Transition Elements and Coordination Chemistry Notes

Transition Elements (d-Block elements)

• Elements where the last orbitals filled are the d orbitals are known as transition elements.

• Located in the middle of the periodic table between electropositive s-block and electronegative p-block elements.

• Elements of group 3 to 12 (IIIB, IVB, VB, VIB, VIIB, VIIIB, IB, and IIB) in the long form of the periodic table are called d-block elements.

• The name transition is given to these d- block elements because of their position between s- and p-block elements.

• All the transition elements are metals and they have a variety of interesting properties.

• Examples:

  • Precious metals such as silver, gold, and platinum.

  • Industrially important metals such as iron, copper, and titanium.

• Most of the common materials contain transition metals or a combination of several transition metals in the form of alloys.

General Characteristics

1. General Electronic Configuration:

   • Transition metals have the electronic configuration $(n-1) d^{1-10}ns^{1-2}$.

   • When electrons fill orbitals, ns-orbital is filled first than $(n -1)d$ orbital.

   • When losing during oxidation, ns electrons are lost first than $(n-1)d$ electrons.

   • Zn, Cd, and Hg (end members of first three series):

     • General electronic configuration $(n -1) d^{10} ns^2$.

     • Do not show properties of transition elements to any appreciable extent.

     • Called non-typical transition elements.

2. Classification

   • Transition elements consist of four series.

3. Metallic Character

   • Transition metals can lose valence electrons and form cations: $M \rightarrow M^{n+} + ne^-$.

   • Possess metallic luster, high density, high tensile strength, hardness, brightness, and are good conductors of heat and electricity.

   • Have all three types of structures: hexagonal close-packed (hcp), face-centered cubic (fcc), and body-centered cubic (bcc) lattices.

   • Both metallic and covalent bonding are present in the atoms of transition elements.

   • Hardness is due to the strong covalent bonding in their d-orbitals.

   • More the unpaired d-electrons, stronger will be the bonding and harder will be the metal.

   • Metallic bonding is due to the possession of one or two electrons in the outermost shell.

   • There is a gradual decrease in electropositive character from left to right.

4. Melting and Boiling Points

   • Due to the strong metallic bond, they have high melting points and boiling points.

   • The melting points of these elements rise to a maximum and then fall with the increase in atomic number.

   • Manganese and technetium have abnormally low melting points.

   • Tungsten has the highest melting point (3683K) among the transition elements.

   • High boiling points and high melting points are due to the strong covalent bonding in their d-orbitals.

5. Atomic and Ionic Radii

   • The atomic and ionic radii of transition elements are smaller than those of s-block elements and larger than p-block elements.

   • Among the elements of a particular transition series, as the atomic number increases:

     • The atomic radii first decrease till the middle.

     • Become almost constant.

     • Then increase towards the end of the period.

   • The decrease in atomic radii in the series in the beginning is due to an increase in nuclear charge, which tends to pull the ns electrons inwards, hence it tends to reduce the size.

$Table 1: Atomic radii of elements of 3d series (in \AA)$

6. Atomic Volume and Density

   • The atomic volume of metals are much lower than those of alkali and alkaline earth metals.

   • As the inner orbital gets filled, the increased nuclear charge pulls the electronic crowd inwards.

   • The atomic volume therefore decreased.

   • The decrease in atomic volume increases the density.

7. Ionization Potential (Ionization Energy) and Electronegativity

   • The ionization potential and electronegativity values of d-block elements lie between s- and p-block elements.

   • The d-block elements are less electropositive than s-block elements and more than p-block elements.

   • There is a small or negligible increase in the ionization potential of d-block elements while moving from left to right in the periodic table.

   • This is due to the fact that the differentiating electron enters in inner d-orbital which does not provide a high shielding effect, while the nuclear charge keeps increasing from left to right.

8. Oxidation States

   • All the transition metals exhibit a great variety of oxidation states.

   • It is related to the electronic configuration of their atoms.

   • The variable oxidation states of transition metals are due to the involvement of $(n-1)d$ and outer ns electrons.

   • While moving from left to right in the periodic table, the number of variable oxidation states is increased at the beginning, reaches the maximum, and then is decreased again.

   • In a lower oxidation state, the compounds formed are ionic, and in a higher oxidation state, they are covalent in nature. For example, CrO is an ionic compound, whereas $CrO_3$ is a covalent compound.

