Transition Elements (d-Block Elements): Properties, Trends, Oxidation States and Complex Formation
Complete Guide for JEE Advanced, NEET, IIT-JAM, GATE, CSIR-NET & BITSAT
Imagine a world without iron bridges, without cobalt-blue pigments, without the platinum catalyst that cleans your car's exhaust, without haemoglobin carrying oxygen in your blood — impossible. The transition elements are the workhorses of chemistry and of life itself. They sit in the middle of the periodic table, bridging the reactive s-block metals and the mostly covalent p-block elements, and they bring along an extraordinary bag of tricks: vivid colours, strong magnetism, variable oxidation states, catalytic genius, and an almost obsessive love of forming complex ions.
This article is your complete companion to mastering the d-block (transition elements) — built directly from the celebrated text Concise Inorganic Chemistry (J.D. Lee), with every concept reshaped for competitive exam success.
1. What Are Transition Elements? (Definition of d-Block Elements) 🔬
2. General Properties of Transition Elements
3. Periodic Trends in d-Block Elements
4. Oxidation States and Variable Valency
5. Complex Formation and Coordination Compounds
6. Chemical Properties of Transition Metals
7. Sc, Ti and V Groups: Important Trends
Three series of elements are formed by progressively filling the 3d, 4d, and 5d shells of electrons. Together these constitute the d-block elements. Their periodic table position is between the s-block (Groups 1–2) and p-block (Groups 13–18), hence the term transitional.
- First row (3d series): Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn (Groups 3–12)
- Second row (4d series): Y, Zr, Nb, Mo, Tc, Ru, Rh, Pd, Ag, Cd
- Third row (5d series): La, Hf, Ta, W, Re, Os, Ir, Pt, Au, Hg
- Incomplete fourth row: Ac and actinides (discussed with f-block)
Why Group 12 (Zn, Cd, Hg) is Different
These elements have a completely filled d-shell (d¹⁰) even in their ionic forms (Zn²⁺, Cd²⁺, Hg²⁺ are all d¹⁰). So they do not show the typical transition-metal behaviour: no variable oxidation states (beyond +II), no colour from d-d transitions, no ferromagnetism. They are the odd ones out.
2. Electronic Configuration — The Root of All Properties
The general electronic configuration of transition elements is: [Noble gas] (n-1)d1–10 ns1–2
Calcium (the s-block element just before the first row) has the configuration: 1s²2s²2p⁶3s²3p⁶4s². From Sc onward, electrons fill the 3d orbital. In theory we expect a perfectly regular filling: 3d¹ → 3d² → 3d³... up to 3d¹⁰. Reality has two famous exceptions:
| Element | Expected Config. | Actual Config. | Reason |
|---|---|---|---|
| Chromium (Cr) | [Ar] 3d⁴ 4s² | [Ar] 3d⁵ 4s¹ | Half-filled d (d⁵) is extra stable |
| Copper (Cu) | [Ar] 3d⁹ 4s² | [Ar] 3d¹⁰ 4s¹ | Fully filled d (d¹⁰) is extra stable |
Electronic Configurations of All First-Row Transition Elements
| Element | Symbol | Configuration | d-electrons |
|---|---|---|---|
| Scandium | Sc | [Ar] 3d¹ 4s² | 1 |
| Titanium | Ti | [Ar] 3d² 4s² | 2 |
| Vanadium | V | [Ar] 3d³ 4s² | 3 |
| Chromium | Cr | [Ar] 3d⁵ 4s¹ | 5 (exception) |
| Manganese | Mn | [Ar] 3d⁵ 4s² | 5 |
| Iron | Fe | [Ar] 3d⁶ 4s² | 6 |
| Cobalt | Co | [Ar] 3d⁷ 4s² | 7 |
| Nickel | Ni | [Ar] 3d⁸ 4s² | 8 |
| Copper | Cu | [Ar] 3d¹⁰ 4s¹ | 10 (exception) |
| Zinc | Zn | [Ar] 3d¹⁰ 4s² | 10 |
3. Variable Oxidation States — The Hallmark Feature
One of the most striking and exam-relevant features of transition metals is their ability to exist in multiple oxidation states. Unlike s-block elements (Na always +1, Ca always +2), transition metals can show anywhere from +1 to +8 or even negative oxidation states!
