Comprehensive Study

The d-block and f-block elements form the heart of inorganic chemistry, displaying a fascinating array of properties due to the involvement of their inner-shell (n-1)d and (n-2)f electrons in bonding.

1. d-Block Elements (Transition Elements)

• Definition: These are elements located between the s-block and p-block elements in the periodic table, specifically from Group 3 to Group 12. A transition element is defined as an element whose atom, either in its ground state or in any one of its common oxidation states, has partially filled (n-1)d orbitals.
  • Note on Zn, Cd, Hg: Zinc (Zn), Cadmium (Cd), and Mercury (Hg) (Group 12) have fully filled (n-1)d10 and ns2 configurations in their ground states and in their common +2 oxidation states. Therefore, they are often not considered true transition elements by this strict definition, although they are still classified within the d-block.
• General Electronic Configuration: (n−1)d1−10ns1−2.
  • ‘n’ represents the outermost shell (period number).
  • (n−1) represents the penultimate shell (one less than the period number), where the d-orbitals are being filled.
  • Important Exceptions to Aufbau Principle:
    • Chromium (Cr, Z=24): Expected: [Ar]3d4 4s2. Actual: [Ar]3d5 4s1. (Achieves stable half-filled d-subshell).
    • Copper (Cu, Z=29): Expected: [Ar]3d9 4s2. Actual: [Ar]3d10 4s1. (Achieves stable fully-filled d-subshell).
    • Similar exceptions occur for other elements in their respective series (e.g., Molybdenum (Mo), Silver (Ag), Gold (Au)) due to the extra stability associated with half-filled or completely filled d-orbitals.
• Metallic Properties: All transition elements are highly metallic. They are typically hard, possess high melting and boiling points, and are excellent conductors of heat and electricity. This strong metallic character arises from the strong metallic bonding facilitated by the extensive delocalization of electrons, involving both the ns and (n-1)d electrons.

