Transcript Periodicity
http://teachable.net/res.asp?r=1494 Periodicity Foundation Chemistry Semester 1 Dr Joanne Nicholson Melting and boiling points of Period 3 Elements The values are plotted in Kelvins (K) to avoid having to use negative values! The metallic structures • Na, Mg, Al • Melting and boiling points rise across the three metals because of the increasing strength of the metallic bonds. • The number of electrons which each atom can contribute to the delocalised "sea of electrons" increases. The atoms also get smaller and have more protons as you go from sodium to magnesium to aluminium. The attractions and therefore the melting and boiling points increase because: • The charge on the metal ions increases from +1 to +3 (with magnesium at +2). • The number of delocalised electrons increases. •The ionic radius decreases as protons are added, meaning a higher charge density. Silicon • Silicon is a metalloid (an element with some of the properties of metals and some of the properties of non-metals). • Silicon has giant covalent bonding. It has a giant lattice structure similar to that of diamond, in which each silicon atom is covalently-bonded to four other silicon atoms in a tetrahedral arrangement. Silicon has a very melting point and boiling point because: • all the silicon atoms are held together by strong covalent bonds ... • which need a very large amount of energy to be broken. The Non-metals • P, S, Cl, Ar • Phosphorus, sulphur, chlorine and argon are simple molecular substances with only van der Waals attractions between the molecules. • Their melting or boiling points will be lower than those of the first four members of the period which have giant structures. • The sizes of the melting and boiling points are governed entirely by the sizes of the molecules. Remember the structures of the molecules: • The strength of the van der Waal’s forces decreases as the size of the molecule decreases ... • So the melting points and boiling points decrease in the order S8 > P4 > Cl2 > Ar Trends in Electrical Conductivity of Period 3 elements The metallic structures Sodium, magnesium and aluminium are all metals. They have metallic bonding, in which positive metal ions are attracted to delocalised electrons. The delocalised electrons are free to move and carry charge. Going from sodium to aluminium: • the number of delocalised electrons increases ... • there are more electrons which can move and carry charge ... • so the electrical conductivity increases. Silicon • Silicon is called a semiconductor because it is able to conduct electricity only at high temperatures and not at room temperature. • At high temperatures, electrons are promoted to higher energy levels and are free to move around. The Non-metals • The remaining elements in Period 3 do not conduct electricity: • In phosphorus, sulphur and chlorine, the outer electrons are not free to move and carry charge because they are held strongly in covalent bonds. • In argon (which exists as single atoms) the outer electrons are not free to move and carry charge because they are held strongly in a stable third energy level.