Unit 6

Review - STPAvogadro's Hypothesis - at the

same temperature and pressure equal volumes of gas have the same number of particles.

273 K (0ºC) and 1 atmosphere (atm)pressure.

Review – Kinetic Molecular Theory There were three assumptions. What were they?

What is the definition of ‘temperature’?

Kinetic = movement; Molecular = looking at the gas at the smallest level possible.

We use this theory to understand and explain the observed behaviour of gas.

How does K-M theory explain the behaviour of gases?

Some properties of gases explained by this theory are:

1) Low density of gases

2) Compression and expansion of gases

3) Diffusion and effusion

1) Low Density of Gases Density = mass per unit volume. (D = m/v)

Gold is 6500 times as dense as chlorine gas. This difference cannot be due solely to the difference in mass of the atoms (3:1)

K-M theory states there is a huge amount of space between gas particles there are fewer chlorine atoms than gold atoms in the same space and the chlorine has a lower density.

This is true of all gases when compared to solids and liq.

2a) Compression To compress something means to ‘squeeze it’.

The large amount of space between gases (defined in the K-M theory) means that a gas can be easily compressed into a smaller volume.

2b) Expansion The random motion of a gas’ particles (defined by the

K-M theory) means that a gas will always move out or ‘expand’ to fill the maximum space of a container.

3a) Diffusion As defined by the K-M theory,

there are no significant forces of attraction between gas particles and they can move easily past each other = diffusion.

When the space into which a gas flows is occupied by another gas, the random motion of gas particles (defined by the K-M theory) causes the gases to mix until they are evenly distributed.

3b) Effusion A process similar to diffusion; a gas escapes through a

tiny opening.

In 1846 Thomas Graham conducted experiments to measure the rate of effusion for different gases at the same temperature.

He discovered an inverse relationship between effusion rates and molecular mass.

Graham’s law of effusion states that the rate of effusion for a gas is inversely proportional to the square root of its molar mass.

Rate of effusion 1/ molar mass

The rate of diffusion depends on the mass of the particles. Lighter particles defuse more rapidly than heavier particles.

Different gases at the same temperature have the same average kinetic energy (defined by the K-M theory). The masses of different gases is different so a lighter gas must have more velocity than heavier particles to have the same kinetic energy (KE = ½ mv2)

Using Graham’s law you can set up a proportion to compare the diffusion rates for two different gases.

RateA = (molar massB)

Rate B molar massA

A worked example Q: Ammonia has a molar mass of 17.0 g/mol;

hydrogen chloride has a molar mass of 36.5 g/mol. What is the ratio of their diffusion rates?

A:

Rate(NH3) = (molar massHCl)

Rate(HCl) molar massNH3

= (17.0 / 38.5) = 1.47

Therefore Ammonia diffuses 1.47 times faster than HCl!

Known Unknown

Molar mass (NH3) = 17.0 g/mol

Ration of diffusion rates = ?

Molar mass (HCl) = 36.5 g/mol

Try BY Yourself! Calculate the ratio 0f effusion rates for Nitrogen (N2)

and neon (Ne).

Calculate the ratio of diffusion rates for carbon monoxide and carbon dioxide.

CHALLENGE! What is the rate of effusion for a gas that has a molar mass twice that of a gas that effuses at a rate of 3.8 mol/min?