CBSE Class 12 Chemistry Board Exam 2026: How to Solve Numericals & Graph Questions Easily

The Class 12 Chemistry board exam is on 28 February 2026, conducted by the Central Board of Secondary Education. This is the final revision phase, and students should focus on smart, exam-oriented preparation. This 5-hour one-shot revision video is designed after carefully analysing previous years’ questions (PYQs). Most exam questions are expected to come directly or indirectly from the concepts covered here. Students are advised to watch the full revision in one go without skipping, as it provides a complete conceptual recap, important reactions, and frequently asked patterns that can significantly improve scores.

30 Most Expected Questions for CBSE Class 12 Chemistry Exam 2026

D-Block Elements Introduction

D-block elements are located between the S-block and P-block in the periodic table, spanning Group 3 to Group 12. They comprise 40 elements whose last electron enters a d-orbital. These elements are organized into four series (3d, 4d, 5d, 6d), with the principal quantum number of the d-series always being one less than the period number (e.g., 4th period contains the 3d series).

Electronic Configuration of D-Block Elements

The general electronic configuration for D-block elements is [Noble Gas] ns¹⁻² (n-1)d¹⁻¹⁰. The (n-1)d orbitals fill after the ns orbital.

The 3d Series: Configuration and Exceptions

The 3d series (Sc to Zn) is crucial for examinations. A mnemonic to remember the elements is "Sky TV Crorepati Mein Feco Nico Jeete" (Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn).

Stability of half-filled and fully-filled orbitals (d⁵ and d¹⁰) is a key principle. This leads to two exceptions in the 3d series:

1. Chromium (Cr, Z=24): Actual Configuration: [Ar] 4s¹ 3d⁵ (instead of expected [Ar] 4s² 3d⁴) for stable half-filled d⁵.

2. Copper (Cu, Z=29): Actual Configuration: [Ar] 4s¹ 3d¹⁰ (instead of expected [Ar] 4s² 3d⁹) for stable fully-filled d¹⁰.

• Self-Assessment: Copper and Chromium show exceptional behavior in electronic configuration.

Transition Elements

D-block elements are called transition elements because their properties bridge S-block metals and P-block non-metals.

A transition element is defined as an element with a partially filled d-orbital (d¹ to d⁹) in its ground state or any common oxidation state.

Differentiating D-Block and Transition Elements

• Zinc (Zn), Cadmium (Cd), and Mercury (Hg): These are D-block elements but NOT transition elements. They have a fully-filled d¹⁰ configuration in their ground state and common oxidation states (e.g., Zn²⁺ is 3d¹⁰), never achieving a partially filled d-orbital.

• Copper (Cu): Copper is a transition element. Its ground state is 4s¹3d¹⁰, but its common Cu²⁺ oxidation state is 3d⁹, a partially filled d-orbital.

Properties of D-Block Elements

1. Metallic Character

All d-block elements are metals, exhibiting high melting/boiling points and good conductivity. Their strong metallic bonding arises from delocalized ns electrons and covalent bonding due to overlap of unpaired (n-1)d orbitals.

• Strength of Metallic Bond: Directly related to the number of unpaired electrons. More unpaired electrons lead to stronger bonds, harder metals, and higher melting points.

• Hard Metals (e.g., Cr, Mo, W): Many unpaired electrons.

• Soft Metals (e.g., Zn, Cd, Hg): Zero unpaired electrons (d¹⁰), resulting in weak metallic bonding.

2. Atomic Radii

Trends in atomic radii are governed by the interplay of increasing nuclear charge and the screening effect.

• Screening Effect: Inner-shell electrons "screen" outer electrons from the full nuclear attraction, reducing the effective nuclear charge.

• Effectiveness Order: s > p > d > f. f-orbitals provide negligible shielding. (Analogy for Shielding: Imagine the nucleus wants to pull a friend (outer electron) into the house. However, your mother (inner electron) is sitting in between and pushes the friend away, preventing them from coming closer. A smaller house (like an s-orbital) makes it easier for the mother to block the friend, representing strong shielding. A larger house (like an f-orbital) makes it harder for the mother to block the friend, representing weak shielding.)

Trend Across a Period (3d Series)

The trend is non-linear:

Trend Down a Group (Top to Bottom)

1. 3d to 4d Series: Atomic size increases due to the addition of a new electron shell.

