Unit 6

Inorganic Chemistry

Class 12 Chemistry

Transition Metals

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NEB/CDC syllabus scope: This chapter covers the characteristics and oxidation states of transition metals, complex ions and metal complexes, shapes of complex ions, a simple crystal-field explanation of d-orbital splitting in octahedral complexes, the origin of colour in transition-metal compounds, and catalytic properties of transition metals.

1. Introduction to Transition Metals

Definition A transition element is an element whose atom has an incompletely filled d-subshell, or which forms at least one stable ion having an incompletely filled d-subshell. Transition metals occur in the d-block of the periodic table.

The first transition series lies in Period 4 and runs from scandium to zinc in the d-block. In school-level chemistry, these elements are commonly discussed together because their properties change gradually across the series. In the stricter transition-element definition, an element must possess an incomplete d-subshell in the atom or in at least one of its common ions.

Important Note Zinc is a d-block element, but Zn and Zn²⁺ have completely filled 3d subshells. Therefore zinc does not show many characteristic transition-metal properties such as variable oxidation state and d–d colour.

Diagram 1: Position of transition metals in the d-block

2. Electronic Configuration of the First Transition Series

The general outer electronic configuration of d-block elements is:

General configuration (n − 1)d¹–¹⁰ ns⁰–²

For the first transition series, the 3d subshell is progressively filled. Chromium and copper show familiar stability-related exceptions to the simplest Aufbau prediction.

Element	Atomic No.	Ground-state electronic configuration	Useful d-count idea
Sc	21	[Ar] 3d¹ 4s²	d¹ atom
Ti	22	[Ar] 3d² 4s²	d² atom
V	23	[Ar] 3d³ 4s²	d³ atom
Cr	24	[Ar] 3d⁵ 4s¹	half-filled d⁵
Mn	25	[Ar] 3d⁵ 4s²	d⁵ atom
Fe	26	[Ar] 3d⁶ 4s²	d⁶ atom
Co	27	[Ar] 3d⁷ 4s²	d⁷ atom
Ni	28	[Ar] 3d⁸ 4s²	d⁸ atom
Cu	29	[Ar] 3d¹⁰ 4s¹	filled d¹⁰ atom
Zn	30	[Ar] 3d¹⁰ 4s²	filled d¹⁰

Remember When first-row transition metals form positive ions, the 4s electrons are generally removed before the 3d electrons. For example: Fe: [Ar] 3d⁶ 4s² → Fe²⁺: [Ar] 3d⁶ Fe³⁺: [Ar] 3d⁵

Diagram 2: Simplified orientation of the five d orbitals

3. Characteristics of Transition Metals

Transition metals show several characteristic properties because the energies of the ns and (n − 1)d orbitals are relatively close and because many of their ions retain partially filled d orbitals.

Characteristic	Main reason	Typical example
Variable oxidation states	Both ns and (n − 1)d electrons can participate in bonding.	Fe²⁺/Fe³⁺, Cu⁺/Cu²⁺, Mn²⁺ to Mn⁷⁺
Complex formation	Small metal ions with high charge density can accept lone pairs from ligands.	[Fe(CN)₆]⁴⁻, [Co(NH₃)₆]³⁺
Coloured ions/compounds	Partially filled d orbitals can absorb visible light through electronic transitions.	Cu²⁺ compounds are commonly blue; many Mn and Cr species are strongly coloured.
Catalytic activity	Variable oxidation states and ability to form intermediate complexes or adsorb reactants.	Fe in Haber process, V₂O₅ in Contact process, Ni in hydrogenation
Magnetic behaviour	Unpaired d electrons produce paramagnetism.	Fe³⁺ (d⁵) has unpaired electrons.
High melting point and density	Strong metallic bonding involving d and s electrons.	Many first-row transition metals are hard, dense metals.
Alloy formation	Similar atomic sizes allow one metal atom to replace another in a lattice.	Fe–Cr–Ni stainless steel

Exam Important When asked “Why do transition metals show variable oxidation states?”, connect the answer directly to the small energy difference between the ns and (n − 1)d orbitals. Electrons from both can take part in bond formation or ion formation.

4. Oxidation States of Transition Metals

Transition metals commonly show more than one oxidation state. Lower oxidation states often involve loss of ns electrons, while higher oxidation states can also involve d electrons.