2. General Characteristics of d-Block Elements

• 1. Metallic Character:
  • Physical Properties: Exhibit typical metallic properties such as high tensile strength, ductility, malleability, and characteristic metallic lustre.
  • High Melting and Boiling Points: They have exceptionally high melting and boiling points due to strong metallic bonds, which are a result of the overlap of valence s and d orbitals leading to strong interatomic interactions. Tungsten (W) has the highest melting point among all metals.
  • Exceptions: Zinc (Zn), Cadmium (Cd), and Mercury (Hg) have significantly lower melting points compared to other transition metals. This is because their d-orbitals are fully filled (d10 configuration), meaning the d-electrons do not participate significantly in metallic bonding, leading to weaker cohesive forces.
• 2. Variable Oxidation States:
  • Most Characteristic Property: This is the hallmark feature of transition elements.
  • Reason: The energy difference between the (n-1)d orbitals and ns orbitals is very small. Consequently, both sets of electrons can participate in bond formation.
  • Trends:
    • The most common oxidation state observed is often +2 (resulting from the loss of two ns electrons).
    • Higher oxidation states are achieved by successively involving electrons from the (n-1)d orbitals.
    • The maximum oxidation state generally increases from Group 3 to Group 7 (Manganese, Mn, shows the highest +7 oxidation state in the 3d series) and then decreases.
    • Stability of Oxidation States:
      • For lighter transition metals (e.g., 3d series), higher oxidation states are generally more stable when combined with highly electronegative elements like oxygen or fluorine (e.g., KMnO4, K2Cr2O7).
      • The stability of higher oxidation states generally decreases down a group (e.g., Fe prefers +3, Ru and Os prefer +8). This is due to increasing size and decreasing electronegativity, making it harder to maintain high positive charge.
      • The stability of lower oxidation states generally increases down a group (related to the inert pair effect, though less pronounced than in p-block).
  • Examples (3d series):
    • Scandium (Sc): Only +3 (d0 configuration in Sc3+ is very stable).
    • Titanium (Ti): +2, +3, +4 (most stable is +4, as in TiO2).
    • Vanadium (V): +2, +3, +4, +5 (most stable is +5, as in V2O5).
    • Chromium (Cr): +2, +3, +4, +5, +6 (most common are +3 and +6, as in Cr2O3, CrO3, K2Cr2O7).
    • Manganese (Mn): +2, +3, +4, +5, +6, +7 (most common are +2, +4, +7, as in MnSO4, MnO2, KMnO4).
    • Iron (Fe): +2, +3 (most common are +2 and +3, as in FeSO4, Fe2O3). +2 is more stable in aqueous solution, +3 in air.
    • Cobalt (Co): +2, +3 (most common are +2 and +3).
    • Nickel (Ni): +2, +3 (most common is +2).
    • Copper (Cu): +1, +2 (most common is +2, as in CuSO4). Cu+ (d10) is stable in solid but disproportionates in aqueous solution.
    • Zinc (Zn): Only +2 (d10 configuration in Zn2+ is very stable).
• 3. Formation of Coloured Ions:
  • Characteristic Property: Most compounds of transition metals are distinctly coloured, both in solid state and in aqueous solutions.
  • Reason (d-d transitions): This property arises due to the presence of partially filled (n-1)d orbitals. When white light falls on transition metal ions, the d-orbitals (which are degenerate in an isolated atom/ion) get split into different energy levels in the presence of ligands (or surrounding ions). The d-electrons can then absorb specific wavelengths of light from the visible region to jump from a lower energy d-orbital to a higher energy d-orbital (d-d transition). The remaining (complementary) wavelengths are transmitted or reflected, giving the compound its observed colour.
  • Colourless Ions: Ions with completely empty (d0 configuration) or completely filled (d10 configuration) d-orbitals are usually colourless because no d-d transitions are possible.