2. 4d to 5d Series (Lanthanoid Contraction): Atomic radii are nearly identical. This is because the filling of the 4f subshell (before 5d) leads to very poor shielding. The increased nuclear charge is not counteracted, pulling outer electrons inward.

• Example: Zirconium (Zr, 160 pm) and Hafnium (Hf, 159 pm) have similar sizes.

• Consequence: Elements like Zr/Hf are difficult to separate due to similar chemical properties.

• PYQ: Assertion (A): The separation of Zirconium (Zr) and Hafnium (Hf) is difficult. Reason (R): Zr and Hf have similar atomic radii due to Lanthanoid Contraction. Answer: Both A and R are true, and R is the correct explanation for A.

Melting and Boiling Point

Generally high, determined by metallic bond strength and thus the number of unpaired electrons.

• Trend: Increases to the middle (Group 6, Cr) then decreases.

• Manganese (Mn): Anomalously low melting point due to stable half-filled 4s²3d⁵ configuration, making electrons less available for bonding.

Enthalpy of Atomization

A direct measure of metallic bond strength, also dependent on the number of unpaired electrons.

• Lowest Value: Zinc (Zn) has the lowest enthalpy of atomization due to zero unpaired electrons (3d¹⁰ 4s²), resulting in weak metallic bonding.

Ionization Energy (IE)

Energy required to remove an electron. Primarily driven by electronic stability (half-filled/fully-filled subshells).

Trend Across a Period (Left to Right)

• Generally increases.

• Irregularities:

• Third IE of Mn is very high: Mn²⁺ (3d⁵) is very stable.

• Second IE of Cr and Cu is very high: Cr⁺ (3d⁵) and Cu⁺ (3d¹⁰) are very stable.

Trend Down a Group (Top to Bottom)

• 3d to 4d Series: IE decreases (increased size).

• 4d to 5d Series: IE increases due to Lanthanoid Contraction (constant size, very high effective nuclear charge). Thus, IE of Hafnium (Hf) > Zirconium (Zr).

Variable Oxidation States

A hallmark of transition elements, enabled by the small energy difference between (n-1)d and ns orbitals, allowing both to participate in bonding.

• Exceptions:

• Scandium (Sc): Only +3 oxidation state.

• Zinc (Zn): Only +2 oxidation state.

• Stability of Oxidation States: Linked to achieving stable electronic configurations:

1. Half-filled d⁵: Fe³⁺ (3d⁵), Mn²⁺ (3d⁵).

2. Fully-filled d¹⁰: Zn²⁺ (3d¹⁰).

3. Inert (Noble Gas) Configuration: Cr⁶⁺, Mn⁷⁺.

4. Half-filled t₂g configuration: Cr³⁺ (3d³). In the presence of ligands, its three d-electrons occupy the three t₂g orbitals singly, providing extra stability.

Stabilizing Higher Oxidation States

Fluorine and Oxygen stabilize higher oxidation states.

• Fluorine: High lattice energy and bond enthalpy compensate for ionization energy.

• Oxygen: Greater ability due to its capacity to form multiple bonds.

Standard Reduction Potential (E°)

E° measures an ion's stability in aqueous solution. D-block metals tend to undergo oxidation (lose electrons).

Oxidation is favored when ionization enthalpy is low, and hydration enthalpy is high.

• Mn²⁺/Mn and Zn²⁺/Zn have very negative E°: Low ionization enthalpy (forming stable 3d⁵ or 3d¹⁰) and high hydration enthalpy make oxidation favorable.

• Copper (Cu) has a positive E° (+0.34 V): Forming Cu²⁺ requires a very high second ionization enthalpy (removing electron from stable 3d¹⁰). This energy is not fully compensated by hydration enthalpy, so Cu does not readily oxidize. Instead, Cu²⁺ readily reduces to Cu(s).

• Stability of Cu²⁺ vs. Cu⁺ in aqueous solution: Cu²⁺ is more stable due to its significantly higher hydration enthalpy, which compensates for the high ionization energy required to form it.

Properties of Transition Elements

Magnetic Properties

• Paramagnetism: Weak attraction to magnetic field due to unpaired electrons.

• Diamagnetism: Weak repulsion due to absence of unpaired electrons.

