Chapter Notes

Here are the comprehensive notes for the chapter The D- and F-Block Elements.

Introduction to d- and f-Block Elements

The periodic table is organized into blocks based on the filling of electron orbitals. The d-block and f-block elements are crucial groups with unique properties.

• The d-Block Elements: These are the elements found in Groups 3-12 in the middle of the periodic table. In these elements, the last electron enters the d orbitals of the penultimate (second to last) energy shell. They are also known as transition elements.

  • There are four series of transition metals:
    1. 3d series: Scandium (Sc) to Zinc (Zn)
    2. 4d series: Yttrium (Y) to Cadmium (Cd)
    3. 5d series: Lanthanum (La), then Hafnium (Hf) to Mercury (Hg)
    4. 6d series: Actinium (Ac), then Rutherfordium (Rf) to Copernicium (Cn)

• The f-Block Elements: These elements are placed in a separate panel at the bottom of the periodic table. In these elements, the last electron enters the f orbitals. They are also known as inner transition elements.

  • There are two series of inner transition elements:
    1. Lanthanoids (4f series): Cerium (Ce) to Lutetium (Lu)
    2. Actinoids (5f series): Thorium (Th) to Lawrencium (Lr)

What Makes an Element a "Transition Element"?

Originally, the name came from their properties being "transitional" between the highly reactive s-block metals and the non-metallic p-block elements.

According to the modern IUPAC definition, a transition metal is an element that has an incompletely filled d subshell either in its neutral atom or in any of its common ions.

THE TRANSITION ELEMENTS (d-BLOCK)

Electronic Configurations of the d-Block Elements

The general outer electronic configuration for d-block elements is $(n-1)d^{1-10} ns^{1-2}$.

• $(n-1)d$ refers to the inner d orbitals.
• $ns$ refers to the outermost s orbital.

There are exceptions to this general rule due to the very small energy difference between the $(n-1)d$ and $ns$ orbitals. The stability of half-filled ($d^5$) and completely filled ($d^{10}$) orbitals is a major factor.

Key Examples from the 3d Series:

• Chromium (Cr, Z=24): Expected configuration is $3d^4 4s^2$. The actual configuration is $3d^5 4s^1$. This is because a half-filled $3d$ subshell ($d^5$) is more stable.
• Copper (Cu, Z=29): Expected configuration is $3d^9 4s^2$. The actual configuration is $3d^{10} 4s^1$. This is because a completely filled $3d$ subshell ($d^{10}$) is more stable.

Solution

Scandium has the ground state configuration $3d^1 4s^2$. It has an incompletely filled $3d$ orbital, so it is a transition element. Zinc has the ground state configuration $3d^{10} 4s^2$. In its common oxidation state ($Zn^{2+}$), the configuration is $3d^{10}$. In both its ground state and its common ion, the d orbital is completely filled. Therefore, it is not regarded as a transition element.

General Properties of the Transition Elements (d-Block)

Physical Properties

• Metallic Character: They exhibit typical metallic properties like high tensile strength, ductility, malleability, metallic lustre, and high thermal and electrical conductivity.
• Hardness and Volatility: They are very hard and have low volatility (except for Zn, Cd, and Hg).
• Melting and Boiling Points: They have very high melting and boiling points. This is due to the involvement of both ns and inner $(n-1)d$ electrons in forming strong interatomic metallic bonds. The melting points generally peak around the middle of the series (at $d^5$), where the number of unpaired electrons available for bonding is maximum.
• Enthalpy of Atomisation: This is the energy required to break the metallic bonds and turn the solid metal into individual gaseous atoms. Transition metals have high enthalpies of atomisation due to strong metallic bonding.
  • The trend is similar to melting points, with a maximum in the middle of the series.
  • Elements of the second (4d) and third (5d) series have higher enthalpies of atomisation than the first (3d) series, leading to more frequent metal-metal bonding in their compounds.

