Coordination compounds also called metal complexes are compounds that contain central metal atom or ion bonded to one or more surrounding ligands. The above definition includes compounds with metal-carbon bonds, called organometallic compounds
These metals could be main group metals or transition metals.
Main group metals are metals in s and p blocks of the periodic table. For example; Mg, Al, Pb.
Transition metals are elements that have partially filled d or f orbitals and exhibit variable oxidation states. They are located in between s and p block on the periodic table.
The d-block metals have 5 orbitals with 10 maximum electrons needed to completely fill them.
However, the f– block metals have 7 orbitals to be maximally filled by 14 electrons.
Chemistry of Coordination Compounds
The d-block, consists of three series: the 3d series (Scandium to Zinc), the 4d series (Yttrium to Cadmium), and the 5d series (Lanthanum to Mercury). These elements have partially filled d orbitals in their atomic or ionic ground state, which gives them unique chemical and physical properties.
Understanding the roles of the ligands even with their different categories in the formation of metal complexes, our focus will only be on the first row members of the transition metals. This will help to pick from the simple aspect of their chemistry because elements of other rows have tendency of have two or more central metals.
Significance of First-Row Transition Metal Complexes
First-row transition metals play crucial roles in various fields, including catalysis, medicine, materials science, and environmental chemistry.
Understanding their chemistry provides insights into their reactivity, bonding, and electronic properties.
First-row Transition Metals
These are the 10 first row transition metals on the Periodic Table:
Scandium (Sc), Titanium (Ti), Vanadium (V), Chromium (Cr), Manganese (Mn), Iron (Fe), Cobalt (Co), Nickel (Ni), and Copper (Cu).
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Periodic Table – 3d Block Elements
| Element | Symbol | Electronic Configuration |
|---|---|---|
| Scandium | Sc | [Ar] 3d1 4s2 |
| Titanium | Ti | [Ar] 3d2 4s2 |
| Vanadium | V | [Ar] 3d3 4s2 |
| Chromium | Cr | [Ar] 3d5 4s1 |
| Manganese | Mn | [Ar] 3d5 4s2 |
| Iron | Fe | [Ar] 3d6 4s2 |
| Cobalt | Co | [Ar] 3d7 4s2 |
| Nickel | Ni | [Ar] 3d8 4s2 |
| Copper | Cu | [Ar] 3d10 4s1 |
| Zinc | Zn | [Ar] 3d10 4s2 |
They are naturally occurring elements.
Each metal has a unique electronic configuration, which influences its chemical behavior.
Ligands and Coordination Chemistry
Ligands are molecules or ions that donate electron pairs to the central metal atom or ion in order to form metal complex.
Classification of Ligands
Ligands are molecules or ions that donate electron pairs to form coordination complexes with transition metal atoms or ions. Ligands can be classified based on their electronic and structural properties, as well as the number of electron pairs they donate to the metal center.
Monodentate Ligands
Monodentate ligands are ligands that donate a single electron pair (or a single atom) to the metal center. They form a single bond with the metal atom and occupy one coordination site.
For examples:
1. Water (H2O): Water is a common monodentate ligand that readily coordinates to metal ions, forming hydrated metal complexes. For example, [Cu(H2O)6]2+ is the hexaaquacopper(II) ion.
2. Chloride (Cl–): Chloride ions are monodentate anionic ligands that can bind to metal centers. An example is [FeCl6]3-, which is the hexachloridoferrate(III) ion.
Bidentate Ligands
Bidentate ligands are ligands that donate two electron pairs to the metal center. They form two bonds with the metal atom and occupy two coordination sites.
For examples:
1. Ethylenediamine (en): Ethylenediamine is a bidentate ligand that has two amino groups (NH2) separated by an ethylene bridge (-CH2CH2-). It can form complexes with metal ions such as [Ni(en)3]2+, which is tris(ethylenediamine)nickel(II) ion.
2. Oxalate (C2O42-): Oxalate is a bidentate anionic ligand that can form complexes with metal ions. An example is [Co(C2O4)3]3-, which is the tris(oxalato)cobaltate(III) ion.
Tridentate Ligands
Tridentate ligands are ligands that donate three electron pairs to the metal center. They form three bonds with the metal atom and occupy three coordination sites.
For example:
1. Terpyridine (tpy): Terpyridine is a tridentate ligand that contains three pyridine rings. It can form complexes with metal ions such as [Fe(tpy)2]2+, which is bis(2,2′:6′,2”-terpyridine)iron(II) ion.
Polydentate Ligands
Polydentate ligands, also known as chelating ligands, donate multiple electron pairs to the metal center. They form multiple bonds with the metal atom and occupy multiple coordination sites, creating a chelate complex.
For examples:
1. EDTA (ethylenediaminetetraacetic acid): EDTA is a hexadentate ligand that contains four carboxylic acid groups (-COOH) and two amino groups (-NH2). It forms strong complexes with metal ions by coordinating through its multiple electron-donating groups. An example is [Cu(EDTA)]2-, which is the copper(II) complex of EDTA.
2. DTPA (diethylenetriaminepentaacetic acid): DTPA is a hexadentate ligand that has five carboxylic acid groups and one amino group. It forms complexes with metal ions and is commonly used in chelation therapy. An example is [Co(DTPA)]–, which is the cobalt (III) complex of DTPA.
Ambidentate Ligands
Ambidentate ligands are ligands that have multiple atoms capable of bonding to a metal center but can only bond through one of those atoms at a time. These ligands can form isomeric complexes depending on which atom is bonded to the metal.
For examples:
1. Nitrite (NO2–): The nitrite ion can coordinate to metal ions through either the nitrogen atom (N-bound) or the oxygen atom (O-bound). For example, [Cu(NO2)2] can exist in two isomeric forms, nitrito-N and nitrito-O.
2. Thiocyanate (SCN-): Thiocyanate can bond to metal ions through either the sulfur atom (S-bound) or the nitrogen atom (N-bound). Isomeric forms include [Fe(NCS)2] and [Fe(SCN)2].
Bridging Ligands
Bridging ligands are ligands that can coordinate to multiple metal centers simultaneously, creating a bridge between them. They can facilitate the formation of coordination compounds with more than one metal atom.
For example:
1. Cyanide (CN–): Cyanide can bridge between two metal centers, forming a metal-metal bond. A well-known example is [Fe(CN)6]4-, which is the hexacyanidoferrate(II) ion.
