Comprehension given below is followed by some multiple choice questions. Each question has one correct option.

NCERT Exemplar · Q12

Q. Comprehension given below is followed by some multiple choice questions. Each question has one correct option. Choose the correct option. In the modern periodic table, elements are arranged in order of increasing atomic numbers which is related to the electronic configuration. Depending upon the type of orbitals receiving the last electron, the elements in the periodic table have been divided into four blocks, viz, s, p, d and f. The modern periodic table consists of 7 periods and 18 groups. Each period begins with the filling of a new energy shell. In accordance with the Aufbau principle, the seven periods (1 to 7) have 2, 8, 8, 18, 18, 32 and 32 elements respectively. The seventh period is still incomplete. To avoid the periodic table being too long, the two series of f-block elements, called lanthanoids and actinoids are placed at the bottom of the main body of the periodic table.

(a) The element with atomic number 57 belongs to

(i) s-block

(ii) p-block

(iii) d-block

(iv) f-block

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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 element with atomic number 57 (lanthanum) is the first element of the lanthanoid series, but its last electron enters a 5d orbital, placing it in the d-block of the periodic table.

The key to solving this lies in understanding how the periodic table blocks are defined. The block an element belongs to is determined by the subshell that receives the last electron during its ground-state electronic configuration — not by its position in a series like the lanthanoids.

Atomic number 57 is lanthanum (La). Many students instinctively place it in the f-block because it sits at the start of the lanthanoid series in the table’s layout. But the actual electronic configuration tells a different story.

Write the configuration up to atomic number 57.

Following the Aufbau principle, the order of filling is:

$1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, \dots$

Up to xenon (atomic number 54), the configuration is:

$[\text{Xe}] = 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6$
Add the next three electrons (55, 56, 57).
- Caesium (55): $[\text{Xe}] 6s^1$
- Barium (56): $[\text{Xe}] 6s^2$
- Lanthanum (57): The next available orbital after $6s$ is $4f$ , but for lanthanum, the $4f$ orbital is higher in energy than $5d$ due to the $(n-2)f$ vs $(n-1)d$ energy crossover. So the 57th electron goes into $5d$ , giving: $[\text{Xe}] 6s^2 5d^1$
Identify the block. …

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