 Properties of Gases  Gases uniformly fill any container  Gases are easily compressed  Gases mix completely with any other gas  Gases exert pressure on their surroundings

 Measuring barometric pressure  The barometer  Invented by Evangelista Torricelli in 1643  Units  mm Hg (torr) 760 torr = 1 atm  newtons/m 2 (pascal (Pa)) 101,325 Pa = 1 atm 101,325 Pa = kPa  Atmospheres 1 atm = standard pressure

Boyle’s law (Robert Boyle) The product of pressure times volume is a constant, provided the temperature and number of moles remains the same  Pressure and volume are inversely related  Volume increases linearly as the pressure decreases (1/P)

Boyle’s Law Continued……  At constant temperature, Boyle’s Law can be used to fine a new volume or pressure  Boyle’s law works best at low pressures  Gases that obey Boyle’s Law are called Ideal Gases

Charles’ Law (Jacques Charles) The volume of a gas increases linearly with temperature provided the pressure and number of moles remain constant.  Temperature and volume are directly proportional

Charles’ Law continued………..  Temperature must be measured in degrees Kelvin  K = o C  0 K is “absolute zero”

Avogadro’s Law (Amedeo Avogadro) For a gas at constant temperature and pressure, the volume is directly proportional to the number of moles.

Combined gas law – (n remaining constant)

 Derived from existing laws…….  V = k/P, V = bT and V = an  V = (k)(b)(a)(Tn/P)  Constants k, b and a are combined into the universal gas constant R and…..

 Limitations of the Ideal Gas Law  Works well at low pressure and high temperatures  Most gases do not behave ideally above 1 atm  Does not work well near the condensation conditions of a gas

Variations of the Ideal Gas Law Density Molar Mass of a gas

At Standard Temperature and Pressure (STP) T = K (0 o C) P = 1 atm (760 torr or kPa) 1 mole of an ideal gas occupies 22.4 L of volume Remember…….. Density = mass/volume and moles (n)= grams/molar mass