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Classification of Ligands
| Monodentate Ligand | IUPAC Name | Formula |
|---|---|---|
| Water | Aqua | H2O |
| Chloride | Chloro | Cl– |
| Ammonia | Amine | NH3 |
| Cyanide | Cyano | CN– |
| Nitrate | Nitrato | NO3– |
| Bidentate Ligand | IUPAC Name | Formula |
| Ethylenediamine | Ethylenediamine | en |
| Oxalate | Oxalato | C2O42- |
| Acetylacetonate | Acetylacetonate | acac– |
| Ethylene glycol | Dihydroxyethane | HOCH2CH2OH |
| Glycinato | Glycinato | gly |
| Tridentate Ligand | IUPAC Name | Formula |
| Terpyridine | Terpyridine | terpy |
| Ethylenediaminetriacetate | Ethylenediaminetriacetato | edta3- |
| Oxamidate | Oxamidate | C2O22- |
| Tris(2-pyridylmethyl)amine | Tris(2-pyridylmethyl)amine | tpa |
| Triethylenetetramine | Triethylenetetramine | trien |
| Polydentate Ligand | IUPAC Name | Formula |
| EDTA (Ethylenediaminetetraacetate) | Ethylenediaminetetraacetato | edta4- |
| DTPA (Diethylenetriaminepentaacetate) | Diethylenetriaminepentaacetato | dtpa5- |
| Tetraazamacrocyclic ligand | Tetraazamacrocyclic ligand | [Ni(C14H28N4)]2+ |
| Phthalocyanine | Phthalocyanine | C32H18N8 |
| Cyclam (1,4,8,11-Tetraazacyclotetradecane) | 1,4,8,11-Tetraazacyclotetradecane | cyclam |
| Ambidentate Ligand | IUPAC Name | Formula |
| Nitrite | Nitrito | NO2– |
| Thiocyanate | Thiocyanato | SCN– |
| Nitro | Nitro | NO2 |
| Nitrosyl | Nitrosyl | NO |
| Isocyanate | Isocyanato | NCO– |
| Bridging Ligand | IUPAC Name | Formula |
| Aquo | Aquo | H2O |
| Cyano | Cyano | CN– |
| Carboxylate | Carboxylato | RCO2– |
| Oxalate | Oxalato | C2O42- |
| Hydroxo | Hydroxo | OH– |
Chelate
Chelates are multidentate ligands that form cyclic structures with a metal ion through multiple donor sites. The ligand, known as a chelating ligand, coordinate to the metal ion through two or more donor atoms, creating a ring-like structure called a chelate ring. The formation of chelate complexes significantly enhances the stability of the coordination compound. For example, copper ion form the complex with EDTA and en to give [Cu(EDTA)]2- and [Cu(en)2]2+ respectively.
In a metal complex containing chelate, the chelating ligand forms a cyclic structure with the metal ion. The coordination bonds between the ligand and the metal ion create a closed loop, known as the chelate ring. The formation of the chelate ring increases the thermodynamic stability of the complex by reducing the number of potential dissociation pathways and minimizing the coordination entropy.
Stability of Metal-Chelate Complexes
Chelate complexes are generally more stable than complexes formed with monodentate or polydentate ligands. This increased stability arises from several factors:
a. Ring Strain: The formation of the chelate ring introduces a degree of strain within the ligand structure. This strain enhances the binding strength between the ligand and the metal ion, making the complex more stable.
b. Entropy Reduction: The formation of the chelate ring reduces the conformational flexibility of the ligand and decreases the number of possible dissociation pathways. As a result, the entropy of the system is lowered, favoring the formation of the chelate complex.
c. Increased Coordination Number: Chelating ligands effectively increase the coordination number of the metal ion by occupying multiple coordination sites. The higher coordination number contributes to increased stability by maximizing the metal-ligand interactions.
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Naming of Metal Complexes or Coordination Compounds
Naming Metal Complex Cations for 3d Metal Series (first transition metals)
Identify the Metal ion
Start by identifying the specific 3d metal ion present in the metal complex. The name of the metal is generally the same as the element’s name.
Specify the Oxidation State
Indicate the oxidation state of the metal ion using Roman numerals in parentheses. This step is crucial for 3d metals as they can exhibit multiple oxidation states. For example, Fe(II) represents iron(II) with a +2-oxidation state, while Fe(III) represents iron(III) with a +3-oxidation state.
Name the Ligands
List the ligands attached to the metal ion in alphabetical order. The ligand names are modified based on the type of ligand:
Anionic ligands: The ligand name is modified by replacing the ending with “-o.” For example, chloride becomes chloro, oxide becomes oxo, and cyanide becomes cyano.
Neutral ligands: The ligand name remains unchanged. Examples include water (aqua), ammonia (ammine), and carbon monoxide (carbonyl).
Cationic ligands: The ligand name is modified by adding the suffix “-ium.” For instance, hydronium becomes hydronium and ammonium becomes ammonium.
Include the Multiplicity
Indicate the charge of the complex cation using Roman numerals in parentheses, followed by the overall charge of the complex in superscript. The overall charge is determined by balancing the charges of the metal ion and the ligands.
For Example: [Co(NH3)6]3+ is hexaamminecobalt(III) ion.
Naming Metal Complex Anions for 3d Metal Series:
The naming of complex anions follows a similar pattern to metal complex cations, but the names end in “-ate.”
Identify the Metal Ion
Identify the specific 3d metal ion present in the metal complex anion, using the same naming conventions as for complex cations.
Specify the Oxidation State
Indicate the oxidation state of the metal ion using Roman numerals in parentheses.
Name the Ligands
List the ligands attached to the metal ion, following the same modifications as for metal complex cations.
Include the Multiplicity
Indicate the charge of the complex anion using Roman numerals in parentheses, followed by the overall charge of the complex in superscript.
For example: [Cr(CN)6]3- is hexacyanochromate(III) ion.
III. Naming Neutral Metal Complexes for 3d Metal Series
Naming neutral complexes follows a similar process, but without specifying the charge.
1. Identify the metal atom or ion.
2. Specify the oxidation state if necessary.
3. Name the ligands in alphabetical order, making appropriate modifications as discussed earlier.
Example: [Ni(en)3] is tris(ethylenediamine)nickel(II).
Prefixes for Ligands and Chelates
| Number | Monodentate Ligand | Chelates |
|---|---|---|
| 1 | Mono | Mono |
| 2 | Di | Bis |
| 3 | Tri | Tris |
| 4 | Tetra | Tetrakis |
| 5 | Penta | Pentakis |
| 6 | Hexa | Hexakis |
| 7 | Hepta | Heptakis |
| 8 | Octa | Octakis |
| 9 | Nona | Nonakis |
| 10 | Deca | Decakis |
| 11 | Undeca | Undecakis |
| 12 | Dodeca | Dodecakis |
Isomers and Prefixes for 3d Metal Complexes
Isomers, which are different compounds with the same molecular formula, are distinguished using prefixes.
Cis- and trans- isomers: These prefixes are used to indicate different geometric arrangements of ligands around a metal atom. Cis– refers to ligands on the same side, while trans- refers to ligands on opposite sides.
Fac- and mer- isomers: These prefixes are used for complexes with three identical ligands. Fac– indicates a facial arrangement where the ligands are adjacent to each other, while mer- indicates a meridional arrangement where the ligands are arranged in a plane.
Example: [Co(en)2Cl2] can exist in both cis-dichloridobis(ethylenediamine)cobalt(III) and trans-dichloridobis(ethylenediamine)cobalt(III) isomers.
Coordination Number and Geometry in 3d Metal Complexes
In coordination chemistry, the coordination number and geometry play a crucial role in describing the spatial arrangement of ligands around a central metal atom in a complex. It is essential to note that these phenomena are part of the distinguishing factors between the molecular formula and the structures of the coordination compounds.
Coordination Number
The coordination number is the number of bonding sites directly bonded to the central metal atom in a metal complex. The bonding sites could come from any of the classes of the ligands mentioned. Coordination number represents the maximum number of donor atoms that can form coordination bonds with the metal center. The coordination number is denoted by the symbol “CN.” In 3d metal complexes, the most common coordination numbers are 4 and 6 because of their stability, although other numbers can also occur.