• Magnetic Moment (μ): Calculated by μ = √[n(n+2)] B.M., where 'n' is the number of unpaired electrons.

• Example: Cr³⁺ ([Ar] 3d³) has n=3. μ = √[3(3+2)] = √15 ≈ 3.87 B.M.

Formation of Coloured Ions

Most transition metal compounds are colored due to d-d transition.

• Mechanism: Unpaired d-electrons absorb visible light to jump between split d-orbitals (t₂g to e₉). The complementary color is observed.

• Condition: Must have one or more unpaired electrons. Ions like Sc³⁺ (3d⁰) and Zn²⁺ (3d¹⁰) are colourless.

Tendency to Form Complex Compounds

Transition metals form complexes due to:

1. Small size of metal ions.

2. High positive charge density.

3. Availability of vacant d-orbitals to accept lone pairs from ligands.

Formation of Interstitial Compounds

Transition metals trap small non-metal atoms (H, C, N) in their crystal lattice.

• Characteristics: Very high melting points, very hard, chemically inert.

Catalytic Property

Transition metals act as catalysts due to their variable oxidation states, which provide alternative reaction pathways with lower activation energy.

• Examples: Fe (Haber's process), Ni (hydrogenation), Pt/Rh (Ostwald's process).

Important Compounds of Transition Elements

Potassium Dichromate (K₂Cr₂O₇)

1. Preparation

From chromite ore (FeO.Cr₂O₃):

1. Oxidation of chromite: FeO.Cr₂O₃ + Na₂CO₃ + O₂ → Na₂CrO₄ + Fe₂O₃ + CO₂

2. Chromate to dichromate: 2Na₂CrO₄ + H₂SO₄ → Na₂Cr₂O₇ + Na₂SO₄ + H₂O

3. K₂Cr₂O₇ formation: Na₂Cr₂O₇ + 2KCl → K₂Cr₂O₇ + 2NaCl

2. Interconversion of Chromate and Dichromate

CrO₄²⁻ (yellow) and Cr₂O₇²⁻ (orange) interconvert based on pH:

• Acidic: 2CrO₄²⁻ + 2H⁺ ⇌ Cr₂O₇²⁻ + H₂O

• Basic: Cr₂O₇²⁻ + 2OH⁻ ⇌ 2CrO₄²⁻ + H₂O

3. Oxidizing Properties

K₂Cr₂O₇ has Cr in +6 oxidation state. Since Cr⁺³ is more stable than Cr⁺⁶, it is a powerful oxidizing agent in acidic medium, reducing to Cr³⁺:

Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O

It oxidizes I⁻ to I₂, Fe²⁺ to Fe³⁺, S²⁻ to S, Sn²⁺ to Sn⁴⁺.

Potassium Permanganate (KMnO₄)

MnO₄⁻ (Mn in +7) is a strong oxidizing agent because Mn⁺⁷ is less stable than Mn⁺², so it readily reduces to Mn⁺².

It oxidizes I⁻ to I₂, Fe²⁺ to Fe³⁺, Sn²⁺ to Sn⁴⁺, C₂O₄²⁻ to CO₂, SO₃²⁻ to SO₄²⁻.

Manganate and Permanganate

Stability and Disproportionation of Manganate (MnO₄²⁻)

MnO₄²⁻ is stable only in strongly alkaline solutions. In neutral or acidic solutions, it undergoes disproportionation (simultaneous oxidation and reduction):

3MnO₄²⁻ + 4H⁺ → 2MnO₄⁻ + MnO₂ + 2H₂O

Mn⁺⁶ disproportionates to Mn⁺⁷ and Mn⁺⁴.

Preparation of Potassium Manganate (K₂MnO₄)

From Pyrolusite ore (MnO₂):

2MnO₂ + 4KOH + O₂ → 2K₂MnO₄ + 2H₂O

Preparation of Potassium Permanganate (KMnO₄)

1. Commercial Method: Electrolytic oxidation of K₂MnO₄ solution.

2. Laboratory Method: Disproportionation of K₂MnO₄ in acidic medium (e.g., with H₂SO₄ or CO₂):
   3K₂MnO₄ + 2H₂SO₄ → 2KMnO₄ + MnO₂ + 2K₂SO₄ + 2H₂O

F-Block Elements (Lanthanoids and Actinoids)

Known as inner transition elements, electrons enter the (n-2)f orbital.

