Write the general outer electronic configuration of s-, p-, d- and f-block elements.

Exercises · 3.29

Q. Write the general outer electronic configuration of s-, p-, d- and f-block elements.

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Concept understanding — Periodic Table Blocks

The Intuition: Why "Blocks" at All?

Imagine you're building a house of cards. Each card has a specific shape and a specific place where it fits. The periodic table is like that house — but instead of cards, we have elements, and instead of shapes, we have electron configurations.

The periodic table is arranged in rows (periods) and columns (groups). But if you look closely, you'll notice that the table isn't a perfect rectangle. There's a detached island of elements (the f-block) floating below, and the main body has a strange "staircase" shape. That shape isn't random — it's dictated by which orbital the last electron enters.

That's the core idea: A block is a set of elements whose last electron enters the same type of orbital (s, p, d, or f).

The Precise Statement

Periodic Table Blocks are regions of the periodic table where elements share the same valence subshell — the subshell being filled as you move across that block.

There are four blocks, named after the four types of atomic orbitals:

Block	Orbital being filled	Location in the table	Number of groups
s-block	$ns$	Leftmost 2 columns (Groups 1 & 2)	2
p-block	$np$	Rightmost 6 columns (Groups 13–18)	6
d-block	$(n-1)d$	Middle 10 columns (Groups 3–12)	10
f-block	$(n-2)f$	Two rows below the main table (Lanthanides & Actinides)	14

Note

The "n" in the orbital notation refers to the principal quantum number (the period number). Notice how for d and f blocks, the orbital being filled has a lower n than the period you're in. That's because of the Aufbau principle — orbitals fill in order of increasing energy, and 4s fills before 3d, etc.

How to Read the Blocks

s-block (Groups 1 & 2)

Last electron enters an $s$ orbital.
Examples: Hydrogen ( $1s^1$ ), Lithium ( $2s^1$ ), Beryllium ( $2s^2$ ).
These are highly reactive metals (except H and He). They lose their $s$ electron(s) easily.

p-block (Groups 13–18)

Last electron enters a $p$ orbital.
Examples: Carbon ( $2p^2$ ), Oxygen ( $2p^4$ ), Chlorine ( $3p^5$ ).
This block contains metals, non-metals, and metalloids — the most chemically diverse block.

d-block (Groups 3–12)

Last electron enters a $d$ orbital — specifically, the $(n-1)d$ subshell.
Examples: Iron ( $3d^6$ ), Copper ( $3d^{10}$ ), Zinc ( $3d^{10}$ ).
These are transition metals. They often have variable oxidation states and form coloured compounds.

f-block (Lanthanides & Actinides)

Last electron enters an $f$ orbital — specifically, the $(n-2)f$ subshell.
Examples: Cerium ( $4f^1$ ), Uranium ( $5f^3$ ).
These are inner transition metals. They are placed below to keep the table from being absurdly wide.

Watch out

A common mistake: thinking that the block tells you the group number. It doesn't. The block tells you the orbital type, not the group. For example, both Carbon (Group 14) and Oxygen (Group 16) are in the p-block, but they're in different groups.

Why This Matters

Knowing the block of an element tells you three things instantly:

Which orbital is being filled — the heart of its electron configuration. …

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The general outer electronic configurations are based on which subshell (s, p, d, or f) is being filled last. For s-block: $ns^{1-2}$ ; p-block: $ns^2 np^{1-6}$ ; d-block: $(n-1)d^{1-10} ns^{0-2}$ ; f-block: $(n-2)f^{1-14} (n-1)d^{0-1} ns^2$ .

The Aufbau principle tells us that electrons fill orbitals in order of increasing energy. The "block" of an element in the periodic table is determined by the subshell that receives the last electron as we build up the atom. This is the key to writing general configurations.

Let’s break it down block by block.

s-block elements (Groups 1 and 2) The last electron enters an s-orbital of the outermost shell (the shell with the highest n). For Group 1, it’s $ns^1$ ; for Group 2, it’s $ns^2$ . So the general configuration is:

ns^{1-2}

Here, $n$ is the principal quantum number of the valence shell (e.g., $n=2$ for Li, $n=3$ for Na).

p-block elements (Groups 13 to 18) The last electron enters a p-orbital of the outermost shell. The s-subshell of that shell is already filled ( $ns^2$ ). The p-subshell can hold 1 to 6 electrons. So:

ns^2 np^{1-6}

For example, carbon (Group 14) has $2s^2 2p^2$ .

d-block elements (Groups 3 to 12, the transition metals) Here, the last electron enters a d-orbital of the second-last shell (n–1). The outermost s-subshell ( $ns$ ) is usually filled first (with 1 or 2 electrons) before the d-orbitals begin filling. So the general configuration is:

(n-1)d^{1-10} ns^{0-2}

Watch out

Two common exceptions: chromium (Cr) has $3d^5 4s^1$ (not $3d^4 4s^2$ ) and copper (Cu) has $3d^{10} 4s^1$ (not $3d^9 4s^2$ ). These arise because half-filled and fully-filled d-subshells are extra stable. The general formula still holds as a range, but these exceptions are exam favourites. …

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Q.Write the general outer electronic configuration of s-, p-, d- and f-block elements.

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