Gases Physical Properties

States of Matter:

Solid: very clumped, little space between molecules, molecules vibrate in place

Liquid: smalk amount of intermolecular space, molecules move about and slide past each other

Gas: lots of intermolecular space, molecules move freely and very fast

Kinetic Molecular Theory:

KMT is based on the idea that particles of matter are always in motion

Volume: most of the volume of gas is empty space

Motion: Move freely and randomly

Collision: No loss of energy due to elastic collisions

Force: No intermolecular forces

Energy: Average kinetic energy is directly related to the temp

KMT is based on ideal gases

identity of ideal gases does not matter as they all act the same due to no intermolecular forces and they don’t take particle size into account

Gases in the real world do not behave ideally, they behave most ideally at low pressure and high temperature

real gases do not perfectly follow the KMT

real gases will behave differently depending on the type of gas because they take particle size into account and intermolecular forces exist.

Characteristics of Gases:

Compressibility: can be compressed or expanses due to lots of empty space

Mass: have a definite mass

shape and volume: have indefinite shape and volume due to fluidity

motion: only change direction due to collision

4 factors that affect gasses:

Amount (n): number of molecules

as the amount increases number of collisions also increases

measures in moles

Volume (V): size of container

size of the container is the amount of space occupied by the gas

Gases will occupy all parts of the container

measured in Liters (L)

Temperature (T): measured in Kelvin (273 + °C)

temperature is the measure of the average kinetic energy

molecules of different gases at the same temperature have the same average kinetic energy

temp increases = kinetic energy increases

temp decreases = kinetic energy decreases

measured in kelvin (K)

Pressure (P): Force exerted over an area

Gas pressure = collisions between molecules and the inside of the container

The more or harder the hits the higher the pressure and vice versa

Effusion: the ability of gas to escape its continee through a tiny hole

Diffusion: Gas soreads through areas and bigger particles move slower

Units: atm (atmospheres) kPa (kilopascal) mmHg(millimeters of Hg) torr

Standard Temperature and Pressure (STP)

temperature: 0°C or 273 K

Pressure: 1 atm, 101.3 kPa, 760 torr / mmHg

Volume: 1 mole = 22.4 liters

Gas Laws:

Boyle’s Law: Volume and Pressure are Inversely related

Constants: Temperature and amount

equation (P)1(V)1 = (P)2(V)2

Charles’ Law: Volume and Temperature are directly related

Constants: Pressure and Amount

(V)1/(T)1 = (V)2/(T)2

Constants: pressure and amount

Gay-Lussac’s Law: Pressure and Temperature are directly related

Constants: Amount and Volume

(P)1/(T)1 = (P)2/(T)2

Avogardo’s Law: Volume and Amount are directly related

Constants: Pressure and Temperature

(V)1 / (n)1 = (V)2 / (n)2

Dalton’s Law: Partial Pressure: pressure exerted by one gaseous component in a mixture

Total Pressure = sum of all partial pressures

(P) total = (P)1 + (P)2 + (P)3 + …

Combined Gas Law: combination of all 4 laws. This can be used when several variables change

Equation: (P)1(V)1/(n)1(T)1 = (P)2(V)2/(n)2(T)2

Ideal Gas Law: volume, pressure, temperature, and the number of moles are interrelated mathematically.

The ideal gas law includes a number called the “ ideal gas law constant”.

P: Pressure

V: Volume (Liters)

n: amount (mols)

R: Ideal Gas Constant

T: Temperature (K)

PV = nRT

8.134 x L x kPa/ mol x K or 0.0821 x L x atm/ mol x K