3.1.3 Bonding - States of matter
| Students should:
The three states of matter
There are three states of matter, solids, liquids and gases. The state of a substance depends on the forces that hold its particles together and their average energy.
We can differentiate the bulk properties of the three states of matter in terms of shape and the volume they occupy:
|Volume and shape||These have a fixed shape and a fixed volume||Have a fixed volume but no fixed shape||Have neither fixed shape nor volume|
These macroscopic properties are a direct result of the forces acting between the microscopic particles and the distances that these particles have between them.
The particles are close together and vibrate about fixed positions.
They cannot move any closer together and so the solid is incompressible.
The particles are held in fixed positions (apart from vibrations), which means that the solid keeps its shape in the macroscopic world unless subject to forces that bend, flatten or otherwise distort the shape.
Solids themselves can vary in properties, some being brittle while others are malleable, some hard while others are soft. These differences can always be explained by studying the underlying structure of the solid and the particles that comprise it.
The particles of liquids are close together but they are not in fixed positions. Their vibrations are enough to shake the whole structure apart and the particles can move with respect to one another.
The result of this freedom of motion is that the liquid flows in the macroscopic world and very small force is required to distort its shape. Gravity is enough to pull the liquid towards the earth and a liquid adopts the shape of the bottom of any container.
Where gravity doesn't act, for example in space, the liquid adopts a spherical shape. This indicates that forces act between the particles pulling them together.
The energy of the particles makes them move rapidly and prevents any appreciable forces of attraction acting between the particles. The spaces between the particles are now very large and the gas is compressible; the particles can be easily pushed closer together. The particles spread out, where possible, by constant collisions and expand to fill the entire space available. The gas adopts the shape of any container.
Changes of state
Melting point (mp)
If the vibrational energy of the particles is less than the force holding them together, the substance is a solid. As the temperature increases so the vibrations increase until eventually the structure breaks apart. We call this temperature the melting point of the solid. If this temperature is approached from above and a liquid gradually cooled, then this is also the temperature at which the liquid solidifies or freezes.
Thus, the definition of melting point is the temperature at which solid and liquid states are in equilibrium (i.e. co-exist). For example, water is a solid below 0ºC and a liquid above 0ºC. At the melting point of ice, both water and ice co-exist in equilibrium with one another.
Example: In what state is nitrogen at 56K, if its melting point and boiling points are 63K and 76K respectively?
56K is below the melting point and therefore the nitrogen is in the solid state at this temperature. The temperature would have to be increased to 63K in order for it to liquefy.
Boiling point (bp)
The same argument cannot be applied to the boiling point, as there are always particles in the body of a liquid that have sufficient energy to escape the forces of attraction within the liquid and escape into the gaseous state. This happens because the energy of the particles is distributed statistically, with some receiving a lot of energy whereas others receive very little. This distribution of energy is shown by the Maxwell - Boltzmann energy distribution curve. The curve is temperature dependent.
In the curve you can see that there are always some particles with high enough energy to overcome the forces of attraction and escape from the rest.
If the particles escaping from a liquid are carried away by currents of air the liquid evaporates.
We call a gas that exists below the normal boiling point of a liquid a 'vapour'. The particles of vapour exert a pressure in the same way as any other gas; this is called the vapour pressure.
The boiling point of a liquid is the temperature at which the vapour pressure of the liquid is equal to the atmospheric pressure and bubbles of gas can spontaneously form within the liquid. This clearly depends on the atmospheric pressure at the time. At the boiling point (or very slightly above it) the gas particles have too much energy to allow forces to pull them back together again.
The graph shows how the vapour pressure of a liquid steadily increases until it reaches atmospheric pressure. At this temperature the liquid boils. We record the boling point and the atmospheric pressure at the time.
Boiling points are useful measures of purity of liquids, as impurities cause a increase in the bp. This is called a colligative property.
|Shape||Fixed||Gravity is enough to pull the liquid into the shape of the bottom of the container||Not fixed, adopts the shape of the container|
|Volume||Fixed||Fixed||Not fixed, expands to fill all available space|
|Forces||Strong||Not strong enough to prevent relative motion of the particles||Negligible|
|Particle proximity||Close together||Close together||Very far apart but with frequent collisions|
|Particle motion||Vibration about a fixed point||Vibration, rotation and translation motion||Rapid translation (also rotation and vibration)|