New sheet - 11/01 11:36

Cheatsheet Content

The d- and f-Block Elements: Overview The d-block elements, also known as Transition Metals , occupy groups 3-12 of the periodic table. Their distinguishing feature is the progressive filling of d-orbitals. The f-block elements, or Inner Transition Metals , consist of the lanthanoids (4f series) and actinoids (5f series), where f-orbitals are progressively filled. These are placed separately below the main body of the periodic table. Transition metals are defined as metals having an incomplete d-subshell in their neutral atom or in their common oxidation states. Elements like Zn, Cd, and Hg (Group 12) have full $d^{10}$ configurations in both ground and common oxidation states, and thus are generally not considered true transition metals , though their chemistry is often studied alongside them due to their position as end members of the transition series. Electronic Configurations of 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. Exceptions: Due to the extra stability of half-filled ($d^5$) and completely filled ($d^{10}$) orbitals, some elements show anomalous configurations. Chromium (Cr): $[Ar]3d^54s^1$ instead of $3d^44s^2$. Copper (Cu): $[Ar]3d^{10}4s^1$ instead of $3d^94s^2$. Palladium (Pd): $[Kr]4d^{10}5s^0$. Table: Electronic Configurations of 1st Series Transition Elements (Ground State) Element Sc Ti V Cr Mn Fe Co Ni Cu Zn Z 21 22 23 24 25 26 27 28 29 30 $4s$ 2 2 2 1 2 2 2 2 1 2 $3d$ 1 2 3 5 5 6 7 8 10 10 General Characteristics of d-Block Elements The d-orbitals protrude to the periphery of an atom, making them significantly influenced by their surroundings and capable of influencing other atoms/molecules. This leads to distinct chemical properties: Variable Oxidation States: Due to the small energy difference between $(n-1)d$ and $ns$ orbitals, electrons from both can participate in bonding. Formation of Coloured Ions: Unpaired d-electrons can undergo d-d transitions by absorbing visible light. Complex Formation: Small size, high ionic charge, and availability of d-orbitals allow them to form stable complexes. Catalytic Properties: Ability to adopt multiple oxidation states and form complexes enables them to act as catalysts. Paramagnetism: Presence of unpaired electrons causes attraction to magnetic fields. Physical Properties of Transition Metals Metallic Character: Most are typical metals with high tensile strength, ductility, malleability, and excellent thermal/electrical conductivity. Exceptions: Zn, Cd, Hg, Mn (less typical). Lattice Structures: Exhibit various metallic structures (hcp, bcc, ccp/fcc). Melting and Boiling Points: Generally high due to strong interatomic metallic bonding involving both $ns$ and $(n-1)d$ electrons. Maxima typically occur around the middle of each series (e.g., Cr, Mo, W) where the number of unpaired electrons is maximized, leading to stronger bonding. Anomalies exist for Mn and Tc. Enthalpy of Atomisation ($\Delta_a H^\circ$): High values, correlating with high melting points. Higher $\Delta_a H^\circ$ generally means stronger interatomic interaction. Metals of 2nd and 3rd series have higher $\Delta_a H^\circ$ than 1st series, leading to more metal-metal bonding in heavier transition metals. Table: Lattice Structures of Transition Metals (Selected) Sc Ti V Cr Mn Fe Co Ni Cu Zn hcp hcp bcc bcc X (bcc, ccp) bcc ccp (hcp) ccp ccp X (hcp) Note: hcp = hexagonal close packed; bcc = body centred cubic; ccp = cubic close packed; X = a