Element	Common oxidation states	Important examples
Sc	+3	Sc³⁺
Ti	+2, +3, +4	Ti³⁺, TiO₂ (Ti⁴⁺)
V	+2, +3, +4, +5	V²⁺, V³⁺, VO²⁺, vanadates
Cr	+2, +3, +6	Cr²⁺, Cr³⁺, Cr₂O₇²⁻
Mn	+2, +3, +4, +6, +7	Mn²⁺, MnO₂, MnO₄⁻
Fe	+2, +3	Fe²⁺, Fe³⁺
Co	+2, +3	Co²⁺, Co³⁺ complexes
Ni	+2 (most common)	Ni²⁺
Cu	+1, +2	Cu⁺, Cu²⁺
Zn	+2	Zn²⁺

Diagram 3: Variation of common oxidation states across the first transition series

How to Find Oxidation State

Charge-balance rule The sum of oxidation numbers of all atoms in a neutral compound is zero. In an ion, the sum equals the overall ionic charge.

Example: KMnO₄

Let oxidation state of Mn = x.

(+1) + x + 4(−2) = 0 x = +7

Example: K₂Cr₂O₇

Let oxidation state of each Cr = x.

2(+1) + 2x + 7(−2) = 0 2x = 12 → x = +6

5. Complex Ions and Metal Complexes

Complex ion A complex ion is a charged coordination entity consisting of a central metal atom or ion bonded to surrounding molecules or ions called ligands through coordinate bonds.

Ligand A ligand is an ion or molecule that donates at least one lone pair of electrons to a central metal atom or ion to form a coordinate bond.

Term	Meaning	Example
Central metal ion	Metal ion to which ligands are directly attached	Fe²⁺ in [Fe(CN)₆]⁴⁻
Ligand	Lone-pair donor bonded to metal	NH₃, H₂O, Cl⁻, CN⁻
Coordination number	Number of donor atoms directly bonded to the metal	6 in [Co(NH₃)₆]³⁺
Complex ion	Charged metal–ligand entity	[Cu(NH₃)₄]²⁺
Coordination compound	Compound containing a coordination entity	K₄[Fe(CN)₆]

Why Transition Metals Form Complexes

Transition-metal ions are often relatively small.
They can possess comparatively high positive charge density.
They have orbitals capable of accepting electron pairs from ligands.
Metal–ligand bonding can stabilize particular oxidation states and structures.

Worked Example: Oxidation state and coordination number in [Fe(CN)₆]⁴⁻

Each CN⁻ ligand has charge −1. Let oxidation state of Fe be x:

x + 6(−1) = −4 x = +2

There are six monodentate CN⁻ ligands, so the coordination number is 6.

6. Shapes of Complex Ions

The shape of a complex depends mainly on the coordination number, the electronic configuration of the metal ion, and the nature of the ligands.

6.1 Octahedral Complex

Six ligand donor atoms surround the central metal ion along the positive and negative x, y, and z directions.

Diagram 4: Octahedral arrangement of six ligands

6.2 Tetrahedral Complex

Four ligands occupy the corners of a tetrahedron around the central metal ion.

Diagram 5: Tetrahedral arrangement of four ligands

6.3 Square-Planar Complex

Four ligands and the metal ion lie approximately in one plane, with bond angles close to 90°.

Diagram 6: Square-planar arrangement of four ligands

Shape	Typical coordination number	Example	Key geometry
Octahedral	6	[Co(NH₃)₆]³⁺	Six ligand positions around metal
Tetrahedral	4	[CoCl₄]²⁻	Four ligands toward tetrahedron corners
Square planar	4	[Ni(CN)₄]²⁻	Four ligands in one plane at about 90°

7. d-Orbitals in Octahedral Complexes: Simple Crystal Field Theory

In a free metal ion, the five d orbitals have the same energy. When six ligands approach a metal ion along the x, y, and z axes to form an octahedral complex, electrostatic interactions remove this degeneracy.

The d orbitals that point directly along the axes experience greater repulsion from approaching ligand electron pairs. The orbitals directed between the axes experience less repulsion.

Octahedral splitting In an octahedral crystal field, the five d orbitals split into:

higher-energy e_g set: d_x²−y² and d_z²
lower-energy t_2g set: d_xy, d_xz and d_yz

The energy separation is called the octahedral crystal-field splitting, commonly written as Δ_o.

Diagram 7: Splitting of d orbitals in an octahedral crystal field

Why e_g Orbitals Have Higher Energy

The d_x²−y² and d_z² orbitals point directly toward ligands approaching along the coordinate axes. Their electron density therefore experiences stronger repulsion. The d_xy, d_xz, and d_yz orbitals point between the axes and experience less repulsion.

Exam Important For an octahedral complex, remember the order: lower t₂g < higher e₉ and the sets: t₂g = dxy, dxz, dyz | e₉ = dx²−y², dz²

8. Why Transition-Metal Compounds Are Coloured

Many transition-metal ions have partially filled d orbitals. In a complex, ligand interactions split the d orbitals into groups of different energies. If visible light provides the correct amount of energy, an electron can be promoted from a lower-energy d level to a higher-energy d level.

Energy absorption idea absorbed light energy = ΔE between split d levels ΔE = hν

Here h is Planck’s constant and ν is the frequency of absorbed light.

