Q.How would you account for the irregular variation of ionisation enthalpies (first and second) in the first series of the transition elements?
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Transition Elements: From Intuition to Definition
Imagine you're building a house with bricks. Most bricks are identical — you stack them in neat rows. But some bricks are special: they have extra slots on their sides where you can attach hooks, magnets, or other bricks. These special bricks can change the shape of the wall, conduct electricity, or even change colour when you heat them.
In the periodic table, transition elements are those special bricks. They are the metals that sit in the middle block — groups 3 to 12 — and they have a unique ability: they can use their inner electrons (not just the outermost ones) to form bonds, change oxidation states, and create colourful compounds.
The Intuition: Why "Transition"?
The word "transition" comes from the idea that these elements form a bridge between the highly reactive metals on the left (like sodium, magnesium) and the less reactive metals / non-metals on the right (like aluminium, silicon). Their properties are not extreme — they are in-between.
But the real reason they are special lies in their electron configuration.
The Precise Definition (IUPAC)
A transition element is an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell.
Let's unpack that.
1. The "d" sub-shell
Electrons are arranged in shells (K, L, M, N...) and sub-shells (s, p, d, f). The d sub-shell can hold a maximum of 10 electrons. In transition elements, the d sub-shell is being filled — but not completely.
For example, consider Iron (Fe):
- Atomic number 26
- Electron configuration: 1s22s22p63s23p64s23d6
- The 3d sub-shell has 6 electrons — it is incomplete (it can hold 10).
So iron is a transition element.
2. The "or" part — cations matter
Some elements have a complete d sub-shell in their neutral atom, but when they lose electrons to form positive ions (cations), the d sub-shell becomes incomplete.
Example: Zinc (Zn)
- Neutral Zn: [Ar]3d104s2 — the 3d sub-shell is full (10 electrons).
- But Zn commonly forms Zn2+: [Ar]3d10 — still full.
- So zinc is NOT a transition element by the IUPAC definition.
Example: Copper (Cu)
- Neutral Cu: [Ar]3d104s1 — 3d is full.
- But Cu2+: [Ar]3d9 — now the 3d sub-shell is incomplete.
- So copper IS a transition element.
A common mistake: thinking that all elements in the d-block (groups 3–12) are transition elements. They are not. Zinc, cadmium, and mercury are d-block elements but NOT transition elements because their common cations have a full d sub-shell.
The "d-block" vs "Transition Elements"
| d-block elements | Transition elements |
|---|---|
| Groups 3 to 12 | Groups 3 to 11 (excluding Zn, Cd, Hg) |
| All have d electrons | Must have incomplete d sub-shell in atom or common cation |
Why this formula?
Transition Element Definition: The "Why" Behind the Definition
The Core Definition
A transition element (IUPAC definition) is an element whose atom has an incomplete d-subshell in its ground state or can form stable ions with an incomplete d-subshell.
Key exam point: This definition covers both the neutral atom and its common ions.
Why This Definition? The Reasoning
1. The d-orbital filling pattern
In the periodic table, transition elements belong to the d-block (Groups 3–12). As we move across a period, electrons fill the (n−1)d orbitals after the ns orbital.
For example, in Period 4:
- Scandium (Sc): [Ar]3d14s2 — has one d-electron → transition element
- Zinc (Zn): [Ar]3d104s2 — d-subshell is full → not a transition element
2. The "incomplete d-subshell" condition
The definition focuses on incompleteness because:
- A full d-subshell (d10) is exceptionally stable (like a noble gas configuration for d-orbitals)
- Elements with d10 configurations do not show the characteristic properties of transition metals (variable oxidation states, coloured compounds, catalytic activity, paramagnetism)
3. Why include ions?
Consider Zinc (Zn):
- Ground state: [Ar]3d104s2 — d-subshell is full → not a transition element
- Common ion: Zn2+: [Ar]3d10 — still full → still not a transition element
Now consider Copper (Cu):
- Ground state: [Ar]3d104s1 — d-subshell is full → by atom definition alone, not a transition element
- But Cu2+: [Ar]3d9 — incomplete d-subshell → is a transition element
Therefore: The definition must include ions to correctly classify elements like Cu, which form stable ions with incomplete d-subshells.