typical metal structure. Atomic and Ionic Sizes Trends Across a Series (d-block): Generally, atomic and ionic radii decrease with increasing atomic number. This is because the new electron enters an $(n-1)d$ orbital, and the effective nuclear charge increases. However, the shielding effect of d-electrons is not as effective as s or p-electrons, so the decrease is not as rapid. Trends Down a Group (d-block): From 3d to 4d series, atomic radii generally increase as expected due to the addition of a new shell. From 4d to 5d series, atomic radii are remarkably similar. This phenomenon is known as Lanthanoid Contraction . Lanthanoid Contraction: Occurs due to the poor shielding effect of 4f-electrons. The 4f-orbitals are filled before the 5d-series begins. This poor shielding causes the effective nuclear charge to increase significantly, pulling the outer electrons closer to the nucleus and resulting in a smaller-than-expected atomic and ionic radii for the 5d-series elements. Consequences of Lanthanoid Contraction: Similar radii of 4d and 5d elements in the same group (e.g., Zr (160 pm) and Hf (159 pm)). Similar chemical properties of 4d and 5d elements within the same group, making their separation difficult. Increased density of 5d series elements compared to 4d series elements. Ionisation Enthalpies ($\Delta_i H$) Trends Across a Series (d-block): Ionisation enthalpy generally increases from left to right due to increasing nuclear charge. However, the increase is less steep than in p-block elements because the added electrons go into an inner $(n-1)d$ orbital, providing some shielding. Irregularities: The first ionisation enthalpy shows an irregular trend due to varying stability of $d^5$ and $d^{10}$ configurations and the energy difference between $(n-1)d$ and $ns$ orbitals. Second and third ionisation enthalpies often show higher values due to the removal of electrons from more stable configurations or more effectively shielded orbitals. For example, high values for Cr and Cu for $M^+$ to $M^{2+}$ due to $d^5$ and $d^{10}$ configurations respectively. Factors Affecting Ionisation Enthalpy: Nuclear Charge: Higher nuclear charge leads to higher ionisation enthalpy. Shielding Effect: Inner d-electrons partially shield outer electrons, reducing ionisation enthalpy. Exchange Energy: Stability associated with half-filled or completely filled degenerate orbitals (e.g., $d^5$, $d^{10}$) due to maximum exchange energy. This makes it harder to remove an electron from such configurations. Oxidation States One of the most notable features of transition elements is their ability to exhibit a wide range of oxidation states. Origin: Arises from the small energy difference between $(n-1)d$ and $ns$ electrons, allowing both to participate in bonding. Range: The greatest number of oxidation states is observed in the middle of the series (e.g., Mn from +2 to +7). Extremes: Early elements (Sc, Ti) show fewer oxidation states due to having fewer d-electrons to lose/share. Late elements (Cu, Zn) show fewer oxidation states due to having many d-electrons, making it harder to access higher valencies. Stability Trends: Maximum oxidation states often correspond to the sum of $ns$ and $(n-1)d$ electrons (e.g., Mn has max +7 from $3d^54s^2$). Stability of higher oxidation states decreases after the middle of the series. Oxidation states typically differ by unity (e.g., V: +2, +3, +4, +5), unlike non-transition elements where they often differ by two. Influence