Chapter 4: The d- and f-Block Elements

Dear Class 12 Student! Welcome to the fascinating world of transition metals. Unlike the predictable s- and p-blocks, the d-block is full of vibrant colors, complex magnetic properties, and catalytic magic due to their partially filled d-orbitals. This chapter is highly conceptual. For the Boards and JEE, mastering the "Why" behind the trends (like Lanthanoid Contraction) and the chemistry of $KMnO_4$ and $K_2Cr_2O_7$ is absolutely essential. Let's delve into the transition series!

1. Introduction and Position in the Periodic Table

The d-block elements occupy the large middle section of the periodic table, flanking between the s-block and p-block (Groups 3 to 12). They are characterized by the gradual filling of the inner $(n-1)d$ orbitals.

The Four Series

• 3d series: Scandium (Sc, 21) to Zinc (Zn, 30) - Most important for syllabus.
• 4d series: Yttrium (Y, 39) to Cadmium (Cd, 48).
• 5d series: Lanthanum (La, 57), Hafnium (Hf, 72) to Mercury (Hg, 80).
• 6d series: Actinium (Ac, 89), Rutherfordium (Rf, 104) to Copernicium (Cn, 112).

2. Electronic Configuration of d-Block Elements

The general outer electronic configuration is $(n-1)d^{1-10} ns^{1-2}$. The $(n-1)$ indicates that the d-orbitals being filled belong to the penultimate (second to last) shell.

3. General Properties of Transition Elements (Trends and Reasons)

NCERT emphasizes the reasoning behind these trends heavily. Master the "Why".

1. Melting and Boiling Points

• Property: Transition metals have very high melting and boiling points (except Group 12).
• Reason (Metallic Bonding): High MPs are due to the involvement of a greater number of electrons from both (n-1)d and ns orbitals in interatomic metallic bonding. The unpaired d-electrons overlap to form covalent-like bonds within the metal lattice.
• The Trend: Melting points rise to a maximum at $d^5$ (middle) because of the maximum number of unpaired d-electrons.
• The Dip (Exception): Manganese (Mn) and Technetium (Tc) show an anomalous dip.
  Reason: Their $d^5$ configuration is so exceptionally stable (due to high exchange energy) that the electrons are held tightly by the nucleus and do not participate effectively in delocalized metallic bonding.

3. Ionization Enthalpies

• General Trend: Generally increases from left to right due to increasing effective nuclear charge ($Z_{eff}$).
• Reason for Irregularity: The trend is irregular because the addition of d-electrons alters the shielding effect. Removing an electron from the $ns$ orbital is influenced by how well the $(n-1)d$ electrons shield it.
• Board Favorite: Why is the third ionization enthalpy of Manganese (Mn) exceptionally high compared to Iron (Fe)?
  Answer: $Mn^{2+}$ has a stable half-filled $[Ar] 3d^5$ configuration. Removing a third electron disrupts this extreme stability (loss of exchange energy), requiring a massive amount of energy. Iron ($Fe^{2+}$) is $3d^6$; removing an electron yields a stable $3d^5$ ($Fe^{3+}$), so its 3rd IE is comparatively lower.

4. Oxidation States

• Transition metals show a vast variety of oxidation states.
• Reason (Small Energy Gap): The energy difference between the $(n-1)d$ and $ns$ orbitals is very small. Therefore, both sets of electrons can be used for bond formation.
• Maximum Oxidation State: Occurs in the middle of the series, corresponding to the sum of all s and d electrons. (e.g., Manganese shows up to +7 in $KMnO_4$).

5. Standard Electrode Potentials ($E^\circ$)

The standard electrode potential ($E^\circ$) of a metal is NOT determined by a single factor. It is the result of a competition between three energy terms:

The trend in $M^{2+}/M$ standard electrode potentials is highly irregular. This is because $E^\circ$ depends on the sum of three enthalpies: Sublimation Enthalpy, Ionization Enthalpy, and Hydration Enthalpy.

4. Important Characteristics and Phenomena

1. Magnetic Properties

• Diamagnetic: Substances repelled by a magnetic field. All electrons are paired.
• Paramagnetic: Substances attracted by a magnetic field. Caused by the presence of unpaired electrons. Transition metals are mostly paramagnetic.

The magnetic moment is calculated using the "Spin-Only" Formula:

2. Formation of Coloured Ions

Most transition metal compounds are highly colored.
Reason: When an anion or ligand approaches the transition metal ion, the 5 degenerate d-orbitals split into two sets of different energy levels. When visible light falls on the ion, an unpaired d-electron absorbs a specific wavelength (energy) and jumps from the lower d-level to the higher d-level. This is called a d-d transition. The transmitted (unabsorbed) light gives the compound its complementary color.

Condition: There MUST be vacant spaces in the d-orbitals for the electron to jump into, and there MUST be an unpaired electron to jump. Ions with $d^0$ (e.g., $Sc^{3+}, Ti^{4+}$) or completely filled $d^{10}$ (e.g., $Zn^{2+}, Cu^+$) cannot undergo d-d transitions and are colorless/white.

3. Formation of Complex Compounds

Transition metals form a vast array of complex coordination compounds (e.g., $[Fe(CN)_6]^{3-}$). They do this because:
1. They have a comparatively smaller ionic size and high nuclear charge density.
2. They have vacant $(n-1)d$ orbitals of appropriate energy to accept lone pairs of electrons donated by ligands.

