Chemistry Chapter 4 Notes of CBSE Class 12

This note covers the d-block and f-block elements of the periodic table, known as transition metals and inner transition metals , respectively. It details their positions, electronic configurations, and general characteristics such as physical properties, atomic and ionic sizes, ionization enthalpies, oxidation states, magnetic properties, and tendency to form coloured ions, complex compounds, interstitial compounds, and alloys. Key concepts like Lanthanoid contraction and Actinoid contraction are explained. The note also describes the preparation and properties of important compounds like Potassium Dichromate (K₂Cr₂O₇) and Potassium Permanganate (KMnO₄) , highlighting their oxidising nature.

1. Position and Definition

• d-block elements are found in Groups 3-12 of the periodic table.
• Their characteristic feature is the progressive filling of the (n-1)d orbitals in each of the four long periods.
• Transition metals are generally referred to as the d-block elements.
• IUPAC Definition: Transition metals are defined as metals that have an incomplete d subshell either in their neutral atom or in their ions.
• Exceptions (Not Transition Metals): Zinc (Zn), Cadmium (Cd), and Mercury (Hg) (Group 12) are generally not regarded as transition metals because they have a full d¹⁰ configuration in their ground state as well as in their common oxidation states. However, their chemistry is often studied alongside transition metals as they are the end members of the 3d, 4d, and 5d series.
• f-block elements are placed in a separate panel at the bottom of the periodic table.
• They consist of elements where the 4f and 5f orbitals are progressively filled.
• They are referred to as inner transition metals .
• The two series are:
• 4f series: Cerium (Ce) to Lutetium (Lu), known as Lanthanoids .
• 5f series: Thorium (Th) to Lawrencium (Lr), known as Actinoids .

2. Electronic Configurations

• Transition Elements (d-block):
• General electronic configuration of outer orbitals: (n-1)d¹ ⁻ ¹⁰ ns ¹ ⁻ ² .
• (n-1)d represents the inner d orbitals with 1 to 10 electrons.
• ns represents the outermost orbital with 1 or 2 electrons.
• Exceptions occur due to the small energy difference between (n-1)d and ns orbitals and the stability of half-filled (d⁵) and completely filled (d¹⁰) orbitals.
• Chromium (Cr): [Ar] 3d⁵ 4s¹ instead of 3d⁴ 4s².
• Copper (Cu): [Ar] 3d¹⁰ 4s¹ instead of 3d⁹ 4s².
• Palladium (Pd): [Kr] 4d¹⁰ 5s⁰ .
• For Zn, Cd, Hg, and Cn, the outer configuration is (n-1)d¹⁰ ns² .
• When forming ions, ns electrons are lost before (n-1)d electrons .
• Example: Scandium (Z=21) is a transition element because its ground state configuration is 3d¹ 4s². Scandium atom has incompletely filled 3d orbitals in its ground state (3d¹).
• Example: Zinc (Z=30) is not a transition element because it has completely filled d orbitals (3d¹⁰) in its ground state (3d¹⁰ 4s²) and in its oxidised state (+2 state is 3d¹⁰).
• Example: Silver (Ag, Z=47) has a ground state configuration of 4d¹⁰ 5s¹. It is considered a transition element because it can exhibit the +2 oxidation state, where it has an incompletely filled d-orbital (4d⁹).

Inner Transition Elements (f-block):

• Lanthanoids:
• Atoms generally have 6s² common, with variable occupancy of the 4f level .
• The most stable oxidation state is +3 .
• Electronic configurations of tripositive ions (Ln³⁺) are generally of the form 4fⁿ (n = 1 to 14).
• Example: Ce³⁺ is [Xe] 4f¹.
• Actinoids:
• Atoms are believed to have 7s² and variable occupancy of the 5f and 6d subshells .
• The 14 electrons are formally added to 5f, starting from Pa, but 5f orbitals are complete at Lr (Z=103).
• Irregularities are related to the stability of f⁰, f⁷, and f¹⁴ configurations in the 5f orbitals.
• Example: Am is [Rn] 5f⁷ 7s²; Cm is [Rn] 5f⁷ 6d¹ 7s².

