Study Notes on d-Block and f-Block Elements

Position and Classification of d-Block and f-Block Elements

• The d-Block defined: The d-block of the periodic table consists of elements in groups 3-12. In these elements, the d orbitals are progressively filled in each of the four long periods.

• The f-Block defined: The f-block consists of elements in which the $4f$ and $5f$ orbitals are progressively filled. They are displayed in a separate panel at the bottom of the periodic table.

• Terminology:

  • Transition Metals: Often used for d-block elements.

  • Inner Transition Metals: Used for f-block elements.

• Transition Metal Series: There are four main series:

  • 3d series: Scandium ($Sc$) to Zinc ($Zn$).

  • 4d series: Yttrium ($Y$) to Cadmium ($Cd$).

  • 5d series: Lanthanum ($La$) and Hafnium ($Hf$) to Mercury ($Hg$).

  • 6d series: Actinium ($Ac$) and elements from rutherfordium ($Rf$) to Copernicium ($Cn$).

• Inner Transition Metal Series:

  • Lanthanoids: The $4f$ series from Cerium ($Ce$) to Lutetium ($Lu$).

  • Actinoids: The $5f$ series from Thorium ($Th$) to Lawrencium ($Lr$).

• Evolution of Definition: Originally, the name "transition" described properties transitional between s-block and p-block elements. The current IUPAC definition states transition metals are metals with an incomplete d subshell in their neutral atom or in their ions.

• The Case of Group 12 (Zn, Cd, Hg):

  • These elements have a full $d^{10}$ configuration in their ground state and common oxidation states.

  • Technically, they are not transition metals by the IUPAC definition.

  • However, they are studied with transition metals as they are the end members of the 3d, 4d, and 5d series.

Electronic Configurations of the d-Block Elements

• General Outer Electronic Configuration: $(n-1)d^{1-10} ns^{1-2}$.

  • $(n-1)$ represents the inner d orbitals (penultimate energy level) containing 1 to 10 electrons.

  • $ns$ represents the outermost orbital containing 1 or 2 electrons.

• Notable Exceptions:

  • Palladium ($Pd$): Electronic configuration is $4d^{10} 5s^0$.

  • Chromium ($Cr$): Exhibits $3d^5 4s^1$ instead of $3d^4 4s^2$ because the energy difference between 3d and 4s is small, and half-filled orbitals ($d^5$) are more stable.

  • Copper ($Cu$): Exhibits $3d^{10} 4s^1$ instead of $3d^9 4s^2$ because completely filled orbitals ($d^{10}$) are more stable.

• Configuration Table (1st Series):

  • Sc ($Z=21$): $3d^1 4s^2$

  • Ti ($Z=22$): $3d^2 4s^2$

  • V ($Z=23$): $3d^3 4s^2$

  • Cr ($Z=24$): $3d^5 4s^1$

  • Mn ($Z=25$): $3d^5 4s^2$

  • Fe ($Z=26$): $3d^6 4s^2$

  • Co ($Z=27$): $3d^7 4s^2$

  • Ni ($Z=28$): $3d^8 4s^2$

  • Cu ($Z=29$): $3d^{10} 4s^1$

  • Zn ($Z=30$): $3d^{10} 4s^2$

General Physical Properties

• Metallic Nature: Transition elements show typical metallic properties including high tensile strength, ductility, malleability, high thermal and electrical conductivity, and metallic lustre.

• Lattice Structures:

  • Many adopt hcp (hexagonal close packed), bcc (body centred cubic), or ccp (cubic close packed) structures.

  • Exceptions: Zn, Cd, Hg, and Mn display unique metallic structures at room temperature.

• Hardness and Volatility: Transition metals (except group 12) are very hard and have low volatility.

• Melting and Boiling Points: These are very high. High melting points are due to the involvement of $(n-1)d$ electrons alongside $ns$ electrons in interatomic metallic bonding.

  • Melting points rise to a maximum at approximately $d^5$ in any row.

  • Anomalously low values are observed for Mn and Tc.

• Enthalpy of Atomisation: Values are high. Maxima occur around the middle of each series, indicating that one unpaired electron per d orbital is ideal for strong interatomic interaction (metallic bonding).

  • Metals of the 4d and 5d series have higher enthalpies of atomisation than those of the 3d series, which explains why metal-metal bonding is more frequent in heavy transition metal compounds.

Trends in Atomic and Ionic Sizes

• Across a series: There is a progressive decrease in radius with increasing atomic number. This happens because the shielding effect of d electrons is poor, resulting in a higher net electrostatic attraction between the nucleus and outermost electrons.