Why Variable Oxidation States?
The energy difference between (n-1)d and ns subshells is very small. So both the ns and (n-1)d electrons can participate in bonding and be lost to form ions. More d electrons = more possible oxidation states.
Fig. 1 — Oxidation state distribution across first-row transition elements (most stable states in bold font in text)
Key Trends in Oxidation States
- First five elements (Sc→Mn): Maximum oxidation state equals group number. Example: Mn (Group 7) shows max +VII in KMnO₄.
- Last five elements (Fe→Zn): After d⁵, tendency of all d electrons to participate in bonding decreases. Fe shows maximum +VI (not +VIII), while Ru and Os show +VIII.
- Higher oxidation states: Found in fluorides and oxides (F⁻ and O²⁻ are small, highly electronegative — they stabilise high oxidation states).
- Lower oxidation states: Stabilised by π-bonding ligands like CO and dipyridyl.
- Isostructural pairs: In highest oxidation state, properties resemble main group. Example: SO₄²⁻ (Group 16) and CrO₄²⁻ (Group 6) are isostructural; SiCl₄ and TiCl₄ are isostructural.
4. Metallic Character — Why They're All Metals
In the d-block, the penultimate (second-to-last) shell of electrons is expanding. This means the d electrons are being added to an inner shell, not the outermost one. The result? All transition elements are metals with:
- High electrical and thermal conductivity (delocalised electrons)
- Metallic lustre
- Hard, strong and ductile nature
- High melting and boiling points (due to strong metallic bonding involving d electrons)
- They form alloys with each other and with other metals
Melting Points — A Dramatic Story
Transition metals typically melt above 1000°C. Ten elements melt above 2000°C, and three melt above 3000°C: Ta (3000°C), W (3410°C) and Re (3180°C). Tungsten holds the record for highest melting point of any element. This contrasts sharply with s-block metals like Li (181°C) and Cs (29°C).
Density — Surprisingly Heavy!
Transition metals have small atomic volumes (nuclear charge is poorly screened by d electrons → all electrons pulled inward → compact atom → high density). Almost all have density > 5 g/cm³. The champions are osmium (22.57 g/cm³) and iridium (22.61 g/cm³) — the densest elements known. A football made of osmium measuring 30 cm in diameter would weigh 320 kg!
5. Size of Atoms and Ions — Lanthanide Contraction Explained
Atomic size follows an interesting pattern across the transition series:
Fig. 2 — Covalent radii decrease from Sc to Ni, then slightly increase at Cu and Zn
Key Size Trends
- Left to right across a row: Radii generally decrease because extra protons are added but d electrons shield the nucleus poorly (d electrons shield less effectively than p or s electrons).
- Near the end of the row: Slight increase at Cu and Zn due to d–d electron repulsion.
- Down a group (Sc→Y→La): Radii increase as expected. But from Group 4 onward, the increase from 2nd row to 3rd row is almost zero because of the Lanthanide Contraction.
Lanthanide Contraction — A Critical Concept
Between the second row (4d) and third row (5d) elements, 14 lanthanide elements (filling 4f orbitals) intervene. The 4f electrons are very poor shielders of nuclear charge. So across the 14 lanthanides, the nuclear charge increases by 14 units while shielding barely increases → each atom shrinks progressively. By the time we reach Hf (next to La), the size has contracted back to nearly the same as Zr (previous row)!
- Zr and Hf are almost identical in size → extremely difficult to separate (most difficult pair in periodic table)
- Nb and Ta have same ionic radii → also very difficult to separate
- 2nd and 3rd row transition elements are more similar to each other than either is to the 1st row
- 3rd row elements are denser, have higher melting points, and show stronger covalent bonding
6. Ionization Energies and Electrode Potentials
Ionization energies of transition metals are intermediate between s-block and p-block. First ionization energies vary from 541 kJ mol⁻¹ (La) to 1007 kJ mol⁻¹ (Hg) — comparable to Li and C respectively. This explains why transition metals are less electropositive than Group 1/2 and can form either ionic or covalent bonds depending on conditions.