3. General Characteristics of Transition Elements (d-Block)

Physical Properties:

• Typical metallic properties: high tensile strength, ductility, malleability, high thermal and electrical conductivity, metallic lustre .
• Except for Zn, Cd, Hg, and Mn, they have typical metallic structures.
• Very hard and have low volatility (except Zn, Cd, Hg).
• High melting and boiling points . This is attributed to the involvement of a greater number of electrons from (n-1)d in addition to ns electrons in interatomic metallic bonding .
• Melting points tend to rise to a maximum at d⁵ configuration in any row.
• High enthalpies of atomisation . Maxima around the middle of each series (d⁵) indicate that one unpaired electron per d orbital is favourable for strong interatomic interaction.
• Heavier transition metals (2nd and 3rd series) generally have greater enthalpies of atomisation than the first series.

Atomic and Ionic Sizes:

• Within a series (horizontal): Generally show a progressive decrease in radius with increasing atomic number for ions of the same charge. Atomic radii also show this trend but with smaller variation. This is because the shielding effect of a d electron is not very effective, increasing the net electrostatic attraction.
• Across series (vertical): Increase from the first (3d) to the second (4d) series, but radii of the third (5d) series are virtually the same as the corresponding members of the second series.
• This phenomenon is the result of Lanthanoid contraction . The filling of 4f orbitals before the 5d series leads to a regular decrease in atomic radii, compensating for the expected size increase.
• Lanthanoid contraction causes second and third series elements (e.g., Zr and Hf) to have similar radii and physical/chemical properties.
• The imperfect shielding of one 4f electron by another is less than that of one d electron by another.
• Density generally increases across a series due to decreasing metallic radius and increasing atomic mass.

Ionisation Enthalpies:

• Increase along each series from left to right due to increasing nuclear charge and ineffective shielding by d electrons.
• Successive ionisation enthalpies do not increase as steeply as in non-transition elements.
• Variation within a series is much less compared to non-transition elements across a period.
• First ionisation enthalpy generally increases, but the increase in second and third ionisation enthalpies is higher.
• Irregular trends in the first ionisation enthalpy of the 3d series can be accounted for by the relative energies of 4s and 3d orbitals and effective shielding.
• Irregularities in second and third ionisation enthalpies break for Mn²⁺ (d⁵) and Fe³⁺ (d⁵), which have stable half-filled configurations.
• Factors influencing ionisation enthalpy: electron attraction to nucleus, electron repulsion, and exchange energy (responsible for stability of half-filled/filled orbitals).
• High third ionisation enthalpies for Cu, Ni, and Zn indicate difficulty in achieving oxidation states greater than +2.

Oxidation States:

• Transition elements exhibit a great variety of oxidation states .
• The lowest common oxidation state is +2 .
• The greatest number of oxidation states occurs in or near the middle of the series (e.g., Manganese exhibits +2 to +7).
• Fewer oxidation states at the ends (Sc, Ti have few electrons; Cu, Zn have many d electrons leaving fewer orbitals for bonding).
• Maximum oxidation states (up to Mn) correspond to the sum of s and d electrons.
• Variability arises from the incomplete filling of d orbitals , with oxidation states often differing by unity (e.g., VII, VIII, VIV, VV). This is in contrast to non-transition elements where oxidation states often differ by two.
• In heavier d-block members (Groups 4-10), higher oxidation states are favoured , opposite to the inert pair effect in p-block. For example, Mo(VI) and W(VI) are more stable than Cr(VI).
• Low oxidation states (like 0 in Ni(CO)₄, Fe(CO)₅) are found in complex compounds with π-acceptor ligands.
• Example: Scandium (Z=21) does not exhibit variable oxidation states. Its common state is +3.

