S1.4 & S1.5: Avogadro’s Law And Ideal Gases Flashcards
What is avogadro’s law?
States that at the same temp and pressure equal volumes of different gases contain the same number of particles
-> at a particular temp and pressure -> 1 mol of any gas -> occupy same volume
Room temp/1 atm -> 24 dm^3
What is boyles law?
Gas in a container -> exert pressure as gas molecules collide with wall of container
Decrease in volume at constant temp -> increased pressure
-> particles squashed, more collision
Boyles law:
Pressure is inversely proportional to volume at constant temp
P∝ 1 /V
PV= a constant
What is Charles law?
Increasing temp at constant pressure -> particle more KE -> more collisions -> molecules move further apart -> increased volume
Charles law:
Volume is directly proportional to the temp (in K) at constant pressure
V∝TorV / T= a constant
What is the relationship between pressure and temperature?
Temp is directly proportional to pressure at a constant volume
P∝TorP / T= a constant
Increasing temp at constant volume -> increased KE -> molecules move faster + collide more frequently -> increased pressure
How is the ideal gas equation constructed?
Boyles + Charles + P/T = all equal a constant
-> PV/T = constant
-> PV = constant x T
Constant is made from 2 components -> number of moles (n) and gas constant (R)
IDEAL GAS EQUATION:
PV = nRT
(P= pa, V=m^3, n=mol, T=Kelvin, R=8.21 (J/K/mol) -> when V:m3 and P:Pa)
What does kinetic theory state? How does it relate to ideal gases?
Gas molecules:
- move fast + randomly
- Have hardly any volume
- Do not attract or repel each other (no intermolecular forces)
No KE lost when gas molecules collide (elastic collisions -> energy conserved and particles colliding strike and move in opposite directions)
Temp of the gas directly proportional to the average KE of molecules
Gases that follow this -> ideal gases
In reality -> gases don’t fit this exact description, but are very close -> real gases
Real gases: ideal gas equation
The real relationship between pressure, volume and temp -> significant deviation from ideal gas equation when temp very low/pressure very high
Réal gases: assumptions about volume
Volume of actual gas molecule -> tiny compared to volume of container -> so it can be ignored
- generally true
- inaccurate at low temp and high pressures
High pressure -> less space -> fraction of space taken up by gas increases
Réal gases: assumptions about attractive forces
Gas molecules far apart -> little interaction between molecules
Gas molecules closer -> intermolecular forces cause attraction -> reduces number of collisions with walls of container
Pressure is less than expected