Variation in Atomic and Ionic Sizes

• Trend Across a Series: As we move from left to right across a transition series, the atomic and ionic radii decrease. This happens because with each step, a new electron is added to a d orbital while the nuclear charge increases by one. The shielding effect of a d electron is poor, so the increased nuclear charge pulls the outermost electrons more strongly, causing the radius to shrink. The variation, however, is small.
• Trend Down a Group (Lanthanoid Contraction):
  • The atomic radius increases from the 3d to the 4d series, as expected.
  • However, the atomic radii of the 5d series elements are virtually the same as their corresponding members in the 4d series (e.g., Zr (160 pm) and Hf (159 pm)).
  • This phenomenon is called the Lanthanoid Contraction.

Lanthanoid Contraction is the steady decrease in the size of the lanthanoid atoms and ions with increasing atomic number.

• Cause: Before the 5d series begins, the 4f orbitals are filled. The 4f electrons provide very poor shielding of the outer electrons from the increasing nuclear charge. This results in a significant contraction in size.
• Consequence: This contraction cancels out the expected increase in size from the 4d to the 5d series. As a result, the 4d and 5d series elements in the same group have very similar sizes and, consequently, very similar chemical properties, making their separation difficult.

Ionisation Enthalpies

• There is a general increase in ionisation enthalpy from left to right across a series due to increasing nuclear charge.
• The increase is not as steep as in main group elements.
• Irregularities in the trend occur due to the extra stability of $d^5$ (half-filled) and $d^{10}$ (fully-filled) configurations. For example, the third ionisation enthalpy of Manganese ($Mn^{2+} \rightarrow Mn^{3+}$) is very high because it involves removing an electron from a stable $d^5$ configuration. Conversely, the third ionisation enthalpy of Iron ($Fe^{2+} \rightarrow Fe^{3+}$) is lower because the resulting $Fe^{3+}$ ion has a stable $d^5$ configuration.

Oxidation States

One of the most characteristic properties of transition metals is their ability to exhibit a great variety of oxidation states in their compounds.

• Reason: The energy difference between the $(n-1)d$ and $ns$ orbitals is very small. Therefore, electrons from both subshells can participate in bonding.
• Trends:
  • The number of oxidation states increases up to the middle of the series and then decreases. Manganese (Mn) shows the maximum number of oxidation states in the 3d series, from +2 to +7.
  • The lower oxidation states (like +2, +3) are usually found in ionic compounds.
  • The higher oxidation states are found in covalent compounds with highly electronegative elements like oxygen and fluorine (e.g., $CrO_4^{2-}$, $MnO_4^-$).
  • Unlike p-block elements, where lower oxidation states are more stable for heavier elements (inert pair effect), in the d-block, higher oxidation states are more stable for heavier elements within a group. For example, W(VI) is more stable than Cr(VI).

Solution

Scandium (Z=21) does not exhibit variable oxidation states. It only shows the +3 oxidation state.

Trends in Standard Electrode Potentials ($E^\ominus$)

• $E^\ominus(M^{2+}/M)$ Trend: The values generally become less negative across the series, indicating a decreasing tendency to be oxidized. However, the trend is irregular.
  • The values for Mn and Zn are more negative than expected due to the stability of the resulting ions ($Mn^{2+}$ with $d^5$ and $Zn^{2+}$ with $d^{10}$ configuration).
  • Copper (Cu) has a unique positive $E^\ominus$ value ($+0.34$ V). This is because the high energy required to atomize Cu and ionize it to $Cu^{2+}$ is not balanced by its hydration enthalpy. This explains why copper does not liberate $H_2$ gas from acids.
• $E^\ominus(M^{3+}/M^{2+})$ Trend: These values give information about the stability of the +3 state relative to the +2 state.
  • The value for $Mn^{3+}/Mn^{2+}$ is highly positive ($+1.57$ V), indicating that $Mn^{3+}$ is a strong oxidizing agent and easily reduces to the very stable $Mn^{2+}$ ($d^5$).
  • The value for $Fe^{3+}/Fe^{2+}$ is less positive ($+0.77$ V), as $Fe^{3+}$ ($d^5$) is also stable.

Solution

• $Cr^{2+}$ ($d^4$) is reducing because it readily loses an electron to become $Cr^{3+}$ ($d^3$). The $d^3$ configuration provides extra stability in an octahedral field, as it corresponds to a half-filled $t_{2g}$ level.
• $Mn^{3+}$ ($d^4$) is oxidising because it readily gains an electron to become $Mn^{2+}$ ($d^5$). The $d^5$ configuration is exceptionally stable because it is a half-filled d-subshell.

