Electronic Configuration in the Four Blocks

Section 7.2 already noted that the modern periodic table's overall shape divides into four blocks; this subsection explains WHY, in terms of electronic configuration. Each block is named after the subshell into which its elements' last electron is placed. The s-BLOCK elements place their last electron in an s subshell; since an s subshell has only one orbital (holding at most 2 electrons), the general outer configuration is $ns^{1-2}$, covering exactly groups 1 and 2, which sit at the left extreme of the table. The p-BLOCK elements place their last electron in a p subshell; with three degenerate p orbitals (holding up to 6 electrons), this covers groups 13 through 18, sitting on the right of the table, and its general outer configuration (from period 2 onward) is $ns^2np^{1-6}$. Group 18 (the noble gases) sits at the p-block's far end; helium is a notable exception, since its valence shell has n=1 and so has no p subshell at all ($1s^2$), yet it is still placed in group 18 of the p-block because, like the other noble gases, its valence shell is completely filled (a duplet rather than an octet). The d-BLOCK elements place their last electron in a d subshell, which exists only for shells with $n \geq 3$; by the (n+l) rule, the ns subshell always fills before the (n-1)d subshell, so the last electron only enters d after ns is already full. With five d orbitals (holding up to 10 electrons), the d-block spans ten groups (3 through 12) at the table's centre, with general outer configuration $ns^{0-2}(n-1)d^{1-10}$ — though some elements deviate from the plain Aufbau-predicted configuration because a half-filled or fully-filled d subshell is extra stable; e.g. chromium (Z=24) is $4s^13d^5$ rather than the 'expected' $4s^23d^4$, since this way both the 4s and 3d subshells end up half-filled. The f-BLOCK elements place their last electron in an f subshell (seven orbitals, up to 14 electrons), giving general outer co …