Standard Electrode Potentials (E⁰):

• E⁰(M²⁺/M) values are generally less negative across the series, related to the increase in the sum of first and second ionisation enthalpies.
• Irregularities are observed, e.g., Mn, Ni, Zn values are more negative than expected.
• The positive E⁰(Cu²⁺/Cu) value (+0.34V) accounts for copper's inability to liberate H₂ from acids. This is due to high enthalpy of atomisation and low hydration enthalpy of Cu²⁺.
• Stability of d⁵ (Mn²⁺) and d¹⁰ (Zn²⁺) configurations and high negative hydration enthalpy (Ni) influence E⁰ values.
• E⁰(M³⁺/M²⁺) values show varying trends.
• Low value for Sc³⁺/Sc²⁺ reflects the stability of Sc³⁺ (noble gas config).
• High value for Zn³⁺/Zn²⁺ due to removing electron from stable d¹⁰ (Zn²⁺).
• High value for Mn³⁺/Mn²⁺ shows stability of Mn²⁺ (d⁵). Mn³⁺ is a strong oxidising agent.
• Low value for Fe³⁺/Fe²⁺ shows extra stability of Fe³⁺ (d⁵). Fe³⁺ is a strong oxidising agent.
• Low value for V³⁺/V²⁺ related to stability of V²⁺ (half-filled t₂g level). V²⁺ is a strong reducing agent.

Stability of Higher Oxidation States:

• Highest oxidation numbers are often achieved in oxides and fluorides . This is because oxygen and fluorine are small and highly electronegative, effectively oxidising the metal to its highest state.
• Highest oxidation state in simple halides: TiX₄ (+4), VF₅ (+5), CrF₆ (+6). Mn shows +4 in MnF₄, but +7 in Mn₂O₇.
• Ability of fluorine to stabilise high oxidation states is due to high lattice energy (CoF₃) or high bond enthalpy (VF₅, CrF₆).
• Ability of oxygen to stabilise high oxidation states exceeds that of fluorine. This is due to oxygen's ability to form multiple bonds to metals.
• Examples of highest oxidation states in oxides/oxoanions: Sc₂O₃ (+3), Mn₂O₇ (+7), CrO₃ (+6), V₂O₅ (+5), (FeO₄)²⁻ (+6), VO₂⁺ (+5), Cr₂O₇²⁻ (+6), MnO₄⁻ (+7).
• Acidity of oxides increases with oxidation number. Lower oxides are basic (V₂O₃, CrO), intermediate are amphoteric (V₂O₄, V₂O₅, Cr₂O₃), higher are acidic (CrO₃, Mn₂O₇).

Chemical Reactivity:

• Vary widely. Many are electropositive and dissolve in mineral acids; a few are 'noble' (unaffected by single acids).
• First series metals (except Cu) are relatively reactive and oxidised by 1M H⁺, though reaction rate can be slow (e.g., Ti, V are passive to dilute non-oxidising acids).
• Ti²⁺, V²⁺, and Cr²⁺ are strong reducing agents ; Mn³⁺ and Co³⁺ are strong oxidising agents . Cr²⁺ is stronger reducing agent than Fe²⁺ because Cr²⁺ (d⁴) changing to Cr³⁺ (d³) results in a stable half-filled t₂g level, whereas Fe²⁺ (d⁶) changing to Fe³⁺ (d⁵) results in a stable half-filled d⁵ configuration.

Magnetic Properties:

• Substances can be diamagnetic (repelled), paramagnetic (attracted), or ferromagnetic (strongly attracted).
• Many transition metal ions are paramagnetic . Ferromagnetism is an extreme form of paramagnetism.
• Paramagnetism arises from the presence of unpaired electrons .
• For first series transition metals, the magnetic moment is primarily determined by the number of unpaired electrons (spin-only formula): µ = √[n(n+2)] BM, where n is the number of unpaired electrons.
• Magnetic moment increases with the increasing number of unpaired electrons.
• Observed magnetic moments generally agree well with calculated spin-only values for the first series, except for some complex cases (Fe²⁺, Co²⁺, Ni²⁺).
• Zn²⁺ (3d¹⁰) and Sc³⁺ (3d⁰) have 0 unpaired electrons and are diamagnetic .
• Mn²⁺ (3d⁵) has 5 unpaired electrons and high paramagnetism.