of Electronegative Elements: Oxygen and Fluorine can stabilize higher oxidation states (e.g., CrF$_6$, Mn$_2$O$_7$) due to their high electronegativity and ability to form multiple bonds. Inert Pair Effect: In d-block, unlike p-block, lower oxidation states are generally less favored for heavier members in a group (e.g., Mo(VI) and W(VI) are more stable than Cr(VI)). Zero Oxidation State: Observed in complex compounds with $\pi$-acceptor ligands (e.g., Ni(CO)$_4$, Fe(CO)$_5$). Table: Common Oxidation States of 1st Row Transition Metals (Bold indicates most common) Sc Ti V Cr Mn Fe Co Ni Cu Zn +1 +1 +2 +2 +2 +2 +2 +2 +2 +2 +2 +2 +3 +3 +3 +3 +3 +3 +3 +3 +3 +3 +4 +4 +4 +4 +4 +4 +4 +4 +5 +5 +5 +5 +6 +6 +6 +6 +7 +7 Standard Electrode Potentials ($E^\circ$) The $E^\circ$ values for $M^{2+}/M$ indicate the ease of reduction of $M^{2+}$ to $M$ (or oxidation of $M$ to $M^{2+}$). A more negative $E^\circ$ indicates a greater tendency to be oxidized. Trends Across a Series: Generally, $E^\circ$ values for $M^{2+}/M$ become less negative (or more positive) across the 3d series, indicating a decreasing tendency to form $M^{2+}$ ions. This correlates with increasing ionisation enthalpies. Anomalies: Mn, Ni, Zn: More negative $E^\circ$ than expected. Mn: Due to the stability of $Mn^{2+}$ ($d^5$ configuration). Zn: Due to the stability of $Zn^{2+}$ ($d^{10}$ configuration). Ni: Due to unusually high negative enthalpy of hydration ($\Delta_{hyd}H^\circ$) of $Ni^{2+}$ ions, which compensates for the high ionisation energy. Copper (Cu): Has a positive $E^\circ$ value ($+0.34V$) for $Cu^{2+}/Cu$. This means copper does not liberate hydrogen from acids. Its high energy required to transform $Cu(s)$ to $Cu^{2+}(aq)$ is not compensated by its hydration enthalpy. $E^\circ$ for $M^{3+}/M^{2+}$: Mn: Very high positive $E^\circ$ for $Mn^{3+}/Mn^{2+}$ ($+1.57V$), indicating $Mn^{3+}$ is a strong oxidising agent (prefers to be reduced to $Mn^{2+}$ which has stable $d^5$). Fe: Relatively low $E^\circ$ for $Fe^{3+}/Fe^{2+}$ ($+0.77V$), indicating $Fe^{3+}$ is less oxidising than $Mn^{3+}$ (due to stable $d^5$ for $Fe^{3+}$). Cr: Negative $E^\circ$ for $Cr^{3+}/Cr^{2+}$ ($-0.41V$), indicating $Cr^{2+}$ is a strong reducing agent (prefers to be oxidised to $Cr^{3+}$ which has stable $d^3$). Chemical Reactivity Most transition metals are electropositive and react with mineral acids. Early transition metals like Ti and V are passive to dilute non-oxidizing acids at room temperature due to oxide layer formation. The reactivity generally decreases across the series. $Ti^{2+}$, $V^{2+}$ and $Cr^{2+}$ are strong reducing agents. For example, $Cr^{2+}$ (d$^4$) can be easily oxidized to $Cr^{3+}$ (d$^3$, stable half-filled $t_{2g}$ orbitals). $Mn^{3+}$ and $Co^{3+}$ are strong oxidizing agents. For example, $Mn^{3+}$ (d$^4$) is easily reduced to $Mn^{2+}$ (d$^5$, stable half-filled d-orbitals). Magnetic Properties Transition metals exhibit either diamagnetism or paramagnetism. Diamagnetism: Substances repelled by a magnetic field, caused by the presence of only paired electrons. Paramagnetism: Substances attracted to a magnetic field, caused by the presence of unpaired electrons. The magnetic moment ($\mu$) is calculated using the "spin-only" formula : $$\mu = \sqrt{n(n+2)}$$ where $n$ is the number of unpaired electrons and $\mu$ is in Bohr magnetons (BM). Ferromagnetism: An extreme form of paramagnetism where substances are very strongly attracted to magnetic