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The melting and boiling points of substances are important as they serve to characterise the purity of a material |
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Melting point
The melting point of a substance is 'the temperature at which the two states, liquid and solid, co-exist in equilibnum'. Or to put it into plain terms, the temperature at which something melts.
The melting point is actually a very important physical property, as it may be used to ascertain the degree of purity of a substance. Pure substances have sharp, well defined melting points, whereas the addition of impurity both lowers and broadens the melting temperature.
Unike boiling points, the melting point remains virtually unaffected by the external air pressure.
Boiling point
The boiling point of a liquid is the temperature at which the vapour pressure of the liquid equals the atmospheric pressure. The atmospheric pressure is variable according to the elevation above sea level and the weather conditions. For this reason boiling points need to be measured at a specific atmospheric pressure. This is 1 atmosphere pressure, 1 x 105 Nm-2.
The boiling point is a measure of the strength of the interparticular forces within the body of a liquid. It is a more convenient comparison for many volatile substances, as their melting points are often too low.
The change of state from solid to liquid takes place when the available energy is sufficient to literally shake the structure apart. In all solids the particles are held together by forces (as we should well know by now!). As energy is given to the particles, they vibrate more about their fixed positions in the solid, until eventually the vibrations are strong enough to tear them from these fixed positions. At this temperature the integrity of the solid structure breaks apart and the solid turns to liquid.
It is appropriate at this point to remind ourselves of the particular chacteristics of solids, liquids and gases in terms of motion and interparticular distance and bulk (real world scale) properties, in terms of shape and volume.
| microscopic | bulk | |||
|---|---|---|---|---|
| particle motion | interparticular distance | shape | volume | |
| solid | only vibration about fixed positions | very close | fixed | fixed |
| liquid | vibration, rotation and limited translation | very close | adopts the shape of the vessel | fixed |
| gas | vibration, rotation and translation | far apart | none, expands and diffuses | expands |
As we can see the particles in a liquid are able to move relative to one another, meaning that the forces that held them in positions in the solid structure have been overcome by the available energy.
The melting temperature can then be taken as a guide to the strength of interparticular forces within the bulk solid structure.
Relationship between structure and change of state
The stronger the forces between the particles in a compound, the higher the melting point and boiling point.
The highest melting points are found in network covalent solids such as diamond, graphite and silicon dioxide. A giant covalent structure has many strong bonds holding it together.
The next highest melting points are found in some metals, notably transitional metals. After these the giant ionic structures and then finally the simple covalent molecules have the lowest melting points of all.
Consideration of the nature of each type of structure allows us to differentiate between and within them.
In giant covalent compounds the term 'melting point' rather loses meaning, the compound must decompose rather than melt, as the bonds actually holding the atoms together have to be broken.
These have high melting points. The strength of inter-ionic attraction depends on the charge on the ion and the size of the ions. Structures with double charged ions have considerably higher melting points than structures with single charged ions.
| single charged ions only | double charged ions only | ||
| compound | m.p. / ºC | compound | m.p. / ºC |
| NaCl | 801 | MgO | 2800 |
| KCl | 776 | CaO | 2572 |
| LiF | 848 | MgS | >2000 |
| LiCl | 605 | CaS | 2400 |
It is possible to also appreciate from the table that the size of the ion has an influence. The smaller the ion the higher the charge density and the stronger the forces between the ions, resulting in a higher melting point. Hence KCl has a lower m.p. than NaCl.
Notice that lithium salts do not follow this trend. According to the theory of electrostatic attraction force, lithium chloride is expected to have a higher melting point than sodium chloride. The fact that it doesn't is explained by the high charge density of lithium ions (due to their small size). These polarise the much larger negative ion's electron shell, creating a degree of covalent character in the lattice that lowers the melting point.
Beryllium chloride is a covalent simple molecule for the same reasons. The high charge density on the beryllium 2+ ion (even smaller than lithium and with a double charge) would repolarise the chloride ions and produce covalent bonds.
The strength of the metallic lattice depends on the charge on the ions within the lattice (as each metal atom provides the same number of delocalised electrons to the sea of electrons as the charge on the ion), and also the ionic radius. There is a third factor, which goes beyond this syllabus, the crystal packing structure of the ions.
| group 1 | group 2 | ||
|---|---|---|---|
| Li | 180 | Be | 1287 |
| Na | 98 | Mg | 650 |
| K | 63 | Ca | 842 |
You can see that group 2 elements in general have much higher melting points associated with much stronger forces of attraction within the metal lattice. The anomalous value for magnesium in group 2 is due to the crystal structures of magnesium and calcium.
The melting points of simple molecular compounds are often very low. As a result is makes more sense to discuss their boiling points.
The definition of the boiling point is the temperature at which the vapour pressure of a liquid equals atomospheric pressure. Clearly, the atmospheric pressure on a given day is variable, so the boiling point is recorded at a pressure of 1 atmosphere, or 100 kPa.
The boiling point of simple molecular substances depends on the nature of the intermolecular forces. This in turn depends on the polarity of the molecules.There are three categories:
- Non-polar molecules
- Polar molecules
- Polar molecules with O-H or N-H groups
Non-polar molecules
The dispersion force is dependent on the size of the molecules, i.e. the relative molecular mass. It is also fine-tuned by the shape of the molecules.
- Increased mass = increased intermolecular attraction = higher boiling point.
- More spherical in shape = lower attractive force = lower boiling point.
Polar molecules
If the molecules are the same size then polarity can differentiate their boiling points. The force is additional to dispersion and the sum of the forces is consequently larger. Polar molecules have higher boiling points than non-polar compounds of the same relative mass.
Polar molecules with O-H or N-H bonds
These have stronger dipole-dipole interactions, called hydrogen bonds, in addition to dispersion forces. Their boiling points are higher than both similiar sized non-polar, or polar molecules.