Standard Electrode Potentials
Noble character (low reactivity with acids) is promoted by:
- High enthalpy of sublimation (strong metallic bonding)
- High ionization energies
- Low solvation enthalpy (large ionic radius → weak hydration)
The platinum group metals (Ru, Rh, Pd, Os, Ir, Pt) and gold are the most noble. Many transition metals dissolve in mineral acids, liberating H₂, but the platinum metals and gold require aqua regia (mixture of concentrated HCl and HNO₃).
7. Colour in Transition Metal Compounds
This is one of the most visually striking and theoretically rich topics in inorganic chemistry — and a favourite in all competitive exams!
Why Are Transition Metal Compounds Coloured?
When light passes through a material, certain wavelengths are absorbed. If absorption occurs in the visible region, the transmitted light has the complementary colour. Transition metal ions with incompletely filled d orbitals can absorb visible light by promoting an electron from a lower d level to a higher d level (d–d transition). This energy jump is small enough to fall in the visible spectrum.
Fig. 3 — Complementary colour chart for transition metal compounds
Factors That Change the Colour of a Complex
- Nature of the ligand: [Ni(NH₃)₆]²⁺ is blue, [Ni(H₂O)₆]²⁺ is green, [Ni(NO₂)₆]⁴⁻ is brown-red.
- Number of ligands
- Shape of the complex
- Oxidation state of the metal
When There Is No Colour
Some transition metal compounds are white/colourless:
- ZnSO₄, TiO₂: Zn²⁺ is d¹⁰ (full d), Ti⁴⁺ is d⁰ (empty d) → no d–d transitions possible.
- Sc³⁺, Ti⁴⁺, V⁵⁺, Cr⁶⁺, Mn⁷⁺: All have empty d orbitals → colourless in principle, but high oxidation states often form oxo-anions (e.g., CrO₄²⁻ yellow, MnO₄⁻ purple) that are intensely coloured due to charge transfer (not d–d).
8. Magnetic Properties — Paired vs. Unpaired Electrons
Magnetism in transition metal compounds is a direct window into their electronic structure — and one of the most numerical-heavy topics in exams.
Types of Magnetism
- Paramagnetism: Substance is attracted into a magnetic field. Caused by unpaired electrons. More unpaired electrons → stronger paramagnetism.
- Diamagnetism: Substance is weakly repelled from a magnetic field. All electron spins are paired. Diamagnetism is much weaker than paramagnetism.
- Ferromagnetism: Very strong magnetism where unpaired electrons on adjacent atoms align parallel. Fe, Co, and Ni are ferromagnetic. They can form permanent magnets.
- Antiferromagnetism: Moments on adjacent atoms align antiparallel → nearly zero net moment. Found in salts of Fe³⁺, Mn²⁺, Gd³⁺.
Calculating Magnetic Moment
For first-row transition elements, orbital contribution is quenched by the crystal field. So we use the spin-only formula:
where n = number of unpaired electrons and μB = Bohr magneton (9.273 × 10⁻²⁴ J T⁻¹).
| n (unpaired e⁻) | μs (B.M.) | Example Ion | Experimental μ (B.M.) |
|---|---|---|---|
| 1 | 1.73 | Ti³⁺, Cu²⁺ | 1.7–2.1 |
| 2 | 2.83 | V³⁺, Ni²⁺ | 2.8–3.5 |
| 3 | 3.87 | Cr³⁺, Co²⁺ | 3.7–5.2 |
| 4 | 4.90 | Cr²⁺, Mn³⁺, Fe²⁺ | 4.8–5.6 |
| 5 | 5.92 | Mn²⁺, Fe³⁺ | 5.7–6.0 |
μ/μB = 797.5 × √(χM × T) = 797.5 × √(2.004 × 10⁻⁸ × 293) = 1.93 BM
Using spin-only: if n=1, μ = √(1×3) = 1.73 BM ≈ experimental. ∴ Cu²⁺ has 1 unpaired electron.
