5.1.1 States of matter

Describe and compare solids, liquids and gases as the three states of matter.

The movement of particles, the attractive forces between particles and the interparticular spacing should be described. A molecular level description of what happens when evaporation, boiling, condensation, melting and freezing occur should be given. Students should understand what is meant by the term diffusion



Solids may be ionic, metallic, simple covalent or giant covalent, but all of these structures have something in common; the smallest part¡cles are not free to move, they are held close together in fixed positions by the forces around them. The only motion allowed is vibration and this is how they absorb energy. A particle in a solid will vibrate more and more as the temperature rises until eventually it has enough energy to break free from the forces holding it in position. This happens at the melting temperature.

Bulk properties:

Solids show a definite shape and a definite volume. Unless forces are used that are not commonly found near the earth’s surface, solids can not be compressed.



Liquids are materials in which the smallest particles are as close together as solids, but the particles can slip over each other to change places. They can vibrate, rotate and translate but the forces that hold the particles of liquid close to each other are greater than the forces due to motion that would force the particles away from each other.

Bulk properties:

Liquids have no fixed shape except for the shape of the container but do have a fixed volume. Liquids can not be compressed under common pressures.



The particles of gas, either atoms or molecules, have too much energy to remain attached to one other. The move by translation, rotation and vibration, but in this case the translational motion is the most important. The particles are on average very far apart and collide incessantly many times a second. Because of the distance between them it is assumed that the forces of attraction between the particles are neglegible.

Bulk properties:

Materials in the gas phase have no fixed shape, that is, they take on the shape of the container. Gases have no fixed volume, the motion of the particles is so great that the forces of attraction between the particles are not able to hold them together. A certain amount of gas at a pressure of one atmosphere and a volume of ten litres could become five litres if the pressure was increased or would become more than ten litres if the pressure was decreased. The gas expands to fill the container. see gas laws


5.1.2 The kinetic theory

Describe kinetic energy in terms of the movement of particles whose average energy is proportional to absolute temperature

Kinetic energy should be interpreted in terms of ideal gases consisting of point masses in random motion whose energy is proportional to absolute temperature. Students should be able to describe what happens when the temperature is changed

The Kinetic Theory of Matter is the statement of how we believe atoms and molecules, particularly in gas form, behave and how it relates to the ways we have to look at the things around us. The Kinetic Theory is a good way to relate the 'micro world' with the 'macro world.'

A statement of the Kinetic Theory is:
1. All matter is made of atoms, the smallest bit of each element. A particle of a gas could be an atom or a group of atoms.
2. Atoms have an energy of motion that we feel as temperature. The motion of atoms or molecules can be in the form of linear motion of translation, the vibration of atoms or molecules against one another or pulling against a bond, and the rotation of individual atoms or groups of atoms.
3. There is a temperature to which we can extrapolate, absolute zero, at which, theoretically, the motion of the atoms and molecules would stop.
4. The pressure of a gas is due to the motion of the atoms or molecules of gas striking the object bearing that pressure. Against the side of the container and other particles of the gas, the collisions are elastic (with no friction).
5. There is a very large distance between the particles of a gas compared to the size of the particles such that the size of the particle can be considered negligible.

In the view at this level, it is useful to look at atoms as if they were close to the hard little balls that Dalton considered.

Changes of state

All of these may be explained by an appreciation of the energy overcoming the forces holding the particles together or becoming insufficient to keep them apart.

melting: solid to liquid phase change

boiling: liquid to gas phase change

evaporation: liquid to gas phase change below the boiling point

solidification: liquid to solid phase change

condensation: gas to liquid phase change

5.1.3 The Maxwell Boltzmann curve

Describe the Maxwell- Boltzmann energy distribution curve.

This describes the distribution of energies of particles in a body of liquid or gas. It also applies to solids although since there is no translation or rotation of the particles we are restricted to vibrational energy only.

Maxwell and Boltzmann explained that in every substance the majority of particles have a mean energy but there are always some particles with high and some with low energies. The total energy of the particles is given by the area under the curve.

5.1.4 : Draw and explain qualitatively the Maxwell- Boltzmann energy distribution curve for different temperatures.

Higher temperature - flatter and broader curve as the spread of energies is greater. The mean energy moves further to the right (higher energy)
Lower temperature curve - bigger hump and not so broad mean energy moves more to the left (lower energy)

For a full explanation of Maxwell Boltzmann click here.

5.1.5 The gas laws:

Describe qualitatively the effects of temperature, pressure and volume changes on a fixed mass of an ideal gas.
State the ideal gas equation PV = nRT

Robert Boyle first identified the dependency of volume on pressure, finding that an increase in pressure caused a corresponding decrease in volume. PV is constant.

Charles investigated the effect of temperature on volume and found a reciprocal relationship between volume and temperature. In other words volume is directly proportional to temperature.

From these two relationships and other experimentation the Ideal Gas Equation was formulated.

The ideal gas equation: PV = nRT


P = pressure (units may be Nm-2, Pa, kPa, atmospheres, mmHg)
V = volume in cm3
n = number of atoms in mol
R = the universal gas constant - usually in SI units 8,314
T = temperature in K

Gases consequently expand when heated at constant pressure, cool when expanded at constant pressure etc.


5.1.7: Apply the ideal gas equation in calculations. Use the relationship between P, V, n and T for gases. Students should be familiar with P1V1/T1 = P2V2/T2 and be able to calculate molar volumes

The above equation can be rearranged as P1V1=P2V2 (temperature constant) or V1/T1=V2/T2 (pressure constant) to calculate volumes, pressures, temperatures etc. under changing conditions.

Avogadro's law

Equal volumes of gas contain equal numbers of particles (at 273K, 1 atm one mole of gas occupies 22.4 dm3)

Dalton's law of partial pressures

The partial pressure of a gas is the pressure the gas would exert if the gas were alone in its container.

Example: If there are Gases A, B and C having a moles, b moles and c moles respectively in a 1dm3 container then the partial pressure of a = (a)/(a+b+c) x total pressure.

Pp = (number of moles of the gas) x total pressure
(total number of moles)



  particular properties bulk properties
  motion distance shape volume
solids Particles vibrate about fixed positions very close together fixed fixed
liquids Translation, rotation and vibration: Translation not so important as particles are very close together very close together not fixed fixed
gases Translation, rotation and vibration: Particles fly about very rapidly and collide many times each second. very far apart in molecular terms not fixed not fixed