Chemistry - Chapter 10: Gas Laws

gas

a substance that has no well-defined boundaries but diffuses rapidly to fill any container in which it is placed in

e.g. NH3 and HCl - diffuse towards each other (high concentration to area of low concentration)

properties of gases

• diffuses quickly

• can be compressed

• no fixed shape/volume

• low density

• increased pressure → volume decreases

• increased temperature → increased volume

• decreased temperature → decreased volume

temperature

uses Kelvin (273K)

pressure

the force that the gas exerts on each unit of area of its container unit

Pascals (Pa/kPa for kilopascals) or Newtons per metre squared (Nm-²)

volume

volume of a sample of gas is the same as the volume of the container in which it is held

*remember:

1L = 1000cm³

1m³ = 100cm³

Boyle’s Law

at constant temperature, the volume of a fixed mass of gas is inversely proportional to its pressure

therefore, the higher the pressure, the lower the volume

// when a fixed mass of gas is kept at a constant temperature, its volume multiplied by its pressure us constant

Charles’ Law

the volume of a fixed mass of gas, kept at a constant pressure, is directly proportional to its temperature measured on the Kelvin scale

combined formula of Boyle and Charles

P1V1/T1 = P2V2/T2

Gay-Lussac’s Law of Combining Volumes

volumes of gases always reacted with each other in simple whole number ratios only if the volumes are measure at the same temperature and pressure

Avogadro’s law

equal volumes of gases contain equal numbers of molecules (under same conditions of temperature and pressure)

assumptions of the kinetic theory of gases

• gases are made up of particles in continuous rapid random motion, colliding with each other and with the walls of the container

• the average kinetic energy of particles in a gas is proportional to the temperature

• there are no attractive or repulsive forces between the molecules

• gas consist of particles that are widely separated in space. most of the volume of a gas is empty space; the gas particles themselves do not occupy much space

• as collisions are elastic, there is no loss of kinetic energy in collisions

limitations

• intermolecular forces exist between molecules in a gas (attractive forces)

• amount of space that gaseous particles occupy is significant especially under high pressures (relatively speaking)

• collisions are not all perfectly elastic as some energy can be lost to surroundings

there, the kinetic theory don’t hold for real gases, but only ideal gases

ideal gases

gases that obey all the assumptions of the kinetic theory of gases under all conditions of temperature and pressure

// gases that obey all gas laws at every temperature and pressure

real gases

differ from ideal gases due to incorrect statements of:

(i) forces of attraction and repulsion do exist between molecules

(ii) collisions are not perfectly elastic as some energy can be lost to surroundings

(iii) amount of space that gaseous particles occupy is significant, especially under high pressures

at s.t.p

100 kPa pressure

22.4L volume

273K temperature

the equation for state of any ideal gas (pg 64 of logsform book)

pV = nRT

• p = pascals

• v = m³

• n = number of moles

• r = 8.31J mol^-1 K^-1 = gas constant

• t = kelvin