• Down a group: Atomic size increases from the 3d series to the 4d series.

• Lanthanoid Contraction:

  • The radii of the 5d series (third series) are virtually identical to the corresponding members of the 4d series (second series).

  • Cause: Filling the $4f$ orbitals before the $5d$ series leads to a regular decrease in atomic radii because shielding by 4f electrons is even less effective than d electrons.

  • Result: Pairs like zirconium ($Zr$, $160\,pm$) and hafnium ($Hf$, $159\,pm$) have nearly identical physical and chemical properties.

• Density: Because atomic radius decreases while atomic mass increases, density increases significantly across the series (e.g., from $3.43\,g\,cm^{-3}$ for Sc to $8.9\,g\,cm^{-3}$ for Cu).

Ionisation Enthalpies

• General Trend: Ionisation enthalpy increases along each series from left to right due to increased nuclear charge.

• Comparison to Non-Transition Elements: Successive enthalpies for d-block elements do not increase as steeply. The variation along a transition series is much less than along a period of s- or p-block elements.

• Mechanics of Ionisation: When d-block elements form ions, $ns$ electrons are lost before $(n-1)d$ electrons.

• Exchange Energy Factors:

  • Exchange energy is proportional to the possible pairs of parallel spins in degenerate orbitals. It stabilizes specific configurations.

  • $Mn^+$ has lower ionisation enthalpy than $Cr^+$ because losing an electron from Manganese results in a stable $d^5$ configuration.

  • $Fe^{2+}$ has lower ionisation enthalpy than $Mn^{2+}$ because $Mn^{2+}$ is already at the highly stable $d^5$ state.

• Second Ionisation Enthalpy: Unusually high for Cr and Cu because removing the second electron breaks stable $d^5$ and $d^{10}$ configurations, respectively.

• Third Ionisation Enthalpy: Generally very high. Breaking the $d^5$ (Mn) and $d^{10}$ (Zn) configurations is extremely difficult.

Oxidation States

• Variability: Transition elements exhibit a great variety of oxidation states due to the participation of both $(n-1)d$ and $ns$ electrons. Oxidation states usually differ by units of one (e.g., $V^{II}, V^{III}, V^{IV}, V^V$).

• Series Trends:

  • Middle of Series: Greatest number of oxidation states (e.g., Mn shows +2 to +7).

  • Extreme Ends: Fewer oxidation states (Sc shows only +3; Zn shows only +2).

• Stability Trends:

  • Early series: Maximum state is often the most stable (e.g., $Ti^{IV}$).

  • Late series: Lower states are more common ($Fe^{II,III}, Cu^{I,II}$).

• Group Trends vs. p-Block: In the d-block, higher oxidation states are more stable in heavier members of a group (e.g., $Mo(VI)$ and $W(VI)$ are more stable than $Cr(VI)$). In the p-block, the lower oxidation state is favored by heavier members (inert pair effect).

• Low Oxidation States: Can be stabilized by ligands with $\pi$-acceptor character (e.g., Carbonyls: $Ni(CO)_4$, oxidation state is zero).

Chemical Reactivity and Electrode Potentials ($E^\ominus$)

• $E^\ominus(M^{2+}/M)$ Trends:

  • The general trend across the series is towards less negative values, meaning a decreasing tendency to form divalent cations.

  • Copper Exception: Cu has a unique positive $E^\ominus$ value ($+0.34\,V$). It cannot liberate $H_2$ from non-oxidising acids because high enthalpy of atomisation and high ionisation enthalpy are not balanced by hydration enthalpy.

  • Anomalous Negativity: Mn, Ni, and Zn values are more negative than expected. For Mn and Zn, this is due to stable $d^5$ and $d^{10}$ configurations. For Ni, it is due to a very high negative enthalpy of hydration.

• $E^\ominus(M^{3+}/M^{2+})$ Values:

  • Low value for Sc reflects the stability of $Sc^{3+}$ (noble gas configuration).

  • High value for Mn reflects the stability of $Mn^{2+}$ ($d^5$).

  • Low value for Fe reflects the stability of $Fe^{3+}$ ($d^5$).

  • $Mn^{3+}$ and $Co^{3+}$ are strong oxidising agents in aqueous solutions.

  • $Ti^{2+}, V^{2+},$ and $Cr^{2+}$ are strong reducing agents.

Chemical Properties of Compounds

• Magnetic Properties:

  • Paramagnetism: Results from unpaired electrons. Spin and orbital angular momenta contribute, though orbital contribution is quenched in the first series.