Coordination Geometries
The coordination geometry describes the spatial arrangement of ligands around the central metal atom. Different coordination numbers can lead to specific geometric shapes. Let’s explore the common coordination geometries observed in 3d metal complexes:
Octahedral Geometry (CN = 6)
Octahedral geometry is observed when the coordination number is 6, indicating that the central metal atoms are bonded to by six donor (bonding) sites from the ligands. These ligands occupy the six corners of an octahedron, resulting in a symmetrical arrangement. The central metal atom resides at the center of the octahedron. Octahedral complexes are commonly encountered in transition metal chemistry due to their stability and structural diversity. Examples of 3d metal complexes with octahedral geometry include: [Co(NH3)6]3+ (hexaamminecobalt(III) ion), [Fe(H2O)6]2+ (hexaaquairon(II) ion), [CrCl6]3- (hexachlorochromate(III) ion), and [Fe(en)3]3+ (tris(ethylenediamine)iron(III) ion) [Ni(en)2Cl2] (dichlorobis(ethylenediamine)nickel(II)), and [Fe(acac)3] (tris(acetylacetonato)iron(III)).
Tetrahedral Geometry (CN = 4)
Tetrahedral geometry arises when the coordination number is 4, signifying that the central metal atoms are bonded to by four donor (bonding) sites from the ligands. The ligands occupy the four corners of a tetrahedron, generating a symmetrical arrangement. The central metal atom is situated at the center of the tetrahedron. Tetrahedral complexes are commonly observed when the metal ion is coordinated by ligands that form strong sigma bonds. Examples of 3d metal complexes with tetrahedral geometry include: [NiCl4]2- (tetrachloronickelate(II) ion), [Zn(NH3)4]2+ (tetraamminedizinc(II) ion), [FeF4]2- (tetrafluoroferrate(II) ion).
Square Planar Geometry (CN = 4)
Square planar geometry has the same coordination number of 4 with the tetrahedral geometry. However, in square planar complexes, the ligands occupy the four corners of a flat square around the central metal atom. The central metal atom lies in the same plane as the ligands. Square planar complexes are commonly observed when the metal ion possesses d8 electronic configuration and is surrounded by strong-field ligands.
Examples of 3d metal complexes with square planar geometry include: [PtCl4]2- (tetrachloroplatinate(II) ion), [Ni(CN)4]2-(tetracyanidonickelate(II) ion), and [Pd(NH3)2Cl2] (diamminedichloropalladium(II)).
Trigonal Bipyramidal Geometry (CN = 5)
Trigonal bipyramidal geometry is observed when the coordination number is 5, indicating that that the central metal atoms are bonded to by five donor (bonding) sites from the ligands. The ligands occupy the corners of an imaginary trigonal bipyramid, with three ligands in the equatorial plane and two ligands along the axial positions. Trigonal bipyramidal complexes often arise from a combination of monodentate, bidentate, or tridentate ligands. Examples of 3d metal complexes with trigonal bipyramidal geometry include:
[CoCl5]2- (pentachlorocobaltate(II) ion), [Fe(CO)5] (pentacarbonyliron(0)) and [MnF5]2- (pentafluoromanganate(II) ion).
Bonding in Metal Complexes
Metal complexes contain the coordinate covalent and covalent bonds.
Coordination covalent bond is formed when is formed when a ligand donates an electron pair to the central metal atom. It is a type of bond that in which the ligand acts as a Lewis base, providing a lone pair of electrons, and the central metal atom acts as a Lewis acid, accepting the electron pair.
In addition to coordinate covalent bond, some ligands can form covalent bonds with the metal center, where electron pairs are shared between the ligand and the metal atom. Covalent bonds involve a more equal sharing of electrons between the two atoms, and they can occur when the ligand has available lone pairs of electrons to interact with the empty or partially filled orbitals of the metal.
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Classification of Metal Complexes for 3d Metals
Metal complexes of 3d transition metals can be classified based on various factors, including the coordination number, geometry, and nature of the ligands. The classification of metal complexes provides insight into their structural and chemical properties specific to the 3d metals.
Mononuclear Metal Complexes
Mononuclear metal complexes of 3d metals consist of a single metal atom or ion coordinated to one or more ligands. The ligands may be monodentate, bidentate, tridentate, polydentate, or ambidentate ligands. Examples of mononuclear complexes of 3d metals include: [Ni(H2O)6]2+ hexaaquonickel(II) ion, and [Co(en)3]3+: tris(ethylenediamine)cobalt(III) ion.
Polynuclear Metal Complexes
Polynuclear complexes of 3d metals contain two or more metal ions that are connected by bridging ligands. These complexes often exhibit interesting magnetic and optical properties due to the metal-metal interactions. Examples of polynuclear complexes of 3d metals include: Prussian Blue ([Fe4[Fe(CN)6]3]), [Mn2(CO)10] (Decacarbonyldimanganese) and Ru3(CO)12.
Homoleptic and Heteroleptic Metal Complexes
Homoleptic complexes of 3d metals contain only one type of ligand coordinated to the metal center. Heteroleptic complexes, on the other hand, consist of two or more different ligands coordinated to the metal ion.
a. [CuCl4]2-: Tetrachloridocuprate(II) ion, a homoleptic metal complex where all four chloride ligands are coordinated to the copper(II) ion.
b. [Co(en)2Cl2]+: Dichloridobis(ethylenediamine)cobalt(III) ion, a heteroleptic complex with two ethylenediamine ligands and two chloride ligands coordinated to a cobalt(III) ion.
Chelate Metal Complexes
Chelate complexes of 3d metals are formed when a multidentate ligand coordinates to a metal ion, creating a cyclic structure called a chelate ring. These complexes exhibit increased stability due to the chelate effect. Examples of chelate complexes of 3d metals include: [Cu(EDTA)]2-: Diamminetetraacetato copper(II) ion, and [Fe(acac)3]: Tris(acetylacetonato)iron(III).
Isomerism in Coordination Compounds
Isomerism refers to the existence of different compounds with the same chemical formula but different arrangements of atoms.
In these 3d metal complexes, the classes of isomerism are structural isomerism and geometric isomerism.
Structural iIsomerism
Structural isomerism can arise from variations in ligand connectivity or the arrangement of ligands around the central metal atom.
Linkage Isomerism
Linkage isomerism occurs when the same ligand can coordinate to the central metal atom through different atoms. The ligand atoms involved in coordination may exhibit different chemical properties, leading to differences in reactivity and behavior of the complex. Examples of linkage isomerism include:
[Co(NH3)5(NO2)]2+ and [Co(NH3)5(ONO)]2+: In these metal complexes, the nitrite (NO2–), an ambidentate ligand can coordinate to the cobalt (Co) center through either the nitrogen or the oxygen atom. The resulting complexes have different structural and electronic properties, influencing their reactivity and biological activities.
Coordination Sphere Isomerism
Coordination sphere isomerism involves the interchange of ligands and/or counterions around the central metal atom while maintaining the same connectivity. The rearrangement of ligands within the coordination sphere can result in different steric effects and electronic environments. An example of coordination sphere isomerism is:
[Co(NH3)4Cl2]Cl and [Co(NH3)4Cl2]Br: These complexes have the same ligand connectivity, with four ammine (NH3) ligands and two chloride (Cl) ligands coordinated to the cobalt (Co) center. However, the counterion is different in each complex, with chloride (Cl) in the first and bromide (Br) in the second. This interchange of counterions leads to different structural and chemical properties.