Magnetic Properties

• Paramagnetism: Arises from the presence of unpaired electrons. Paramagnetic substances are weakly attracted by a magnetic field. Most transition metal ions and their compounds are paramagnetic.
• Diamagnetism: Arises when all electrons are paired. Diamagnetic substances are weakly repelled by a magnetic field.
• The magnetic moment of a paramagnetic substance can be calculated using the spin-only formula: $\mu = \sqrt{n(n+2)}$ where n is the number of unpaired electrons and μ is the magnetic moment in units of Bohr Magneton (BM). The greater the number of unpaired electrons, the higher the magnetic moment.

Given

• Atomic number, Z = 25 (This is Manganese, Mn)
• The ion is divalent, so it is $Mn^{2+}$.

To Find

The magnetic moment, $\mu$.

Formula

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

Solution

The electronic configuration of Mn (Z=25) is $[Ar] 3d^5 4s^2$. For the divalent ion $Mn^{2+}$, two electrons are removed from the 4s orbital. The configuration of $Mn^{2+}$ is $[Ar] 3d^5$. The number of unpaired electrons, $n = 5$.

Substitute this into the formula: $\mu = \sqrt{5(5+2)} = \sqrt{5 \times 7} = \sqrt{35} \approx 5.92 \text{ BM}$

Final Answer The magnetic moment is $5.92$ BM.

Formation of Coloured Ions

Most compounds of transition metals are coloured in the solid state or in solution.

• Reason: The colour arises from the absorption of energy from visible light to excite an electron from a lower energy d orbital to a higher energy d orbital. This is known as a d-d transition.
• The colour we see is the complementary colour of the light that is absorbed.
• Ions with completely filled ($d^{10}$, e.g., $Zn^{2+}$, $Cu^+$) or completely empty ($d^0$, e.g., $Sc^{3+}$, $Ti^{4+}$) d orbitals are colourless because no d-d transition is possible.

Formation of Complex Compounds

Transition metals form a large number of coordination compounds or complexes.

• Reason: This is due to:
  1. The relatively small size of the metal ions.
  2. Their high ionic charges.
  3. The availability of empty d orbitals for bond formation with ligands.

Catalytic Properties

Transition metals and their compounds are excellent catalysts.

• Reason:
  1. Variable Oxidation States: They can easily switch between oxidation states, allowing them to provide a new pathway with lower activation energy for a reaction.
  2. Ability to Form Complexes: They can form intermediate complexes with reactants, bringing them closer and weakening their bonds.
• Examples:
  • Iron (Fe) in the Haber's process for ammonia synthesis.
  • Vanadium(V) oxide ($V_2O_5$) in the Contact process for sulphuric acid manufacture.
  • Nickel (Ni) in the hydrogenation of oils.

Formation of Interstitial Compounds

Transition metals can trap small atoms like hydrogen (H), carbon (C), or nitrogen (N) in the empty spaces (interstitial sites) of their crystal lattices.

• Properties of Interstitial Compounds:
  • They are very hard and rigid.
  • They have high melting points, even higher than the pure metals.
  • They retain metallic conductivity.
  • They are chemically inert.
  • They are usually non-stoichiometric (e.g., $TiH_{1.7}$, $VH_{0.56}$).

Alloy Formation

Alloys are homogeneous solid solutions of two or more metals.

• Reason: Transition metals have similar atomic radii and other characteristics, which allows the atoms of one metal to easily substitute the atoms of another in its crystal lattice.
• Examples:
  • Brass: an alloy of copper and zinc.
  • Bronze: an alloy of copper and tin.
  • Stainless Steel: an alloy of iron with chromium and nickel.

Some Important Compounds of Transition Elements

Potassium Dichromate ($K_2Cr_2O_7$)

This is an important orange-red crystalline compound and a powerful oxidizing agent.

