Gas is omnipresent: in space, on Mars, dissolved in blood and soda.
We are constantly surrounded by an ocean of gas.
Gases can be theoretically, experimentally, and mathematically understood when they behave ideally.
Boyle’s Law
First described the relationship between pressure and volume in a gas.
In a closed system, decreasing volume increases pressure and vice versa.
Boyle’s Law: Pressure (P) × Volume (V) = Constant (k) (at constant temperature and gas amount).
Historically, Robert Boyle popularized the law, but the experiments were by Richard Townley and Henry Power.
Development of Gas Laws
Jacques Charles: Volume/Temperature = Constant (at constant pressure).
Amedeo Avogadro: Volume/Amount of Gas (moles) = Constant (at constant pressure and temperature).
Combined, these form the Ideal Gas Law:
PV = nRT
Pressure (P) × Volume (V) = number of moles (n) × gas constant (R) × Temperature (T).
Components of the Ideal Gas Law
Pressure (P): Measured in pascals (Pa) or atmospheres (atm).
Volume (V): Space available for gas particles.
Amount of Gas (n): Number of moles.
Gas Constant (R): 8.3145 L·kPa/K·mol.
Temperature (T): Measure of kinetic energy of particles; higher temperature means higher kinetic energy.
Demonstrating the Ideal Gas Law
Example with a soda can: Heating turns water into vapor inside the can, pressure changes when cooled in ice water, and pressure differential crushes the can.
The law allows calculation of one variable if three others are known.
Non-Ideal Behavior and Conditions
Gases deviate from ideal behavior at low temperatures and high pressures.
STP (Standard Temperature and Pressure): 0°C and 100 kPa.
Absolute Zero: 0 K (-273.15°C), where all particle motion stops.
Summary
The evolution of the Ideal Gas Law involved multiple scientists beyond Boyle.
The Ideal Gas Law is a pivotal tool in chemistry for understanding gas properties.
Real-world deviations and further complexities to be explored in future lessons.