The "Formula" — A Decision Tree
The definition can be expressed as a logical condition:
Transition element⟺(Atom has d1−9)∨(Stable ion has d1−9)
Where:
- d1−9 means incomplete d-subshell (1 to 9 electrons)
- d0 or d10 means complete (empty or full) → not a transition element
Common Exam Exceptions …
The key idea is that first ionisation enthalpy rises only mildly and irregularly across the series (d-electrons shield the 4s electrons somewhat, but not perfectly), while the real, sharp irregularity shows up in the second ionisation enthalpy.
Reasoning:
- From Sc to Zn, nuclear charge increases steadily, so first ionisation enthalpy (ΔiH1) generally rises — but only slightly and with minor bumps, because the added 3d electrons shield the outer 4s electron from the growing nuclear charge almost as effectively as it shields itself, keeping the rise gentle (unlike the steep rise across a normal period of non-transition elements).
- The second ionisation enthalpy (ΔiH2) shows the real irregularity: it is unusually high for Cr and Cu, because their singly-charged ions (Cr+=3d5, Cu+=3d10) already have an extra-stable, half-filled or fully-filled d-subshell — removing a second electron means breaking into that stable subshell, which costs much more energy. …
The first ionisation enthalpy rises only slightly and irregularly across the first transition series, because the added 3d electrons imperfectly shield the 4s electrons from the growing nuclear charge. The sharper irregularity is in the second ionisation enthalpy: it is unusually high for Cr and Cu (whose M+ ions have the extra-stable 3d5/3d10 configurations, so removing a further electron breaks that stability) and comparatively low for Mn and Zn (whose M+ ions still carry one loosely-held 4s electron beyond a stable d5/d10 core).
The Real Data (Table 4.2)
| Element | ΔiH1 | M+ configuration | ΔiH2 |
|---|---|---|---|
| Sc | 631 | 3d14s1 | 1235 |
| Ti | 656 | 3d24s1 | 1309 |
| V | 650 | 3d34s1 | 1414 |
| Cr | 653 | 3d5 | 1592 |
| Mn | 717 | 3d54s1 | 1509 |
| Fe | 762 | 3d64s1 | 1561 |
| Co | 758 | 3d74s1 | 1644 |
| Ni | 736 | 3d84s1 | 1752 |
| Cu | 745 | 3d10 | 1958 |
| Zn | 906 | 3d104s1 | 1734 |
Step-by-Step Reasoning
1. First ionisation enthalpy — a gentle, only mildly irregular rise.
Across Sc→Zn, nuclear charge rises by one unit each step, so ΔiH1 generally increases (631 → 906). But the rise is much gentler than across a normal (non-transition) period, because each new electron is added to an inner 3d orbital rather than the outer shell — a 3d electron shields the 4s electrons from the nucleus almost as effectively as another 4s electron would, so the effective nuclear charge felt by the valence electron increases only slowly. This is why the chapter describes the first-ionisation-enthalpy trend as "irregular... though of little chemical significance," without pinning the irregularity to any one specific element.
2. Second ionisation enthalpy — the real, well-defined break.
ΔiH2 removes an electron from the singly-charged ion M+, and here the electronic configuration of M+ matters directly:
- Chromium: neutral Cr is 3d54s1 (the well-known half-filled-stability exception), so Cr+ is already 3d5 — a stable, half-filled d-subshell. Removing a second electron means breaking into this stable arrangement, so ΔiH2 for Cr (1592) is unusually high.
- Copper: neutral Cu is 3d104s1, so Cu+ is 3d10 — a stable, fully-filled subshell. Breaking into it likewise makes Cu's ΔiH2 (1958) the highest in the row.