   • The minimum oxidation state = number of s electrons.

   • The maximum oxidation state = ns electrons + unpaired $(n-1)d$ electrons.

9. Color

   • The color exhibited by transition metal ions is due to the presence of unpaired electrons in d-orbitals which permits the d-d electron transition in their compounds, which lies in the visible region.

10. Magnetic Properties

    • Most of the transition metal elements are paramagnetic in nature. This is due to the presence of unpaired electrons in their d-orbital.

    • Thus, the transition elements having all the electrons paired are diamagnetic.

    • The substance having unpaired electrons is attracted by the magnetic field and is said to show paramagnetism.

    • The substance having no unpaired electrons repelled by the magnetic field and said to show diamagnetism.

11. Catalytic Properties

    • Most of the transition metals, their alloys, and compounds are used as catalysts.

    • Some common examples are Pt, Ni, Fe, $Cr, V<em>2O</em>5$, etc.

    • The catalytic power of these metals is due to:

      • The presence of partly filled d-orbitals.

      • Exhibiting various oxidation states, or the formation of intermediate complexes with reactants, thus lowering the energy of activation.

12. Formation of Alloys

    • d-block elements have a strong tendency to form alloys since their atomic sizes are very similar, and in the crystal lattice, one metal can be readily replaced by another.

    • Examples: Brass, bronze, and various types of steels.

    • The alloys so formed are relatively hard and possess a high melting point.

13. Complex Formation

    • They are renowned to form a large number of complex compounds, mainly due to:

      • Small atomic size and higher nuclear charge.

      • Presence of partly filled or vacant orbitals.

    • Examples: $[Fe(CN)<em>6]^{3-}$, $[Ni(CO)</em>4]$, $K<em>4[Fe(CN)</em>6]$ etc.

General Characteristics of Transition Elements

• Almost all the transition elements have typical metallic properties such as ductility, high tensile strength, and high thermal and electrical conductivity.

• They have high melting and boiling points.

• They have high densities.

• Most of them are colored compounds.

• They are electropositive in nature.

• They have a tendency to form complexes.

• They show variable oxidation states.

• The first ionization energy of these elements is higher than those of s-block elements but lesser than those of p-block elements.

• Their compounds are generally paramagnetic in nature.

• They can act as catalysts in different chemical reactions, etc.

Complex Ions and Metal Complexes

• A compound in which a metal atom or metal ion is coordinated into two or more ligands is called a metal complex or coordinated compound.

Different Terms Used in Metal Complexes

1. Ligand or Coordinating Groups:

   • The molecular or ionic species which get attached directly to the central metal atom or ion during the formation of a complex compound are called ligands.

   • Ligands are commonly negative ions ($Cl^-$, $Br^-$, $CN^-$ etc.) or neutral molecules (such as $NH_3$, CO, NO etc.) containing a lone pair of electrons.

   • Ligands donate one or more electron pairs to the central metal ion, forming an $L \rightarrow M$ coordinate bond.

   • Examples:

     • In $[Fe(CN)_6]^{4-}$ ion, $CN^-$ is a ligand.

     • In $[Ag(NH<em>3)</em>2]^+$ ion, $NH_3$ is a ligand.

Types of Ligands

• a) Monodentate Ligand:

  • The ligand which can share only one pair of electrons is called a monodentate ligand.

  • Example: $F^-$ (Fluoro), $Cl^-$ (Chloro), $CN^-$ (Cyano), $OH^-$ (Hydroxo), $H<em>2O$ (Aqua), $NH</em>3$ (ammine), CO (Carbonyl), NO (Nitrosyl), etc.

• b) Bidentate Ligand:

  • The ligand which can share two pairs of electrons is called a bidentate ligand.

  • Examples: $CO<em>3^{2-}$ (Carbonato), $C</em>2O<em>4^{2-}$ (oxalate), $SO</em>4^{2-}$ (sulphato), $NH<em>2-CH</em>2-CH<em>2-NH</em>2$ (ethylenediamine, en), etc.

• c) Polydentate Ligand:

  • The ligand which can share many electron pairs is called a polydentate ligand.

  • Examples:

    • Diethyltriamine (dien) = can share 3 pairs = tridentate ligand.

    • EDTA (ethylenediamine tetraacetate) = can share 6 pairs = hexadentate ligand.

