Which one of the following general electronic configurations correctly represent a transition metal element?

1. Atomic Structure: Shells and Subshells (basic)

To understand how the periodic table is organized, we must first look at the internal architecture of an atom. At its core, an atom consists of a nucleus containing protons and neutrons, which represents the positive central portion of the atom Environment and Ecology, Majid Hussain, Major Crops and Cropping Patterns in India, p.100. Electrons, which carry a negative charge, do not just float randomly around this nucleus; they are organized into specific energy levels called shells.

Think of shells as concentric layers or floors in a building. These shells are labeled as K, L, M, N, and so on. For instance, the smallest shell, K, can hold only 2 electrons, which is why Helium (with 2 electrons) has a completely filled K shell Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.60. As we move further from the nucleus, the shells get larger and can hold more electrons according to the formula 2n² (where 'n' is the shell number). When an atom like Sodium loses an electron from its outer M shell, the L shell beneath it becomes the new outermost shell, creating a stable octet Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.46.

However, to master advanced chemistry, we must look even closer. Each shell is further divided into subshells, labeled s, p, d, and f. These subshells describe the specific shapes of the regions where electrons are likely to be found. For example, the first shell (K) has only one subshell (1s), while the third shell (M) has three (3s, 3p, and 3d). This hierarchy—Shells → Subshells → Orbitals—is the fundamental logic that dictates where an element sits on the periodic table and how it reacts with others.

Shell Label (n)	Max Electrons (2n²)	Subshells Present
K (n=1)	2	s
L (n=2)	8	s, p
M (n=3)	18	s, p, d
N (n=4)	32	s, p, d, f

Sources: Environment and Ecology, Majid Hussain, Major Crops and Cropping Patterns in India, p.100; Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.60; Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.46

2. The Modern Periodic Law and Table Layout (basic)

To understand the Modern Periodic Table, we must first understand the concept of periodicity. Much like the phases of the Moon or the changing of seasons repeat in a predictable cycle, the chemical and physical properties of elements also repeat at regular intervals Science, Class VIII, Keeping Time with the Skies, p.178. This led Henry Moseley to formulate the Modern Periodic Law, which states that the properties of elements are a periodic function of their atomic numbers (the number of protons), rather than their atomic weights. This shift was revolutionary because it aligned the table with the internal electronic structure of atoms.

The layout of the table is organized into 18 vertical columns called Groups and 7 horizontal rows called Periods. The elements are categorized into four main blocks based on the orbital being filled by the last electron: s, p, d, and f. Elements in the same group share the same outer electronic configuration, which is why they exhibit similar chemical behaviors. For example, the reactivity series shows how metals like Potassium (K) and Sodium (Na) are grouped near each other because they are highly reactive and belong to the same s-block family Science, class X, Metals and Non-metals, p.45.

A significant portion of the table is occupied by the d-block (Groups 3 to 12), known as the transition metals. These elements are unique because they fill their inner d-orbitals while their outer s-shell is already partially or fully occupied. The general electronic configuration for these elements is (n – 1) d¹⁻¹⁰ ns⁰⁻². Here, 'n' represents the outermost shell, and '(n – 1)' indicates that the electrons are entering a shell one level deeper. This configuration explains why transition metals often have variable oxidation states and form colorful compounds.

Feature	Groups (Columns)	Periods (Rows)
Total Number	18	7
Significance	Identical outer shell electron count; similar reactivity.	Represents the highest energy level (shell) being filled.

Sources: Science, Class VIII, Keeping Time with the Skies, p.178; Science, class X, Metals and Non-metals, p.45

3. Rules for Filling Electrons (Aufbau Principle) (intermediate)

To understand how the periodic table is organized, we must first understand the Aufbau Principle (from the German word Aufbau, meaning 'building up'). This principle dictates that in the ground state of an atom, electrons fill atomic orbitals in order of increasing energy levels. While basic chemistry often describes electron distribution in simple shells like K, L, M, and N Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.47, the Aufbau principle provides the precise roadmap for how subshells (s, p, d, f) are populated.