fields (e.g., Fe, Co, Ni). Orbital Angular Momentum: For 1st series transition metals, the contribution of orbital angular momentum to magnetic moment is often quenched and can be ignored. Table: Calculated and Observed Magnetic Moments (BM) for 1st Row Transition Metal Ions Ion Configuration Unpaired electrons (n) $\mu$ (Calculated) $\mu$ (Observed) Sc$^{3+}$ $3d^0$ 0 0 0 Ti$^{3+}$ $3d^1$ 1 1.73 1.75 Ti$^{2+}$ $3d^2$ 2 2.84 2.76 V$^{2+}$ $3d^3$ 3 3.87 3.86 Cr$^{2+}$ $3d^4$ 4 4.90 4.80 Mn$^{2+}$ $3d^5$ 5 5.92 5.96 Fe$^{2+}$ $3d^6$ 4 4.90 5.3 - 5.5 Co$^{2+}$ $3d^7$ 3 3.87 4.4 - 5.2 Ni$^{2+}$ $3d^8$ 2 2.84 2.9 - 3.4 Cu$^{2+}$ $3d^9$ 1 1.73 1.8 - 2.2 Zn$^{2+}$ $3d^{10}$ 0 0 0 Formation of Coloured Ions Most transition metal ions form coloured compounds in solid or aqueous solution. Mechanism: When white light falls on a transition metal compound, electrons in the lower energy d-orbitals absorb specific wavelengths of light to get excited to higher energy d-orbitals (d-d transitions). The remaining transmitted light, which is the complementary colour of the absorbed light, is observed. Conditions: For d-d transitions to occur, the ion must have partially filled d-orbitals (i.e., $d^1$ to $d^9$ configurations). Ions with $d^0$ or $d^{10}$ configurations (e.g., Sc$^{3+}$, Ti$^{4+}$, Zn$^{2+}$) are colourless because d-d transitions are not possible. Factors Affecting Colour: The colour depends on the nature of the metal ion, its oxidation state, and the ligands surrounding it. Table: Colours of Some 1st Row Transition Metal Ions (Aquated) Configuration Example Colour $3d^0$ Sc$^{3+}$ colourless $3d^0$ Ti$^{4+}$ colourless $3d^1$ Ti$^{3+}$ purple $3d^1$ V$^{4+}$ blue $3d^2$ V$^{3+}$ green $3d^3$ V$^{2+}$ violet $3d^3$ Cr$^{3+}$ violet $3d^4$ Mn$^{3+}$ violet $3d^4$ Cr$^{2+}$ blue $3d^5$ Mn$^{2+}$ pink $3d^5$ Fe$^{3+}$ yellow $3d^6$ Fe$^{2+}$ green $3d^6, 3d^7$ Co$^{3+}$, Co$^{2+}$ blue/pink $3d^8$ Ni$^{2+}$ green $3d^9$ Cu$^{2+}$ blue $3d^{10}$ Zn$^{2+}$ colourless Formation of Complex Compounds Transition metals readily form complex compounds (coordination compounds) where the metal ion acts as a Lewis acid and ligands (anions or neutral molecules) act as Lewis bases. Reasons: Small Size and High Charge: The relatively small size and high positive charge of transition metal ions allow them to attract electrons from ligands effectively. Availability of d-orbitals: Presence of vacant d-orbitals of appropriate energy allows them to accept electron pairs from ligands for bond formation. Examples: $[Fe(CN)_6]^{3-}$, $[Fe(CN)_6]^{4-}$, $[Cu(NH_3)_4]^{2+}$, $[PtCl_4]^{2-}$. Catalytic Properties Many transition metals and their compounds act as effective catalysts in various chemical reactions. Reasons: Variable Oxidation States: They can readily change their oxidation states, forming unstable intermediates and providing an alternative reaction pathway with lower activation energy. Large Surface Area: Many transition metals provide a suitable surface for reactants to adsorb and react (e.g., finely divided metals). Formation of Complexes: They can form intermediate complexes with reactants, facilitating the reaction. Examples: $V_2O_5$ in Contact Process (oxidation of $SO_2$ to $SO_3$). Fe in Haber's Process (synthesis of ammonia). Ni in catalytic hydrogenation of oils. $PdCl_2$ in Wacker process (oxidation of ethyne to ethanal). $Fe^{3+}$ catalysing reaction between $I^-$ and $S_2O_8^{2-}$ ($2Fe^{3+} + 2I^- \rightarrow 