The d- and f-Block Elements — NEET notes

d-Block: General Properties

The d-block elements (Groups 3–12) have partially filled d-orbitals in their ground state or common oxidation states. Their general electronic configuration is (n−1)d¹⁻¹⁰ ns¹⁻². Cr ([Ar]3d⁵4s¹) and Cu ([Ar]3d¹⁰4s¹) are exceptions due to extra stability of half-filled and fully-filled d-subshells. All are metals with high melting points, densities, and hardness. They form colored ions, show variable oxidation states, and exhibit catalytic activity.

General config: (n-1)d^ 1-10 ns^ 1-2

Position & Electronic Configuration

d-Block elements occupy Groups 3–12 in periods 4–7. Their distinguishing electron enters the (n−1)d subshell. Chromium and copper have exceptional configurations due to the extra stability of half-filled (d⁵) and fully-filled (d¹⁰) d-subshells.

The d-block elements are placed between the s-block and p-block in the periodic table — Groups 3 through 12, spanning periods 4, 5, 6, and 7. The differentiating electron enters the (n-1)d subshell, NOT the outermost shell.

• General electronic configuration: (n-1)d^ 1-10 ns^ 1-2
• 3d series (Period 4): Sc ( Z=21 ) to Zn ( Z=30 ) — most important for NEET
• 4d series (Period 5): Y ( Z=39 ) to Cd ( Z=48 )
• 5d series (Period 6): La ( Z=57 ), Hf ( Z=72 ) to Hg ( Z=80 )
• The 4s orbital fills before 3d but empties first during ionization (lower principal quantum number has higher nuclear penetration after filling)

Transition Elements vs d-Block Elements

A transition element must have a partially filled d-subshell in its ground state OR in any common oxidation state. Zinc (3d¹⁰4s²) is a d-block element but NOT a transition element because its d-subshell is always full.

IUPAC defines a transition element as one that has an incomplete d-subshell either in its ground state or in one or more of its common oxidation states. This is a stricter criterion than simply being in the d-block.

• Sc ( 3d^1 ) and Ti ( 3d^2 ) are transition elements — partially filled d in ground state
• Cu ( 3d^ 10 4s^1 ) IS a transition element — Cu^ 2+ has 3d^9 (partially filled)
• Zn ( 3d^ 10 4s^2 ) is NOT a transition element — Zn^ 2+ still has 3d^ 10 (completely filled)
• This distinction explains why Zn compounds are colorless and diamagnetic while Cu compounds are colored

Metallic Character & Physical Properties

Transition metals have high melting points, boiling points, enthalpies of atomization, and densities due to strong metallic bonding involving both (n−1)d and ns electrons. They are good conductors and form alloys readily.

Transition metals are harder, denser, and have higher melting points than s-block metals. Why? Because they have more electrons available for metallic bonding. In s-block metals, only 1–2 s -electrons contribute. In transition metals, both ns and (n-1)d electrons participate in metallic bonding.

• Melting points: Generally high; maximum near the middle of each series (Groups 5–6)
• Mn is an exception — anomalously low m.p. for its position (weak metallic bonding due to half-filled d^5 stability)
• Zn, Cd, Hg have low m.p. — fully filled d^ 10 electrons don't participate in bonding
• Enthalpy of atomization follows the same trend — peaks in middle, dips at Mn and Zn
• Density increases from Sc to Cu, with a maximum near the middle of the 3d series

Alloy formation is characteristic of transition metals. Because transition metals have similar atomic sizes and same crystal structures, one metal can substitute for another in the lattice. Examples: steel (Fe + C), brass (Cu + Zn), bronze (Cu + Sn), stainless steel (Fe + Cr + Ni).

Quick recall

Why do transition metals have higher enthalpies of atomization than Group 1 or 2 metals?

Because both ns and (n-1)d electrons contribute to metallic bonding, leading to stronger interatomic forces. More bonding electrons → stronger metallic bonds → higher energy needed to separate atoms.

Oxidation States & Ionization

Transition metals show variable oxidation states because both ns and (n−1)d electrons have comparable energies and can be removed during bonding. The +2 state arises from losing ns² electrons, while higher states involve d-electron removal. Maximum oxidation state increases from Sc (+3) to Mn (+7), then decreases. Ionization enthalpies increase gradually across the 3d series but less steeply than p-block due to poor d-electron shielding. Stability of +2 state linked to exchange energy and d-electron configuration.