Diagram 8: d–d transition and the origin of colour

Important Observations

The colour depends on the metal ion, its oxidation state, the ligands, and the geometry of the complex.
Ions with d⁰ or d¹⁰ configurations often do not show ordinary d–d colour because there is no suitable partially filled d-level transition.
Some compounds can still be coloured by other electronic processes; therefore “d⁰/d¹⁰ means always colourless” is too absolute.

Common Mistake Do not simply write “transition metals are coloured because they have d electrons.” The better answer is that ligand fields split partially filled d orbitals, and visible light can be absorbed to promote electrons between these levels.

9. Catalytic Properties of Transition Metals

Transition metals and their compounds are widely used as catalysts. Two important reasons are their ability to change oxidation state and their ability to form temporary intermediate complexes or adsorb reactant molecules on a metal surface.

Variable Oxidation State

A transition-metal species can accept electrons in one step and donate them in another, providing an alternative reaction pathway.

Adsorption / Intermediate Formation

Metal surfaces can hold reactant molecules close together and weaken bonds, while metal ions can form short-lived intermediates.

Catalyst	Process / reaction	Role
Fe	N₂ + 3H₂ ⇌ 2NH₃	Catalyst in the Haber process
V₂O₅	2SO₂ + O₂ ⇌ 2SO₃	Catalyst in the Contact process
Ni	Hydrogenation of C=C bonds	Metal-surface catalyst
MnO₂	2H₂O₂ → 2H₂O + O₂	Speeds decomposition of hydrogen peroxide

Diagram 9: A catalyst provides a lower-activation-energy pathway

Exam Important A catalyst changes the reaction pathway and lowers activation energy. It does not change the overall thermodynamic energy difference between reactants and products.

10. Worked Examples

Worked Example 1: Electronic configuration of Fe³⁺

Neutral iron:

Fe = [Ar] 3d⁶ 4s²

Remove two 4s electrons first, then one 3d electron:

Fe³⁺ = [Ar] 3d⁵

Answer: Fe³⁺ is a d⁵ ion.

Worked Example 2: Oxidation state of Co in [Co(NH₃)₆]³⁺

NH₃ is a neutral ligand. Let oxidation state of Co = x:

x + 6(0) = +3 x = +3

Coordination number: 6. Shape: octahedral.

Worked Example 3: Oxidation state of Co in [CoCl₄]²⁻

Each Cl⁻ carries −1 charge:

x + 4(−1) = −2 x = +2

Coordination number: 4. This complex is commonly represented as tetrahedral.

Worked Example 4: Why is Zn²⁺ usually not coloured by d–d transition?

Zn²⁺ has configuration:

Zn²⁺ = [Ar] 3d¹⁰

The d subshell is completely filled. There is no partially filled d-level arrangement for an ordinary d–d transition, so Zn²⁺ compounds are commonly colourless unless another type of electronic transition causes colour.

Worked Example 5: Why does Fe act as a catalyst?

Iron can provide a metal surface for adsorption and can interact with reactants in ways that lower the activation-energy barrier. In the Haber process it accelerates:

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

The catalyst speeds both forward and reverse reactions without being consumed overall.

11. Quick Comparison Tables

Transition Metal vs Typical Main-Group Metal

Property	Transition metals	Typical main-group metals
Oxidation states	Often variable	Often fewer common values
Complex formation	Very common	Generally less characteristic
Colour	Many ions/complexes coloured	Many simple ions colourless
Catalytic activity	Common	Less characteristic as a group
d electrons	Partially filled d orbitals are often important	d orbitals not the defining feature

Octahedral vs Tetrahedral vs Square Planar

Geometry	Coordination number	Ligand arrangement	Example
Octahedral	6	Six directions around metal	[Co(NH₃)₆]³⁺
Tetrahedral	4	Four corners of tetrahedron	[CoCl₄]²⁻
Square planar	4	Four ligands in one plane	[Ni(CN)₄]²⁻

12. Common Exam Mistakes

Calling every d-block element a transition element without considering whether its atom or ions have an incomplete d subshell.
Removing 3d electrons before 4s electrons when writing first-row transition-metal cations.
Forgetting ligand charges while calculating oxidation state in a complex ion.
Confusing coordination number with the oxidation state of the metal.
Writing that all four-coordinate complexes have the same shape. Four-coordinate complexes may be tetrahedral or square planar.
Reversing the octahedral crystal-field order. In an octahedral field, t₂g is lower and e₉ is higher.
Explaining colour only by saying “d electrons are present” instead of explaining splitting and visible-light absorption.
Saying a catalyst changes the equilibrium constant or overall reaction energy. A catalyst changes the pathway and rate, not the thermodynamic equilibrium position.