2Fe^{2+} + I_2$; $2Fe^{2+} + S_2O_8^{2-} \rightarrow 2Fe^{3+} + 2SO_4^{2-}$). Interstitial Compounds These are formed when small non-metal atoms (H, C, N, B) get trapped in the interstitial voids (spaces) within the crystal lattices of transition metals. Characteristics: High melting points (higher than pure metals). Very hard (some borides approach diamond in hardness). Retain metallic conductivity. Chemically inert. Non-stoichiometric (e.g., $TiC$, $Mn_4N$, $Fe_3H$, $VH_{0.56}$). Alloy Formation Transition metals readily form alloys (homogeneous solid solutions of two or more metals). Reasons: Transition metals have similar atomic sizes (radii within 15% of each other) and similar chemical properties. Characteristics: Alloys are generally harder and have higher melting points than their constituent metals. Examples: Ferrous alloys: Stainless steel (Fe, Cr, Ni), tool steels (Fe, W, V, Mo). Brass (Cu, Zn), Bronze (Cu, Sn). Mischmetall: An alloy of lanthanoid metals (~95%) and iron (~5%), used for lighter flints and in Mg-based alloys. Oxides and Oxoanions of Metals Formation: Most transition metals form oxides upon heating with oxygen. Oxidation States: Highest oxidation state in oxides often coincides with the group number (e.g., $Sc_2O_3$ to $Mn_2O_7$). Beyond group 7, higher oxides are less stable. Acidic/Basic Nature: Lower oxidation state oxides are generally basic (e.g., $V_2O_3$). Intermediate oxidation state oxides are amphoteric (e.g., $V_2O_4$, $Cr_2O_3$). Higher oxidation state oxides are acidic (e.g., $CrO_3$, $Mn_2O_7$). Ionic/Covalent Character: As oxidation number increases, covalent character increases (e.g., $Mn_2O_7$ is a covalent green oil). Potassium Dichromate ($K_2Cr_2O_7$) Preparation: From chromite ore ($FeCr_2O_4$): Fusion with alkali ($Na_2CO_3$ or $K_2CO_3$) in presence of air. $$4FeCr_2O_4 + 8Na_2CO_3 + 7O_2 \rightarrow 8Na_2CrO_4 + 2Fe_2O_3 + 8CO_2$$ Conversion of sodium chromate ($Na_2CrO_4$) to sodium dichromate ($Na_2Cr_2O_7$) by acidification. $$2Na_2CrO_4 + 2H^+ \rightarrow Na_2Cr_2O_7 + 2Na^+ + H_2O$$ Crystallization of $K_2Cr_2O_7$ by adding $KCl$ to $Na_2Cr_2O_7$ solution (less soluble). $$Na_2Cr_2O_7 + 2KCl \rightarrow K_2Cr_2O_7 + 2NaCl$$ Chromate-Dichromate Interconversion: In acidic medium, chromate ions ($CrO_4^{2-}$, yellow) convert to dichromate ions ($Cr_2O_7^{2-}$, orange): $$2CrO_4^{2-} + 2H^+ \rightleftharpoons Cr_2O_7^{2-} + H_2O$$ In alkaline medium, dichromate ions convert to chromate ions: $$Cr_2O_7^{2-} + 2OH^- \rightleftharpoons 2CrO_4^{2-} + H_2O$$ Structure: Chromate ($CrO_4^{2-}$) is tetrahedral. Dichromate ($Cr_2O_7^{2-}$) consists of two tetrahedra sharing one corner (Cr-O-Cr bridge), with a bond angle of $126^\circ$. Oxidising Agent: $K_2Cr_2O_7$ is a strong oxidising agent in acidic medium. $$Cr_2O_7^{2-} + 14H^+ + 6e^- \rightarrow 2Cr^{3+} + 7H_2O \quad (E^\circ = 1.33V)$$ It oxidises $I^-$ to $I_2$, $S^{2-}$ to $S$, $Sn^{2+}$ to $Sn^{4+}$, $Fe^{2+}$ to $Fe^{3+}$. Potassium Permanganate ($KMnO_4$) Preparation: From pyrolusite ore ($MnO_2$): Fusion with $KOH$ and an oxidising agent ($KNO_3$ or air) produces green potassium manganate ($K_2MnO_4$). $$2MnO_2 + 4KOH + O_2 \rightarrow 2K_2MnO_4 + 2H_2O$$ Disproportionation of manganate in neutral or acidic solution: $$3MnO_4^{2-} + 4H^+ \rightarrow 2MnO_4^- + MnO_2 + 2H_2O$$ Electrolytic oxidation of manganate in alkaline solution: $$MnO_4^{2-} \xrightarrow{\text{electrolytic oxidation}} MnO_4^-$$ In laboratory: Oxidation of $Mn^{2+}$ salt by