Formation of Coloured Ions:

• Many transition metal ions are coloured in solid state and aqueous solutions.
• Colour is attributed to the presence of unpaired d electrons and d-d transitions .
• When an electron absorbs energy (from visible light) to get excited from a lower energy d orbital to a higher energy d orbital, the observed colour is the complementary colour of the light absorbed.
• Ions with d⁰ or d¹⁰ configurations are usually colourless (e.g., Sc³⁺ (3d⁰), Ti⁴⁺ (3d⁰), Zn²⁺ (3d¹⁰), Cu⁺ (3d¹⁰)).
• Ions with unpaired electrons are typically coloured (e.g., Ti³⁺ (3d¹)-purple, V⁴⁺ (3d¹)-blue, V³⁺ (3d²)-green, V²⁺ (3d³)-violet, Cr³⁺ (3d³)-violet, Cr²⁺ (3d⁴)-blue, Mn³⁺ (3d⁴)-violet, Mn²⁺ (3d⁵)-pink, Fe³⁺ (3d⁵)-yellow, Fe²⁺ (3d⁶)-green, Co³⁺ (3d⁶)-blue, Co²⁺ (3d⁷)-pink, Ni²⁺ (3d⁸)-green, Cu²⁺ (3d⁹)-blue).

Formation of Complex Compounds:

• Transition metals form a large number of complex compounds .
• Reasons: comparatively smaller sizes of metal ions, high ionic charges , and availability of d orbitals for bond formation with ligands.
• Examples: [Fe(CN)₆]³⁻, [Fe(CN)₆]⁴⁻, [Cu(NH₃)₄]²⁺.

Catalytic Properties:

• Transition metals and their compounds exhibit catalytic activity .
• Reasons: ability to adopt multiple oxidation states and form complexes .
• Mechanism: involves formation of bonds between reactants and catalyst surface atoms (using d and s electrons), increasing reactant concentration and weakening bonds. Ability to change oxidation states also makes them effective catalysts.
• Examples: V₂O₅ (Contact Process), finely divided Fe (Haber's Process), Ni (Hydrogenation), Fe(III) (Iodide-Persulphate reaction).

Formation of Interstitial Compounds:

• Formed when small atoms (H, C, N) are trapped inside the crystal lattices of metals.
• Usually non-stoichiometric .
• Properties: high melting points (higher than pure metals), very hard (some borides near diamond hardness), retain metallic conductivity , chemically inert .

Alloy Formation:

• Transition metals readily form alloys due to similar radii and other characteristics.
• Alloys can be homogeneous solid solutions . Formed by atoms with metallic radii within about 15% of each other.
• Properties: often hard and have high melting points .
• Examples: Ferrous alloys (steel with Cr, V, W, Mo, Mn), stainless steel, brass (Cu-Zn), bronze (Cu-Sn).

4. Important Compounds of Transition Elements

Potassium Dichromate (K₂Cr₂O₇):

• Prepared from chromite ore (FeCr₂O₄) by fusion with alkali carbonate (e.g., Na₂CO₃) in air, followed by acidification and treatment with KCl.
• FeCr₂O₄ + Na₂CO₃ + O₂ ® Na₂CrO₄ + Fe₂O₃ + CO₂
• Na₂CrO₄ (yellow) + H⁺ ® Na₂Cr₂O₇ (orange)
• Na₂Cr₂O₇ + KCl ® K₂Cr₂O₇ (orange crystals) + NaCl
• Chromates (CrO₄²⁻, tetrahedral) and dichromates (Cr₂O₇²⁻, two tetrahedra sharing a corner with Cr-O-Cr angle 126°) are interconvertible based on pH.
• CrO₄²⁻ (yellow) ⇌ Cr₂O₇²⁻ (orange) (Acidic ⇌ Alkaline)
• Oxidation state of Cr is +6 in both.
• Strong oxidising agents, especially in acidic solution .
• Oxidising action (acidic): Cr₂O₇²⁻ + 14H⁺ + 6e⁻ ® 2Cr³⁺ + 7H₂O (E⁰ = 1.33V).
• Oxidises I⁻ to I₂, H₂S to S, Sn(II) to Sn(IV), Fe(II) to Fe(III).
• Used in leather industry, as oxidant in organic chemistry, and as a primary standard in volumetric analysis.