Maximum O.S. = number of ns + (n-1)d electrons (up to Mn)

Variable Oxidation States

Transition metals exhibit multiple oxidation states because the energy difference between ns and (n−1)d orbitals is small enough that varying numbers of d-electrons can participate in bonding. The maximum oxidation state reaches +7 at Mn, then decreases.

The hallmark of transition metals is their ability to show multiple oxidation states. This happens because ns and (n-1)d electrons have very similar energies — so a transition metal can lose 2, 3, 4, or even 7 electrons depending on the reaction partner.

• Minimum common O.S. = +2 (loss of two 4s electrons)
• Maximum O.S. increases from Sc (+3) to Mn (+7) — each element uses one more d -electron
• After Mn, maximum O.S. decreases — too many d -electrons hold each other tightly (increased nuclear charge + pairing)
• Sc shows only +3; Zn shows only +2 — both have fixed oxidation states
• Mn shows the widest range: +2 to +7 (all O.S. from +2 through +7 are known)

Key oxidation state data for the 3d series: - Sc: +3 only - Ti: +2, +3, +4 - V: +2, +3, +4, +5 - Cr: +2, +3, +6 - Mn: +2, +3, +4, +6, +7 - Fe: +2, +3, +4, +6 - Co: +2, +3 - Ni: +2, +3 - Cu: +1, +2 - Zn: +2 only (Most stable O.S. in bold)

Stability of Oxidation States

The relative stability of +2 vs +3 (and higher) states depends on d-electron configuration, exchange energy, electrode potential, and the nature of the reaction partner. Half-filled d⁵ (Mn²⁺, Fe³⁺) and fully-filled d¹⁰ (Cu⁺, Zn²⁺) configurations confer extra stability.

The +2 state comes from losing two 4s electrons — this is universal across the 3d series. The stability of the +2 state relative to higher states depends on how easy or hard it is to also remove 3d electrons.

• Mn^ 2+ ( 3d^5 , half-filled) is very stable — resists oxidation to Mn^ 3+
• Fe^ 3+ ( 3d^5 , half-filled) is more stable than Fe^ 2+ ( 3d^6 ) in acidic aqueous solution
• Ti^ 4+ ( 3d^0 ) is most stable for Ti — noble gas core, no d -electrons to destabilize
• Cu^ + ( 3d^ 10 ) is expected to be stable (filled d ), but disproportionates in water: 2Cu^+ arrow Cu^ 2+ + Cu
• Higher O.S. are stabilized by electronegative ligands (F, O) — e.g., CrO_4^ 2- (Cr in +6), MnO_4^- (Mn in +7)

Quick recall

Why is Fe^ 3+ more stable than Fe^ 2+ even though removing one more electron costs energy?

Because Fe^ 3+ achieves a half-filled 3d^5 configuration, which has extra stability due to high exchange energy (maximum number of parallel-spin d -electrons = 10 exchange pairs). This stabilization energy outweighs the third ionization energy cost.

Ionization Enthalpy Trends

Ionization enthalpies of 3d elements increase slowly from left to right due to poor shielding by d-electrons. The increase is not as steep as in p-block because added d-electrons poorly shield each other from nuclear charge.

Across the 3d series (Sc → Zn), ionization enthalpies increase gradually from left to right. The increase is slow and irregular — not a smooth climb like in the p-block.

IE_1 (3d series): Sc (631) < Ti (656) < V (650) < Cr (653) < Mn (717) < Fe (762) < Co (758) < Ni (737) < Cu (745) < Zn (906 kJ/mol)

• General trend: gradual increase from Sc to Zn (nuclear charge increases, but d -electrons shield poorly)
• Mn shows a bump — removing an electron from half-filled 3d^5 costs extra energy
• Zn has the highest IE₁ in the 3d series — fully filled 3d^ 10 4s^2 is very stable
• IE₁ values are generally higher than s-block but lower than p-block elements in the same period
• IE_2 is important: Cr has abnormally high IE_2 because removing the second electron breaks into the stable 3d^5 core

Color, Magnetic & Catalytic Properties

Transition metal compounds are colored due to d-d electronic transitions when light is absorbed and an electron jumps between split d-orbitals. They are paramagnetic when unpaired d-electrons are present — magnetic moment is calculated by the spin-only formula μ = √(n(n+2)) BM. They act as catalysts by providing variable oxidation states and surfaces for reactant adsorption. They also form interstitial compounds and alloys.

