AP Chemistry

Explore

Pages

Notes

5. Gases

Pressure

Torricelli

pressure: how often a gas hits the side of a container

⁠

760 mm Hg (millimeters of mercury) = 760 torr = 1 atm (standard atmospheres) = 101325 Pa (pascals)

number of newtons per meter squared

Boyle’s law: pressure and volume

⁠

example: A sample of 1.53 L of SO₂ is at pressure of 5.6 × 10³ Pa. What happens to the volume if the pressure is changed to 1.5 × 10⁴ Pa?

(5.6 × 10³ Pa)(1.53 L) = (1.5 × 10⁴ Pa)(V₂)

V₂ = 0.57 L

Charles’s law: volume and temperature

⁠

kelvin (K): unit of temperature where 0 K is absolute zero

⁠

example: A sample of gas at 15℃ and 1 atm has a volume of 2.58 L. What volume will the gas have at 38℃ and 1 atm?

2.58 L / (15 + 273.15 K) = V₂ / (38 + 273.15 K)

V₂ = 2.79 L

Gay-Lussac’s law: pressure and temperature

⁠

Avogadro’s law: volume and moles

⁠

combined gas law: combination of the other laws

⁠

Ideal gases

ideal gas: gas that you hope all gases follow

⁠

ideal gas law (”pervert”)

P: pressure (atm)

V: volume (L)

n: moles (mol)

T: temperature (K)

R: gas constant (0.08206 atm × L / mol × K)

Dalton’s law of partial pressures

⁠

total pressure is the pressures of the gases added together

example: A container contains 1.2 atm H₂ and 0.8 atm He. What is the total pressure?

1.2 atm + 0.8 atm = 2.0 atm

mole fraction (χ)

⁠

example: 2.0 mol H₂ and 1.0 mol of He. What is the mole fraction of He?

χ = 1.0 mol / 3.0 mol = 0.3...

mole fractions can be used to determine partial pressures of gases

χ × total pressure = partial pressure

example: The total pressure is 9.0 atm. What is the pressure of helium?

0.3... × 9.0 atm = 3.0 atm

STP

standard temperature and pressure (STP): 0℃ and 1 atm

molar volume at STP

PV = nRT at STP

(1.000 atm)(1.000 L) = n(0.0206 atm L / mol K)(273.15 K)

22.42 L/mol at STP

Gas stoichiometry

whenever you see STP: use 22.42 L/mol (gases only)

example: CaCO₃(s) → CaO(s) + CO₂(g) Calculate the volume of CO₂ at STP produced from 152g CaCO₃

⁠

Density of a gas

⁠

“dirty pee”

d: density

R: gas constant

T: temperature (K)

P: pressure (atm)

example: the density of a gas was measured at 1.50 atm and 27℃ and found to be 1.95 g/L; calculate molar mass

⁠

Kinetic molecular theory (KMT)

rules

gas particles are so small that they have no volume

particles are always moving; the collisions cause pressure

particles never react with each other

average kinetic energy in the gas is directly proportional to the temperature

⁠

T: temperature (K)

KE: kinetic energy (J)

⁠

Joule (J): unit of energy

⁠

root mean square velocity

⁠

T: temperature (K)

M: molar mass (kg/mol)

meters/second

big → slow

small → fast

Real gases

diffusion: one gas goes through another gas

effusion: one gas goes through a hole

ideal gases work best with high temperature and low pressure

real gases: have volume and stick to each other

van der Waals:

⁠

why wouldn’t a gas follow ideal gas law? because it’s real

Kinetic molecular theory (KMT)

Real gases

Gallery

Want to print your doc?
This is not the way.

Try clicking the ⋯ next to your doc name or using a keyboard shortcut (

CtrlP

) instead.