7. Real Gases Do Not Behave Ideally
Real gases DO experience inter-molecular attractions
Real gases DO have volume
Real gases DO NOT have elastic collisions

8. Deviations from Ideal Behavior
Likely to behave nearly ideally
Gases at high temperature and low pressure
Small non-polar gas molecules
Likely not to behave ideally
Gases at low temperature and high pressure
Large, polar gas molecules

9. Variables that Describe a Gas
Pressure (atm, mmHg, kPa)
Volume (L)
Temperature (˚C or K)
Number of moles (mol)

10. The Ideal Gas Law
Allows us to solve for a property of an ideal gas when properties are constant!
PV=nRT
P=pressure
V=volume (Liters!!)
T=temperature (Kelvin!!)
n=number of moles
R= 8.31 L x kPa or L x atm
K x mol K x mol

11. Lets Practice
Determine the volume occupied by mol of a gas at 15˚C if the pressure is 81.1 kPa.
Identify what you know
n=0.582 mol
T=15˚C = 288 K
P=81.1 kPa.
R= 8.31 L x kPa / K x mol
V=?

12. PV=nRT 81.1 Kpa x V=(0.582 mol) (8.31 L x kPa / K x mol) (288K)
81.1 KPa x V= L x kPa
V=17.2 L

13. Two rules! Temp= Kelvin, Volume =Liters, and pressure= atm or kPa
If you are given grams convert into moles.
36 grams of H2O x 1 mole H2O =2 mol H2O
18 g H2O

14. Dalton’s Law of Partial Pressures
the total pressure of a mixture of gases is equal to the sum of the partial pressures of all the gases present. (at constant Temp & Volume)
Formula:
Ptotal= P1 + P2 + P3 + P4…

15. Let’s Practice
1) What is the total pressure of a gas mixture if it contains CO2 at 40.8 torr and O2 at torr & H2 at torr? 2) What is the pressure of Ne gas if the total pressure of the gas is atm, and the mixture contains 40.4 atm of He and 22.6 atm of HCl?
PT=P1 +P2+P3
PT= = torr
PT=P1 +P2+P3
100.6 atm = PNe atm atm = 37.6 atm

16. Water Displacement
Gases produced in the laboratory are often collected over water. Thus, the gas is not pure but is always mixed with water vapor.
The collection process makes it so the total pressure inside the bottle would be the same as the atmospheric pressure.
Patm=Pgas + PH2O

17. Lets practice – (Hw check: pg 367 #1)
3) Some hydrogen gas is collected over water at 20.0 °C. The levels of water inside and outside the gas-collection bottle are the same. The partial pressure of hydrogen is torr (mmHg).
What is the barometric pressure at the time the gas is collected?
Formula:Patm=Pgas + PH2O (to find the P use table A-8 pg. 859)
Patm= torr torr= torr

18. Lets Practice Continued….
4. Some CO2 gas is collected over water at 15.0 °C. The levels of water inside and outside the gas-collection bottle are the same. The barometric pressure at the time the gas is collected is 780 mmHg. What is the partial pressure of CO2?
Formula: Patm=Pgas + PH2O
780 mmHg = PCO mmHg= 767.2mmHg

19. Dalton’s Law- not on notesheet
Example #1: What is the partial pressure of oxygen at 101.3kPa of total pressure if the partial pressure of nitrogen is kPa and carbon dioxide is 0.040kPa?

20. Solution: Ptotal= Po+ PN2 + PCO2 Po = Ptotal – (PN2 + PCO2 )
Po = kPa – (79.10kPa kPa)
=22.16 kPa

21. Dalton Example #2
What is the partial pressure of each gas if you have a mixture of 2.0 moles of O2 and 2.0 moles of CO2 that exert at total pressure of 700 torr?
O2= 2.0 mol O2 x 700 torr = 350 torr
4.0 total mol
CO2= 2.0 mol CO2 x 700 torr = 350 torr
4.0 total mol

22. Now you try! Show all your work on worksheet-pg 16
Ideal Gas Law: PV=nRT
Make sure:
1) Temp is in Kelvin (C+273= K)
2) Volume is in L (1000mL=1L) 350 mL x 1 L .
3) Pressure is in atm or kPa mL
4) 101.3Kpa=1 atm=760 mmHg=760 torr
5) grams-> mol (use molar mass)
36 grams of H2O x 1 mole H2O =2 mol H2O
18 g H2O

23. When to use the equations
If the conditions of the gas change
Then use the combined gas law
If the conditions of the gas are fixed
Then use the ideal gas law