The energy of an orbital is determined by the (n + l) rule, where 'n' is the principal quantum number (shell) and 'l' is the azimuthal quantum number (subshell type). Orbitals with a lower (n + l) value are filled first. If two orbitals have the same (n + l) value, the one with the lower 'n' is filled first. This explains a common point of confusion: why the 4s orbital is filled before the 3d orbital. For 4s, n+l is 4+0=4; for 3d, n+l is 3+2=5. Since 4 is less than 5, the 4s subshell is lower in energy and gets its electrons first.

As we move into the transition metals, we encounter a unique filling pattern. Here, electrons begin to occupy the (n – 1)d subshell. This means the electrons are actually filling an 'inner' shell even though the 'outer' s-shell already has electrons. This is why the general configuration for d-block elements is written as (n – 1) d¹⁻¹⁰ ns⁰⁻². It reflects that the d-orbital can hold up to 10 electrons and is technically one energy level 'below' the outermost valence shell (n). This transition in filling subshells is what defines the different blocks (s, p, d, f) of the modern periodic table.

Sources: Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.47

4. Division of the Periodic Table into Blocks (intermediate)

To understand the Periodic Table deeply, we must look beyond just rows and columns and see it as a map of electron addresses. The table is divided into four distinct blocks—s, p, d, and f. Each block is named after the specific subshell (orbital) that is being filled by the last electron entering the atom. This organization isn't just for convenience; it tells us exactly how an element will behave chemically. For instance, while we often look at basic shell configurations (K, L, M, N) to see valence electrons Science, Class X, Metals and Non-metals, p.47, the block system provides a more surgical view of the atom's energy levels.

The d-block (Groups 3 to 12) is particularly fascinating. These elements are known as transition metals because they represent a bridge or 'transition' between the highly reactive metals of the s-block and the non-metals of the p-block. Their general electronic configuration is represented as (n–1) d¹⁻¹⁰ ns⁰⁻². The (n-1) notation is the most important part of this formula—it tells us that as we move across the period, electrons are actually filling an inner d-orbital belonging to an energy level lower than the outermost shell. While most metals are solid at room temperature and have high melting points Science, Class X, Metals and Non-metals, p.39, it is this unique d-electron structure that gives transition metals their characteristic properties like variable oxidation states and catalytic activity.

Here is a quick comparison of how the blocks are structured:

Block	Groups	Characteristic
s-block	1 and 2	Valence electrons enter the outermost s-orbital.
p-block	13 to 18	Valence electrons enter the outermost p-orbital (includes metals, metalloids, and non-metals).
d-block	3 to 12	Electrons fill the inner (n-1) d-subshell; known as transition elements.
f-block	Bottom two rows	Electrons fill the (n-2) f-subshell; known as inner-transition elements.

Sources: Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.47; Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.39

5. Characteristics of s-block and p-block Elements (intermediate)

To understand the architecture of the Periodic Table, we first look at the s-block and p-block elements, collectively known as the Representative Elements. The s-block consists of Group 1 (Alkali Metals) and Group 2 (Alkaline Earth Metals). These elements are characterized by the filling of the s-orbital in their outermost shell (valence shell). They are typically soft, highly reactive metals with low ionization enthalpies, meaning they lose electrons easily to achieve a stable noble gas configuration. For instance, alkali metals like Sodium (Na) and Potassium (K) are so soft they can be cut with a knife Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.40. Because they lose electrons so readily, they form strong ionic bonds and their oxides are generally basic in nature Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.41.

In contrast, the p-block (Groups 13 to 18) is the most diverse region of the periodic table. It is the only block that contains metals, non-metals, and metalloids. As we move across a period in the p-block, the character shifts from metallic to non-metallic. For example, while Aluminium is a metal, Carbon and Oxygen are non-metals, and Silicon acts as a metalloid. The p-block also contains the Noble Gases (Group 18), which have a completely filled valence shell (ns² np⁶), making them chemically inert Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.46. Interestingly, many p-block elements show allotropy—the ability to exist in different physical forms, such as Carbon existing as both the hardest natural substance, Diamond, and the electrical conductor, Graphite Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.40.