= n(n+2) BM

Color of Compounds

When visible light hits a transition metal ion, an electron in a lower-energy d-orbital absorbs a photon and jumps to a higher-energy d-orbital (d-d transition). The color we see is the complementary color of the light absorbed. Ions with d⁰ or d¹⁰ configurations are colorless.

In an isolated transition metal ion, all five d -orbitals have the same energy (degenerate). But when surrounded by ligands (in a compound or solution), the d -orbitals split into two groups with different energies. An electron can absorb visible light and jump from the lower set to the upper set — this is a d-d transition.

• The absorbed wavelength depends on the crystal field splitting energy ( )
• The color seen is the complementary color of the light absorbed
• Cu^ 2+ ( 3d^9 ) absorbs red light → appears blue-green
• Ti^ 3+ ( 3d^1 ) absorbs yellow-green light → appears purple
• Mn^ 2+ ( 3d^5 , all spins parallel) has a very pale pink color — d-d transitions are spin-forbidden (weak absorption)

Quick recall

Why is KMnO_4 intensely purple even though Mn is in the +7 state ( d^0 )?

The color of KMnO_4 is NOT from d-d transitions (Mn⁷⁺ has no d-electrons!). It's from a charge transfer transition — an electron transfers from oxygen's filled orbital to manganese's empty d -orbital. Charge transfer bands are very intense, giving the deep purple color.

Magnetic Properties

Transition metal ions with unpaired d-electrons are paramagnetic (attracted by magnetic field). The spin-only magnetic moment μ = √(n(n+2)) BM, where n = number of unpaired electrons. Ions with all electrons paired (d⁰, d¹⁰, low-spin d⁶) are diamagnetic.

Each unpaired electron acts as a tiny magnet. More unpaired electrons → stronger paramagnetism. The spin-only formula gives the magnetic moment:

= n(n+2) BM (Bohr Magneton)

where n = number of unpaired electrons.

• n = 0 = 0 BM → diamagnetic (repelled by magnetic field)
• n = 1 = 3 = 1.73 BM
• n = 2 = 8 = 2.83 BM
• n = 3 = 15 = 3.87 BM
• n = 4 = 24 = 4.90 BM
• n = 5 = 35 = 5.92 BM (maximum — all five d -orbitals singly occupied)

Examples of magnetic moment calculation: - Fe^ 2+ ( 3d^6 ): ↑↓ ↑ ↑ ↑ ↑ → n = 4 → = 24 = 4.90 BM - Fe^ 3+ ( 3d^5 ): ↑ ↑ ↑ ↑ ↑ → n = 5 → = 35 = 5.92 BM - Ni^ 2+ ( 3d^8 ): ↑↓ ↑↓ ↑↓ ↑ ↑ → n = 2 → = 8 = 2.83 BM - Zn^ 2+ ( 3d^ 10 ): ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ → n = 0 → = 0 (diamagnetic)

Catalytic Activity & Other Properties

Transition metals are excellent catalysts because they offer variable oxidation states (to form intermediates) and have large surface areas for adsorption. They also form interstitial compounds (with H, C, N, B trapped in crystal lattice voids) which are hard and have high melting points.