- Manganese: neutral Mn is 3d54s2, so Mn+ is 3d54s1 — the stable 3d5 core is already intact, with one "spare" 4s electron still to remove. Taking that easy 4s electron gives Mn a comparatively low ΔiH2 (1509) — a dip just below Cr's spike, not a peak. …
Method: Configuration-and-Data Analysis (Table 4.2)
This method explains the irregular variation by pairing each element's electronic configurations (atom and M+ ion) with the actual printed data, instead of relying on remembered "dip" rules.
Steps
Step 1: Write the configurations that matter
ΔiH1 removes an electron from the neutral atom M; ΔiH2 removes one from the M+ ion. So write both configurations for each element (Table 4.2's M and M+ rows).
Step 2: Check the first ionisation enthalpy against the data
Table 4.2 (kJ/mol): Sc 631, Ti 656, V 650, Cr 653, Mn 717, Fe 762, Co 758, Ni 736, Cu 745, Zn 906. The rise is gentle and only mildly irregular — each added 3d electron shields the 4s electrons from the growing nuclear charge, so the effective nuclear charge climbs slowly. Note there is no clean dip at Cr or Cu: Cr (653) is slightly above V (650), and Cu (745) is slightly above Ni (736).
Step 3: Explain the second ionisation enthalpy — where the sharp irregularity lives
- Cr⁺ is 3d5 (stable, half-filled) — removing another electron breaks into it, so ΔiH2 is unusually high (1592).
- Cu⁺ is 3d10 (stable, fully filled) — breaking into it makes Cu's ΔiH2 the highest in the row (1958).
- Mn⁺ is 3d54s1 — the second electron comes from the spare 4s, leaving the stable 3d5 intact, so ΔiH2 is comparatively low (1509).
- Zn⁺ is 3d104s1 — same pattern, so Zn's ΔiH2 (1734) sits below Cu's.
Step 4: Summarise (all values from Table 4.2, kJ/mol)
| Element | ΔiH1 | ΔiH2 | Reason |
|---|---|---|---|
| Cr | 653 — no dip (just above V's 650) | 1592 — unusually high | Cr⁺ is stable 3d5; a second removal breaks it |
| Mn | 717 — above Cr | 1509 — comparatively low | Mn⁺ is 3d54s1; the spare 4s electron goes easily |
Here is a breakdown of the common mistakes students make when explaining the irregular variation of ionisation enthalpies in the first transition series, along with how to avoid each.
The Core Concept (The "Why")
The irregular variation is attributed to the varying degrees of stability of the different 3d configurations (e.g., d0, d5, d10 are exceptionally stable). Two things follow from the real data (Table 4.2):
- First ionisation enthalpy (ΔiH1): rises gently and only mildly irregularly across Sc → Zn (631 → 906 kJ/mol), because each added 3d electron partially shields the 4s electrons from the growing nuclear charge. There is no clean dip at Cr or Cu.
- Second ionisation enthalpy (ΔiH2): shows the sharp, well-defined irregularity — unusually high for Cr (1592) and Cu (1958), whose M+ ions are the stable 3d5 and 3d10; comparatively low for Mn (1509) and Zn (1734), whose M+ ions still carry one spare, easily-removed 4s electron.
Common Mistake #1: Claiming the First Ionisation Enthalpy "Dips" at Cr and Cu
The Mistake: Students reason "Cr and Cu each have only one 4s electron, and removing it gives a stable d5/d10 ion — so their ΔiH1 must dip below their neighbours'."
Why it's wrong: Table 4.2's own numbers refute it: Cr's ΔiH1 (653) is slightly higher than V's (650), and Cu's (745) is slightly higher than Ni's (736). The stable-configuration effect shows up cleanly in ΔiH2, not as first-ionisation dips.
How to Avoid: Quote the data, not the remembered rule. If asked about ΔiH1, say the rise is gentle and only mildly irregular; save the d5/d10 stability argument for ΔiH2.
Common Mistake #2: Telling the Same Story for ΔiH1 and ΔiH2
The Mistake: Students say "Cr has low ΔiH1 and also low ΔiH2" (or high for both), applying one blanket rule to both quantities.
The Correct Logic: The two quantities remove electrons from different species:
- ΔiH1 (M → M⁺) depends on the neutral atom.