Potassium Permanganate (KMnO₄):

• Prepared by fusion of MnO₂ (pyrolusite ore) with alkali hydroxide (KOH) and oxidising agent (KNO₃ or O₂) to form manganate (K₂MnO₄, green), followed by disproportionation in neutral/acidic solution or electrolytic oxidation in alkaline solution.
• MnO₂ + KOH + O₂ ® K₂MnO₄ + H₂O
• MnO₄²⁻ (green) + H⁺ ® MnO₄⁻ (purple) + MnO₂ + H₂O (Disproportionation)
• Dark purple (almost black) crystals.
• Manganate ion (MnO₄²⁻, green) is paramagnetic (one unpaired electron).
• Permanganate ion (MnO₄⁻, purple) is diamagnetic (absence of unpaired electron).
• Strong oxidising agent, depends on pH.
• Reduction in different media:
• Acidic: MnO₄⁻ + 8H⁺ + 5e⁻ ® Mn²⁺ + 4H₂O (E⁰ = +1.52V). Oxidises oxalates, Fe(II), nitrites, iodides, H₂S, sulphites. Reaction with HCl is unsatisfactory as HCl is oxidised to Cl₂.
• Neutral/Faintly alkaline: MnO₄⁻ + 2H₂O + 3e⁻ ® MnO₂ + 4OH⁻ (E⁰ = +1.69V in acidic, but effectively lower here). Oxidises iodide to iodate, thiosulphate to sulphate, Mn(II) to MnO₂.
• Strongly alkaline: MnO₄⁻ + e⁻ ® MnO₄²⁻ (E⁰ = +0.56V).
• Uses: analytical chemistry, oxidant in organic chemistry, bleaching, decolourisation of oils.

5. The Inner Transition Elements (f-Block)

Lanthanoids (4f series):

• Fourteen elements following Lanthanum (La). La is often included in discussions.
• Resemble one another more closely than transition elements in a series.
• Lanthanoid Contraction: Gradual decrease in atomic and ionic radii from Lanthanum (La) to Lutetium (Lu).
• Cause: Imperfect shielding of one 4f electron by another as nuclear charge increases along the series. Shielding by 4f electrons is less effective than d electrons.
• Consequence: Radii of third transition series members are very similar to corresponding second series members (e.g., Zr and Hf). This accounts for their co-occurrence and difficulty in separation.
• Oxidation States: Predominantly +3 . Occasionally exhibit +2 and +4 due to stability of f⁰, f⁷, and f¹⁴ configurations.
• Ce(IV) is favoured by f⁰ config but is a strong oxidant (E⁰ = +1.74V) reverting to +3.
• Pr, Nd, Tb, Dy show +4 only in oxides (MO₂).
• Eu(II) (f⁷) is a strong reducing agent reverting to +3.
• Yb(II) (f¹⁴) is a reductant.
• Tb(IV) (f⁷) is an oxidant.
• Sm exhibits +2 and +3.
• General Characteristics: Silvery white soft metals (hardness increases with Z). Typical metallic structure, good conductors. Melting points 1000-1200 K (except Sm at 1623 K).
• Colour: Many trivalent lanthanoid ions are coloured . Attributed to f electrons, but absorption bands are narrow due to excitation within f level. Ions with f⁰ (La³⁺, Ce⁴⁺) and f¹⁴ (Yb²⁺, Lu³⁺) configurations are colourless.
• Magnetic Properties: Ions other than f⁰ and f¹⁴ types are paramagnetic .
• Ionisation Enthalpies: First and second are comparable to Calcium. Variations in third I.E. show stability of f⁰, f⁷, f¹⁴ (low values for La, Gd, Lu).
• Chemical Behaviour: Earlier members reactive like calcium; heavier members behave more like aluminium. E⁰(Ln³⁺/Ln) values -2.2 to -2.4V (Eu -2.0V). React with hydrogen, carbon (carbides), dilute acids (liberate H₂), halogens (halides). Form oxides (M₂O₃) and basic hydroxides (M(OH)₃).
• Uses: Alloy steels (mischmetall), catalysts (petroleum cracking), phosphors in television screens.