Ionization Isomerism
Ionization isomerism arises when the ligand and counterion interchange their positions within the complex. This isomerism is particularly observed in metal complexes where the ligand can act as both a ligand and a counterion. An example of ionization isomerism is:
[Co(NH3)5SO4]Br and [Co(NH3)5Br]SO4: In these metal complexes, the sulfate (SO4) and bromide (Br) ions exchange their positions with respect to the central cobalt (Co) atom. This results in different structural and chemical properties due to the different nature of the counterions.
Hydrate Isomerism
Hydrate isomerism involves the presence of water molecules as ligands, and the isomers differ in the arrangement of water molecules around the central metal atom. This isomerism is particularly relevant when water molecules are coordinated directly to the metal center or when they are present as lattice water molecules in the crystal lattice of the complex.
Examples of hydrate isomerism include:
[Cr(H2O)6]Cl3 and [Cr(H2O)5Cl]Cl2(H2O): In the first metal complex, six water molecules are coordinated to the central chromium (Cr) atom, while in the second metal complex, five water molecules are coordinated, and an additional water molecule is present as a lattice water molecule. These isomers have different physical properties, such as solubility and color, due to the different arrangements of water molecules.
Geometric Isomerism
Geometric isomerism is a type of isomerism that arises from the different spatial arrangements of ligands around a central metal atom in a coordination compound. It occurs when there is a restricted rotation around a bond or a rigid ligand framework, resulting in distinct isomeric forms. In 3d metal complexes, geometric isomerism can manifest in two main types: cis-trans isomerism and facial-meridional isomerism.
Cis-Trans Isomerism
Cis-trans isomerism, also known as geometric isomerism, occurs when ligands are arranged differently around a metal center due to the restricted rotation around a bond or a rigid ligand framework. This type of isomerism is commonly observed in metal complexes with bidentate or polydentate ligands. For example:
a. cis-[Co(NH3)4Cl2]+ and trans-[Co(NH3)4Cl2]+:
In the cis isomer, the two chloride ligands (Cl) are adjacent to each other, while in the trans isomer, the chloride ligands are on opposite sides of the coordination sphere. The isomerism arises due to the restricted rotation around the coordination bond between the cobalt (Co) atom and the chloride ligands.
b. cis-[Pt(NH3)2Cl2] (cis) and trans-[Pt(NH3)2Cl2]:
In the cis isomer, the two chloride ligands (Cl) are on the same side of the coordination sphere, while in the trans isomer, the chloride ligands are on opposite sides. The isomerism arises due to the restricted rotation around the coordination bond between the platinum (Pt) atom and the chloride ligands.
c. cis-[Ru(en)2Cl2] and trans-[Ru(en)2Cl2]:
In the cis isomer, the two ethylenediamine (en) ligands are adjacent to each other, while in the trans isomer, the ethylenediamine ligands are on opposite sides of the coordination sphere. The isomerism arises due to the restricted rotation around the coordination bond between the ruthenium (Ru) atom and the ethylenediamine ligands.
Facial-Meridional Isomerism
Facial-meridional isomerism, also known as facial-mer isomerism, occurs when three ligands of a coordination compounds are arranged either in a facial (fac) or meridional (mer) orientation around a central metal atom. This type of isomerism is commonly observed in metal complexes with tridentate or hexadentate ligands. For examples,
fac-[Fe(CO)3(NO)3] (fac) and mer-[Fe(CO)3(NO)3]:
In the fac-isomer, the three carbonyl (CO) ligands are arranged in a facial manner, while in the mer-isomer, the carbonyl ligands are arranged in a meridional manner around the iron (Fe) atom. The isomerism arises due to the restricted rotation around the coordination bonds between the iron atom and the ligands.
Some other examples are fac-[Ru(en)3]2+ and mer-[Ru(en)3]2+; fac-[Co(en)3]3+ (fac) and mer-[Co(en)3]3+
Optical Isomerism
Optical isomerism, also known as enantiomerism, is a type of stereoisomerism that arises when a molecule or metal complex exhibits chirality. Chirality refers to the property of an object or molecule that cannot be superimposed onto its mirror image. In 3d metal complexes, optical isomerism occurs when the arrangement of ligands around the central metal atom results in the presence of a chiral center.
In optical isomerism, two distinct isomers, known as enantiomers, exist for a given metal complex. Enantiomers have identical chemical and physical properties, except for their interaction with plane-polarized light. One enantiomer rotates the plane of polarized light clockwise (dextrorotatory, designated as “+”), while the other enantiomer rotates it counterclockwise (levorotatory, designated as “-“).
The presence of chiral ligands or an asymmetrical arrangement of ligands around the metal center gives rise to optical isomerism in 3d metal complexes. For example:
[Cr(en)2Cl2]+ (dichlorobis(ethylenediamine)chromium(III) ion) has the chromium ion (Cr3+) coordinated to two bidentate ethylenediamine (en) ligands and two chloride ions. The presence of the chelating ligand, ethylenediamine, leads to the formation of two enantiomers.
[Pt(NH3)2Cl2] (diamminedichloroplatinum(II)), also, features a central platinum ion (Pt2+) coordinated to two ammonia (NH3) ligands and two chloride ions. The arrangement of ligands around the platinum center gives rise to optical isomerism, resulting in two enantiomers.
Factors that Affect Isomerism in 3d Metal Complexes
The isomerism of 3d metal complexes can be influenced by various factors. Understanding these factors is important in predicting and explaining the occurrence of different isomeric forms. Here are some of these key factors:
Ligand Type and Geometry
The type and geometry of ligands play a significant role in determining the isomeric forms of 3d metal complexes. Different ligands exhibit different steric and electronic effects, which can influence the arrangement of ligands around the central metal atom. Ligands with different sizes, shapes, and coordination modes can lead to distinct isomeric forms.
Coordination Number
Different coordination numbers give rise to different possibilities for ligand arrangements and bond angles, resulting in the formation of isomeric forms. For example, complexes with a coordination number of 4 may exhibit tetrahedral or square planar isomers as earlier explained, the ones with coordination number of 6 may exhibit octahedral or trigonal prismatic isomers while with coordination number of 5, the geometry of metal complexes is trigonal bipyramidal.
Steric Effects
Steric effects, which arise from the spatial hindrance caused by bulky ligands, can significantly influence the isomerism of 3d metal complexes. Bulky ligands may restrict the rotation around coordination bonds or impose specific ligand orientations, leading to the formation of specific isomeric forms. Steric effects can also influence the preference for cis– or trans- arrangements in complexes with bidentate or polydentate ligands.
Chelate Effect
The chelate effect refers to the enhanced stability and geometric constraints observed in metal complexes with chelating ligands compared to metal complexes with monodentate ligands. Chelating ligands can form a ring structure by coordinating to the central metal atom through multiple donor atoms. This ring structure restricts the rotation around coordination bonds, leading to the formation of specific isomeric forms.