    • Examples: Sc3+ (d0), Ti4+ (d0), Cu+ (d10), Zn2+ (d10).
  • Examples of Coloured Ions: Cu2+ (blue), Fe3+ (yellow/brown), MnO4- (purple), Cr2O7^2- (orange), Ni2+ (green).
• 4. Paramagnetic Behaviour:
  • Characteristic Property: Most transition metal ions and their compounds are paramagnetic (they are weakly attracted into a magnetic field).
  • Reason: Paramagnetism is due to the presence of unpaired electrons in their (n-1)d orbitals. The spin of an unpaired electron generates a magnetic moment.
  • Calculation of Magnetic Moment: The magnetic moment (μ) is measured in Bohr Magnetons (BM) and is calculated using the “spin-only” formula:
    μ=n(n+2)​ BM Where ‘n’ is the number of unpaired electrons.
  • Diamagnetic Ions: Ions with no unpaired electrons are diamagnetic (they are weakly repelled by a magnetic field).
    • Examples: Sc3+ (d0, n=0), Ti4+ (d0, n=0), Zn2+ (d10, n=0), Cu+ (d10, n=0).
• 5. Catalytic Properties:
  • Excellent Catalysts: Many transition metals and their compounds (both solid and ionic forms) act as excellent catalysts in various chemical reactions.
  • Reasons:
    • Variable Oxidation States: Their ability to exhibit variable oxidation states allows them to readily change their oxidation states and form unstable intermediate compounds, providing new reaction pathways with lower activation energies.
    • Large Surface Area: In finely divided states, they provide a large surface area for reactants to adsorb onto, facilitating the reaction.
    • Ability to Form Complexes: They can form coordination compounds with reactants, providing transient intermediates that facilitate the reaction.
  • Examples:
    • Vanadium(V) oxide (V2O5): Used as a catalyst in the Contact Process for the manufacture of sulfuric acid.
    • Finely divided Iron (Fe): Used in Haber’s Process for the synthesis of ammonia.
    • Nickel (Ni) or Palladium (Pd): Used in the hydrogenation of oils.
    • Titanium trichloride (TiCl3): Used in Ziegler-Natta polymerization of ethene.
• 6. Formation of Interstitial Compounds:
  • Definition: Transition metals form non-stoichiometric compounds with small non-metal atoms (such as Hydrogen (H), Boron (B), Carbon (C), Nitrogen (N)) that get trapped in the interstitial voids (empty spaces or holes) within the metal crystal lattice.
  • Properties: These compounds generally exhibit:
    • High melting points (higher than pure metals).
    • Are very hard (e.g., carbides of Ti, W, Fe).
    • Retain metallic conductivity.
    • Are chemically inert.
  • Example: TiC, Mn4N, Fe3H, ZrC.
• 7. Alloy Formation:
  • High Tendency: Transition metals readily form a large number of alloys among themselves and with non-transition metals.
  • Reason: Their atomic radii are very similar, allowing atoms of one metal to easily substitute for atoms of another metal in the crystal lattice without much distortion.
  • Examples:
    • Steel: An alloy of Iron (Fe) with Carbon (C), Manganese (Mn), Chromium (Cr), etc.
    • Brass: An alloy of Copper (Cu) and Zinc (Zn).
    • Bronze: An alloy of Copper (Cu) and Tin (Sn).
    • Nichrome: An alloy of Nickel (Ni), Chromium (Cr), Iron (Fe), and sometimes Manganese (Mn).
• 8. Formation of Complex Compounds (Coordination Compounds):
  • Strong Tendency: Transition metals have a very strong tendency to form complex compounds (or coordination compounds).
  • Reasons:
    • Small Size of Metal Ions: Allows for strong electrostatic attraction with ligands.
    • High Nuclear Charge: Provides a strong center for coordination.
    • Availability of Empty (n-1)d Orbitals: These orbitals are of suitable energy to accept lone pairs of electrons from ligands (Lewis acid-base interaction).
  • Examples: [Fe(CN)6]4-, [Cu(NH3)4]^2+, [Cr(H2O)6]3+, [Ni(CO)4].