Feature	s-block Elements	p-block Elements
Nature	Exclusively reactive metals.	Includes metals, non-metals, and gases.
Oxides	Primarily basic (e.g., Na₂O).	Acidic or Amphoteric (e.g., Al₂O₃).
Valence Electrons	1 or 2 (ns¹ or ns²).	3 to 8 (ns² np¹⁻⁶).

From a geographical and industrial perspective, these two blocks dominate the Earth's crust. Oxygen and Silicon (both p-block) are the two most abundant elements by weight in the crust, followed by Aluminium (p-block) and Iron (d-block) Physical Geography by PMF IAS, Earth's Interior, p.53. Understanding these blocks is crucial because their reactivity—the tendency to attain a completely filled valence shell—dictates how they form the compounds that make up our world Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.46.

Sources: Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.40; Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.41; Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.46; Physical Geography by PMF IAS, Earth's Interior, p.53

6. Inner Transition Elements (f-block) (exam-level)

The Inner Transition Elements, also known as the f-block elements, consist of two series of elements—the Lanthanides and the Actinides—that are usually displayed in two separate rows at the bottom of the periodic table. They are called 'inner transition' because the electrons being added fill the f-orbitals of the anti-penultimate shell (the shell two levels below the outermost valence shell). While representative elements like Sodium (Na) or Magnesium (Mg) focus on filling the outermost shell to reach an octet Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.47, these elements possess a more complex internal structure.

The general electronic configuration for these elements is (n-2)f¹⁻¹⁴ (n-1)d⁰⁻¹ ns². Here, 'n' represents the outermost shell. Because the f-orbitals are buried deep within the atom's electronic structure, they are shielded by the outer electrons. This shielding leads to a unique phenomenon known as Lanthanide Contraction, where the atomic and ionic radii decrease steadily as the atomic number increases, significantly impacting the chemical behavior of subsequent elements in the periodic table.

The two series have distinct characteristics that are vital for both chemistry and environmental science. The Actinides, for instance, are primarily known for their radioactivity. Elements like Thorium and Uranium spontaneously emit alpha and beta particles or gamma rays as their nuclei disintegrate Environment, Shankar IAS Academy, Environmental Pollution, p.82. This makes them central to discussions regarding both nuclear energy and environmental safety.

Feature	Lanthanides (4f series)	Actinides (5f series)
Shell Filled	4f subshell	5f subshell
Radioactivity	Mostly non-radioactive (except Promethium)	All are radioactive
Oxidation States	Principally +3	Variable (+3, +4, +5, +6, +7)

Sources: Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.47; Environment, Shankar IAS Academy, Environmental Pollution, p.82

7. The Nature of Transition Metals (d-block) (exam-level)

To understand the transition metals, we must look at the 'heart' of the periodic table—specifically Groups 3 to 12. These elements are known as the d-block because their distinguishing electrons are added to the inner d-orbitals. Unlike the representative elements where the outermost shell is filling up, transition metals are unique because they fill a penultimate shell—that is, the energy level just inside the outermost one. This is represented by the general electronic configuration (n – 1) d¹⁻¹⁰ ns⁰⁻². Here, 'n' represents the outermost valence shell, but the chemical 'action' is happening in the (n-1) level. This internal complexity is why elements like iron (Fe) and copper (Cu) exhibit such diverse chemical personalities, including multiple oxidation states and the ability to form colorful compounds Science, Class VIII, Nature of Matter, p.123.

While we often call the entire d-block 'transition metals,' IUPAC (the global authority on chemistry) provides a stricter definition: a transition element is one that has an incomplete d-subshell either in its neutral state or in one of its common ions. This is a crucial distinction for the UPSC aspirant. For instance, Zinc (Zn) is in the d-block, but because its d-orbitals are completely full (d¹⁰) in both its metallic and ionic (Zn²⁺) forms, it technically doesn't meet the strict IUPAC definition of a transition metal. However, it is still studied alongside them because it shares many metallic characteristics, such as the ability to form amphoteric oxides like ZnO, which react with both acids and bases Science, Class X, Metals and Non-metals, p.41.