Electronic Effects
Electronic effects, such as the electronic configuration of the metal ion and ligand-field effects, can influence the isomerism of 3d metal complexes. Changes in the electronic environment can affect the bonding and coordination preferences of ligands, resulting in different isomeric forms. Ligands with different electronic properties can exhibit different affinities for specific metal-ligand bond geometries.
Solvent Effects
The choice of solvent can also influence the isomerism of 3d metal complexes. Solvents can interact with the metal-ligand bonds and affect their strength and orientation, leading to changes in the isomeric forms. Different solvents may stabilize or destabilize specific isomeric forms, depending on their polarity, coordination ability, and interactions with the metal complex.
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Theories of Electronic Structure of 3d Metal Complexes
The electronic structure of 3d metal complexes can be understood through various theories that aim at providing insight into the bonding and electronic properties of these metal complexes. For now, the valence bond, crystal field, and the ligand field theories will provide the needed explanation.
Valence Bond Theory
The valence bond theory, which incorporates the concept of hybridization, is a fundamental approach used to understand the electronic structures of transition metal complexes. In the context of 3d metal complexes, the valence bond theory explains the formation of coordination bonds based on hybridized orbitals.
For octahedral complexes, such as [Co(NH3)6]3+, the valence bond theory suggests the involvement of d2sp3 hybrid orbitals of the metal atom. It is important to note that the d orbitals utilized by first-row transition metals can either be 3d or 4d orbitals. Originally, these structures were described as covalent and ionic complexes, corresponding to the usage of 3d and 4d orbitals, respectively. The terms “hyperligated” and “hypoligated” were introduced as alternatives to describe these metal complexes, also referred to as inner orbital (using 3d) and outer orbital (using 4d) complexes.
The choice of which d orbitals are involved in bonding depends on the number of unpaired electrons present, which can be determined by observing the magnetic behavior of the compounds. The valence bond theory introduces the concepts of low spin and high spin configurations to describe the electronic configurations of d4 through d7 ions.
Consider the example of Fe(III), which has five unpaired electrons as an isolated ion, with one electron occupying each of the 3d orbitals. In an octahedral coordination compound, Fe(III) can exhibit either one or five unpaired electrons. When there is one unpaired electron, the ligand electrons exert a strong enough influence to force the pairing of the metal’s d electrons, resulting in the availability of two 3d orbitals for hybridization and bonding. Conversely, when there are five unpaired electrons, the ligands do not exert sufficient bonding strength to compel the pairing of the 3d electrons.
Another example is Co(II), which requires seven electrons. In this case, Co(II) can exhibit either one or three unpaired electrons. In the low-spin scenario with one unpaired electron, the seventh electron occupies a higher orbital. In the high-spin scenario with three unpaired electrons, the bonding involves the 4d or outer orbital hybrid, while the metal electrons remain in the 3d levels.
Similar considerations arise when dealing with eight or nine electrons, as seen in complexes like Ni(II) and Cu(II). These metal complexes often adopt tetrahedral or square-planar geometries. It’s important to note that while the valence bond theory has played a significant role in understanding bonding in coordination compounds, it is less frequently employed today, except for discussions on hybrid orbitals used in bonding. The theory’s reliance on the high-energy 4d orbitals raises doubts, and it struggles to provide a satisfactory explanation for the electronic spectra of complexes. Since electronic spectra provide valuable experimental data, this limitation represents a significant drawback of the valence bond theory.
For example: In [Co(NH3)6]3+ (hexaamminecobalt(III) ion), the central cobalt (Co3+) ion undergoes sp3d2 hybridization, resulting in the formation of six hybrid orbitals directed towards the six ammonia (NH3) ligands. Each NH3 ligand donates a pair of electrons to the Co3+ ion, resulting in the formation of six sigma bonds. The valence bond theory explains the coordination and bonding in this complex based on the hybridization of the cobalt d-orbitals and the formation of sigma bonds.
Another example, [Fe(CN)6]4- (hexacyanoferrate(II) ion) has the central iron (Fe2+) ion undergoes sp3d2 hybridization, resulting in the formation of six hybrid orbitals. The six cyanide (CN–) ligands act as pi acceptor ligands and form pi back-donation bonds with the Fe2+ ion. The valence bond theory can explain the formation of both sigma bonds and pi back-donation bonds, taking into account the hybridization and orbital overlap.
Also, [Cu(NH3)4]2+ (tetraamminecopper(II) ion) features a central copper (Cu2+) ion that undergoes sp3 hybridization, resulting in the formation of four hybrid orbitals directed towards the four ammonia (NH3) ligands. The valence bond theory describes the coordination and bonding in this complex by considering the overlap of the hybrid orbitals with the ligand orbitals, leading to the formation of sigma bonds.
Valence Bond Theory
| Hybridization | Geometry | Coordination Number | Example of 3d Metal Complex |
|---|---|---|---|
| d2sp3 | Octahedral | 6 | [Co(NH3)6]3+ |
| d2sp3 | Octahedral | 6 | [Cr(H2O)6]3+ |
| d2sp3 | Octahedral | 6 | [FeCl6]3- |
| sp3 | Tetrahedral | 4 | [NiCl4]2- |
| sp3 | Tetrahedral | 4 | [Zn(NH3)4]2+ |
| sp3 | Tetrahedral | 4 | [FeF4]2- |
| d2sp2 | Square Planar | 4 | [PtCl4]2- |
| d2sp3 | Trigonal Bipyramidal | 5 | [CoCl5]2- |
| d2sp3 | Trigonal Bipyramidal | 5 | [Fe(CO)5] |
Crystal Field Theory
Crystal field theory (CFT) provides a useful framework for understanding the electronic structure of metal ions in crystalline environments, such as coordination compounds.
In CFT, the primary focus is on the electrostatic interactions between metal ions and surrounding anions, such as oxide ions. These anions create an electrostatic field, which depends on the symmetry of the crystal structure. The d orbitals of the metal ions are affected by this field, leading to their splitting into different energy levels. While CFT was originally developed for describing ionic crystals, it neglected the consideration of covalent bonding.
It was later realized that the same arrangement of electron pair donor species around metal ions exists in both crystals and coordination compounds. This insight paved the way for the development of a more comprehensive theory called molecular orbital theory, which offers a deeper understanding of the electronic structure in metal complexes. As interest in coordination chemistry grew, these theories gained wider usage.
Consider the complex [Co(NH3)6]3+. Here, the cobalt (Co) ion is surrounded by six ammonia (NH3) ligands, forming an octahedral coordination environment. According to crystal field theory, the d orbitals of the cobalt ion split into two sets of energy levels due to the repulsion from the ligands. The dz2 and dx2–y2 orbitals (axial orbitals) (eg), which point directly at the ligands, experience increased repulsion and are consequently raised in energy. Conversely, the dxy, dxz, and dyz orbitals (non-axial orbitals) (t2g), which point between the ligands, are less affected by the field and remain at lower energy levels.
This splitting of the d orbitals gives rise to distinctive colors observed in transition metal complexes. In the case of [Co(NH3)6]3+, the energy difference between the raised d orbitals and the lower-energy orbitals is denoted as Δo (delta-o) or 10Dq (ten times the symbol Dq) in older references. The magnitude of Δo determines the energy spacing between the different d orbitals, influencing the absorption of specific wavelengths of light and thus giving the complex its characteristic color.