Actinoids (5f series):

• Fourteen elements following Actinium (Ac). Ac is often included.
• All are radioactive . Latter members have very short half-lives, making study difficult.
• Actinoid Contraction: Gradual decrease in size of atoms or M³⁺ ions across the series. Similar to lanthanoid contraction but greater from element to element due to poorer shielding by 5f electrons.
• Electronic Configurations: Have 7s² and variable 5f/6d occupancy. Irregularities linked to f⁰, f⁷, f¹⁴ stability. 5f orbitals are less buried than 4f and 5f electrons participate in bonding to a greater extent.
• Oxidation States: Show a greater range than lanthanoids. +3 is the common state.
• First half: frequently exhibit higher oxidation states . Maximum state increases from +4 (Th) to +7 (Np).
• Latter half: stability of higher states decreases.
• Distribution of oxidation states is uneven and differs between former and later elements.
• General Characteristics and Comparison with Lanthanoids:
• Silvery appearance, variety of structures (due to irregularities in radii).
• Highly reactive metals when finely divided. Attack boiling water and HCl. Passivated by HNO₃. Unaffected by alkalies.
• Magnetic properties more complex; values lower than lanthanoids.
• Ionisation enthalpies of early actinoids are lower than early lanthanoids. This is because 5f electrons penetrate less into the inner core and are less effectively shielded, making them more available for bonding.
• Behaviour similar to lanthanoids (close similarities, gradual property variation without oxidation state change) is more evident in the second half of the series.
• The chemistry is more complex than lanthanoids due to the wider range of oxidation states and radioactivity.

6. Comparison: Transition Elements vs Non-Transition Elements

7. Comparison: Lanthanoids vs Actinoids

8. Applications of d- and f-Block Elements

• Iron and steels (construction materials, alloys with Cr, Mn, Ni, W, Mo, V).
• TiO₂ (pigment industry).
• MnO₂ (dry battery cells).
• Zn and Ni/Cd (battery industry).
• Coinage metals (Cu, Ag, Au).
• Alloys (Cu/Ni in coins, brass Cu-Zn, bronze Cu-Sn).
• Catalysts (V₂O₅ for SO₂ oxidation, TiCl₄+Al(CH₃)₃ Ziegler catalyst for polyethylene, Fe for Haber process, Ni for hydrogenation, PdCl₂ for Wacker process, Ni complexes for polymerisation).
• AgBr (photographic industry).
• Inner transition elements: Th, Pa, U (sources of nuclear energy).
• Lanthanoids (mischmetall alloy for bullets, shells, lighter flints; mixed oxides as catalysts; individual oxides as phosphors).

Frequently Asked Questions (FAQs)

1. Why are Zinc, Cadmium, and Mercury not considered transition metals?
• They have a complete d¹⁰ electronic configuration in their ground state.
• They also have a complete d¹⁰ configuration in their common oxidation state of +2.
• According to IUPAC definition, transition metals must have an incomplete d subshell in either the neutral atom or its ions.
2. What is Lanthanoid Contraction and its main consequence?
• It is the gradual decrease in the atomic and ionic radii of Lanthanoids with increasing atomic number.
• It is caused by the poor shielding effect of 4f electrons.
• The main consequence is that elements in the third transition series have very similar radii and properties to their corresponding elements in the second transition series.
3. Why do transition metals exhibit variable oxidation states?
• This property arises from the incomplete filling of their (n-1)d orbitals.
• Electrons from both the ns and (n-1)d orbitals can participate in bonding.
• The small energy difference between these orbitals allows for the loss of a variable number of electrons.