peroxodisulphate ($S_2O_8^{2-}$). $$2Mn^{2+} + 5S_2O_8^{2-} + 8H_2O \rightarrow 2MnO_4^- + 10SO_4^{2-} + 16H^+$$ Properties: Dark purple (almost black) crystals, isostructural with $KClO_4$. Sparingly soluble in water. Decomposes on heating. Structure: Both manganate ($MnO_4^{2-}$, green) and permanganate ($MnO_4^-$, purple) ions are tetrahedral. Manganate is paramagnetic (1 unpaired electron), permanganate is diamagnetic (no unpaired electrons). Oxidising Agent: $KMnO_4$ is a very strong oxidising agent. Acidic Medium: $$MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O \quad (E^\circ = 1.52V)$$ Oxidises $I^-$ to $I_2$, $Fe^{2+}$ to $Fe^{3+}$, oxalate to $CO_2$, $H_2S$ to $S$, $SO_3^{2-}$ to $SO_4^{2-}$, $NO_2^-$ to $NO_3^-$. Neutral or Faintly Alkaline Medium: $$MnO_4^- + 2H_2O + 3e^- \rightarrow MnO_2 + 4OH^-$$ Oxidises $I^-$ to $IO_3^-$, $S_2O_3^{2-}$ to $SO_4^{2-}$, $Mn^{2+}$ to $MnO_2$. The f-Block Elements (Inner Transition Elements) Consist of Lanthanoids (4f series) and Actinoids (5f series). Lanthanoids (4f Series) Elements from Cerium (Ce, Z=58) to Lutetium (Lu, Z=71). They are placed after Lanthanum (La, Z=57) in the periodic table, which itself is a d-block element but is often discussed with lanthanoids due to chemical similarities. Electronic Configurations of Lanthanoids General configuration: $[Xe]4f^{1-14}5d^{0-1}6s^2$. Lanthanum (La): $[Xe]5d^16s^2$. Cerium (Ce): $[Xe]4f^15d^16s^2$. Gadolinium (Gd): $[Xe]4f^75d^16s^2$ (due to stable half-filled $4f^7$). Lutetium (Lu): $[Xe]4f^{14}5d^16s^2$. Most stable ion: $Ln^{3+}$ (tripositive ions) with configuration $4f^{n}$. Atomic and Ionic Sizes (Lanthanoid Contraction) Definition: The gradual decrease in atomic and ionic radii of lanthanoids from La to Lu. Cause: Poor shielding effect of 4f-electrons. As the atomic number increases, the nuclear charge increases, but the 4f-electrons provide very poor shielding, leading to a stronger effective nuclear charge that pulls the outer electrons closer. Consequences: (Already discussed in d-block section, but applies here as the origin of the effect). Oxidation States of Lanthanoids Most Common: +3 oxidation state is the most stable and common for all lanthanoids. Other States: +2 and +4 oxidation states are sometimes observed, particularly when they lead to stable $f^0$ (empty), $f^7$ (half-filled), or $f^{14}$ (fully filled) configurations. $Ce^{4+}$ ($4f^0$): Strong oxidising agent, reverts to $Ce^{3+}$. $Eu^{2+}$ ($4f^7$): Strong reducing agent, reverts to $Eu^{3+}$. $Yb^{2+}$ ($4f^{14}$): Strong reducing agent, reverts to $Yb^{3+}$. $Tb^{4+}$ ($4f^7$): Strong oxidising agent, reverts to $Tb^{3+}$. General Characteristics of Lanthanoids Appearance: Silvery white, soft metals (Sm is steel hard). Tarnish rapidly in air. Melting Points: Range from $1000K$ to $1200K$ (Sm melts at $1623K$). Density: Increases with increasing atomic number. Colour: Most $Ln^{3+}$ ions are coloured in solid state and aqueous solutions due to f-f transitions (except $La^{3+}$ ($4f^0$) and $Lu^{3+}$ ($4f^{14}$) which are colourless). Magnetic Properties: Most $Ln^{3+}$ ions are paramagnetic (except $La^{3+}$ and $Lu^{3+}$). Reactivity: Early members are quite reactive, similar to Calcium. Reactivity decreases across the series. React with hydrogen, halogens, acids, and form oxides ($Ln_2O_3$) and hydroxides ($Ln(OH)_3$). Hydroxides are basic. Ionisation Enthalpies: First IE is around $600 kJ/mol$, second IE around $1200 kJ/mol$. Irregularities