It’s important to note that crystal field theory has certain limitations. Firstly, it solely focuses on the electrostatic repulsion between orbitals and ligands and does not account for potential bonding interactions. Molecular orbital theory, on the other hand, considers bonding, leading to the formation of both higher and lower energy molecular orbitals. Secondly, CFT does not provide a complete picture of the electronic structure since it overlooks lower-energy (bonding) molecular orbitals.
To further illustrate the concepts, let’s explore another example. Consider the complex [Fe(CN)6]3-. Here, the iron (Fe) ion is surrounded by six cyanide (CN–) ligands, forming an octahedral coordination environment. According to crystal field theory, the d orbitals of the iron ion split due to the repulsion from the ligands. The dz2 and dx2–y2 (eg)orbitals are raised in energy, while the dxy, dxz, and dyz orbitals (t2g) remain relatively unchanged.
These energy-level differences in the d orbitals directly impact the electronic and magnetic properties of the complex. For instance, in the case of [Fe(CN)6]3-, the d-orbital splitting contributes to the compound’s paramagnetic behavior and its ability to undergo ligand exchange reactions.
Crystal Field Stabilization Energy
The energy difference between the actual electron distribution in a 3d metal complex and the distribution that would occur in a hypothetical uniform field is known as the crystal field stabilization energy (CFSE). CFSE plays a crucial role in determining the stability and electronic properties of coordination complexes.
In crystal field theory, the CFSE arises from the interaction between the metal’s d orbitals and the ligands surrounding it. When the metal ion is placed in a ligand field, the d orbitals experience repulsion from the ligands, resulting in their splitting into different energy levels. The extent of this splitting is determined by the nature and arrangement of the ligands.
Let’s consider an example using a 3d metal complex: [Co(H2O)6]2+. In this metal complex, the cobalt (Co) ion is coordinated by six water (H2O) ligands in an octahedral arrangement. The d orbitals of the cobalt ion split into two sets of energy levels due to the repulsion from the ligands. The dz2 and dx2–y2 orbitals (axial, eg orbitals), which point directly at the ligands, experience increased repulsion and are raised in energy. On the other hand, the dxy, dxz, and dyz orbitals (non-axial, t2g orbitals), which point between the ligands, are less affected by the field and remain at lower energy levels.
The CFSE can be determined by calculating the energy difference between the actual electron distribution in the metal complex and the distribution that would occur in the absence of ligand field effects. The CFSE value is equal in magnitude to the ligand field stabilization energy (LFSE), which quantifies the stabilization of the metal complex due to the ligand field. CFSE and LFSE are often represented by the symbol ΔE, and their values can be determined experimentally or calculated using theoretical methods.
For example, in [Co(H2O)6]2+, the CFSE represents the energy stabilization achieved by the cobalt ion in the presence of the surrounding water ligands. This stabilization arises from the repulsion between the ligands and the cobalt’s d orbitals. By quantifying the CFSE, we can better understand the overall stability and electronic properties of the metal complex.
It is worth noting that CFSE and LFSE are essential in predicting and explaining various properties of coordination complexes, including their spectroscopic behavior, magnetic properties, and reactivity. These concepts provide valuable insights into the electronic structure and stability of 3d metal complexes and are widely used in coordination chemistry research and applications.
Ligand Field Theory (LFT)
Ligand field theory (LFT) is a comprehensive theory that builds upon the foundation of crystal field theory (CFT) to provide a more detailed understanding of the electronic structure and bonding in coordination complexes. While CFT focuses primarily on the electrostatic interactions between metal d orbitals and ligands, LFT incorporates additional considerations such as σ and π bonding interactions.
In LFT, the ligands surrounding the metal ion are treated as molecular entities with their own molecular orbitals. These molecular orbitals interact with the metal’s d orbitals to form a set of hybrid orbitals, known as ligand field orbitals. These hybrid orbitals arise from the combination of metal d orbitals and ligand molecular orbitals, facilitating bonding and antibonding interactions.
When a metal ion is coordinated by ligands, the ligand field orbitals split into different energy levels. The magnitude of this splitting is influenced by various factors, including the nature and geometry of the ligands, as well as the metal’s oxidation state and electronic configuration.
Let’s consider an example to illustrate ligand field theory in the context of 3d metal complexes. Take the complex [Fe(CN)6]4-. Here, the iron (Fe) ion is surrounded by six cyanide (CN–) ligands, arranged in an octahedral geometry. In ligand field theory, the d orbitals of the iron ion interact with the molecular orbitals of the cyanide ligands, resulting in the splitting of the ligand field orbitals.
Due to the interaction between the iron d orbitals and the cyanide ligand molecular orbitals, the d orbitals split into two sets of energy levels. The dx2-y2 and dz2 orbitals, which experience greater repulsion from the ligands, are raised in energy, while the dxz, dyz, and dxy orbitals, which exhibit less repulsion, remain at lower energy levels. This splitting is often represented by the symbol Δ (delta).
In ligand field theory, the energy difference between these two sets of orbitals, represented as Δ, reflects the strength of the bonding and antibonding interactions between the metal and ligands. This energy splitting influences various properties of the complex, including its color, magnetism, and reactivity.
LFT provides a more comprehensive understanding of the electronic structure and bonding in metal complexes compared to CFT. By considering not only the electrostatic repulsion but also the bonding interactions between metal and ligands, LFT offers insights into the stability, spectroscopic behavior, and chemical reactivity of 3d metal complexes.
Spectrochemical Series
The spectrochemical series is a useful tool in coordination chemistry that ranks ligands based on their ability to cause ligand field splitting in a metal complex. It provides a qualitative measure of ligand strength, which determines the energy difference between the metal ion’s d orbitals in the presence of different ligands.
In the spectrochemical series, ligands are arranged in increasing order of their ligand field splitting strength, denoted as Δ. A ligand with a higher splitting strength will result in a larger energy difference between the metal ion’s d orbitals, leading to a greater separation in energy levels. This energy splitting directly affects the electronic and spectroscopic properties of the metal complex.
The expression for the strengths of some ligands in increasing order is as follows:
I– < Br– < S2- < SCN– < Cl– < NO3– < N3- < F– < OH– < C2O42- < H2O < NCS– < pyridine (py) < NH3 < en < bipy < phen < NO2– < PPh3 < CN– < CO
Let’s examine the ligand series and understand why ligands are arranged in this order.
At the lower end of the series, we have iodide (I–) and bromide (Br–) ions. These ligands are relatively weak in terms of ligand field splitting and have a small impact on the energy difference between the metal ion’s d orbitals i.e causing high spin of the d orbital electrons.
Moving up the series, we encounter ligands such as sulfide (S2-), thiocyanate (SCN–), and chloride (Cl–). These ligands are stronger than iodide and bromide ions but weaker than other ligands found later in the series. They induce a moderate splitting of the d orbitals.
As we progress further, ligands like nitrate (NO3–), azide (N3-), and fluoride (F–) exhibit a stronger ligand field splitting effect. Their ability to donate electron density to the metal ion increases, resulting in a larger energy difference between the metal’s d orbitals.
Continuing along the series, hydroxide (OH–) and oxalate (C2O42-) ligands have a higher splitting strength than previous ligands due to their stronger electron-donating ability.
Water (H2O) is a moderately strong ligand that produces a significant ligand field splitting effect on the metal’s d orbitals. It is often considered a benchmark ligand in the spectrochemical series.