- ΔiH2 (M⁺ → M²⁺) depends on the M+ ion:
- Cr⁺ (3d5) → Cr²⁺ (3d4): destroys stable d5 → high ΔiH2.
- Cu⁺ (3d10) → Cu²⁺ (3d9): destroys stable d10 → high ΔiH2.
- Mn⁺ (3d54s1) → Mn²⁺ (3d5): removes the spare 4s electron, leaving stable d5 → low ΔiH2.
- Zn⁺ (3d104s1) → Zn²⁺ (3d10): same pattern → ΔiH2 below Cu's.
How to Avoid: Write the configurations of M, M⁺, and M²⁺ for each element. Ask: "Is the electron being removed a spare 4s electron (easy), or does it break a stable d5/d10 core (hard)?"
Common Mistake #3: Writing Wrong Ground-State Configurations
The Mistake: Students write Cr as 3d44s2 or Cu as 3d94s2, and then cannot explain the irregularities.
The Correct Logic: The actual ground-state configurations are the two exceptions:
- Cr: [Ar]3d54s1 (not 3d44s2)
- Cu: [Ar]3d104s1 (not 3d94s2)
How to Avoid: Memorise the two exceptions (Cr and Cu). For every other 3d element, the general 3dn4s2 rule holds.
Common Mistake #4: Blaming Only "Nuclear Charge" or "Shielding"
The Mistake: Students say "ionisation enthalpy increases because nuclear charge increases" and stop there. That explains the general rise, not the irregularities. …
Showing the 12 most recent of 36 on this concept.
- CBSE 2026Set 56/1/11 markMCQQ.Which of the following is not a transition metal ? (A) Sc (B) Ag (C) Hg (D) Cu
›Reveal solutionSolution
Transition metals are defined by having partially filled d-orbitals in their common oxidation states. Mercury (Hg) has a full d¹⁰ configuration in both its elemental and common +2 state, so it is not a transition metal. The correct answer is (C).
Why This Definition Matters
The classification of transition metals isn't about being a metal or having d-electrons — it's about partially filled d-orbitals in the atom or in any common ion. This is the IUPAC definition. If an element's d-subshell is completely full (d¹⁰) in all its common forms, it doesn't qualify, even if it sits in the d-block of the periodic table.
Let's check each option against this rule.
-
Scandium (Sc, Z=21)
Electronic configuration: [Ar]3d14s2.
Its common ion is Sc³⁺: [Ar] — the 3d orbital is empty.
But the atom has a partially filled d-orbital (3d¹). The definition says "in the atom or in any common oxidation state." Since the atom itself has an incomplete d-subshell, Sc is a transition metal.
-
Silver (Ag, Z=47)
Configuration: [Kr]4d105s1.
Common ion: Ag⁺ → [Kr]4d10 — full d-subshell.
However, silver also forms Ag²⁺ (e.g., in AgF₂), which has 4d9 — a partially filled d-orbital. Because at least one common oxidation state (Ag²⁺) has an incomplete d-subshell, Ag qualifies as a transition metal.
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Mercury (Hg, Z=80)
Configuration: [Xe]4f145d106s2.
Common ions: Hg²⁺ → [Xe]4f145d10 (full d¹⁰); Hg₂²⁺ (dimeric) also has each Hg with a full d¹⁰ core.
Mercury does not form any stable ion with a partially filled d-orbital. Its d-subshell is always full. Therefore, Hg is not a transition metal.
-
Copper (Cu, Z=29)
Configuration: [Ar]3d104s1. …
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- CBSE 2026Set 56/2/11 markMCQQ.Assertion (A) : Zinc, cadmium and mercury are not considered as transition elements. Reason (R) : These elements have completely filled orbitals in their ground state as well as in their common oxidation states.
›Reveal solutionSolution
The key idea is that transition elements are defined by having partially filled d-orbitals in their ground state or common oxidation states. Since Zn, Cd, and Hg have completely filled d¹⁰ configurations in both, they fail this criterion. Thus, both Assertion (A) and Reason (R) are true, and (R) correctly explains (A).