High Spin vs. Low Spin Complexes
In an octahedral crystal field, the five d-orbitals split into two sets:
- t₂g (lower energy): dxy, dxz, dyz
- eg (higher energy): dx²-y², dz²
Fig. 4 — Crystal field splitting of d-orbitals in an octahedral field into t₂g and eg sets
- Strong field ligands (large Δ₀): CN⁻, CO, NO⁺ → electrons pair in t₂g → low spin → fewer unpaired electrons
- Weak field ligands (small Δ₀): F⁻, H₂O, OH⁻ → electrons fill all orbitals first → high spin → more unpaired electrons
I⁻ < Br⁻ < S²⁻ < Cl⁻ < NO₃⁻ < F⁻ < OH⁻ < H₂O < NCS⁻ < NH₃ < en < bpy < NO₂⁻ < CN⁻ < CO
"I Brought Some Carrots, Not For Our Home, Never Asked. Enough! No Nicely, Bye Noisy Colleagues!"
9. Complex Formation — Transition Metals as Superb Lewis Acids
Transition metals have an unparalleled tendency to form coordination compounds with Lewis bases (ligands). A ligand is any neutral molecule (like NH₃) or anion (like Cl⁻ or CN⁻) that can donate a lone pair of electrons to the metal.
Fe²⁺ + 6CN⁻ → [Fe(CN)₆]⁴⁻
Why Do Transition Metals Form So Many Complexes?
- They have small, highly charged ions → strong electrostatic attraction for ligands
- They have vacant low-energy d orbitals to accept lone pairs from ligands
- The +III oxidation state is more stable in complexes than the +II state
- Cobalt forms more complexes than any other element except carbon!
Class-a vs. Class-b Acceptors (Hard-Soft Acid-Base Theory)
| Class | Metal Examples | Preferred Ligand Donor Atoms | HSAB Term |
|---|---|---|---|
| Class-a (hard) | Sc, Ti, V, Cr, Mn, Fe, Co (early/first row) | N, O, F (small, electronegative) | Hard acid |
| Class-b (soft) | Rh, Ir, Pd, Pt, Ag, Au, Hg | P, S, I (large, polarizable) | Soft acid |
| Borderline | Fe, Co, Ni, Cu, Ru, Os | Both types | Intermediate |
Coordination Numbers and Shapes
- CN = 6, Octahedral — by far the most common
- CN = 4, Tetrahedral — common for d¹⁰ metals (Zn²⁺) and large ligands
- CN = 4, Square Planar — characteristic of d⁸ metals (Ni²⁺, Pd²⁺, Pt²⁺, Au³⁺)
- CN = 7 or 8 — rare in first row, but much more common in second and third row (e.g., [ZrF₇]³⁻, [ZrF₈]⁴⁻)
Fig. 5 — Schematic of an octahedral [ML₆] transition metal complex
10. Catalytic Properties — Transition Metals as Nature's Catalysts
Transition metals and their compounds are among the most important catalysts in both industry and biological systems. They catalyse reactions by:
- Using their variable valency to form unstable intermediate compounds
- Providing a suitable reaction surface (heterogeneous catalysis)
Key Industrial Catalysts
| Catalyst | Process | Reaction |
|---|---|---|
| TiCl₃ (Ziegler-Natta) | Polymerisation of ethylene | n(CH₂=CH₂) → (–CH₂–CH₂–)ₙ |
| V₂O₅ | Contact Process (H₂SO₄ manufacture) | SO₂ + ½O₂ → SO₃ |
| Fe (promoted) | Haber-Bosch Process (NH₃) | N₂ + 3H₂ ⇌ 2NH₃ |
| MnO₂ | Decomposition of KClO₃ | 2KClO₃ → 2KCl + 3O₂ |
| PdCl₂ | Wacker Process | C₂H₄ + H₂O + PdCl₂ → CH₃CHO + 2HCl + Pd |
