Electronic Configurations and Types of Elements

Period comes from the outermost shell. The period number of an element equals the principal quantum number $n$ of the outermost (valence) shell being filled. An element whose highest-$n$ electrons sit in $n=4$ is in period 4, regardless of which subshell within that shell is being built.

Block comes from the last-filled subshell. Elements split into four blocks depending on which subshell type receives the differentiating electron:

• $s$-block — last electron enters an $s$ subshell; outer configuration $ns^{1}$ or $ns^{2}$; Groups 1 and 2 (helium too, by configuration $1s^2$, though it is drawn with the noble gases). Soft, reactive metals.
• $p$-block — last electron enters a $p$ subshell; outer configuration $ns^{2}np^{1-6}$; Groups 13 to 18. Contains metals, non-metals, metalloids, and the exceptionally stable noble gases at $ns^2np^6$.
• $d$-block — last electron enters a $(n-1)d$ subshell while the outer shell is $ns^{1-2}$; outer configuration $(n-1)d^{1-10}ns^{1-2}$; Groups 3 to 12, the transition elements, sitting between the $s$- and $p$-blocks in periods 4–7.
• $f$-block — last electron enters an $(n-2)f$ subshell; outer configuration $(n-2)f^{1-14}(n-1)d^{0-1}ns^{2}$; the lanthanoids (after lanthanum) and actinoids (after actinium), together the inner-transition elements, shown as a detached strip below the main table.

Group number for the $d$-block is not just the outer-$s$ electron count — it needs the penultimate shell too:

where: $G$ is the group number, $n_{(n-1)d}$ is the number of electrons in the $(n-1)d$ subshell, and $n_{ns}$ is the number of electrons in the outer $ns$ subshell. For $s$- and $p$-block elements the group is read directly from the outer $ns^{a}np^{b}$ electron count instead ($G = a$ for $s$-block; $G = 10+a+b$ for $p$-block).

The classification is not an arbitrary label — it falls straight out of the Aufbau filling order and the shape of the periodic table itself, for three linked reasons.

One, the table is built by filling subshells in a fixed sequence. Aufbau order ($1s,2s,2p,3s,3p,4s,3d,4p,\dots$) means that as atomic number increases, new electrons keep entering the same type of subshell for a run of consecutive elements — two elements while $ns$ fills, six while $np$ fills, ten while $(n-1)d$ fills, fourteen while $(n-2)f$ fills. Grouping elements by "which subshell is currently filling" is therefore identical to grouping them by column-runs in the table: that is exactly what blocks are.

Two, the last electron, not the whole configuration, decides chemistry. Even though a $d$-block atom has full $ns$ and inner shells, its outermost, most loosely bound electrons available for bonding are the ones in the partially filled $(n-1)d$ subshell alongside $ns$ — so calling it "$d$-block" correctly flags where its distinctive chemistry (variable oxidation states, coloured ions, catalytic behaviour) comes from, none of which is required knowledge here but explains why this filling-based label, not a numerically-last-written subshell, is used.

Three, stability exceptions are still explained by the same last-electron logic. A half-filled ($d^5$, $f^7$) or fully filled ($d^{10}$, $f^{14}$) subshell is unusually stable (extra exchange energy), so some atoms promote one $ns$ electron into the $(n-1)d$ or $(n-2)f$ subshell to reach that stable arrangement — chromium ($3d^5\,4s^1$ instead of the naive $3d^4\,4s^2$) and copper ($3d^{10}\,4s^1$ instead of $3d^9\,4s^2$) are the two you must know. The block assignment does not change: the last electron in both still enters a $d$ subshell, so both remain $d$-block.

Question: Classify chromium ($Z = 24$) and zinc ($Z = 30$) into their blocks, and state each one's period and group.

Step 1. Write chromium's configuration by Aufbau order and apply the stability exception: the naive prediction $1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2\,3d^4$ shifts one $4s$ electron into $3d$ to reach the extra-stable half-filled subshell, giving the actual configuration $1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^1\,3d^5$.

Step 2. Identify chromium's last-filled subshell and block: the differentiating electron is in $3d$, so chromium is a $d$-block element. Period: outermost principal shell used is $n=4$, so period 4. Group from equation (1): $G = n_{(n-1)d} + n_{ns} = 5 + 1 = 6$, so Group 6.

Step 3. Write zinc's configuration by Aufbau order, with no exception needed: $1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2\,3d^{10}$.

Step 4. Identify zinc's last-filled subshell and block: the last electron completes $3d^{10}$, a fully filled $d$ subshell, so zinc is still classified $d$-block by filling logic even though its $d$ subshell is now complete. Period: $n=4$, so period 4. Group from equation (1): $G = 10 + 2 = 12$, so Group 12 — matching zinc's known position at the end of the transition row. Answer: chromium is $d$-block, period 4, Group 6 (configuration $[\text{Ar}]\,3d^5\,4s^1$); zinc is $d$-block, period 4, Group 12 (configuration $[\text{Ar}]\,3d^{10}\,4s^2$).

• Reading the block off the numerically highest subshell written in the configuration instead of the one that received the electron last in Aufbau order — for $d$- and $f$-block atoms these can point to different subshells.
• Forgetting that $d$-block group numbers need $(n-1)d$ electrons added to $ns$ electrons together, not the outer $s$ count alone.
• Treating chromium's and copper's shifted configurations as errors to "correct" rather than the expected extra-stability exceptions you are required to know and apply.
• Placing helium in the $p$-block by its visual position above the noble gases, when its configuration $1s^2$ makes it structurally $s$-block.

Ask the tutor: want to practice classifying a few more elements — say vanadium, silver, or a lanthanoid — into block, period, and group? Ask the AI tutor for extra worked examples.