Many of the most important industrial catalysts are transition metals or their compounds. Two properties make them exceptional catalysts: 1. Variable oxidation states — they can accept and donate electrons, forming reaction intermediates 2. Ability to adsorb reactants — their surfaces bind reactant molecules, weakening bonds and lowering activation energy

• Fe — Haber process ( N_2 + 3H_2 arrow 2NH_3 ), finely divided iron at 500°C
• V₂O₅ — Contact process ( 2SO_2 + O_2 arrow 2SO_3 )
• Pt/Pd/Rh — Catalytic converters in cars (reduce NO _x , oxidize CO)
• Ni — Hydrogenation of oils (vegetable oil → vanaspati ghee)
• MnO₂ — Decomposition of KClO_3 and H_2O_2
• TiCl₄ + Al(C₂H₅)₃ — Ziegler-Natta catalyst (polymerization of ethene)

Interstitial compounds form when small atoms (H, C, N, B) get trapped in the voids (interstices) of the transition metal crystal lattice. Examples: TiC, steel ( Fe_3C ), TiH. These compounds are: - Very hard (TiC used in cutting tools) - High melting point - Chemically inert - Retain metallic conductivity

Quick recall

Why is finely divided iron used (not a lump) in the Haber process?

Finely divided iron has a much larger surface area for adsorbing N_2 and H_2 molecules. More surface area → more reactant molecules adsorbed simultaneously → faster reaction. This is heterogeneous catalysis — the catalyst and reactants are in different phases.

Important Compounds: KMnO₄

Potassium permanganate (KMnO₄) is a dark purple crystalline solid and one of the most powerful oxidizing agents. Mn is in the +7 oxidation state. It is prepared by fusion of MnO₂ with KOH and oxidant to give K₂MnO₄ (manganate), which is then oxidized to KMnO₄. In acidic medium, MnO₄⁻ is reduced to Mn²⁺ (+2). In neutral medium, it forms MnO₂ (+4). In alkaline medium, it forms MnO₄²⁻ (+6). KMnO₄ is used in titrations (self-indicator), Baeyer's test for unsaturation, and water purification.

MnO_4^- + 8H^+ + 5e^- arrow Mn^ 2+ + 4H_2O

Preparation of KMnO₄

KMnO₄ is prepared industrially by fusing MnO₂ (pyrolusite ore) with KOH in the presence of O₂ to form green K₂MnO₄ (potassium manganate), which is then oxidized electrolytically or by Cl₂/ozone to give purple KMnO₄.

The preparation of KMnO_4 is a two-step process: Step 1: Fusion — MnO_2 (pyrolusite ore) is fused with KOH in the presence of air/oxygen. This produces green potassium manganate ( K_2MnO_4 , Mn in +6 state).

2MnO_2 + 4KOH + O_2 arrow 2K_2MnO_4 + 2H_2O

Step 2: Oxidation — The green manganate solution is then oxidized to permanganate. This can be done: - Electrolytically (industrial method) - By passing Cl_2 or O_3 - Or by disproportionation in acidic/neutral solution

3MnO_4^ 2- + 4H^+ arrow 2MnO_4^- + MnO_2 + 2H_2O

• Raw material: MnO_2 (pyrolusite) — most common Mn ore
• K_2MnO_4 (manganate) is green — Mn is in +6 state
• KMnO_4 (permanganate) is purple — Mn is in +7 state
• The +6 → +7 oxidation requires a mild oxidizing agent (electrolysis, Cl_2 , O_3 )

Oxidizing Properties of KMnO₄

KMnO₄ is a versatile oxidizing agent. In acidic medium (H₂SO₄), Mn goes from +7 to +2 (Mn²⁺) — gaining 5 electrons. In neutral medium, Mn goes to +4 (MnO₂). In strongly alkaline medium, Mn goes to +6 (MnO₄²⁻). The acidic medium gives the strongest oxidizing power.

The reduction product of MnO_4^- depends entirely on the pH of the solution. This is one of the most-tested concepts in NEET chemistry.

Acidic: MnO_4^- + 8H^+ + 5e^- arrow Mn^ 2+ + 4H_2O

Neutral: MnO_4^- + 2H_2O + 3e^- arrow MnO_2 + 4OH^-

Alkaline: MnO_4^- + e^- arrow MnO_4^ 2-

• In acidic medium: oxidizes Fe^ 2+ to Fe^ 3+ , C_2O_4^ 2- to CO_2 , SO_3^ 2- to SO_4^ 2-
• The reaction with oxalic acid ( H_2C_2O_4 ) is autocatalytic — Mn^ 2+ produced acts as catalyst
• HCl is NOT used to acidify KMnO_4 — Cl^- would get oxidized to Cl_2 , consuming extra KMnO_4
• Always use H_2SO_4 to acidify KMnO_4 (sulfate ion is not oxidizable)

KMnO₄ Titrations

KMnO₄ acts as its own indicator in titrations because the purple color disappears when MnO₄⁻ is reduced to colorless Mn²⁺ in acidic medium. The endpoint is marked by the first persistent pale pink/purple color. The reaction with oxalic acid is the most important NEET titration.