4. Catalytic Properties

Many transition metals and their compounds act as excellent catalysts (e.g., Iron in the Haber process for ammonia, $V_2O_5$ in the Contact process for sulfuric acid).
Reasons:
1. Their ability to show variable oxidation states allows them to form unstable intermediate compounds with reactants, creating a low-energy pathway.
2. Their solid lattice provides a large surface area with free valencies for reactant molecules to adsorb onto, bringing them closer together.

5. Formation of Interstitial Compounds

Transition metal crystal lattices have tiny empty spaces (interstices) between the large metal atoms. Very small non-metal atoms like Hydrogen (H), Carbon (C), or Nitrogen (N) get trapped in these spaces, forming non-stoichiometric compounds.
Examples: $TiC, Mn_4N, Fe_3H, VH_{0.56}, TiH_{1.7}$.
Properties: They are extraordinarily hard (some approach diamond's hardness), have extremely high melting points (higher than pure metal), retain metallic conductivity, and become chemically inert.

6. Alloy Formation

An alloy is a solid solution of two or more metals. Transition metals readily form alloys with each other (like Brass = Cu+Zn, Bronze = Cu+Sn).
Reason: Because their atomic radii are very similar (difference is less than 15%), atoms of one transition metal can easily replace atoms of another transition metal in the solid crystal lattice without severely distorting it.

5. Important Compounds of Transition Elements (High Priority)

A. Potassium Dichromate ($K_2Cr_2O_7$)

It is a bright orange crystalline solid, widely used as a primary standard and powerful oxidizing agent.

Chromate-Dichromate pH Equilibrium:
Chromate ($CrO_4^{2-}$, Yellow) and Dichromate ($Cr_2O_7^{2-}$, Orange) are interconvertible depending on the pH of the solution. The oxidation state of Chromium is +6 in both! No redox happens here.
In Acidic medium (Low pH): $2CrO_4^{2-} + 2H^+ \rightleftharpoons Cr_2O_7^{2-} + H_2O$ (Turns Orange)
In Basic medium (High pH): $Cr_2O_7^{2-} + 2OH^- \rightleftharpoons 2CrO_4^{2-} + H_2O$ (Turns Yellow)

B. Potassium Permanganate ($KMnO_4$)

A dark purple, almost black crystalline solid. It is an exceptionally powerful oxidizing agent.

6. The f-Block Elements: Lanthanoids

The f-block elements are the "Inner Transition Elements". The first series involves the filling of the 4f orbitals and consists of 14 elements from Cerium (Z=58) to Lutetium (Z=71), following Lanthanum.

• Electronic Configuration: $[Xe] 4f^{1-14} 5d^{0-1} 6s^2$.
• Mastery Table 5: Lanthanoids (Configurations and Radii)

  Element	Symbol	Configuration (Ln)	Radius Ln³⁺ (pm)
  Lanthanum	La	$5d^1 6s^2$	106
  Cerium	Ce	$4f^1 5d^1 6s^2$	103
  Europium	Eu	$4f^7 6s^2$	95
  Gadolinium	Gd	$4f^7 5d^1 6s^2$	94
  Lutetium	Lu	$4f^{14} 5d^1 6s^2$	86

• Oxidation States: The most common and stable oxidation state for ALL lanthanoids is +3. Occasionally, +2 and +4 states are observed.
  Reason: These anomalous states only occur when they lead to extra stable configurations like $f^0$ (empty, e.g., $Ce^{4+}$), $f^7$ (half-filled, e.g., $Eu^{2+}$), or $f^{14}$ (fully-filled, e.g., $Yb^{2+}$). Note that $Ce^{4+}$ is a strong analytical oxidizing agent because it desperately wants to gain an electron to return to the stable +3 state.

Lanthanoid Contraction

As previously discussed, this is the steady decrease in atomic and ionic radii of the lanthanoids with increasing atomic number.

• Consequences:
  1. Similarity in size of 4d and 5d transition series elements (e.g., Zr/Hf).
  2. Difficulty in separating lanthanoids from each other in chemistry because their radii (and thus chemical properties) are nearly identical.
  3. Decrease in Basic Strength of Hydroxides: As the size of the $Ln^{3+}$ ion decreases from $La^{3+}$ to $Lu^{3+}$, the covalent character of the $Ln-OH$ bond increases (Fajans' rule). Therefore, $La(OH)_3$ is the most basic, while $Lu(OH)_3$ is the least basic.

Chemical Reactivity

Earlier members of the series are quite reactive, similar to calcium. They react with water to release $H_2$ gas, burn in oxygen to form $Ln_2O_3$, and react with halogens to form $LnX_3$.
Mischmetall: An important alloy consisting of $\approx 95\%$ lanthanoid metals, $5\%$ iron, and traces of S, C, Ca, and Al. It is used to make bullets, shells, and the sparkling flints for lighters.

7. The f-Block Elements: Actinoids

The second inner transition series involves the filling of the 5f orbitals. It consists of 14 elements from Thorium (Z=90) to Lawrencium (Z=103).

• Electronic Configuration: $[Rn] 5f^{1-14} 6d^{0-1} 7s^2$.
• General Characteristics: ALL actinoids are radioactive. The earlier elements have long half-lives, but elements beyond Uranium are highly unstable synthetics with half-lives ranging from a day to 3 minutes (Lawrencium).
• Oxidation States: Unlike lanthanoids, actinoids show a very wide range of oxidation states (up to +7 for Np and Pu).
  Reason: The 5f, 6d, and 7s energy levels are of comparable energies. Therefore, electrons from all these subshells can be excited and participate in bonding.
• Actinoid Contraction: There is a steady decrease in size across the actinoid series. However, the contraction from element to element is greater than the lanthanoid contraction.
  Reason: The 5f electrons provide even poorer shielding from the nuclear charge than the 4f electrons.