Preparation

It is prepared from chromite ore ($FeCr_2O_4$) in three steps:

1. Conversion of chromite ore to sodium chromate: The ore is fused with sodium carbonate ($Na_2CO_3$) in the presence of air. $4FeCr_2O_4 + 8Na_2CO_3 + 7O_2 \rightarrow 8Na_2CrO_4 + 2Fe_2O_3 + 8CO_2$
2. Conversion of sodium chromate to sodium dichromate: The yellow solution of sodium chromate is acidified with sulphuric acid. $2Na_2CrO_4 + 2H^+ \rightarrow Na_2Cr_2O_7 + 2Na^+ + H_2O$
3. Conversion of sodium dichromate to potassium dichromate: A hot, concentrated solution of sodium dichromate is treated with potassium chloride (KCl). Potassium dichromate is less soluble and crystallizes out on cooling. $Na_2Cr_2O_7 + 2KCl \rightarrow K_2Cr_2O_7 + 2NaCl$

Chromate-Dichromate Equilibrium

In solution, chromate ($CrO_4^{2-}$, yellow) and dichromate ($Cr_2O_7^{2-}$, orange) ions are interconvertible, depending on the pH.

• In acidic solution (low pH): Chromate ions convert to dichromate ions. $2CrO_4^{2-} + 2H^+ \rightleftharpoons Cr_2O_7^{2-} + H_2O$
• In alkaline solution (high pH): Dichromate ions convert to chromate ions. $Cr_2O_7^{2-} + 2OH^- \rightleftharpoons 2CrO_4^{2-} + H_2O$

Oxidising Action

In acidic solution, potassium dichromate is a powerful oxidizing agent. The chromium is reduced from +6 to +3. $Cr_2O_7^{2-} + 14H^+ + 6e^- \rightarrow 2Cr^{3+} + 7H_2O \quad (E^\ominus = 1.33 \text{ V})$ It can oxidize:

• Iodide to iodine: $6I^- \rightarrow 3I_2 + 6e^-$
• Iron(II) to Iron(III): $6Fe^{2+} \rightarrow 6Fe^{3+} + 6e^-$
• $H_2S$ to Sulphur: $3H_2S \rightarrow 6H^+ + 3S + 6e^-$

Potassium Permanganate ($KMnO_4$)

This is a dark purple crystalline compound and a very strong oxidizing agent.

Preparation

It is prepared from pyrolusite ore ($MnO_2$).

1. Conversion of $MnO_2$ to potassium manganate: The ore is fused with an alkali (KOH) and an oxidizing agent like $KNO_3$ or in the presence of air. This produces dark green potassium manganate ($K_2MnO_4$). $2MnO_2 + 4KOH + O_2 \rightarrow 2K_2MnO_4 + 2H_2O$
2. Oxidation of manganate to permanganate: The green manganate solution can be oxidized in two ways:
   • Disproportionation in acidic/neutral solution: $3MnO_4^{2-} + 4H^+ \rightarrow 2MnO_4^- + MnO_2 + 2H_2O$
   • Electrolytic oxidation (commercial method): The manganate solution is electrolyzed.

Oxidising Action

Potassium permanganate is a versatile oxidizing agent that works in acidic, alkaline, and neutral solutions. The product of reduction depends on the pH of the medium.

• In acidic solution (strongest oxidizing action): Permanganate ($MnO_4^-$) is reduced to $Mn^{2+}$ (colourless). $MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O \quad (E^\ominus = 1.52 \text{ V})$ It oxidizes $Fe^{2+}$ to $Fe^{3+}$, oxalate ($C_2O_4^{2-}$) to $CO_2$, and iodide ($I^-$) to $I_2$.
• In neutral or faintly alkaline solution: Permanganate is reduced to manganese dioxide ($MnO_2$), a brown precipitate. $MnO_4^- + 2H_2O + 3e^- \rightarrow MnO_2 + 4OH^-$
• In strongly alkaline solution: Permanganate is reduced to manganate ($MnO_4^{2-}$, green). $MnO_4^- + e^- \rightarrow MnO_4^{2-}$

THE INNER TRANSITION ELEMENTS (f-BLOCK)

These elements are characterized by the filling of the inner 4f (Lanthanoids) and 5f (Actinoids) orbitals.

