d-Block & f-Block Elements – Chemistry Study Notes

General Characteristics and Electronic Configurations

As you move across the periodic table, the d-block spans groups 3 to 12. It is divided into three main transition series corresponding to the filling of the 3d, 4d, and 5d orbitals, alongside an incomplete 4f/5f series for the f-block.

The general outer electronic configuration for transition elements is represented as $(n-1)d^{1-10}ns^{1-2}$. Understanding these configurations is vital because they dictate the chemical and physical behaviors of these metals, such as variable valency and complex formation.

There are distinct anomalies in the electronic configuration of transition elements, particularly due to the extra stability associated with half-filled ($d^5$) and completely filled ($d^{10}$) subshells. For example, chromium ($\text{Cr}$) and copper ($\text{Cu}$) exhibit unexpected configurations like $3d^5 4s^1$ and $3d^{10} 4s^1$ respectively, deviating from the standard Aufbau filling sequence.

Examiners frequently test these exceptions, making it critical to memorize the ground-state configurations of the first transition series from Scandium to Zinc.

• Scandium ($\text{Sc}, Z=21$): $[Ar] 3d^1 4s^2$
• Chromium ($\text{Cr}, Z=24$): $[Ar] 3d^5 4s^1$ (Half-filled stability)
• Copper ($\text{Cu}, Z=29$): $[Ar] 3d^{10} 4s^1$ (Fully-filled stability)
• Zinc ($\text{Zn}, Z=30$): $[Ar] 3d^{10} 4s^2$ (Strictly speaking, zinc and its group homologs do not show typical transition metal properties due to filled $d$ orbitals).

Metallic Radii, Oxidation States, and Electrode Potentials

In any transition series, the atomic and metallic radii decrease gradually from left to right as nuclear charge increases. However, the decrease becomes very small across the middle of the series. This happens because the added $d$-electrons provide poor screening or shielding effect for the outer $s$-electrons against the increasing nuclear charge.

Furthermore, the radii of the second and third transition series elements are very similar due to the famous lanthanoid contraction—the cumulative effect of poor shielding by $4f$ electrons in lanthanides.

Transition elements exhibit a wide variety of oxidation states because the energy gap between the $(n-1)d$ and $ns$ orbitals is very small. This allows electrons from both subshells to participate in bonding. Manganese ($\text{Mn}$), for instance, exhibits oxidation states ranging from $+2$ to $+7$ in compounds like $\text{KMnO}_4$.

The highest oxidation states are typically observed in oxides and fluorides (e.g., $\text{OsO}_4$ or $\text{RuO}_4$) due to the high electronegativity and small size of oxygen and fluorine atoms.

“Standard electrode potentials ($E^\circ$) for transition metal systems ($M^{2+}/M$) do not follow a smooth regular trend. This irregularity is a complex consequence of the interplay of enthalpy of sublimation, ionization enthalpy, and hydration enthalpy.”

Colored Ions and Magnetic Properties

One of the most visually striking properties of d-block transition metal ions is their ability to form colored ions in solid and solution states. When ligands approach a transition metal ion, the degenerate $d$-orbitals split into different energy levels via crystal field splitting. Electrons can then absorb specific wavelengths of visible light to undergo $d\text{-}d$ transitions, transmitting the complementary color to our eyes.

Ions with completely empty ($d^0$) or completely filled ($d^{10}$) configurations, such as $\text{Sc}^{3+}$ or $\text{Zn}^{2+}$, are colorless because $d\text{-}d$ transitions are impossible.

Most transition metal compounds are paramagnetic, meaning they are attracted by a magnetic field due to the presence of one or more unpaired electrons. The magnetic moment of these species is calculated using the spin-only formula:

$$\mu = \sqrt{n(n+2)} \text{ BM}$$

Here, $n$ represents the number of unpaired electrons, and $\text{BM}$ stands for Bohr Magneton. As the number of unpaired electrons increases from $1$ to $5$ (in a $d^5$ system), the magnetic moment increases correspondingly, providing a direct experimental method to determine the oxidation state and coordination geometry of transition complexes.

Inner-Transition Elements: Lanthanoids and Actinoids

The f-block elements consist of the lanthanoids ($4f$-series) and actinoids ($5f$-series). Lanthanoids span from Cerium ($\text{Ce}, Z=58$) to Lutetium ($\text{Lu}, Z=71$). The most characteristic oxidation state for lanthanoids is $+3$, although some elements also exhibit $+2$ or $+4$ states to achieve empty, half-filled, or fully-filled $f$-orbital configurations (e.g., $\text{Ce}^{4+}$ or $\text{Eu}^{2+}$).

The steady decrease in atomic and ionic radii across the lanthanoid series is known as the lanthanoid contraction.

The actinoids span from Thorium ($\text{Th}, Z=90$) to Lawrencium ($\text{Lr}, Z=103$). Unlike lanthanoids, actinoids exhibit a much broader range of oxidation states (such as $+3, +4, +5, +6,$ and $+7$ for uranium, neptunium, and plutonium) because the $5f$, $6d$, and $7s$ orbitals have comparable energy levels.

Furthermore, all actinoids are radioactive, and elements beyond uranium are synthetic or transuranic elements produced via nuclear reactions.

• Lanthanoids are predominantly non-radioactive (except Promethium, $\text{Pm}$).
• Actinoids are intensely radioactive and often display actinide contraction, which is even more pronounced than lanthanoid contraction due to poorer shielding by $5f$ electrons.
• Compounds of actinoids are frequently colored, mirroring d-block behavior via $f\text{-}f$ transitions.

Important Facts / Formulas

Key Points to Remember

• Zinc, cadmium, and mercury ($\text{Zn}, \text{Cd}, \text{Hg}$) have $d^{10}$ configurations and are typically excluded from true transition metal definitions despite being in the d-block.
• Chromium ($\text{Cr}$) and Copper ($\text{Cu}$) show anomalous electronic configurations to achieve extra stability from symmetrical half-filled and full-filled shells.
• Transition metals and their compounds act as excellent catalysts due to variable oxidation states and their capacity to form interstitial complexes.
• Alloys are readily formed by transition metals because atoms of similar radii can easily substitute for one another in crystal lattices (e.g., brass, bronze, steel).
• $\text{KMnO}_4$ and $\text{K}_2\text{Cr}_2\text{O}_7$ are powerful oxidizing agents in acidic media involving manganese and chromium transitions.
• The color of transition metal complexes is fundamentally caused by $d\text{-}d$ electronic transitions facilitated by crystal field splitting.
• Actinoids show greater variability in oxidation states than lanthanoids because $5f$ electrons are less tightly bound than $4f$ electrons.

Quick Revision Summary

• General Config: $(n-1)d^{1-10}ns^{1-2}$ governs the chemistry of d-block transition series.
• Exceptions: $\text{Cr}$ and $\text{Cu}$ anomalies stem from the extra stability of half-filled and full-filled $d$-orbitals.
• Variable Valency: Small energy difference between $ns$ and $(n-1)d$ orbitals enables multiple oxidation states.
• Magnetic Moment: Calculated via $\mu = \sqrt{n(n+2)}$ where $n$ is the number of unpaired electrons.
• Color: Arises from $d\text{-}d$ electronic transitions when ligand fields split $d$-orbital degeneracies.
• Lanthanoid Contraction: Poor shielding of $4f$ electrons results in a progressive decrease in ionic radii.
• Catalysis: Transition metals possess variable valencies and large surface areas, making them premier catalysts.
• Actinoids: Radioactive inner-transition elements exhibiting higher oxidation states due to accessible $5f$ and $6d$ levels.