Electrical Conductivity Analysis

Table 5 shows the electrical conductivity of each element in the periodic table. This is known to be the amount of electrical current a material can carry. All of the elements that do not show a value on this table are known to be non-metals which should not be able to conduct electricity. Some of them do however show a value (some very small) that means they do conduct electricity. The general trend down the groups 1, 2, and 3 shows the electrical conductivity to be increasing as you down the groups. Although it seems that the opposite is happening for groups 4, 5 and 6 as it seems with this group that as you go down the overall value is becoming larger but towards the middle of the groups the value seems to spike in the period 3 elements.

Argon is a noble gas and so does not form a molecule with its other like atoms and so there is very little Van der Waals forces between the atoms and so it has the lowest melting point of the period 3 elements. The metals (sodium, magnesium, and aluminium) are all joined together with metallic bonding and so have a very strong attraction between the ions and the delocalized electrons which requires a lot of energy to overcome these bonds which results in a high melting point. The reason that the melting point increases between them is because the sodium ion has a charge of positive one and therefore has one delocalized electron for each sodium ion within the lattice structure but the magnesium ion has a charge of positive two and therefore has two delocalized electrons for every magnesium ion within the lattice structure. This means that the metallic lattice for magnesium is stronger than the lattice structure of sodium because there is a larger number of a delocalized electron in the structure which requires more energy to break the attraction between them. This means that magnesium will have a higher melting point than

The general trend here is that as you go across the group the temperature decreases. We can see from the table though that the temperature drops significantly from silicon to phosphorous by nearly 1600°C. This can be related back to the bonding between the silicon and the oxygen. The bonding here is ionic bonding which is a very strong form of bonding. The oxygen has taken the the four electrons from silicon’s outer shell and split them equally between the two oxygen atoms taking two each to complete their outer shells. The silicon ion ends up with a charge of positive four and each oxygen with a charge of negative two but both ions are very stable. This means that a lot of energy is required to break this kind of