13. Exam-Important Points

Know the general d-block electronic configuration.
Remember Cr = [Ar]3d⁵4s¹ and Cu = [Ar]3d¹⁰4s¹.
Explain variable oxidation state using close ns and (n−1)d energies.
Be able to calculate oxidation state inside complex ions.
Know the meanings of ligand, central metal ion, coordination number, and complex ion.
Draw octahedral, tetrahedral, and square-planar shapes correctly.
For octahedral CFT: t₂g is lower and e₉ is higher.
Explain colour using d-orbital splitting and absorption of visible light.
Explain catalysis using variable oxidation states and/or adsorption/intermediate-complex formation.
Know common catalyst examples: Fe, V₂O₅, Ni, and MnO₂.

14. Important Exam Questions

Short-Answer Questions

Define a transition element.
Write the general electronic configuration of d-block elements.
Why do transition metals show variable oxidation states?
Why do transition metals readily form complex ions?
Define ligand and coordination number.
Find the oxidation state of Fe in [Fe(CN)₆]⁴⁻.
What is the coordination number of Co in [Co(NH₃)₆]³⁺?
Name the two groups of d orbitals formed in an octahedral field.
Which d orbitals belong to t₂g and e₉ sets?
Why are many transition-metal compounds coloured?
Why are Zn²⁺ compounds commonly colourless?
State two reasons for the catalytic activity of transition metals.
Give two industrial examples of transition-metal catalysts.
Why are Cr and Cu electronic configurations considered exceptions to the simple Aufbau pattern?

Long-Answer Questions

Explain the important characteristics of transition metals with suitable examples.
Discuss the variable oxidation states of the first transition series.
Explain complex ions and metal complexes, defining central metal ion, ligand, and coordination number.
Draw and explain octahedral, tetrahedral, and square-planar complex-ion geometries.
Using simple crystal field theory, explain the splitting of d orbitals in an octahedral complex.
Explain the origin of colour in transition-metal compounds on the basis of d-orbital splitting.
Explain the catalytic properties of transition metals and give important examples.
Compare t₂g and e₉ orbitals in an octahedral crystal field.

Calculation / Application Questions

Calculate the oxidation state of Mn in KMnO₄.
Calculate the oxidation state of Cr in K₂Cr₂O₇.
Calculate the oxidation state and coordination number of Fe in [Fe(CN)₆]⁴⁻.
Calculate the oxidation state and coordination number of Co in [CoCl₄]²⁻.
Write the d-electron configuration of Fe²⁺, Fe³⁺, Cu²⁺, and Zn²⁺.
Classify [Co(NH₃)₆]³⁺, [CoCl₄]²⁻, and [Ni(CN)₄]²⁻ by coordination number and shape.

Diagram Questions

Show the position of the d-block in the periodic table.
Draw the five d-orbitals schematically.
Draw an octahedral complex and label the central metal and ligands.
Draw tetrahedral and square-planar complex shapes.
Draw the octahedral crystal-field splitting diagram and label t₂g, e₉, and Δₒ.
Draw a simple energy-level diagram showing how absorption of visible light can cause a d–d transition.

15. One-Minute Revision

Transition metals belong to the d-block and characteristically involve incomplete d subshells.
The general d-block configuration is (n−1)d¹–¹⁰ ns⁰–².
Cr is [Ar]3d⁵4s¹ and Cu is [Ar]3d¹⁰4s¹.
When forming first-row cations, 4s electrons are generally removed before 3d electrons.
Transition metals commonly show variable oxidation states.
They readily form complex ions with ligands that donate lone pairs.
Coordination number counts donor atoms directly bonded to the metal.
Common complex shapes include octahedral, tetrahedral, and square planar.
In an octahedral field, t₂g orbitals are lower in energy and e₉ orbitals are higher.
Octahedral splitting energy is represented by Δₒ.
Many transition-metal compounds are coloured because visible light can promote electrons between split d levels.
d⁰ and d¹⁰ ions usually lack ordinary d–d transitions.
Transition metals often act as catalysts because of variable oxidation states and intermediate formation/adsorption.
Fe catalyses ammonia synthesis; V₂O₅ catalyses SO₂ oxidation; Ni catalyses hydrogenation.
Zinc is a d-block element but does not display many characteristic transition-metal properties in Zn²⁺ because it is d¹⁰.

16. Diagram Practice

Students should practice these diagrams for the NEB examination:

Position of the d-block / first transition series in the periodic table.
Five d-orbitals — schematic orientation.
Common oxidation states across the first transition series.
Octahedral complex.
Tetrahedral complex.
Square-planar complex.
Octahedral d-orbital splitting: t₂g, e₉, and Δₒ.
d–d electronic transition responsible for colour.
Energy-profile diagram showing the lower activation-energy pathway of a catalyst.

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