in third IE due to stability of $f^0, f^7, f^{14}$. Actinoids (5f Series) Elements from Thorium (Th, Z=90) to Lawrencium (Lr, Z=103). They are placed after Actinium (Ac, Z=89). Electronic Configurations of Actinoids General configuration: $[Rn]5f^{1-14}6d^{0-1}7s^2$. All actinoids are radioactive. The 5f-orbitals are not as deeply buried as 4f-orbitals and can participate in bonding to a greater extent. Similar irregularities in configurations due to stability of $f^0, f^7, f^{14}$ (e.g., Am, Cm). Atomic and Ionic Sizes (Actinoid Contraction) Definition: Gradual decrease in atomic and ionic radii across the actinoid series. Magnitude: Actinoid contraction is greater from element to element than lanthanoid contraction. Cause: Poor shielding effect of 5f-electrons. 5f-electrons are less effective at shielding than 4f-electrons, leading to a more pronounced contraction. Oxidation States of Actinoids Exhibit a greater range of oxidation states than lanthanoids. Common: +3 is the most common oxidation state. Higher States: Early actinoids (Th, Pa, U, Np) exhibit higher oxidation states (+4, +5, +6, +7). This is because the 5f, 6d, and 7s orbitals are of comparable energies, allowing more electrons to participate in bonding. Stability: Higher oxidation states are more stable for early actinoids, while +3 becomes more stable for later actinoids. General Characteristics of Actinoids Appearance: Silvery in appearance, but display a variety of structures due to irregular metallic radii. Reactivity: Highly reactive metals, especially when finely divided. React with boiling water, acids, and non-metals at moderate temperatures. Magnetic Properties: More complex than lanthanoids, with lower values. Ionisation Enthalpies: Lower than lanthanoids, especially for early actinoids, because 5f electrons are less firmly held and can participate more easily in bonding. Comparison of Lanthanoids and Actinoids Feature Lanthanoids Actinoids Electronic Configuration $4f$ orbitals filled $5f$ orbitals filled Oxidation States Mainly +3; few show +2, +4 (stable $f^0, f^7, f^{14}$) Mainly +3; exhibit higher oxidation states (+4, +5, +6, +7) especially for early members Atomic/Ionic Radii Lanthanoid contraction (less pronounced) Actinoid contraction (more pronounced) Chemical Reactivity Less reactive; early members like Ca More reactive; react with non-metals readily Magnetic Properties Simpler; spin-only formula often works More complex Colour Most $Ln^{3+}$ ions are coloured due to f-f transitions Most $An^{3+}$ ions are coloured Complex Formation Less tendency to form complexes Greater tendency to form complexes Radioactivity Only Promethium ($Pm$) is radioactive All are radioactive Applications of d- and f-Block Elements Iron and Steel: Most important construction materials. Used in various alloys like stainless steel. Catalysts: Many transition metals and their compounds (e.g., $V_2O_5$, Fe, Ni, Pd) are used as catalysts in industrial processes. Coinage Metals: Copper, silver, and gold are used for making coins due to their durability and aesthetic appeal. Electrical Conductors: Copper and silver are excellent electrical conductors, widely used in wiring and electronic components. Pigments: Compounds of transition metals are often brightly colored and used as pigments in paints, ceramics, and glass (e.g., chromium compounds for green, cobalt compounds for blue). Batteries: Metals like nickel and cadmium are used in various types of rechargeable batteries.