KMnO_4 is a self-indicator in titrations. During the titration, each drop of KMnO_4 is immediately decolorized as it reacts. At the endpoint, there is no more reducing agent left to decolorize the permanganate — so the very next drop imparts a permanent pale pink/purple color to the solution.

2KMnO_4 + 5H_2C_2O_4 + 3H_2SO_4 arrow 2MnSO_4 + K_2SO_4 + 10CO_2 + 8H_2O

• No external indicator needed — KMnO_4 is purple, Mn^ 2+ is colorless
• Endpoint: first persistent pale pink color that doesn't fade within 30 seconds
• The reaction is slow initially (no catalyst). But once Mn^ 2+ is produced, it catalyzes the reaction — this is called autocatalysis
• The titration is done at 60–70°C (warm) to ensure fast reaction rate
• n -factor of KMnO_4 in acid = 5 (gains 5 electrons); n -factor of H_2C_2O_4 = 2 (loses 2 electrons per C)

Quick recall

In the KMnO_4 vs oxalic acid titration, what is the molar ratio?

2 mol KMnO_4 : 5 mol H_2C_2O_4 . Equivalence: KMnO_4 gains 5e^- per molecule (Mn: +7 → +2), oxalic acid loses 2e^- per molecule (2C: +3 → +4 each). By equivalence: 2 5 = 5 2 = 10 electrons transferred.

Important Compounds: K₂Cr₂O₇

Potassium dichromate (K₂Cr₂O₇) is an orange crystalline solid where Cr is in the +6 oxidation state. Prepared from chromite ore (FeCr₂O₄) through fusion with Na₂CO₃, oxidation, and acidification. CrO₄²⁻ (chromate, yellow) and Cr₂O₇²⁻ (dichromate, orange) are pH-interconvertible. In acidic medium, Cr₂O₇²⁻ is reduced to Cr³⁺ (green), gaining 6 electrons per formula unit (3 per Cr). Used in chromyl chloride test for Cl⁻ detection, leather tanning, and as a primary standard in titrations.

Cr_2O_7^ 2- + 14H^+ + 6e^- arrow 2Cr^ 3+ + 7H_2O

Preparation of K₂Cr₂O₇

K₂Cr₂O₇ is prepared from chromite ore (FeCr₂O₄) by fusion with Na₂CO₃ in the presence of air, producing sodium chromate (Na₂CrO₄, yellow). Acidification converts chromate to dichromate. Treatment with KCl gives K₂Cr₂O₇.

The preparation of K_2Cr_2O_7 starts from chromite ore ( FeCr_2O_4 , also written as FeO Cr_2O_3 ). The process has three key steps.

Step 1: Fusion with Na_2CO_3 + air → produces sodium chromate (yellow)

4FeCr_2O_4 + 8Na_2CO_3 + 7O_2 arrow 8Na_2CrO_4 + 2Fe_2O_3 + 8CO_2

Step 2: Acidification → converts yellow chromate to orange dichromate

2Na_2CrO_4 + H_2SO_4 arrow Na_2Cr_2O_7 + Na_2SO_4 + H_2O

Step 3: Conversion to K salt → Na_2Cr_2O_7 is treated with KCl

Na_2Cr_2O_7 + 2KCl arrow K_2Cr_2O_7 + 2NaCl

Chromate-Dichromate Equilibrium

CrO₄²⁻ (chromate, yellow, tetrahedral) and Cr₂O₇²⁻ (dichromate, orange) exist in a pH-dependent equilibrium. Adding acid shifts to dichromate (orange); adding base shifts to chromate (yellow). Both have Cr in +6 state.

Chromate ( CrO_4^ 2- , yellow) and dichromate ( Cr_2O_7^ 2- , orange) are interconvertible depending on pH. This is NOT a redox reaction — Cr stays at +6 in both forms.

