S- and P-Block Elements

The table is divided into four blocks by the orbital that receives the differentiating electron. The s-block is Groups 1 and 2, whose valence electrons occupy an $ns$ orbital, and it also takes helium, because both of helium's electrons sit in the $1s$ orbital even though helium is printed at the head of the noble gases. The p-block runs from Group 13 to Group 18 as the $ns^2np^{1\text{ to }6}$ series. The d-block is the transition series, and the f-block is the pair of detached rows. The paper treats this as a spotter rather than a definition: among sulphur, carbon, potassium and neon the s-block element is potassium, and among lithium, caesium, iron and helium the one that does not belong is iron, because iron is filling a $3d$ orbital while helium, for all its shyness, still belongs to the s-block through its $1s^2$ configuration.

Every trend in the chapter descends from two pulls working against each other, the nuclear charge climbing across a period while each new row adds a shell that shields it. Atomic radius falls across a period, because electrons enter the same shell while the protons grow, and rises down a group, because each step adds a shell. Ionic radius follows the same shape but not the same size: a cation is smaller than its parent atom, since losing electrons either strips a shell or lets the nucleus tighten its hold on the rest, and an anion is larger, because the added electron dilutes the grip of the protons. Covalent radius is half the distance between two bonded identical atoms. Ionisation energy, the energy needed to pull an electron off a gaseous atom, climbs across a period and falls down a group, and electronegativity moves with it, which is why the Group IIA order from least to most electronegative runs $\text{Ba} < \text{Sr} < \text{Ca} < \text{Mg} < \text{Be}$.

Electron affinity is the one trend with a hitch worth holding onto, because it does not fall cleanly. Down the halogens it decreases from chlorine to iodine as the atom enlarges and the added electron is held further from the nucleus, but fluorine does not top the list. Fluorine sits below chlorine because the fluorine atom is so small that an incoming electron meets strong repulsion from the electrons already crowded into the second shell, so chlorine ends with the more negative electron affinity. Bond energy and bond length close the set, and they run in opposite directions: a bigger atom makes a longer bond and a longer bond is a weaker one, so hydrogen fluoride carries the strongest hydrogen halogen bond and hydrogen iodide the weakest.