The Lanthanoids

The 14 elements following Lanthanum (La), from Cerium (Ce) to Lutetium (Lu), are called lanthanoids.

Electronic Configurations

• The general configuration is $[Xe] 4f^{1-14} 5d^{0-1} 6s^2$.
• The most stable and common oxidation state for all lanthanoids is +3.
• The tripositive ions ($Ln^{3+}$) have the configuration $[Xe] 4f^n$.

Atomic and Ionic Sizes (Lanthanoid Contraction)

There is a steady and gradual decrease in atomic and ionic radii across the series from La to Lu.

• Cause: The 4f electrons have a very poor shielding effect. As the atomic number and nuclear charge increase, the effective nuclear charge experienced by the outer electrons increases, pulling them closer to the nucleus.
• Consequences:
  1. Causes the radii of the 5d transition series elements to be very similar to the 4d series.
  2. Leads to very similar chemical properties among the lanthanoids, making their separation difficult.
  3. The basicity of their hydroxides decreases across the series.

Oxidation States

• +3 is the most common and stable oxidation state.
• Some elements show +2 and +4 oxidation states, which are associated with the extra stability of empty ($f^0$), half-filled ($f^7$), or completely filled ($f^{14}$) f-subshells.
  • Cerium (Ce) shows a +4 state ($Ce^{4+}$) as it achieves a noble gas configuration ($f^0$). It is a strong oxidizing agent.
  • Europium (Eu) shows a +2 state ($Eu^{2+}$) as it achieves a stable half-filled configuration ($f^7$). It is a strong reducing agent.
  • Ytterbium (Yb) shows a +2 state ($Yb^{2+}$) as it achieves a stable fully-filled configuration ($f^{14}$).
  • Terbium (Tb) shows a +4 state ($Tb^{4+}$) to achieve a half-filled configuration ($f^7$).

General Characteristics

• They are silvery-white, soft metals that tarnish in air.
• They are good conductors of heat and electricity.
• Many trivalent lanthanoid ions ($Ln^{3+}$) are coloured due to f-f transitions.
• Most $Ln^{3+}$ ions are paramagnetic (except $La^{3+}$ and $Lu^{3+}$).
• They are chemically reactive, similar to alkaline earth metals. They react with acids to liberate hydrogen gas.

Uses

• The most important use is in the production of alloy steels.
• Mischmetall: An alloy of lanthanoid metals (~95%) and iron (~5%). It is used in Mg-based alloys to make bullets, shells, and lighter flints.

The Actinoids

The 14 elements following Actinium (Ac), from Thorium (Th) to Lawrencium (Lr), are called actinoids.

Key Features

• All actinoids are radioactive.
• The later members have very short half-lives and are prepared only in nanogram quantities, making their study difficult.

Electronic Configurations

• The general configuration is $[Rn] 5f^{1-14} 6d^{0-1} 7s^2$.
• The energies of 5f, 6d, and 7s orbitals are very close.
• The 5f electrons are not as deeply buried as 4f electrons and can participate more in bonding.

Oxidation States

• Actinoids show a greater range of oxidation states than lanthanoids. This is because the 5f, 6d, and 7s orbitals have comparable energies.
• The most common oxidation state is +3, but it is not always the most stable.
• The maximum oxidation state increases from +4 for Th up to +7 for Np, and then decreases.

Actinoid Contraction

Similar to lanthanoid contraction, there is a gradual decrease in atomic and ionic size across the series. The actinoid contraction is greater from element to element than the lanthanoid contraction due to the poorer shielding by 5f electrons.

Comparison of Lanthanoids and Actinoids

Applications of d- and f-Block Elements

• Iron and Steel: The backbone of construction and manufacturing.
• Catalysts: $V_2O_5$ (sulphuric acid), Fe (ammonia), Ni (hydrogenation), $TiCl_4$ (polythene production).
• Pigments: $TiO_2$ is a brilliant white pigment used in paints.
• Batteries: $MnO_2$ is used in dry cells; Zn, Ni, and Cd are used in various batteries.
• Coinage Metals: Cu, Ag, Au (though Ag and Au are now mostly for collection).
• Photography: Silver bromide (AgBr) is light-sensitive and crucial for photographic film.