Chemical Periodicity

One period, one lap of the whole table

Period 3 runs $\text{Na} \rightarrow \text{Mg} \rightarrow \text{Al} \rightarrow \text{Si} \rightarrow \text{P} \rightarrow \text{S} \rightarrow \text{Cl} \rightarrow \text{Ar}$: three metals, then silicon (a semiconductor, between metal and non-metal), then four non-metals. Periodicity means that a property changes with the same shape across every period. If you learn the shape across Period 3, you can expect the same shape across Period 2 or Period 4.

Atomic radius: falls all the way across

You met this in the AS Atomic Structure note (“Atomic and ionic radius”). Crossing Period 3, each element has one more proton in the nucleus and one more electron in the same outer shell (n = 3). The inner shells do not change, so the shielding stays about the same. The larger nuclear charge therefore pulls the outer shell in more strongly, and the atomic radius falls from sodium to chlorine.

Mark schemes want both halves: nuclear charge increases AND shielding is similar, so there is stronger attraction between the nucleus and the outer electrons.

Ionic radius: two separate shrinking runs, not one

Ionic radius does not follow one smooth line. The metals (Na, Mg, Al) lose their outer electrons and keep only two shells (the neon arrangement, 10 electrons). The non-metals (P, S, Cl) gain electrons and keep three shells (the argon arrangement, 18 electrons). So the ions form two separate groups. Ions in a group with the same number of electrons are called isoelectronic.

• $\text{Na}^+, \text{Mg}^{2+}, \text{Al}^{3+}$: all 10 electrons. The nuclear charge rises (11, 12, 13 protons) on the same electrons, so the radius shrinks from $\text{Na}^+$ to $\text{Al}^{3+}$.
• $\text{P}^{3-}, \text{S}^{2-}, \text{Cl}^-$: all 18 electrons. Same reason: the radius shrinks from $\text{P}^{3-}$ to $\text{Cl}^-$.

Every anion is bigger than every cation, because the anions have one more occupied shell than the cations. So the pattern is: shrink across the cations, a big jump up, then shrink across the anions.

A quick check on an invented pair: compare $\text{Mg}^{2+}$ and $\text{S}^{2-}$. Magnesium loses 2 electrons and keeps 2 shells; sulfur gains 2 electrons and has 3 shells. $\text{S}^{2-}$ is much bigger, because it has one more shell, even though sulfur has more protons.

Melting point and conductivity: the same three structures decide both

Both properties come from the structure and bonding of each element, so build that picture once and use it twice.

• Na, Mg, Al — giant metallic. Positive ions sit in a lattice surrounded by a "sea" of delocalised electrons. From Na to Mg to Al, each atom gives more delocalised electrons (1, then 2, then 3) and forms a smaller ion with a bigger charge. Both make the metallic bonding stronger, so the melting point rises from Na to Al. Conductivity also rises, because there are more delocalised electrons to carry charge.
• Si — giant covalent (mark schemes also say giant molecular). Every silicon atom is joined to four others by strong covalent bonds all through the solid, like diamond in the AS States of Matter note. Melting means breaking a huge number of strong covalent bonds, so silicon has the highest melting point in the period. Its bonding electrons are held in fixed covalent bonds, not delocalised, so silicon conducts very poorly. It is a semiconductor.
• P, S, Cl, Ar — simple molecular (argon is single atoms). They exist as small molecules, $\text{P}_4$, $\text{S}_8$ and $\text{Cl}_2$, held to each other only by weak van der Waals' forces. The strong covalent bonds inside each molecule do not break on melting. So the melting points are low. They have no delocalised electrons and no ions, so they do not conduct.

Within the simple molecular group, the order is S > P > Cl > Ar. $\text{S}_8$ is the biggest molecule, with the most electrons, so it has the strongest van der Waals' forces. $\text{P}_4$ comes next, then $\text{Cl}_2$, then single Ar atoms.