  • Spin-only formula: $\mu = \sqrt{n(n+2)}$, where $n$ is the number of unpaired electrons and $\mu$ is the magnetic moment in Bohr magnetons (BM).

• Formation of Coloured Ions: Color arises from d-d transitions. Light corresponding to specific frequencies is absorbed in the visible region, and the complementary color is observed.

  • $f^0$ (e.g., $Sc^{3+}, Ti^{4+}$) and $f^{10}$ (e.g., $Zn^{2+}, Cu^+$) configurations are colourless.

• Complex Formation: Transition metals form numerous complexes due to small ionic size, high ionic charge, and available d orbitals for bonding with ligands.

• Catalytic Property: Attributed to the ability to adopt multiple oxidation states and form complexes. Examples include $V_2O_5$ (Contact Process), Fe (Haber’s Process), and Ni (Hydrogenation).

• Interstitial Compounds: Formed when small atoms (H, C, N) are trapped in metal lattices.

  • Characteristics: Non-stoichiometric (e.g., $TiC, Mn_4N, Fe_3H$), higher melting points than pure metals, very hard, retain metallic conductivity, chemically inert.

• Alloy Formation: Easy to form because transition metal radii are similar (within 15%). Examples: Steels (Fe with Cr, Mn, V, etc.), Brass ($Cu-Zn$), Bronze ($Cu-Sn$).

Important Compounds: K2Cr2O7 and KMnO4

• Potassium Dichromate ($K_2Cr_2O_7$):

  • Preparation: Fusion of chromite ore ($FeCr_2O_4$) with $Na_2CO_3$ in air to get yellow sodium chromate. Acidification with $H_2SO_4$ yields orange sodium dichromate ($Na_2Cr_2O_7 \cdot 2H_2O$). Treatment with $KCl$ yields orange crystals of $K_2Cr_2O_7$.

  • Reversible equilibrium: Chromate ($CrO_4^{2-}$) and dichromate ($Cr_2O_7^{2-}$) interconvert based on pH ($2CrO_4^{2-} + 2H^+ \rightleftharpoons Cr_2O_7^{2-} + H_2O$).

  • Structure: Chromate is tetrahedral; dichromate is two tetrahedra sharing a corner ($Cr-O-Cr$ angle is $126^\circ$).

  • Oxidising Agent: Acidified dichromate converts $I^- \rightarrow I_2$, $Sn^{2+} \rightarrow Sn^{4+}$, $H_2S \rightarrow S$, and $Fe^{2+} \rightarrow Fe^{3+}$.

• Potassium Permanganate ($KMnO_4$):

  • Preparation: Fusion of $MnO_2$ with hydroxide and an oxidiser ($KNO_3$) yields green potassium manganate ($K_2MnO_4$), which disproportionates in neutral/acidic solution to purple permanganate.

  • Structure: Both manganate ($MnO_4^{2-}$) and permanganate ($MnO_4^-$) are tetrahedral.

  • Magnetic nature: Manganate is paramagnetic (one unpaired electron); permanganate is diamagnetic.

  • Oxidising reactions (Acidic): Oxidises oxalate to $CO_2$, $Fe^{2+}$ to $Fe^{3+}$, $NO_2^-$ to $NO_3^-$, and $I^-$ to $I_2$.

  • In Basic solution: Oxidises iodide to iodate ($IO_3^-$) and thiosulphate ($S_2O_3^{2-}$) to sulphate ($SO_4^{2-}$).

The Inner Transition Elements (f-Block)

• Lanthanoids ($4f$ series):

  • Electronic Configuration: Most stable state is Tripositive ions ($Ln^{3+}$), generally of the form $4f^n$.

  • Atomic Size: Shows regular decrease (Lanthanoid Contraction). Shielding of one $4f$ electron by another is very poor.

  • Oxidation States: Dominant state is +3. Occasionally +2 (e.g., $Eu^{2+}, Yb^{2+}$) and +4 (e.g., $Ce^{IV}$) occur due to stability of empty, half-filled, or filled f-subshells.

  • Reactivity: Earlier members react like Calcium; later members react like Aluminium. They form oxides ($M_2O_3$), hydroxides ($M(OH)_3$), and liberate $H_2$ from acids.

• Actinoids ($5f$ series):

  • Radiochemistry: All are radioactive. Earlier members have long half-lives; later ones are very short-lived.

  • Comparison to Lanthanoids: The actinoid contraction is greater than the lanthanoid contraction due to poorer shielding by $5f$ electrons.

  • Oxidation States: Greater range of states (+2 to +7) because $5f$, $6d$, and $7s$ levels have comparable energies.