2.3. Important Compounds of d-Block Elements

• 1. Potassium Dichromate (K2Cr2O7):
  • Properties: A bright orange-coloured crystalline solid. It is a very powerful oxidizing agent.
  • Preparation (from Chromite Ore, FeCr2O4):
    1. Oxidation of Chromite Ore to Sodium Chromate: Finely powdered chromite ore is fused with sodium carbonate (or potassium carbonate) in the presence of air (oxidizing atmosphere) in a reverberatory furnace. 4FeCr2O4(chromiteore)+8Na2CO3+7O2→8Na2CrO4(sodiumchromate,yellowsolution)+2Fe2O3+8CO2
    2. Conversion of Sodium Chromate to Sodium Dichromate: The yellow solution of sodium chromate is filtered, acidified with sulfuric acid, which converts the chromate ions (CrO4^2-) to dichromate ions (Cr2O7^2-). 2Na2CrO4(aq,yellow)+2H+(aq)→Na2Cr2O7(aq,orange)+2Na+(aq)+H2O(l) (This equilibrium is pH-dependent: in acidic medium, chromate converts to dichromate; in alkaline medium, dichromate converts back to chromate).
    3. Conversion of Sodium Dichromate to Potassium Dichromate: Sodium dichromate is more soluble than potassium dichromate. So, potassium chloride is added to the concentrated solution of sodium dichromate, causing K2Cr2O7 to crystallize out. Na2Cr2O7(aq)+2KCl(aq)→K2Cr2O7(s)+2NaCl(aq)
  • Oxidizing Nature: K2Cr2O7 is a very strong oxidizing agent in acidic medium. The dichromate ion (Cr2O7^2-) gets reduced to Cr3+ ion. Cr2O72−+14H++6e−→2Cr3++7H2O (Standard electrode potential E° = +1.33 V)
    • Examples of Oxidation Reactions:
      • Oxidizes Iodides to Iodine: K2Cr2O7+6KI+7H2SO4→Cr2(SO4)3+4K2SO4+3I2+7H2O
      • Oxidizes Ferrous (Fe2+) to Ferric (Fe3+): Cr2O72−+6Fe2++14H+→2Cr3++6Fe3++7H2O
      • Oxidizes SO2 to SO4^2-: Cr2O72−+3SO2+2H+→2Cr3++3SO42−+H2O
• 2. Potassium Permanganate (KMnO4):
  • Properties: A dark purple-coloured crystalline solid. It is also a powerful oxidizing agent.
  • Preparation (from Pyrolusite Ore, MnO2):
    1. Conversion of MnO2 to Potassium Manganate: Powdered pyrolusite ore (MnO2) is fused with an alkali (KOH) and an oxidizing agent (like KNO3 or air/O2). This forms green-coloured potassium manganate (K2MnO4). 2MnO2+4KOH+O2→2K2MnO4(green)+2H2O
    2. Oxidation of Potassium Manganate to Potassium Permanganate:
       • Electrolytic Oxidation: The green K2MnO4 solution is electrolytically oxidized. Manganate ions (MnO4^2-) are oxidized to permanganate ions (MnO4-) at the anode. 2K2MnO4+2H2Oelectrolyticoxidation​2KMnO4(purple)+2KOH+H2
       • Chemical Oxidation (Disproportionation): Manganate disproportionates in acidic or neutral medium to form permanganate and manganese dioxide. 3MnO42−+4H+→2MnO4−+MnO2+2H2O
  • Oxidizing Nature: KMnO4 acts as a very strong oxidizing agent in various media, with different reduction products for Mn.
    • In Acidic Medium (H+): MnO4- gets reduced to Mn2+ (oxidation state changes from +7 to +2, gain of 5 electrons). Standard electrode potential E° = +1.51 V. MnO4−+8H++5e−→Mn2++4H2O
      • Examples of Oxidation Reactions:
        • Oxidizes Ferrous (Fe2+) to Ferric (Fe3+): MnO4−+5Fe2++8H+→Mn2++5Fe3++4H2O
        • Oxidizes Oxalate (C2O4^2-) to CO2: 2MnO4−+5C2O42−+16H+→2Mn2++10CO2+8H2O
        • Oxidizes Iodide (I-) to Iodine (I2): 2MnO4−+10I−+16H+→2Mn2++5I2+8H2O
        • Oxidizes Hydrogen Sulfide (H2S) to Sulfur (S): 2MnO4−+5H2S+6H+→2Mn2++5S+8H2O
    • In Neutral or Weakly Alkaline Medium: MnO4- gets reduced to Manganese Dioxide (MnO2) (brown precipitate, oxidation state changes from +7 to +4, gain of 3 electrons). MnO4−+2H2O+3e−→MnO2(s)+4OH−
    • In Strongly Alkaline Medium: MnO4- gets reduced to Manganate ion (MnO4^2-) (green, oxidation state changes from +7 to +6, gain of 1 electron). MnO4−+e−→MnO42−

3. f-Block Elements (Inner Transition Elements)

• Definition: These are elements where the (n-2)f orbitals are being progressively filled. They are placed separately at the bottom of the periodic table in two series to maintain the periodic table’s structure.
• Two Series:
  • Lanthanoids (4f series): Cerium (Ce, Z=58) to Lutetium (Lu, Z=71). These are elements that follow Lanthanum (La, Z=57).
  • Actinoids (5f series): Thorium (Th, Z=90) to Lawrencium (Lr, Z=103). These are elements that follow Actinium (Ac, Z=89).
• General Electronic Configuration: (n−2)f1−14(n−1)d0−1ns2. (Typically, one electron occupies the (n-1)d orbital, but sometimes it is empty, particularly after f-filling is complete or for very stable f-configurations like f0, f7, f14).