The nature of these metals makes them indispensable in engineering and industry. Because of their unique electronic structure, they can easily mix with other elements to form alloys. A prime example is stainless steel, an alloy of iron that is far more durable and corrosion-resistant than pure iron Science, Class VIII, Nature of Matter, p.129. Their ability to switch between oxidation states also makes them excellent catalysts, speeding up chemical reactions without being consumed themselves. When you see a blue-green solution after reacting copper oxide with hydrochloric acid, you are witnessing the transition metal copper forming a complex ion (CuCl₂), a direct result of its available d-orbitals Science, Class X, Acids, Bases and Salts, p.21.

Sources: Science, Class VIII, NCERT, Nature of Matter: Elements, Compounds, and Mixtures, p.123, 129; Science, Class X, NCERT, Metals and Non-metals, p.41; Science, Class X, NCERT, Acids, Bases and Salts, p.21

8. General Electronic Configuration of d-block (exam-level)

In our journey through the periodic table, we reach the "bridge" between the highly reactive s-block metals and the p-block elements: the d-block elements. These elements, spanning Groups 3 to 12, are primarily known as transition metals. The defining characteristic of these elements is that their last electron enters the d-orbital of the penultimate (second-to-last) shell. This is why we use the notation (n – 1), where 'n' represents the outermost valence shell. For instance, in the 4th period, electrons fill the 3d orbital after the 4s orbital is occupied.

The general electronic configuration for the d-block is represented as (n – 1) d¹⁻¹⁰ ns⁰⁻². In this formula:

• (n – 1) d¹⁻¹⁰: The penultimate d-subshell is being progressively filled with 1 to 10 electrons.
• ns⁰⁻²: The outermost s-subshell typically holds 1 or 2 electrons, though it can be 0 in specific cases like Palladium (Pd).

You encounter these elements frequently in practical science. For example, Iron (Fe), Nickel (Ni), and Chromium (Cr) are transition metals known for their high electrical resistivity when used in alloys like Nichrome Science, class X (NCERT 2025 ed.), Electricity, p.179. Furthermore, the chemical behavior of these elements is often dictated by their ability to lose electrons from both the ns and (n – 1)d shells, leading to multiple oxidation states, such as Chromium VI (Hexavalent Chromium) used in steel hardening Environment, Shankar IAS Academy (ed 10th), Environmental Pollution, p.93.

Sources: Science, class X (NCERT 2025 ed.), Electricity, p.179; Environment, Shankar IAS Academy (ed 10th), Environmental Pollution, p.93

9. Solving the Original PYQ (exam-level)

Now that you have mastered the Aufbau principle and the structure of the periodic table, you can see how those building blocks converge in this question. The core concept here is the filling of the penultimate shell. In transition metals, the last electron enters the d-orbital of the shell just below the outermost valence shell. This is why we use the notation (n-1)d. As you recall from our study of electronic shells, the energy level of the (n-1)d orbital is slightly higher than the ns orbital but lower than the np orbital, which is why the d-block sits where it does in the periodic table.

To arrive at the correct answer, (D) (n – 1) d1–10 ns0–2, you must look closely at the electron ranges. While the d-subshell naturally holds 1 to 10 electrons, the outermost s-subshell is more variable. Why is it 0-2 and not just 2? Remember the exceptions we discussed: elements like Chromium and Copper have only 1 electron in their 4s orbital due to half-filled or fully-filled stability, and Palladium (Pd) has 0. This specific range is a common detail UPSC uses to test if a student truly understands the anomalous configurations within the d-block, as explained in NCERT Class 11 Chemistry.

UPSC often creates "distractor" options that look technically sophisticated but are fundamentally incorrect. Option (B) is a classic trap; the (n-2)f notation refers to inner transition elements (lanthanides and actinides), not transition metals. Option (A) incorrectly uses (n-2) for the d-block, which doesn't follow the energy level rules. Option (C) ignores the standard filling order entirely. By identifying that transition metals are defined by the (n-1)d filling process, you can quickly eliminate the others and focus on the correct valence shell representation.