Elements in Period 3 of Periodic Table

Period 3 Elements – Summary of Physical Properties

Advanced Chemistry Book + website

The elements in Period 3 of the periodic table are the elements sodium to argon. They represent the most straightforward trend in properties.
As the atomic number of the elements increases across a period, the trend is from metallic to a non-metallic character.
Electronegativity is a measure of an element's ability to attract a shared electron pair to itself. Metals have a smaller electronegativity than 1.7 and non metals greater than 2.4.
As the size of an atom decreases the electronegativity increases; so the value increases across a period and decreases down a group.
Melting Points: The melting points across Period 3 rise with increasing atomic number until silicon after which they fall dramatically.

The metallic structures
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" increase. The atoms also get smaller and have more protons as you go from sodium to magnesium to aluminum. The attractions and therefore the melting and boiling points increase because the nuclei of the atoms are getting more positively charged.
Silicon: silicon is a giant covalent structure, formed like a diamond, where the atoms are bonded covalently. Each atom is bonded to four other silicon atoms. Breaking all these bonds takes a great amount of energy and that's why silicon has the highest melting point of 1410°C.
Phosphorus, Sulphur and Chlorine are simple covalent molecules while Argon is held together simply by dispersion forces and that is why argon has a very low melting point.
Effects of atomic size
The value of the atomic radius of an element gives a measure of its size. The size of an atom has an influence on its ionization energy which is the minimum energy required to remove one or more of the outermost electrons. In period 3 the atom radius decreases as atomic number increases as the increase in nuclear charge pulls all electrons closer to the nucleus. The first ionization energy increases as the valence electron becomes closer to the nucleus. This is true for all the period 3 elements except for the two “dips” in the group 3 and 4 elements aluminum and sulphur. The dip in these two elements have given evidence that there are sub-levels called orbitals, because this is where some electrons are found thus needing less ionization energy as they will not be as close to the nucleus as they would be if there weren't any sub-levels.