2CrO_4^ 2- _ (yellow) + 2H^+ leftharpoons Cr_2O_7^ 2- _ (orange) + H_2O

• Add acid ( H^+ ): Equilibrium shifts right → orange dichromate forms
• Add base ( OH^- ): Equilibrium shifts left → yellow chromate forms
• CrO_4^ 2- is tetrahedral (Cr at center, 4 O around it)
• Cr_2O_7^ 2- has two tetrahedra sharing one corner O (Cr–O–Cr bridge, angle ~126°)
• Oxidation state of Cr = +6 in BOTH species — this is purely an acid-base equilibrium

Oxidizing Properties & Chromyl Chloride Test

In acidic medium, Cr₂O₇²⁻ is reduced to Cr³⁺ (green), gaining 6e⁻ per formula unit. It oxidizes Fe²⁺, I⁻, and S²⁻. The chromyl chloride test detects Cl⁻ by forming red vapors of CrO₂Cl₂ when heated with K₂Cr₂O₇ and conc. H₂SO₄.

In acidic solution, Cr_2O_7^ 2- is a strong oxidizing agent. Each Cr atom goes from +6 to +3, gaining 3 electrons. Since there are 2 Cr atoms per formula unit, the total gain is 6 electrons.

Cr_2O_7^ 2- + 14H^+ + 6e^- arrow 2Cr^ 3+ _ (green) + 7H_2O

• Oxidizes Fe^ 2+ to Fe^ 3+ : Cr_2O_7^ 2- + 6Fe^ 2+ + 14H^+ arrow 2Cr^ 3+ + 6Fe^ 3+ + 7H_2O (orange → green)
• Oxidizes I^- to I_2 : Cr_2O_7^ 2- + 6I^- + 14H^+ arrow 2Cr^ 3+ + 3I_2 + 7H_2O
• Oxidizes Sn^ 2+ to Sn^ 4+
• n -factor of K_2Cr_2O_7 = 6 (2 Cr atoms, each gaining 3 electrons)
• Solution turns from orange (dichromate) to green ( Cr^ 3+ ) on reduction

Chromyl Chloride Test — a qualitative test for detecting chloride ions ( Cl^- ): When a chloride salt is heated with K_2Cr_2O_7 and concentrated H_2SO_4 , deep red vapors of chromyl chloride ( CrO_2Cl_2 ) are evolved.

K_2Cr_2O_7 + 4NaCl + 6H_2SO_4 arrow 2CrO_2Cl_2 _ (red) + 2KHSO_4 + 4NaHSO_4 + 3H_2O

f-Block: Lanthanoids

The lanthanoids (La to Lu, Z=57 to 71) are the first inner transition series where the 4f orbitals are progressively filled. Their general configuration is [Xe]4f¹⁻¹⁴5d⁰⁻¹6s². The dominant oxidation state is +3 (due to removal of 5d¹6s²), with Ce⁴⁺ (noble gas [Xe] core) and Eu²⁺/Yb²⁺ (half-filled/fully-filled 4f) as exceptions. Lanthanoid contraction — the steady decrease in ionic radii from La³⁺ to Lu³⁺ — is caused by poor shielding of 4f electrons and has far-reaching consequences for 5d element chemistry.

General config: [Xe] 4f^ 1-14 5d^ 0-1 6s^2

Electronic Configuration & Oxidation States

Lanthanoids have the general configuration [Xe]4f¹⁻¹⁴5d⁰⁻¹6s². The +3 state is universally the most stable because removing 5d¹6s² electrons is energetically straightforward. Anomalous states: Ce⁴⁺ (f⁰), Eu²⁺ (f⁷), Yb²⁺ (f¹⁴) are stabilized by empty/half-filled/fully-filled f-subshells.