Why this approach works
The definition of a transition element is not arbitrary — it stems from the unique chemistry of d-block elements. The IUPAC defines a transition element as an element whose atom has a partially filled d-subshell, or which can give rise to cations with an incomplete d-subshell. This means we must check two things: the ground state electron configuration of the neutral atom, and the configurations of its common oxidation states. If either has a partially filled d-orbital, the element qualifies. If both are completely filled, it does not.
Zinc, cadmium, and mercury sit at the end of their respective d-block series (Group 12). Their neutral atoms have the configuration (n−1)d10ns2 — the d-subshell is full. Their common oxidation state is +2, formed by losing the two s-electrons, leaving (n−1)d10 — still full. No partially filled d-orbital appears anywhere. Hence, they are not transition elements.
Let’s walk through the reasoning step by step.
-
Recall the defining criterion for transition elements.
A transition element must have an atom or a common ion with an incomplete d-subshell. This is the official IUPAC definition. Elements that have completely filled d-orbitals in both the ground state and all common oxidation states are excluded.
-
Examine the ground state configurations of Zn, Cd, and Hg.
- Zinc (Z=30): [Ar]3d104s2
- Cadmium (Z=48): [Kr]4d105s2
- Mercury (Z=80): [Xe]4f145d106s2 In each case, the d-subshell is completely filled (d10). No partially filled d-orbital exists in the neutral atom.
-
Check their common oxidation states.
The most stable and common oxidation state for all three is +2. For example:
- Zn loses its two 4s electrons to form Zn2+: [Ar]3d10
- Cd loses its two 5s electrons to form Cd2+: [Kr]4d10
- Hg loses its two 6s electrons to form Hg2+: [Xe]4f145d10 In every case, the d-subshell remains completely filled. No partially filled d-orbital appears in any common oxidation state. …
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- CBSE 2026Set DZ1 markMCQQ.Which of the following is not a transition metal?(a) Cu(b) Cr(c) Fe(d) Na
›Reveal solutionSolution
A transition metal has a partially filled d-subshell (in the atom or a common ion). Cu, Cr and Fe qualify; Na does not — so option (d).
A transition element is defined as one whose atom or one of its stable ions has an incompletely filled d-orbital.
- Cu ([Ar]3d104s1): forms Cu2+ (3d9) — partially filled d → transition metal.
- Cr ([Ar]3d54s1) — partially filled d → transition metal. …
- CBSE 2026Set ANNUAL1 markMCQQ.Transition element among the following is(a) Zinc(b) Cadmium(c) Cerium(d) Rutherfordium
›Reveal solutionSolution
A transition element is defined as one having a partially (incompletely) filled d-subshell in the ground state or in any of its common oxidation states.
- Zinc (3d10 4s2) and Cadmium (4d10 5s2) have a completely filled d-subshell in the atom AND in their only common +2 ion, so they are NOT classified as transition elements (they are 'post-transition'/group 12 metals).
- Cerium is an f-block (lanthanide) element, not d-block. …
- CBSE 2026Set ANNUAL1 markQ.Why is Zn not considered as transition element? (Z = 30)
›Reveal solutionSolution
Transition elements must have partly filled d-orbitals in the elemental or common ionic state; zinc's d-subshell is always completely full, so it does not qualify.
Zinc (Z = 30) has the electronic configuration [Ar] 3d10 4s2. When it forms its common oxidation state, Zn2+, it loses the two 4s electrons, giving [Ar] 3d10 - the d-subshell remains completely filled (all 10 electrons present) in both the metal and its ion.
…
- CBSE 2026Set ANNUAL1 markMCQQ.Chromium belongs to which block?(a) s-block(b) p-block(c) d-block(d) f-block
›Reveal solutionSolution
Chromium (Z=24) has the configuration [Ar] 3d⁵ 4s¹, placing it in the d-block.
Chromium has atomic number 24 and electronic configuration [Ar] 3d⁵ 4s¹ (an exception to the expected 3d⁴4s², arising from the extra stability of a half-filled d-subshell). Since its last-entering electron occupies a …
- CBSE 2026Set ANNUAL1 markQ.______ block elements are called transition elements.