• Lanthanoids (4f series): Ce (Z=58) to Lu (Z=71).

• Actinoids (5f series): Th (Z=90) to Lr (Z=103).

General Electronic Configuration

Noble Gasf¹⁻¹⁴ (n-1)d⁰⁻¹ ns²

The 4f and 5d orbitals for lanthanoids are very close in energy, leading to exceptions for Ce, Gd, and Lu where a 5d¹ electron is retained for stability.

Comparison of Lanthanoids and Actinoids

Uses of F-Block Elements

• Lanthanoids: Used in alloy steels, e.g., Mischmetal (~95% lanthanoid, ~5% Fe) in bullets, shells, lighter flints.

• Actinoids: Primarily radioactive, used in nuclear energy and research.

Coordination Compounds

Introduction to Molecular and Addition Compounds

Formed by combining simple salts.

• Double Salts (e.g., Mohr's Salt): Lose their identity in solution, dissociating completely into constituent ions.

• Coordination Compounds (e.g., K₄[Fe(CN)₆]): Retain their identity in solution, forming a complex ion.

• K₄[Fe(CN)₆] → 4K⁺ + [Fe(CN)₆]⁴⁻ (not Fe²⁺ and CN⁻).

Basic Terminology

• Central Metal Atom/Ion: Electron-pair acceptor (Lewis acid).

• Ligands: Ions or molecules bound to the central metal, acting as electron-pair donors (Lewis bases).

• Coordinate Bond: Bond formed by ligand donating electron pair to metal.

Classification of Ligands

1. Based on Charge: Anionic (Cl⁻), Cationic (NO⁺), Neutral (H₂O).

2. Based on Denticity: Number of donor atoms.

• Monodentate: One donor (H₂O, NH₃).

• Bidentate: Two donors (Ethylenediamine 'en', Oxalate 'ox').

• Polydentate: More than two donors (EDTA⁴⁻ is hexadentate).

3. Based on Bonding Mode:

• Chelating Ligands: Bidentate/polydentate ligands forming stable ring structures with a metal (chelation increases stability).

• Ambidentate Ligands: Monodentate ligands with more than one potential donor atom, but bind through only one at a time (e.g., NO₂⁻: Nitro -NO₂ or Nitrito -ONO). (Practice Question: Oxalate (C₂O₄²⁻) is bidentate, not ambidentate.)

• Flexidentate Ligands: Exhibit variable denticity (e.g., EDTA can be hexa-, penta-, or tetradentate).

Coordination Number

The total number of ligand donor atoms directly bonded to the central metal atom.

CN = Σ (Number of a specific ligand × its Denticity)

• [Fe(CN)₆]⁴⁻: CN = 6 × 1 = 6.

• [Co(Cl)₂(en)₂]⁺: CN = (2 × 1) + (2 × 2) = 6.

• [Fe(EDTA)]⁻: CN = 1 × 6 = 6.

Coordination Sphere and Ionization Sphere

• Coordination Sphere: Metal + ligands, enclosed in []. Remains intact in solution.

• Ionization Sphere: Counter-ions outside []. Dissociates in solution.

• In K₄[Fe(CN)₆]: [Fe(CN)₆]⁴⁻ is coordination sphere, K⁺ is ionization sphere.

Calculating the Oxidation State of the Central Metal Atom

Sum of oxidation states equals the complex's overall charge.

• K₄[Fe(CN)₆]: 4(+1) + x + 6(-1) = 0 → x = +2.

Homoleptic and Heteroleptic Complexes

• Homoleptic: Metal bound to only one type of ligand (e.g., [Co(NH₃)₆]³⁺).

• Heteroleptic: Metal bound to more than one type of ligand (e.g., [Co(NH₃)₄Cl₂]⁺).

IUPAC Nomenclature of Coordination Compounds

1. Cation-Anion Order: Cation named first, then anion.

2. Coordination Sphere: Ligands named first (alphabetical order, ignoring prefixes), then metal.

3. Metal Naming:

• Cationic/Neutral Complex: Metal's name as is (Cobalt, Platinum).

• Anionic Complex: Metal name ends with -ate (Cobaltate, Ferrate).

