Notes and Assignments: Gas Laws, Chapter 13
Homework on KMT test day: Wednesday, 4/22
Read Section 13.1A.
Do Example and Practice 13.1. Don’t forget to check your answer in the green section.
Memorize pressure equivalents on page 444.
Units of Pressure (Page 444):
1atm = 101.3 kPa = 760 torr = 760 mm Hg = 14.7 psi
Atmospheric pressure is measured with a barometer (Picture on page 443)
Pressure in a closed container is measure with a manometer (Picture on page 445)
Units of Temperature
Usually measured in oC (normal freezing point of water = 0 oC & normal boiling point of water = 100 oC)
SI unit = Kelvin (an absolute scale)
0 K = absolute zero, coldest possible temperature, particle motion stops = no KE
K = oC + 273 (not exactly, impacts sigfigs)
Assumptions About Properties of Gases – Kinetic Theory p.475-477
Small particles with lots of space between = particle volume is insignificant
Random motion and straight paths
Elastic collisions
Particles don’t attract or repel each other
Average KE is directly proportional to Kelvin temperature
Variables That Describe Gases
Pressure (P)
Volume (V)
Temperature (T)
Moles (n)
What is an Ideal Gas? Conforms to KMT
No such gas – however, real gases often behave as ideal gases
Real gases can be liquefied and sometimes solidified – ideal gases cannot
Real vs Ideal p.478
Kinetic theory assumption – no attraction among particles
But they do! Van der Waal’s intermolecular forces.
At high temperatures with lots of KE, these weak forces have little effect.
At low temperatures (near the substance’s boiling point), the intermolecular forces matter.
Real vs Ideal p.478
Kinetic theory assumptions – particles have no volume
But they do! At low pressure, there’s lots of empty space between the particles.
At high pressure, the particles are squished together and their volume matters.
Real vs Ideal p.478
So when do real gases behave most like an ideal gas?
high temperature (above boiling point) and low pressure
Graham’s Law
The rate of effusion is inversely proportional to the square root of the gases’ molar mass
Effusion = escape from a container through a tiny hole.
Rate A/Rate B = √(MM of B)/√(MM of A)
Gas Laws Problem Solving
•Define variables
•Choose a law
•Solve for missing variable
•Substitute (check units)
•Calculate
•Check
Boyle’s Law p.447
For a fixed amount of gas at constant temperature, pressure and volume vary inversely.
P1V1 = P2V2
Charles’ Law p.451
For a fixed amount of gas at constant pressure, temperature and volume vary directly.
V1/T1 = V2/T2 or T1/V1 = T2/V2
Gay-Lussac’s Law
For a fixed amount of gas at constant volumne, pressure and absolute temperature vary directly.
P1/T1 = P2/T2 or T1/P1 = T2/P2
Combined Gas Law p. 464
Combines Boyle’s, Charles’ and Gay-Lussac
Allows for calculations where none of the variables is constant
P1V1/T1 = P2V2/T2 or T1/P1V1 = T2/P2V2
Homework: Page 480, #7c, 8c, 10, 13c, 17, 29. Charles Law bonus lab
STP (standard temperature and pressure) (1 atm, 0oC)
Avogadro’s Principle p.455
Equal volumes of gases at the same temperature and pressure contain equal numbers of particles
One mole (6.02 x 1023 particles) of an ideal gas at STP occupies a volume of 22.4 L
Ideal Gas Law p.458
PV = nRT
“R” is known as the ideal gas constant
R = 0.08206 atm×L/(mol×K)
Also note: You know how mass and moles are related for any given substance.
Homework: Page 481, #23, 25
Partial pressure = pressure gas would have had if it were alone in the container.
Dalton’s Law of Partial Pressures page 464
At constant volume and temperature the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures
P1 + P2 + P3…= Ptotal
Homework: Page 481, #33
Collecting gases over water
Picture Page 467
Vapor pressure of water: Table 13.2
Combine Avogadro and Dalton: n/nT = P/PT mole ratio = pressure ratio
Homework: Page 481, #35, 36