Moving towards the stronger ligands, we encounter ligands such as pyridine (py), ammonia (NH3), ethylenediamine (en), bipyridine (bipy), and phenanthroline (phen). These ligands have strong π-donor capabilities and create a substantial ligand field splitting effect on the metal’s d orbitals.
Towards the end of the series, we have ligands like nitrite (NO2–), triphenylphosphine (PPh3), cyanide (CN–), and carbon monoxide (CO). These ligands are highly efficient at ligand field splitting due to their strong π-acceptor properties. They exert a significant influence on the energy separation of the metal ion’s d orbitals causing low spins of the d orbital electrons.
Crystal Field Splitting of Octahedral 3d Metal Complexes
d1 Electron Configuration
In this case, we have a single electron in one of the five degenerate (having the same energy) d orbitals. When ligands approach the metal ion, they create an electrostatic field that repels the electron. As a result, the d orbital that points directly at the ligands (usually dx2–y2) experiences stronger repulsion, causing it to be raised in energy. Simultaneously, the remaining d orbitals (dxy, dxz, dyz, dz2) that are not directly pointing at the ligands are less affected and remain at a lower energy level. This energy difference between the raised and unaltered d orbitals is represented by Δo (delta-o). For d1 complexes, the electron occupies the lower-energy d orbital, resulting in a low-spin configuration.
For example: [Ti(H2O)6]3+ (hexaaquatitanium(III) ion); the titanium (Ti) ion is surrounded by six water (H2O) ligands. The d1 electron configuration implies that there is only one electron occupying one of the five degenerate d orbitals. When the ligands approach the titanium ion, they create an electrostatic field that repels the electron in the d orbitals. Due to this repulsion, the d orbital that points directly at the ligands (usually dx2–y2 orbital) experiences stronger repulsion, causing it to be raised in energy. On the other hand, the remaining d orbitals (dxy, dxz, dyz, dz2) that are not directly pointing at the ligands are less affected and remain at a lower energy level.
This energy difference between the raised dx2–y2 orbital and the unaltered dxy, dxz, dyz, and dz2 orbitals is represented by Δo (delta-o), which is known as the crystal field splitting parameter. The magnitude of Δo depends on the nature and coordination geometry of the ligands.
This splitting of the d orbitals in d1 metal complexes is responsible for the characteristic colors observed in transition metal complexes. The energy difference between the raised and unaltered orbitals affects the absorption and reflection of light, resulting in the appearance of specific colors.
d2 Electron Configuration
There are two electrons occupying the d orbitals. As the ligands approach the metal ion, they create an electrostatic field that affects the energies of the d orbitals. The d orbitals split into two sets: the lower-energy set (t2g) and the higher-energy set (eg).
For the d2 configuration, both electrons occupy the same t2g orbitals (axial orbitals) (dxy, dxz, dyz) with parallel spins, following Hund’s rule. This results in a low-spin configuration since the energy required to pair up the electrons in the eg set is higher than the energy difference between the t2g and eg sets.
The magnitude of the splitting, or the energy difference between the t2g and eg sets, depends on the nature of the ligands. The Spectrochemical Series provides a ranking of ligands based on their ability to create a splitting. Ligands higher in the Spectrochemical Series, such as CN– and CO, are strong π-acceptors, leading to a larger splitting (Δo) between the t2g and eg sets.
Considering the example: [Fe(CN)6]4- (hexacyanoferrate(II) ion), the iron (Fe) ion is surrounded by six cyanide (CN–) ligands. The d2 electron configuration implies that there are two electrons occupying two of the five d orbitals as earlier explained.
When the ligands approach the iron ion, they create an electrostatic field that affects the energies of the d orbitals. The d orbitals split into two sets: the lower-energy set called t2g (which includes the dxy, dxz, and dyz orbitals) and the higher-energy set called eg (which includes the dx2–y2 and dz2 orbitals).
Surely, both electrons occupy the same t2g orbital with parallel spins, following Hund’s rule. This results in a low-spin configuration because the energy required to pair up the electrons in the eg set is higher than the energy difference between the t2g and eg sets.
In this example, the d2 electron configuration of the iron ion causes both electrons to occupy the same t2g orbital with parallel spins. The cyanide ligands, being strong π-acceptors and high in the spectrochemical series, create a significant splitting (Δo) between the t2g and eg sets. As a result, we have a low-spin configuration with the two electrons paired in the t2g set, while the eg set remains unoccupied.
d3 Electron Configuration
In a d3 electronic configuration, there are three electrons occupying the five degenerate (having the same energy) d orbitals of a transition metal ion. When ligands approach the metal ion, they create an electrostatic field that affects the energies of the d orbitals. According to Crystal Field Theory (CFT), the d orbitals split into two sets: the lower-energy set (t2g) and the higher-energy set (eg).
Considering [Fe(H2O)6]3+ (hexaaquairon(III) ion), the iron (Fe) ion is surrounded by six water (H2O) ligands.
When the water ligands approach the iron ion, they generate an electrostatic field that influences the energies of the d orbitals. Since water ligands are weak field ligands, they cause a smaller splitting (Δo) between the t2g and eg sets compared to strong field ligands.
In the d3 configuration, all three d orbitals are occupied by electrons. According to CFT, the three electrons will prefer to occupy the lower-energy t2g set to minimize electron-electron repulsion, following Hund’s rule. As a result, in the [Fe(H2O)6]3+ complex, the three electrons will occupy the t2g orbitals with parallel spins, resulting in a low-spin configuration.
d4 Electron Configuration.
In a d4 electronic configuration, there are four electrons occupying the five degenerate (having the same energy) d orbitals of a transition metal ion. As ligands approach the metal ion, they create an electrostatic field that affects the energies of the d orbitals according to Crystal Field Theory (CFT).
In the case of d4 high-spin complexes, the d orbitals split into two sets: a lower-energy set (eg) and a higher-energy set (t2g) due to the repulsion from the ligand field. The energy difference between the eg and t2g sets is denoted as Δo (delta-o).
Considering [Mn(H2O)6]2+ (hexaqua manganese(II) ion), the manganese (Mn) ion is surrounded by six water (H2O) ligands.
As the water ligands approach the manganese ion, they generate an electrostatic field that influences the energies of the d orbitals. Since water ligands are weak field ligands, they cause a smaller splitting (Δo) between the eg and t2g sets compared to strong field ligands.
In the d4 high-spin configuration, all five d orbitals are singly occupied by electrons, following Hund’s rule. This occurs because the energy required to pair up two electrons in the same orbital (forming a low-spin configuration) is higher than the energy difference between the eg and t2g sets. Therefore, in high-spin d4 complexes, such as [Mn(H2O)6]2+, we find one electron in each of the five d orbitals, resulting in a high-spin configuration.
d5 Electronic Configuration
In the high-spin d5 configuration, the five electrons occupy each of the five available d orbitals with parallel spins, following Hund’s rule. This arrangement maximizes the number of unpaired electrons, resulting in a high-spin configuration. In this case, the energy difference (Δo) between the t2g and eg sets is relatively small compared to low-spin complexes.