| Pt/PtO (Adams) | Reductions (hydrogenation) | R–C≡C–R + H₂ → R–CH₂–CH₂–R |
| Pt/Rh | Ostwald Process (HNO₃) | 4NH₃ + 5O₂ → 4NO + 6H₂O |
| Ni (Raney) | Hydrogenation of alkenes, fats | –CH=CH– + H₂ → –CH₂–CH₂– |
Metalloenzymes — Nature's Own Transition Metal Catalysts
Many biological enzymes require a transition metal ion as a cofactor. These are called metalloenzymes. Examples from living systems:
- Fe (Haemoglobin): O₂ transport in blood
- Fe (Cytochrome): Electron transfer in respiration
- Fe & Mo (Nitrogenase): Biological fixation of N₂ → NH₃
- Co (Ribonucleotide reductase): DNA biosynthesis
- Cu (Tyrosinase): Skin pigmentation
- Zn (Carbonic anhydrase): CO₂ regulation (pH control in blood)
- Mn (Arginase): Urea formation
11. Nonstoichiometry — An Unusual Feature
Another fascinating property of transition metal compounds is their tendency to form nonstoichiometric compounds — compounds with variable, non-integer ratios of atoms. This arises directly from the variable oxidation states of transition metals.
Iron(II) oxide (FeO) is the textbook example. It should be written as FeO, but the actual formula varies between Fe0.94O and Fe0.84O! The formula is written with a bar: FeO (with overbar) to indicate this variability. The structure contains a mixture of Fe²⁺ and Fe³⁺ ions.
Vanadium and selenium form a series: VSe0.98 → VSe1.0 → VSe1.6 — all related by the variable valency of vanadium.
12. Differences Between First Row and Second/Third Row Transition Elements
This topic is very important for IIT-JAM and CSIR-NET examinations.
Metal–Metal Bonding and Cluster Compounds
Metal-metal (M–M) bonding is quite rare in first-row elements — it occurs mainly in carbonyls like Fe₂(CO)₉ and Co₄(CO)₁₂. In the second and third row, M–M bonds are much more common:
- Carbonyls: Ru₃(CO)₁₂, Os₃(CO)₁₂, Rh₄(CO)₁₂, Ir₄(CO)₁₂ — also the unique Rh₆(CO)₁₆ (no first-row analogue)
- Halide cluster ions: [Re₂Cl₈]²⁻ and [Mo₂Cl₈]⁴⁻ have M–M bonds
- Cluster compounds: Nb and Ta form octahedral 6-metal-atom clusters, e.g. [Nb₆Cl₁₂]²⁺ — 6 Nb atoms at corners of an octahedron with 12 bridging Cl atoms along edges
Stability of Oxidation States
- First row: +II and +III states are most important. Simple M²⁺ and M³⁺ ions are very common.
- Second and third row: Higher oxidation states are more stable. Example: CrO₄²⁻ is a strong oxidiser, but MoO₄²⁻ and WO₄²⁻ are stable. MnO₄⁻ is a strong oxidiser, but TcO₄⁻ and ReO₄⁻ are stable.
- Some second/third row compounds have no first-row counterpart: WCl₆, ReF₇, RuO₄, OsO₄, PtF₆.
Magnetism Differences
Second and third row elements tend to give low-spin complexes regardless of the ligand field strength. The crystal field splitting Δ is larger for heavier elements (diffuse 4d/5d orbitals → better overlap with ligands → larger splitting). So electrons always pair up in the lower energy levels. This makes the spin-only formula less applicable for second and third row elements.
