General Trends in Physical Properties Across Period 3 of the Periodic Table
Atomic Number (Z) increases across the period from left to right.
The modern Periodic Table is arranged in order of increasing atomic number.
As you go across the period from left to right, 1 more proton is added to the nucleus of the atom which increases its atomic number (Z) by 1.
Electron configuration increases by 1 across the period from let to right.
As one moves from one element to another on the right, one more proton (positive charge) is found in the nucleus, so 1 more electron (negative charge) must be added to an electron 'shell' in order to balance the charge.
The electron is added to the same electron 'shell' (energy level). For this reason, all the elements in Period 3 have the first electron 'shell' (energy level) filled with 2 electrons and the second electron 'shell' (energy level) filled with 8 electrons (the electronic configuration of Neon). Sodium begins a new electron 'shell' ( 3rdenergy level) with 1 electron, magnesium has 2 electrons in the third electron 'shell' (energy level), aluminium has 3 electrons in the third electron 'shell' (energy level) etc, until finally the third electron 'shell' (energy level) is filled with 8 electrons and the stable electronic configuration of the Noble Gas Argon is reached (2,8,8).
Atomic radius of the elements decrease across the Period from left to right.
As we move from left to right across the period one more proton is added to the nucleus of each successive atom, and one more electron is added to the same electron 'shell' (energy level) of each successive atom. The increased positive charge in the nucleus of each successive atom attracts all the electrons in the atom more strongly, so they are drawn in more closely towards the nucleus.
1st Ionization Energy increases across the Period from left to right.
First ionisation energy is the energy required to remove an electron from the gaseous atom:
M(g) → M+(g) + e-
The further away from the positively charged nucleus that a negatively charged electron is, the less strongly the electron is attracted to the nucleus and so the more easily that electron can be removed. So, as the atomic radius decreases from left to right across the Period so the 1st Ionization Energy increases.
Electronegativity increases across the Period from left to right.
Electronegativity is the relative tendency shown by a bonded atom to attract electrons to itself.
Typically, metals have low electronegativity, little ability to attract electrons, while non-metals have high electronegativity, greater ability to attract electrons.
Elements become less metallic in nature (more non-metallic) across the period from left to right.
In general metals are hard (EXCEPT Group 1 (IA) metals which are quite soft), have metallic lustre, high melting and boiling points (Except for mercury which is a liquid at room temperature, and the Group 1 (IA) metals which have low melting/boiling points compared to other metals) and good electrical conductivity.
In general, non-metals are dull, brittle, have low melting and boiling points and are electrical insulators (non-conductors of electricity).
Elements to the left of Period 3 exhibit metallic properties, elements to the right show non-metallic properties.
Silicon is a semi-metal (metalloid).
General Trends in Properties of Oxides and Chlorides Across Period 3
The bonding in oxides and chlorides of period 3 elements becomes more covalent across the period.
Group 1 and 2 metallic elements form ionic oxides and chlorides.
Aluminium (Group 13) forms an ionic oxide, but a covalent chloride.
Silicon (Group 14 semi-metal or metalloid) forms covalent chlorides and oxides, but the oxide forms a giant 3-dimensional covalent network similar to that of diamond.
The oxides and chlorides of the non-metals to the right of the period are all covalent, and the species exist as small, discrete, covalent molecules.
Melting point of oxides and chlorides decrease from left to right across Period 3.
Ionic oxides of the Group 1, 2 and 13 metallic elements have higher melting points than the covalent molecular oxides of non-metals on the right hand side of the period.
Note that the oxide of the semi-metal (metalloid) silicon is high because it forms a giant covalent network rather than discrete covalent molecules.
Ionic chlorides of Group 1 and 2 metallic elements have higher melting points than the covalent chlorides of the rest of the period 3 elements.
Note that the anhydrous chloride of aluminium is covalent and has a melting point lower than that of the ionic chlorides to the left of it in the period.
Molten ionic oxides and chlorides conduct electricity.
When molten, the ions in the ionic compounds are free to move and therefore the liquid can conduct electricity.
Note that solid ionic compounds will not conduct electricity.
The covalent oxides and chlorides will not conduct electricity when solid or liquid because there are no "charge carriers" available.
The exception is aluminium chloride because, when molten, there will be a few ions available to move, so it can conduct electricity but only very poorly.
Oxides become more acidic from left to right across period 3.
Ionic oxides of Group 1 and 2 metallic elements are basic.
Aluminium oxide is amphoteric, it is capable of acting as either an acid or a base.
Oxide of the semi-metal (metalloid) silicon is only weakly acidic.
Oxides of the non-metallic elements on the right of the period are all acidic.
Chlorides become more acidic across period 3 from left to right.
Ionic chlorides of Group 1 and 2 metallic elements are neutral (neither acidic nor basic).