• d) Ambidentate Ligand:

  • A ligand that has more than one type of donor atoms is called an ambidentate ligand.

  • Example: CN

1. Complex Ion:

   • A complex ion is a positively, negatively, or neutrally charged species which contains a central metal atom and a suitable number of ligands attaching the central metal atom or ion.

   • The complex ion should be closed by a square bracket.

   • Depending on the nature of the charge on the complex ion, it may be a cationic or anionic complex ion.

   • A cationic complex ion has a positive charge, while an anionic complex ion has a negative charge, and a complex which has no charge on it is neutral.

   • Thus, $[Co(NH<em>3)</em>6]^+$, $[Cu(NH<em>3)</em>6]^{2+}$ etc. are cationic complex ions, and $[Ni(CN)<em>4]^{4-}$, $[Fe(CN)</em>6]^{4-}$ etc. are anionic complex ions, and $[Ni(CO)<em>4]^0$, $[Co(NH</em>3)<em>3Cl</em>3]^0$ etc. are neutral complexes.

2. Coordination Number (CN) of Central Metal Atom or Cation:

   • The coordination number of the central metal atom or cation in a given complex compound is equal to the total number of donor atoms which are actually attached to the central metal atom.

   • Examples:

     • C.N. of Cu in $[Cu(NH<em>3)</em>6]^{2+}$ is 6.

     • C.N. of Co in $[Co(en)_3]^{3+}$ is 6 as the en is a bidentate ligand, and each shares two lone pairs.

   • In the case of complex compounds which contain only monodentate ligands, the coordination number of the central metal atom or ion is equal to the number of monodentate ligands coordinated to the metal atom.

   • This rule does not hold good for the complexes containing polydentate (bidentate, tridentate, tetradentate, etc.) ligands.

3. Molecular or Addition Compound:

   • If the solution of a mixture of two stable compounds in a molecular proportion is allowed to crystallize, it may give crystals of a new substance, which is called a molecular compound.

   • $KCl + MgCl<em>2 + 6H</em>2O \rightarrow KCl.MgCl<em>2.6H</em>2O$ (Carnallite)

4. Double Salt:

   • The molecular compound which gets dissociated completely in water giving individual constituent ions is called a double salt.

   • $KCl.MgCl<em>2.H</em>2O \rightarrow K^+(aq.) + Mg^{2+}(aq.) + Cl^-(aq.)$

Naming of Coordination Complexes

IUPAC Nomenclature of Coordination Compounds

• From left to right.

• Cation first, then anion.

• Coordination sphere = negative = add "ate" to CMA.

Example 1:

• $K<em>4[Fe(CN)</em>6]$

• Potassium hexacyanoferrate (ii)

• Oxidation state calculation:

  • $4(+1) + Fe + 6(-1) = 0$

  • $Fe = 2$

Example 2:

• $[Ag(NH<em>3)</em>2]Cl$

• Diammine silver (i) chloride

• Oxidation state calculation:

  • $Ag + 2(NH_3) + Cl = 0$

  • $Ag + 2(0) + (-1) = 0$

  • $Ag = 1$

Example 3:

• $K<em>2[NiCl</em>4]$

• Potassium tetrachloro nickelate (ii)

• Oxidation state:

  • Ni = 2

Example 4:

• $[FeF_6]$

• Hexaflouro ferrate (ii) ion

• Oxidation state:

  • Fe = 2

Example 5:

• $[Cr(H<em>2O)</em>4Cl<em>2]NO</em>3$

• Tetraaqua dichlorochromium (iii) nitrate

• Oxidation state:

  • Cr = 3

Example 6:

• $[Co(en)<em>2Cl</em>2]Cl$

• Dichloro bis(ethan-1,2-diamine) cobolt (iii) chloride

• Oxidation state:

  • Co = 3

Rules of Naming the Coordination Complexes by IUPAC Nomenclature System

1. The complex compounds are named by giving the name of the cation followed by the anion. There is space between the name of the cation and anion.

2. Naming the complex ion: Whether the complex ion is cation or anion, the name of the ligand must be written first followed by the name of the metal. The ligands must be written in their alphabetical order.

   • Ligands (in alphabetical order) + metal + oxidation state of metal

   • If the ligand contains di, tri, etc. in their names, the prefix like Bis (for two), Tris (for three), Tetrakis (for 4), etc. are used before the name of the ligand.