13. Abundance and Occurrence
Three first-row transition metals are very abundant in Earth's crust:
- Iron (Fe): 60,000 ppm — 4th most abundant element, most abundant transition metal
- Titanium (Ti): 6,320 ppm — 9th most abundant element
- Manganese (Mn): 1,060 ppm — 12th most abundant
The first row elements as a group make up 6.79% of Earth's crust. The second and third row elements are mostly rare (Technetium does not occur naturally at all!). Of the last six elements in 2nd and 3rd rows (Tc, Ru, Rh, Pd, Ag, Cd; Re, Os, Ir, Pt, Au, Hg), none exceeds 0.16 ppm in Earth's crust.
14. Important Reaction Equations — Exam Ready
Reactions of Transition Metals with Acids
Fe + H₂SO₄(dil.) → FeSO₄ + H₂↑
Fe + 4HNO₃(dil.) → Fe(NO₃)₃ + NH₄NO₃ + 3H₂O
Cr + 6HCl(conc.) → 2CrCl₃ + 3H₂↑ (Cr passivated by HNO₃)
Reactions of KMnO₄ (Acidic Medium)
2KMnO₄ + 5H₂C₂O₄ + 3H₂SO₄ → 2MnSO₄ + K₂SO₄ + 10CO₂ + 8H₂O
2KMnO₄ + 5FeSO₄ + 4H₂SO₄ → 2MnSO₄ + K₂SO₄ + 5Fe₂(SO₄)₃ + 4H₂O
Reactions of K₂Cr₂O₇ (Acidic Medium)
K₂Cr₂O₇ + 6FeSO₄ + 7H₂SO₄ → K₂SO₄ + Cr₂(SO₄)₃ + 3Fe₂(SO₄)₃ + 7H₂O
K₂Cr₂O₇ + 3SO₂ → K₂SO₄ + Cr₂(SO₄)₃
Interconversion of Cr Species
Cr₂O₇²⁻ + 2NaOH → 2Na₂CrO₄ + H₂O
Cu²⁺ and Fe³⁺ Reactions
Fe³⁺ + SCN⁻ → [Fe(SCN)]²⁺ (blood red — qualitative test)
Cu²⁺ + 4NH₃ → [Cu(NH₃)₄]²⁺ (deep blue/Schweitzer's reagent)
Disproportionation Reactions
3MnO₄²⁻ + 4H⁺ → 2MnO₄⁻ + MnO₂ + 2H₂O (Manganate disproportionation)
15. Molecular Structures by SVG
Structure of Permanganate Ion (MnO₄⁻)
Fig. 6 — Tetrahedral structure of permanganate ion MnO₄⁻
Structure of Chromate vs. Dichromate Ions
Fig. 7 — Structures of chromate (CrO₄²⁻) and dichromate (Cr₂O₇²⁻) ions
16. Special Topics for Advanced Exams (IIT-JAM / GATE / CSIR-NET)
The Curie and Curie-Weiss Laws
Pierre Curie showed that the paramagnetic susceptibility χM of most materials varies inversely with absolute temperature:
Many systems deviate and obey the Curie-Weiss Law:
where θ (the Weiss constant) is an empirical correction factor. This deviation often signals ferromagnetic or antiferromagnetic interactions between neighbouring atoms.
Russell-Saunders Coupling (for Lanthanides)
For lanthanide elements (and in some heavy transition metals), the spin contribution S and orbital contribution L couple together to give a total angular momentum quantum number J. The magnetic moment must be calculated using:
where g = Landé g-factor = 1 + [S(S+1) − L(L+1) + J(J+1)] / 2J(J+1)
Orbital Contribution in Magnetism
The orbital contribution to magnetism exists when a t₂g orbital can be transformed into an equivalent orbital by rotation (i.e., the orbital is not quenched). In octahedral complexes, orbital contribution is significant for:
For these configurations, the measured μ is higher than the spin-only value. Co²⁺ (d⁷) is a classic example — it has the (t₂g)⁵(eg)² configuration, giving an orbital contribution and an experimental μ of 4.3–5.2 BM, significantly above the spin-only value of 3.87 BM.