The extent of splitting, Δo, depends on the nature of the ligands involved. Ligands lower in the spectrochemical series, such as fluoride (F–) and chloride (Cl–), possess weak field strengths. These weak field ligands do not strongly interact with the metal ion and induce a small splitting. Consequently, the electrons in the d orbitals can occupy higher energy orbitals, maintaining a high-spin configuration.
Let’s consider the example of [Fe(H2O)6]2+ (hexaaquairon(II) ion). In this complex, the iron (Fe) ion is surrounded by six water (H2O) ligands. Since water ligands are weak field ligands, they induce a small splitting (Δo) between the t2g and eg sets. As a result, the five d electrons of iron (Fe) occupy each of the five d orbitals with parallel spins, leading to a high-spin configuration.
d6 Electronic Configuration
The term “d6 high spin electronic configuration” refers to the arrangement of electrons in the d orbitals of a transition metal ion when it has a total of 6 electrons in those orbitals. In this configuration, the electrons occupy the d orbitals in a way that maximizes their spin, resulting in unpaired electrons. This configuration occurs when the crystal field splitting energy (the energy difference between the lower and higher energy d orbitals) is relatively small, and the energy required to pair up electrons exceeds the energy gained from the crystal field splitting.
For example. the hexaquairon(II) ion, [Fe(H2O)6]2+, the iron (Fe) ion is in the +2 oxidation state, resulting in a d6 electronic configuration. The d6 high spin electronic configuration for the Fe(II) ion can be represented as:
dxy↑ dxz↑ dyz↑ dz2↑↑ dx2-y2↑
In this configuration, the three lower energy d orbitals (dxy, dxz, dyz) are singly occupied with one electron each, and the two higher energy d orbitals (dz2 and dx2-y2) have two unpaired electrons each.
The hexaquairon(II) ion is surrounded by six water ligands (H2O) that interact with the d orbitals of the iron ion. The water ligands generate a weak crystal field, resulting in a small energy splitting between the lower and higher energy d orbitals.
Since the crystal field splitting energy is relatively small, the energy required to pair up the unpaired electrons in the higher energy orbitals exceeds the energy gained from the crystal field splitting. As a result, the unpaired electrons occupy the higher energy orbitals, leading to the d6 high spin electronic configuration.
The hexaquairon(II) ion is pale green in color and exhibits paramagnetic behavior due to the presence of unpaired electrons. The unpaired electrons can interact with an external magnetic field, leading to a weak attraction.
d7 Electronic Configuration
In this configuration, the electrons occupy the d orbitals in a way that maximizes their spin, resulting in unpaired electrons. This configuration occurs when the crystal field splitting energy (the energy difference between the lower and higher energy d orbitals) is relatively small, and the energy required to pair up electrons exceeds the energy gained from the crystal field splitting.
Let’s consider an example of a coordination compound that exhibits a d7 high spin electronic configuration. One such example is the heptacyanomanganate(III) ion, [Mn(CN)7]4- with the manganese (Mn) ion is in the +3 oxidation state, resulting in a d7 electronic configuration.
In this configuration, the three lower energy d orbitals (dxy, dxz, dyz) are singly occupied with one electron each, and the two higher energy d orbitals (dz2 and dx2-y2) have two unpaired electrons each. The dy2 orbital, which is perpendicular to the xy-plane, also has one unpaired electron.
The heptacyanomanganate(III) ion is surrounded by seven cyanide ligands (CN–) that interact with the d orbitals of the manganese ion. The cyanide ligands generate a weak crystal field, resulting in a small energy splitting between the lower and higher energy d orbitals.
The heptacyanomanganate(III) ion is dark purple in color and exhibits paramagnetic behavior due to the presence of unpaired electrons. The unpaired electrons can interact with an external magnetic field, leading to a weak attraction.
d8 High Spin Electronic Configuration
In this configuration, the electrons occupy the d orbitals in a way that maximizes their spin, resulting in unpaired electrons. This configuration occurs when the crystal field splitting energy (the energy difference between the lower and higher energy d orbitals) is relatively small, and the energy required to pair up electrons exceeds the energy gained from the crystal field splitting.
In the octahedral complex [Ni(H2O)6]2+, the nickel (Ni) ion is in the +2 oxidation state, resulting in a d8 electronic configuration.
The three lower energy d orbitals (dxy, dxz, dyz) are singly occupied with one electron each, and the two higher energy d orbitals (dz2 and dx2-y2) have three unpaired electrons each. The dyz and dxz orbitals each have an unpaired electron originating from the 4s orbital. The dxy and dx2-y2 orbitals each have one unpaired electron originating from the 3d orbital.
The nickel ion in [Ni(H2O)6]2+ is surrounded by six water ligands (H2O) that interact with the d orbitals of the nickel ion. The water ligands generate a weak crystal field, resulting in a small energy splitting between the lower and higher energy d orbitals.
Since the crystal field splitting energy is relatively small, the energy required to pair up the unpaired electrons in the higher energy orbitals exceeds the energy gained from the crystal field splitting. As a result, the unpaired electrons occupy the higher energy orbitals, leading to the d8 high spin electronic configuration.
The [Ni(H2O)6]2+ complex is usually pale green in color and exhibits paramagnetic behavior due to the presence of unpaired electrons. The unpaired electrons can interact with an external magnetic field, leading to a weak attraction.
d9 High Spin Electronic Configuration
The electrons in this configuration occupy the d orbitals in a way that maximizes their spin, resulting in unpaired electrons.
Let’s consider an example of a coordination compound that exhibits a d9 high spin electronic configuration.
In the octahedral complex [Cu(NH3)4(H2O)2]2+, the copper (Cu) ion is in the +2-oxidation state, resulting in a d9 electronic configuration.
In this configuration, the three lower energy d orbitals (dxy, dxz, dyz) are singly occupied with one electron each, and the two higher energy d orbitals (dz2 and dx2-y2) have four unpaired electrons each. The dxz and dyz orbitals each have an unpaired electron originating from the 4s orbital. The dxy, dz2, and dx2-y2 orbitals each have one unpaired electron originating from the 3d orbital.
The [Cu(NH3)4(H2O)2]2+ complex consists of four ammonia ligands (NH3) and two water ligands (H2O) surrounding the copper ion. These ligands interact with the d orbitals of the copper ion, generating a weak crystal field.
The [Cu(NH3)4(H2O)2]2+ complex is usually pale blue in color and exhibits paramagnetic behavior due to the presence of unpaired electrons. The unpaired electrons can interact with an external magnetic field, leading to a weak attraction.
d10 Electronic Configuration
In this configuration, the electrons occupy the available orbitals with parallel spins before pairing occurs. This configuration is commonly observed for transition metal ions in the +2 oxidation state.
An example of a coordination compound that exhibits the d10 electronic configuration] is [Zn(NH3)4]2+. In this complex, the zinc (Zn) ion is in the +2 oxidation state, and it has a d10 electronic configuration.
In this configuration, all five d orbitals (dxy, dxz, dyz, dz2, dx2-y2) are fully occupied with two electrons each, resulting in a total of 10 electrons.
The [Zn(NH3)4]2+ complex consists of four ammonia ligands (NH3) surrounding the zinc ion. These ligands do not cause any crystal field splitting or energy difference between the d orbitals since the Zn(II) ion does not have any unpaired electrons. Therefore, the complex does not exhibit any color or magnetic properties associated with d-d transitions or magnetic moments.