4. Oxidation State: Roman numerals in parentheses after metal.

5. Ligand Naming:

• Anionic Ligands: -ide → -ido (Chlorido), -ite → -ito (Nitrito), -ate → -ato (Sulfato).

• Neutral Ligands: H₂O (aqua), NH₃ (ammine), CO (carbonyl), NO (nitrosyl).

• Numerical Prefixes: di-, tri-, tetra- for simple ligands; bis-, tris-, tetrakis- if ligand name already has prefix (e.g., ethylenediamine).

6. Counter-ions: No numerical prefixes.

• Example 1: [Pt(NH₃)₂(Cl)₂]²⁺ → diamminedichloridoplatinum(IV) ion

• Example 2: [Co(NH₃)₅(NO₂)]Cl₂ → pentaamminenitritocobalt(III) chloride

Coordination Polyhedra (Geometry)

Spatial arrangement of ligands. Determined by coordination number.

• CN 4: Tetrahedral or Square Planar.

• CN 6: Octahedral (most common).

• CN 5: Trigonal Bipyramidal or Square Pyramidal.

Bonding in Coordination Compounds: Valence Bond Theory (VBT)

Explains bonding via hybridization of metal orbitals, predicting geometry and magnetic properties.

• Hybridization and Geometry:

• CN 4: sp³ (Tetrahedral), dsp² (Square Planar).

• CN 6: sp³d² (Octahedral, outer orbital complex), d²sp³ (Octahedral, inner orbital complex).

Ligand Classification and the Spectrochemical Series

Crucial for VBT.

• Strong Field Ligands (SFL): Cause strong interaction, force electrons to pair up (violating Hund's rule). Form low spin complexes.

• Weak Field Ligands (WFL): Cause weak interaction, follow Hund's rule. Form high spin complexes.

The Spectrochemical Series

Ligands ordered by increasing field strength (ability to cause d-orbital splitting):

"I Br**ought **S**ome **C**olorful **S**weets **F**rom **O**ffice **C**ontaining **Water, NCS, EDTA, Ammonia, en, CN⁻, CO."

(I⁻ < Br⁻ < SCN⁻ < Cl⁻ < S²⁻ < F⁻ < O²⁻/OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < EDTA⁴⁻ < NH₃ < en < CN⁻ < CO)

Applying Valence Bond Theory (VBT): Examples

1. [NiCl₄]²⁻: Ni²⁺ (3d⁸). Cl⁻ is WFL. Hybridization: sp³ (Tetrahedral). Magnetic: Paramagnetic (2 unpaired e⁻).

2. [Ni(CO)₄]: Ni(0) (4s²3d⁸). CO is SFL. 4s and 3d electrons pair into 3d¹⁰. Hybridization: sp³ (Tetrahedral). Magnetic: Diamagnetic (0 unpaired e⁻).

3. [CuCl₄]²⁻: Cu²⁺ (3d⁹). Cl⁻ is WFL. Hybridization: sp³ (Tetrahedral). Magnetic: Paramagnetic (1 unpaired e⁻).

4. [Ni(CN)₄]²⁻: Ni²⁺ (3d⁸). CN⁻ is SFL. 3d electrons pair, one 3d orbital is empty. Hybridization: dsp² (Square Planar). Magnetic: Diamagnetic (0 unpaired e⁻).

Complexes with Coordination Number 6

• Low Spin Complexes: SFL, electrons pair up, minimum unpaired electrons, d²sp³ (inner orbital).

• High Spin Complexes: WFL, Hund's rule followed, maximum unpaired electrons, sp³d² (outer orbital).

• Example: [Cr(NH₃)₆]³⁺: Cr³⁺ (3d³). NH₃ is SFL. Hybridization: d²sp³ (Octahedral). Magnetic: Paramagnetic (3 unpaired e⁻).

Limitations of Valence Bond Theory (VBT)

1. No explanation for color.

2. Arbitrary rules for SFL/WFL and pairing.

3. No clear reason for inner/outer orbital complexes.

4. Cannot interpret spectra.

Crystal Field Theory (CFT)

Purely electrostatic model. Focuses on repulsion between metal d-electrons and ligand lone pairs, which causes crystal field splitting.

The d-Orbitals: Shapes and Classification

• t₂g: dxy, dyz, dzx (lobes between axes).

