gases

A collective term for molecules & Individual atoms

By using:

Macroscopic properties: is it condensed, is it fluid

Microscopic properties: particle separation, particle movement

Translation (a->b)

Spin (rotation)

Vibration (bend or stretch)

Heat capacity: heat capacity is higher with more types of motion

Spectroscopy: a beam of light is shone at the sample. Photons which have the right energy to match the difference between 2 quantised energy levels are absorbed. This causes movement. The light that passes through the sample is the missing wavelength that were absorbed. This creates a unique pattern of dark lines or peaks (called the spectrum). Each peak corresponds to a specific motion (vibrations in infrared and rotation in microwaves). By analysing the pattern, you can tell what the molecule is and what its motion is.

Ordered (regular repeating structure): crystals, also liquids can be in order (liquid crystals)

Unordered: amorphous of glasses

Some material can be in an ordered state sometimes and disordered state other times eg silicon dioxide (crystal quartz or amorphous silica glass)

By using diffraction:

Pass X-rays through. If ordered, some rays will be diffracted by a precise angle defined by the repeating structure. Diffraction pattern can be used to determine structure.

Less gaps, more dense, increase in melting point

Optical properties:

Birefringence : crystal materials experience birefringence, where different polarisations of light have different reflection angles.

A beam of light splits into 2 and “direction of vibration” (polarization) determines which way the light goes.

Liquid crystals with electron field are ordered and unordered without. This changes optical properties eg you can see the material when the electronic field is on where it was invisible previously.

Gases: particles are far apart so interparticle forces are weaker. They are negligible but gases do collide.

Liquids & solids: Stronger interparticle forces than gases.

Electrostatic

Van Der Waals

Hydrogen bonds (bond between hydrogen and electronegative atom in other molecules).

If 2 like charges are together

ALSO If particles are too close, repulsive forces will push them apart.

The range of temperatures and pressures at which each phase is stable.

Triple point (TP): shows the pressure and temperature at which all 3 phases are stable (can happen)

Critical point (CP): As temp and pressure increases until the gases and the liquids have the same density. At this point (CP) the phase boundary disappears and you can't distinguish between liquid and gas. This forms a supercritical fluid (liquid density & gas diffusion/flow)

All gases act ideally at low density

There are no interparticle forces (attractive forces weaken as separation increases)

There is no volume to the particles (volume of particles will be very small compared to total volume if density is low)

p in pa, v in m3, R = gas constant=

8.3145JK-1mol-1 = Boltzmann constant x avogadros constant, t in K

N = number of particles

Vm in m3mol-1 = molar volume= v/n

T(k) = T(°c) + 273.15

T(k) = (5/9)T(°f) + 255.372

1L = 1dm3 = 10^-3 m3

1ml =1 cm³ = 10^-6 m3

pa = Nm-2

Atm =101325 pa

Bar = 100000 pa

760 Torr =101.325 kPa = 1 atm

Psi = 6894.8 pa

The pressure of a mixture of gases is equal to the sum of the partial pressures (the pressure a component gas would have if it occupied the same volume alone) of each gas

It works because:

No interparticle forces mean that each gas behaves as if it’s the only one in the container. Ie they move independently of eachother.

Negligable particle volume: The minimal size of gas particles ensures they occupy insignificant space, allowing independent motion without impeding one another.

Both: conserve momentum (mass x velocity)

Elastic: Atomic gases have elastic collisions. Particles bounce off of eachother and there is no change to total KE.

Inelastic: Happens in most molecular gases. Either KE is converted to rotational or vibrational energy OR if vibrating or rotating particles collide, some energy is transfered into KE. Overall, since both are equally likely, there is no overall change to KE (elastic)

THEREFORE EVEN IF INDIVIDUAL COLLISIONS ARE INELASTIC, THE AVERAGE ELASTIC BEHAVIOUR DETERMINES MACROSCOPIC PROPERTIES.

Everything also applies to collisions with walls

At high temperatures and pressures the ideal assumptions of a gas (no particle volume and no interparticle interactions) don’t apply.

Gases that aren’t ideal gases are called real gases.

We use the compression factor (z):

pv/nrt = 1 for ideal gases

pv/nrt= z for real gases

z>1 it is harder to compress than an ideal gas

z<1 it is easier to compress than an ideal gas

For air, z=0.9999 so we can treat it as an ideal gas

Pv= nrt doesnt work for real gases as they don’t follow the assumptions of ideal gases so we use corections for pressure and volume.

Volume: we use the volume that the particles can be compressed into. V(total) - v(particles).

V(cor) = V-nb where b is the molar volume in m3mol-1

Pressure: We need to take into account attractive forces as attraction means that particles hit the walls less frequently and/or with less force

a is in units m6PaMol-2

Pressure correction: at low density a (the intermolecular forces) = 0 so p+ a/vn2 = p

Volume correction: at low density Vm is big so Vm- b ≈ Vm

The same as van der waals but the pressure correction takes into account the temperature

B (interaction between 2 particles) & c (interaction between 3 particles) are dependent on temperature.

If you measure z at different conditions, you can work out the coefficients

Can be expressed as a force (f) or as potential energy (u)

Liquidification can’t happen in ideal gases because there are no interparticle interactions but it can happen in real gases because there are interparticle interactions.

There are discs that spin. Discs spin at different rates which allow different speeds of particles to be selected. If you measure the number of particles passing through, this gives the distribution of speeds.

It shows the distribution of speeds in a gas. I.e. the probability of a particle having a particular speed within the gas.

It is asymmetrical and there is 1 peak

V = speed

m = mass of particle (kg). If in gmol-1, divide by 1000 and multiply by avogadros number

Can also use M (molar mass) but then use R instead of Kb

It uses mass and temperature. The ratio of m:T determines the distribution of speeds.

Maxwell’s doesn’t use pressure or volume showing that 2 samples of the same gas at the same temp will have the same distribution of speeds regardless of the vessel they are in.

It is the energy associated with translational motion. It is = to 1/2mv2.

It depends on the temperature not the mass of the particles. This means that all particles that have the same temperature will have the same distribution of KE.

This helps to explain why the ideal gas equation works as the values in the ideal gas equation p&v depend on KE which always has the same distribution not mass (or the speed distribution) which isn’t the same for every particle.

Collision frequency: the number of collisions an individual particle has per second

Collision density: how many collisions happen for all particles in 1 m³ per second

Mean free path: how far on average a particle travels between collisions

The mean relative speed of the particles:

The size of the particles:

For collisions between the same kind of particle divided by 2 to avoid counting the same collision twice

Through the second form of the equation, we can tell that mean free path does not depend on the temperature

Gas molecules escaping through a small hole in the wall of a vessel with a vacuum on the other side. It's rate, depends on the rate of collisions with the walls. This means it is also inversely proportional to mass^1/2.

As Zw Doesn't depend on the cross-section neither does the rate of effusion. (Mass matters, size doesn’t)