17. High-Yield Exam Tips and Tricks
- Cr and Cu exceptions: [Ar]3d⁵4s¹ and [Ar]3d¹⁰4s¹ — non-negotiable for JEE/NEET
- Highest melting point: Tungsten (W), 3410°C
- Densest elements: Os and Ir (~22.6 g/cm³)
- Only liquid transition metal: Mercury (Hg), m.p. = −38°C
- Mn oxidation states: +II to +VII (+VII in KMnO₄)
- Maximum oxidation state > group number: Ru and Os can show +VIII (RuO₄, OsO₄) — Fe (same group) only shows +VI
- MnO₄⁻ in acidic medium: → Mn²⁺ (colourless); alkaline: → MnO₄²⁻ (green); neutral: → MnO₂ (brown)
- CrO₄²⁻ ⇌ Cr₂O₇²⁻: Add acid → dichromate (orange); add base → chromate (yellow)
- Spin-only formula: μ = √[n(n+2)] BM — valid for first-row only
- KMnO₄ is d⁰ yet purple: Colour due to charge transfer, not d-d transition
- Ziegler-Natta catalyst: TiCl₃ + AlEt₃ → stereoregular polythene (Nobel Prize 1963)
- Haber-Bosch: Fe catalyst, K₂O promoter, Al₂O₃ support, 450°C, 200 atm
- Lanthanide contraction causes: Zr ≈ Hf in size; Nb ≈ Ta in size
- Class-b (soft) metals: Rh, Ir, Pd, Pt, Ag, Au, Hg — prefer P, S, I ligands
- Square planar geometry: Characteristic of d⁸ metal ions (Ni²⁺, Pd²⁺, Pt²⁺, Au³⁺)
- Don't assume all transition metals show variable oxidation states — Sc³⁺ is the only stable state for Sc
- Don't confuse "spin-only" formula with Russell-Saunders formula — spin-only works for first row only
- Don't say KMnO₄ is purple due to d-d transition — Mn is +VII (d⁰) in KMnO₄; it's charge transfer
- Zn, Cd, Hg are d-block but NOT transition metals by the strict definition
- Cu⁺ (d¹⁰) is colourless in solution but Cu²⁺ (d⁹) is blue — don't mix them up
18. Quick Revision Summary
| Property | Trend / Key Fact |
|---|---|
| Electronic configuration | [Noble gas] (n-1)d¹⁻¹⁰ ns¹⁻² |
| Exceptions (config.) | Cr: 3d⁵4s¹; Cu: 3d¹⁰4s¹ |
| Oxidation states | Multiple; max = group no. for Sc→Mn; decreases after Mn |
| Melting/boiling points | Very high (d electrons in metallic bond); exceptions: Zn, Cd, Hg |
| Density | >5 g/cm³ for most; Os and Ir densest (22.6 g/cm³) |
| Atomic size (across row) | Decreases Sc→Ni, slightly increases Cu, Zn |
| Lanthanide contraction | 2nd and 3rd row have almost same size → Zr≈Hf, Nb≈Ta |
| Colour | d-d transitions; absent for d⁰ and d¹⁰ ions (unless charge transfer) |
| Magnetism | μ = √[n(n+2)] BM for first row; orbital contribution for some |
| Complex formation | Small, charged ions + vacant d orbitals → stable complexes |
| Catalysis | Variable valency + surface activity → industrial and biological catalysts |
| Nonstoichiometry | Variable composition due to variable oxidation state (e.g. FeO) |
| M–M bonding | Rare in 1st row; common in 2nd/3rd row; forms cluster compounds |
References & Further Reading: J.D. Lee, Concise Inorganic Chemistry, 5th Ed., Blackwell Science; Atkins, Shriver & Atkins' Inorganic Chemistry; IUPAC Recommendations for Nomenclature of Inorganic Chemistry (2005). All molecular geometries, bond angles, and oxidation states are consistent with IUPAC guidelines and experimental crystallographic data.
 (Sc, Ti and V Groups) : Properties, Trends, Oxidation States and Complex Formation Explanation For PGT, TGT, IIT-JAM, GATE, BITSAT, CSIR-NET, JEE Advanced, And BHU Exams){kind=link}