›Reveal solutionSolution
d-block elements are called transition elements.
The periodic table is divided into four blocks (s, p, d, f) based on which subshell receives the last (differentiating) electron. Elements in which the last electron enters a d-orbital form the d-block, spanning groups 3 to 12. These are called transition elements because most of them have partially filled d-orbitals in their elemental or common ionic states, giving them properties 'transitional' between the highly reactive s-b …
- CBSE 2026Set ANNUAL1 markMCQQ.Transition metals have incomplete(a) s-orbital(b) p-orbital(c) d-orbital(d) f-orbital
›Reveal solutionSolution
Transition metals have incomplete d-orbitals.
Transition elements are defined as those whose atoms or stable ions have partially filled (incomplete) d-orbitals. This partly filled d-subshell is responsible for their characteristic properties: var …
- CBSE 2026Set ANNUAL1 markMCQQ.Total number of elements are in d-block is ______(a) 39(b) 40(c) 38(d) 36
›Reveal solutionSolution
Four transition series of 10 elements each -> 40 d-block elements.
The d-block (transition elements) consists of elements in which the last electron enters a d-subshell. There are four such series:
- 3d series (Sc to Zn),
- 4d series (Y to Cd),
- 5d series (La/Hf to Hg),
- 6d series (Ac/Rf onward). …
- CBSE 2026Set ANNUAL1 markQ.Fill in the blank: Zn, Cd and Hg generally do not considered as a ______ elements.
›Reveal solutionSolution
Zn, Cd and Hg are not true transition elements because their d-orbitals are completely filled.
A transition element is defined as one that has a partially filled d-subshell in its atomic state or in one of its common oxidation (ionic) states. Zinc, cadmium and mercury have the configuration (n-1)d10 ns2, and even in their common +2 ions they are (n-1)d10 - the d-subshell is completely filled. Since they …
- CBSE 2025Set 56/6/11 markMCQQ.Which of the following is the softest metal ? (A) Zn (B) Sc (C) Cu (D) Fe
›Reveal solutionSolution
Softness in metals is determined by how easily their atoms slide past one another; among the transition elements listed, scandium (Sc) has the weakest metallic bonding due to having only one d-electron available for bonding, making it the softest. The answer is (B).
Understanding Metallic Softness
Hardness in metals arises from the strength of metallic bonding—the electrostatic attraction between the "sea" of delocalized electrons and the positive metal ions. The more electrons available for delocalization (especially from d-orbitals in transition metals), and the smaller the atomic radius, the stronger the bonding and the harder the metal.
Conversely, a soft metal has weaker metallic bonds. Its atoms can slide past one another more easily under stress, making it malleable and easy to deform.
Comparing the Four Metals
Let's examine each candidate by looking at their electronic configurations and the number of electrons contributing to metallic bonding:
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Zinc (Zn): Electronic configuration [Ar]3d104s2
All ten d-electrons are paired in filled orbitals. While the two 4s electrons participate in bonding, the filled d10 subshell contributes relatively little to directional bonding. Zinc is moderately soft but not the softest here.
-
Scandium (Sc): Electronic configuration [Ar]3d14s2
Only one d-electron is available. This is the first transition metal in the series, with minimal d-orbital participation in bonding. The metallic bond is weak because there are fewer electrons to delocalize and the d-orbitals are only just beginning to fill. Scandium is notably soft and can be cut with a knife.
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Copper (Cu): Electronic configuration [Ar]3d104s1 …
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- CBSE 2025Set D1 markMCQQ.Lutetium (Lu) is(a) a transition element(b) a lanthanide(c) an actinide(d) a p-block element
›Reveal solutionSolution
Lutetium (Z = 71) is the last of the fourteen lanthanide (4f-series) elements.
The lanthanoids (lanthanides) are the fourteen elements from cerium (Z = 58) to lutetium (Z = 71) in which the 4f subshell is progressively filled. Lutetium has the configuration [Xe] 4f14 5d1 6s2, comp …
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