3. The name of the neutral coordination compound is given in a one-word way.

Shape of Complex Ion

• The ligands which satisfy the secondary valency have a directional nature, i.e., directed towards the fix position in the space around the central atom.

• Therefore, the secondary valency or coordination number determines the shape or geometry of the complex ion.

• Examples:

  • If the coordination number is 4, the complex ion is either tetrahedral or square planar.

  • Similarly, the complexes with coordination number 6 are octahedral.

Crystal Field Theory (CFT)

• Proposed by Hans Bethe and modified by Van Vleck.

• Assumes that the attraction between the central metal atom and the ligand in the complex is purely electrostatic, i.e., ionic.

• Successful in interpreting many important properties, i.e., color, magnetic properties, etc., of the complexes.

Assumptions of CFT

1. This theory considers a complex as a combination of a central metal atom ion surrounded by suitable ligands which are treated as point charges or point dipoles.

2. The bonding between the metal cation and ligand arises due to the electrostatic attraction between the nucleus of the cation and the partial negative charge present on the ligand. Thus, the bond between the metal and ligand is purely ionic.

3. The CFT does not consider the overlap between the metal orbitals and ligand orbitals.

4. The interaction between the electrons of the cation and those of the ligand is entirely repulsive. These repulsive forces are responsible for the splitting of d-orbitals.

5. Due to the repulsion of negative ligands with the valency d-orbital of the metal cation, the degeneracy of d-orbitals is destroyed, resulting in the splitting of d-orbitals to attain stabilization.

Shape of d-orbitals

• In tetrahedral complexes, d-orbitals are split into $t<em>{2g}$ and $e</em>g$ sets where $t<em>{2g}$ forms higher energy sets, and $e</em>g$ forms lower energy sets.

Color of Transition Metal Compounds

• When the transition metal ion is isolated, $(n-1)d$ orbitals are degenerate.

• But when the ion is surrounded by anions or ligands, the degeneracy is lost, and d-orbitals split up into $e<em>g$ and $t</em>{2g}$ sets.

• Jumping of electrons in between these sets of d-orbitals is associated with the absorption and emission of energy of definite wavelength, which corresponds to a particular color in the visible range.

• This transition of d- electrons between $e<em>g$ and $t</em>{2g}$ sets is called d-d electron transition.

• The jumping of electrons from the sets of higher energy to that of the lower energy corresponds to the emission of a particular wave in the visible range.

• Therefore, the compounds of transition metals and complex ions are colored.

Catalytic Properties of the Transition Metals

• Transition metals and their compounds act as catalysts in various processes.

• Some examples are:

  1. Hydrogenation of alkenes, alkynes, oils, etc. [Finely divided nickel]

  2. Decomposition of $H<em>2O</em>2$. [Manganese oxide $MnO_2$]

  3. Manufacture of $H<em>2SO</em>4$ in the contact process. [Vanadium pentoxide, $V<em>2O</em>5$]

  4. Manufacture of $NH_3$ in Haber’s process. [Iron in presence of metal oxides]

• Transition metals have surface-active sites, which act as surfaces for gaseous reactions. These metals act as surface catalysts as they adsorb the gaseous molecules on their surface to speed up the reaction.

• Example: Hydrogenation of alkenes and alkynes requires Ni or Pd as surface catalysts.

• In some compounds of transition metals, they show catalytic properties due to the various oxidation states of the metal.

• The compound with one oxidation state combines with the reactive species to form an intermediate in which the metal exhibits different oxidation states.

• Now, the intermediate changes itself into a product, and the catalyst in which the metal gains the original oxidation state is regenerated.

Transition elements are metals found in groups 3 to 12 of the periodic table, characterized by the filling of d orbitals. They possess unique properties, including variable oxidation states and the ability to conduct electricity and heat. Their general electronic configuration is $(n-1)d^{1-10}ns^{1-2}$. Key characteristics include:

• Metallic luster, high density, and melting points due to strong metallic bonding.

• Small atomic and ionic radii compared to s-block elements, becoming smaller and then larger within a series.

• Variable oxidation states influenced by the electronic configuration, often forming both ionic and covalent compounds.

• Color and paramagnetic nature attributed to unpaired electrons in d-orbitals.

• Tend to form complex ions and alloys due to similar atomic sizes.

• Catalytic properties are prominent, enhanced by oxidation state variability and partly filled d-orbitals.