The 14 lanthanoids from Ce ( Z=58 ) to Lu ( Z=71 ) (with La as the prototype) involve the progressive filling of 4f orbitals. The general electronic configuration is: [Xe] 4f^ 1-14 5d^ 0-1 6s^2

• Most lanthanoids have 5d^0 — the electron goes into 4f directly
• Exceptions with 5d^1 : La ( 4f^0 5d^1 ), Ce ( 4f^1 5d^1 ), Gd ( 4f^7 5d^1 ), Lu ( 4f^ 14 5d^1 )
• Gd has 4f^7 5d^1 instead of 4f^8 5d^0 — half-filled f^7 stability pulls one electron to 4f
• The +3 state is universally stable — formed by removing 6s^2 + 5d^1 (or one 4f electron if 5d^0 )

Anomalous Oxidation States: - Ce⁴⁺ ( 4f^0 ) — achieves noble gas core [Xe] , extra stable. CeO₂ is a powerful oxidizing agent. - Eu²⁺ ( 4f^7 ) — half-filled f -subshell, extra stable. Eu²⁺ is a strong reducing agent (wants to stay at +2). - Yb²⁺ ( 4f^ 14 ) — fully-filled f -subshell, extra stable.

Lanthanoid Contraction

Lanthanoid contraction is the steady decrease in atomic and ionic radii from La³⁺ (103.2 pm) to Lu³⁺ (86.1 pm). Caused by poor shielding of 4f electrons, leading to a gradual increase in effective nuclear charge. This makes 4d and 5d elements in the same group nearly identical in size.

As we move from La ( Z=57 ) to Lu ( Z=71 ), each element has one more proton AND one more 4f electron. But 4f electrons are terrible at shielding — their diffuse, inner orbital shape means they don't effectively block the nuclear charge from outer electrons.

Z_ eff (across lanthanoids) atomic/ionic radius

• Ionic radii decrease: La^ 3+ (103.2 pm) → Lu^ 3+ (86.1 pm) — a total drop of ~17 pm across 14 elements
• The decrease is steady but small — about 1–1.5 pm per element
• Cause: poor shielding by 4f electrons → effective nuclear charge increases gradually
• Shielding ability order: s > p > d > f — f -electrons are the worst shielders

Major Consequences of Lanthanoid Contraction: 1. 4d and 5d elements in the same group have nearly equal radii — e.g., Zr (160 pm) ≈ Hf (159 pm), Nb ≈ Ta, Mo ≈ W 2. Similar chemical properties of 4d and 5d pairs — making their separation extremely difficult 3. Basicity of lanthanoid hydroxides decreases from La(OH)₃ (most basic) to Lu(OH)₃ (least basic) 4. Difficulty in separation of lanthanoids from each other — similar radii mean similar properties

Quick recall

What is the order of basicity among La(OH)₃, Nd(OH)₃, and Lu(OH)₃?

La(OH)_3 > Nd(OH)_3 > Lu(OH)_3 . Basicity decreases as ionic radius decreases (lanthanoid contraction). Smaller Ln^ 3+ holds OH^- more tightly → harder to release OH^- → less basic.

General Properties of Lanthanoids

Lanthanoids are silvery-white, soft metals with high melting points. They are highly electropositive and react with water (slowly cold, rapidly hot). They tarnish in air forming oxide layers. Their compounds are generally ionic and paramagnetic (except La³⁺ and Lu³⁺ which are diamagnetic).

Lanthanoids are silvery-white, soft, malleable metals that tarnish rapidly in air. They are highly reactive — more electropositive than most transition metals. Their chemistry is dominated by the +3 oxidation state and ionic bonding.

• All are highly electropositive — react with water (slowly cold, vigorously hot): 2Ln + 6H_2O arrow 2Ln(OH)_3 + 3H_2
• Burn in oxygen to form Ln_2O_3 — used in lighter flints (mischmetal: Ce + La + other Ln)
• Form Ln³⁺ ions with large ionic radii → compounds are predominantly ionic
• Most Ln³⁺ ions are paramagnetic (unpaired f -electrons), except La^ 3+ ( f^0 ) and Lu^ 3+ ( f^ 14 ) — diamagnetic
• Color: Many Ln³⁺ ions are colored due to f-f transitions (pale colors, much weaker than d-d transitions)

f-Block: Actinoids

The actinoids (Ac to Lr, Z=89 to 103) are the second inner transition series with progressive filling of 5f orbitals. General configuration: [Rn]5f¹⁻¹⁴6d⁰⁻¹7s². They show a much wider range of oxidation states than lanthanoids (+2 to +7, with +3 to +6 common) because 5f, 6d, and 7s energy levels are very close. All actinoids are radioactive; only Th, Pa, and U occur naturally. They show greater tendency for complex formation than lanthanoids due to higher charge and smaller size in higher oxidation states.