• e₉: dx²-y², dz² (lobes along axes).

Crystal Field Splitting of d-Orbitals

When ligands approach, d-orbitals split into different energy levels.

• Octahedral Complexes (CN 6): Ligands approach along axes. e₉ orbitals increase in energy, t₂g orbitals decrease. Energy difference is Δo.

• e₉ destabilized by +0.6Δo, t₂g stabilized by -0.4Δo.

• Tetrahedral Complexes (CN 4): Ligands approach between axes. Splitting is inverted: t₂g orbitals increase, e₉ orbitals decrease. Energy difference is Δt.

• Δt = (4/9)Δo.

Electronic Configuration in Octahedral Complexes

Electron filling depends on Δo vs. Pairing Energy (P).

• Weak Field Ligands (High Spin): Δo < P. Electrons fill singly first (Hund's rule).

• Strong Field Ligands (Low Spin): Δo > P. Electrons pair up in t₂g first.

Calculation of Crystal Field Splitting Energy (CFSE)

CFSE = (No. of electrons in T₂g × -0.4Δ₀) + (No. of electrons in E₉ × +0.6Δ₀)

• For D¹: CFSE = -0.4Δ₀.

• For D⁷ (Low Spin): CFSE = -1.8Δ₀.

• PYQ: For D⁴ ion in an octahedral complex where Δo > P (SFL), the configuration is t₂g⁴ e₉⁰.

Werner's Coordination Theory

Metals exhibit two valencies:

• Primary Valency: Corresponds to oxidation state, satisfied by negative ions, shown by dotted lines (---).

• Secondary Valency: Corresponds to coordination number, satisfied by ligands, shown by solid lines (—).

• Example: [Co(NH3)5Cl]Cl2. Primary Valency = 3, Secondary Valency = 6.

• PYQ: [Co(NH3)5(SO4)]Cl with AgNO₃ gives white precipitate because Cl⁻ is a counter-ion.

Isomerism in Coordination Compounds

Same molecular formula, different properties.

I. Structural Isomerism

Different connectivity of atoms.

• Ionization Isomerism: Produce different ions in solution (e.g., [Co(NH3)5(SO4)]Br vs. [Co(NH3)5Br](SO4)).

• Hydrate Isomerism: Different number of water molecules inside/outside coordination sphere (e.g., [Cr(H2O)6]Cl3 vs. [Cr(H2O)5Cl]Cl2·H2O).

• Coordination Isomerism: Ligand interchange between cationic and anionic complex ions (e.g., [Co(NH3)6][Cr(CN)6] vs. [Cr(NH3)6][Co(CN)6]).

• Linkage Isomerism: Ambidentate ligand bonds through different donor atoms (e.g., [Co(NH3)5(NO2)]²⁺ vs. [Co(NH3)5(ONO)]²⁺).

II. Stereoisomerism

Different spatial arrangement of ligands.

• Geometrical Isomerism (Cis/Trans):

• Square Planar (CN 4): MA₂B₂, MA₂BC, MABCD show GI. MA₄, MA₃B do not. Tetrahedral complexes do not show GI.

• Octahedral (CN 6): MA₄B₂ (cis/trans). MA₃B₃ (Fac/Mer). [Co(en)₂Cl₂]⁺ shows cis/trans.

• Optical Isomerism (Enantiomers): Non-superimposable mirror images. Requires absence of plane of symmetry.

• Trans isomers often have a plane of symmetry, thus optically inactive.

• Cis isomers often lack symmetry, thus optically active.

• Square planar complexes are always optically inactive.

Color of Coordination Compounds

Explained by d-d transitions. Electron absorbs visible light to jump between split d-orbitals. We see the complementary color.

• Conditions for Color:

1. Presence of Ligands: To cause d-orbital splitting.

2. Partially Filled d-Orbitals: To allow electron transitions (d⁰ or d¹⁰ complexes are colorless).

Organometallic Compounds

Contain at least one direct metal-carbon (M-C) bond.

• Metal Carbonyls: CO ligands (e.g., Ni(CO)₄). Exhibit synergic bonding:

1. σ bond: CO donates lone pair to vacant metal d-orbital.

2. π bond (Back-bonding): Metal donates electrons from filled d-orbital to vacant CO π* orbital, strengthening the M-C bond.