3.1. Lanthanoids (4f-Series): Ce to Lu (14 elements)

• Electronic Configuration: General configuration is [Xe] 4f1-14 5d0-1 6s2. Specific configurations can vary due to stability of f7 or f14.
• Common Oxidation State: The most stable and predominant oxidation state for all lanthanoids is +3.
  • Some elements also exhibit +2 (e.g., Europium (Eu2+, 4f7) and Ytterbium (Yb2+, 4f14)) or +4 (e.g., Cerium (Ce4+, 4f0), Praseodymium (Pr4+), Terbium (Tb4+, 4f7)) states. These alternate oxidation states are more stable when they result in an empty (f0), half-filled (f7), or completely filled (f14) f-orbital configuration.
• Lanthanoid Contraction:
  • Definition: This refers to the gradual and regular decrease in the atomic and ionic radii (specifically for M3+ ions) as one moves across the lanthanoid series from Lanthanum (La) to Lutetium (Lu).
  • Cause: The primary cause is the poor shielding effect of the 4f electrons. As the atomic number increases across the series, the nuclear charge increases. However, the 4f electrons are very ineffective at shielding the outer electrons from this increased nuclear pull. This results in a stronger attraction between the nucleus and the valence electrons, leading to a progressive contraction in size.
  • Consequences (Very Important for JEE/NEET):
    • Similarity in Atomic Radii of 4d and 5d Elements: This is the most significant consequence. Elements of the second transition series (4d, e.g., Zirconium, Zr) and the third transition series (5d, e.g., Hafnium, Hf) that follow the lanthanoids in the same group have almost identical atomic radii. This makes their chemical properties remarkably similar, making their separation very difficult. (e.g., Zr and Hf, Nb and Ta, Mo and W).
    • Slight Increase in Ionization Enthalpy for 5d Elements: Due to the increased effective nuclear charge and smaller size caused by the contraction.
    • Increased Density of 5d Elements: Their atomic masses increase significantly, but their atomic volumes decrease (due to contraction), leading to higher densities.
• Colour: Many lanthanoid ions are coloured in both solid state and solution. This colour arises from f-f transitions, which are relatively weak because the 4f orbitals are well-shielded. Ions with f0, f7, or f14 configurations (e.g., La3+ (4f0), Gd3+ (4f7), Lu3+ (4f14)) are generally colourless.
• Magnetic Properties: Most lanthanoid ions are paramagnetic due to the presence of unpaired electrons in their 4f orbitals. The magnetic moments are generally calculated from orbital contributions as well, not just spin-only, but for NEET/JEE, emphasis is on paramagnetism due to unpaired electrons.
• Chemical Reactivity: All lanthanoids are highly reactive metals, comparable in reactivity to calcium. Reactivity generally increases down the series. They typically react with acids, halogens, and oxygen.
• Basicity of Hydroxides: The basicity of their hydroxides (Ln(OH)3) decreases slightly from La(OH)3 to Lu(OH)3 due to lanthanoid contraction (decreasing ionic radius leads to increasing covalent character). La(OH)3 is the most basic.

3.2. Actinoids (5f-Series): Th to Lr (14 elements)

• Electronic Configuration: General configuration is [Rn] 5f1-14 6d0-1 7s2.
• Common Oxidation State: +3 is the most common and stable oxidation state for actinoids, similar to lanthanoids.
  • However, actinoids exhibit a wider and more varied range of oxidation states compared to lanthanoids. Elements like Uranium (U), Neptunium (Np), Plutonium (Pu) show +3, +4, +5, +6, and even +7 oxidation states. This is because the 5f orbitals are not as deeply buried (less shielded) as the 4f orbitals, and thus they participate more readily in bonding.
• Actinoid Contraction:
  • A gradual decrease in atomic and ionic radii occurs across the actinoid series due to the poor shielding by 5f electrons. This contraction is generally more pronounced than the lanthanoid contraction.
• Colour: Most actinoid ions are coloured, both in solid state and solution, due to f-f transitions.
• Magnetic Properties: Actinoid ions are generally paramagnetic due to unpaired electrons in their 5f orbitals.
• Radioactivity: All actinoids are radioactive. Uranium (U) and Thorium (Th) are relatively long-lived, but most other actinoids are short-lived synthetic elements.
• Chemical Reactivity: Highly reactive metals, especially in finely divided forms. Their reactivity generally decreases across the series.
• Complex Formation: They have a greater tendency to form complexes than lanthanoids. This is attributed to their higher charge density and the availability of 5f, 6d, and 7s orbitals of comparable energy for bonding.

3.3. Comparison between Lanthanoids and Actinoids (Key Distinctions)