General config: [Rn] 5f^ 1-14 6d^ 0-1 7s^2

Electronic Configuration & Oxidation States

Actinoids have the general configuration [Rn]5f¹⁻¹⁴6d⁰⁻¹7s². Unlike lanthanoids, they show a wide range of oxidation states because the energy gap between 5f, 6d, and 7s orbitals is very small. U shows +3 to +6, Pu shows +3 to +7. The early actinoids (Th–Am) show the greatest variety.

The actinoids (Ac, Z=89 to Lr, Z=103 ) involve filling of the 5f orbitals. Their electronic configurations are more irregular than lanthanoids because the energy difference between 5f, 6d, and 7s orbitals is very small.

• General configuration: [Rn] 5f^ 1-14 6d^ 0-1 7s^2
• Configurations are irregular — e.g., Th ( 5f^0 6d^2 7s^2 ), Pa ( 5f^2 6d^1 7s^2 ), U ( 5f^3 6d^1 7s^2 )
• Oxidation states: Much more variable than lanthanoids — +3 to +7 for early actinoids
• +3 is common for ALL actinoids (like lanthanoids), but +4, +5, +6, +7 also exist for early members
• U: +3, +4, +5, +6 ( UO_2^ 2+ , uranyl ion, most stable for U)
• Pu: +3, +4, +5, +6, +7 (widest range)

Radioactivity & Natural Occurrence

Every actinoid is radioactive. Only Th (Z=90), Pa (Z=91), and U (Z=92) occur naturally in significant amounts. Elements beyond U (transuranium elements: Np, Pu, Am...) are all synthetic, produced in nuclear reactors or particle accelerators.

A critical difference between lanthanoids and actinoids: all actinoids are radioactive. Their nuclei are unstable due to the large number of protons. Only three actinoids occur naturally in significant quantities.

• Natural actinoids: Th ( Z=90 , t_ 1/2 = 1.4 10^ 10 years), Pa ( Z=91 ), U ( Z=92 , t_ 1/2 = 4.5 10^9 years)
• Transuranium elements ( Z > 92 ): Np, Pu, Am, Cm, Bk, Cf... all are synthetic (man-made)
• Produced by nuclear reactions: neutron bombardment, particle accelerator collisions
• Radioactivity makes actinoid chemistry difficult to study — handling requires special facilities
• Pm ( Z=61 ) is the only radioactive lanthanoid (all others are stable)

Quick recall

How many naturally occurring actinoids exist, and which are they?

Three: Thorium (Th, Z=90 ), Protactinium (Pa, Z=91 ), and Uranium (U, Z=92 ). All elements beyond uranium ( Z > 92 ) are synthetic — called transuranium elements.

Actinoid Contraction & Comparison with Lanthanoids

Like lanthanoid contraction, actinoid contraction is the steady decrease in ionic radii across the actinoid series due to poor shielding by 5f electrons. The contraction is slightly greater than lanthanoid contraction because 5f electrons shield even more poorly than 4f electrons.

Just as the lanthanoid series shows a steady decrease in ionic radii (lanthanoid contraction), the actinoid series shows a similar actinoid contraction. The cause is identical — poor shielding by f -electrons leading to increasing Z_ eff .

• Actinoid contraction is the steady decrease in ionic radii from Ac³⁺ to Lr³⁺
• 5f electrons are even poorer at shielding than 4f electrons → contraction is slightly greater
• Consequence: 5d and 6d elements in the same group DON'T have as similar sizes as 4d/5d pairs
• The combined effect of lanthanoid + actinoid contraction makes elements beyond Period 6 behave uniquely

Key Differences: Lanthanoids vs Actinoids Summary - Oxidation states: Ln → mainly +3 | An → +3 to +7 (wider range) - Radioactivity: Ln → stable (except Pm) | An → ALL radioactive - Complex formation: Ln → less | An → more (higher charge density in high O.S.) - Magnetic properties: Ln → spin-orbit coupling important | An → even more complex, orbital contribution significant - Compounds: Ln → mostly ionic | An → both ionic and covalent (especially UF₆)