Coordination Compounds Class 12 Chemistry Chapter 5 Notes

September 26, 2023 / By / CBSE Class 12 Chemistry Notes

Coordination Compounds Class 12 Chemistry Chapter 5 Notes

Werner’s Theory of Coordination Compounds

Alfred Werner’s groundbreaking work in the field of coordination chemistry laid the foundation for our understanding of the structures and bonding in coordination compounds. Here are some key points about his contributions and the principles he proposed:

  1. Alfred Werner: Alfred Werner was a Swiss chemist who lived from 1866 to 1919. He is renowned for his pioneering work in coordination chemistry, particularly in understanding the structures of coordination compounds.
  2. Primary and Secondary Valence: Werner introduced the concept of primary and secondary valence for metal ions in coordination compounds. Primary valence represents the ionizable valence of the metal, which is satisfied by negative ions. Secondary valence, on the other hand, is non-ionizable and is satisfied by neutral molecules or negative ions.
  3. Coordination Number: The coordination number is equal to the secondary valence and represents the number of groups (atoms or molecules) directly bonded to the central metal ion. It is fixed for a particular metal in a given compound.
  4. Coordination Entities: Werner proposed that the species within square brackets in chemical formulas represent coordination entities or complexes. These complexes have characteristic spatial arrangements, which we now refer to as coordination polyhedra.
  5. Isomerism: Werner’s work also led to the recognition of isomerism in coordination compounds. Isomers are compounds with the same empirical formula but distinct properties due to differences in the spatial arrangement of atoms or groups around the central metal ion.
  6. Geometrical Shapes: Werner noted that octahedral, tetrahedral, and square planar geometrical shapes are more common in coordination compounds of transition metals. For example, [Co(NH3)6]3+ and [CoCl(NH3)5]2+ are octahedral entities, while [Ni(CO)4] and [PtCl4]2– are tetrahedral and square planar, respectively.

Werner’s coordination theory revolutionized the understanding of coordination compounds, and his ideas continue to be fundamental in the field of inorganic chemistry. Coordination compounds are vital in various applications, including catalysis, materials science, and bioinorganic chemistry.

Difference between a double salt and a complex

The main difference between a double salt and a complex lies in how they behave when dissolved in water:

  1. Double Salt:
    • Double salts are compounds formed by the combination of two different salts or ionic compounds in a fixed stoichiometric ratio.
    • When dissolved in water, double salts dissociate completely into their constituent ions. In other words, they undergo complete ionization in aqueous solution.
    • The dissociation of double salts results in the formation of simple ions in solution.
    • Examples of double salts include carnallite (KCl.MgCl2.6H2O), Mohr’s salt (FeSO4.(NH4)2SO4.6H2O), and potash alum (KAl(SO4)2.12H2O).
  2. Complex:
    • Complexes are compounds formed by the coordination of metal ions with ligands (molecules or ions) through coordinate covalent bonds.
    • When dissolved in water, complexes do not dissociate into their constituent metal ions and ligands. Instead, they remain intact as complex ions.
    • Complex ions are a single entity with a central metal ion bonded to one or more ligands. They do not break down into separate metal ions and ligands in solution.
    • An example of a complex is [Fe(CN)6]4– in the compound K4[Fe(CN)6]. When this complex is dissolved in water, it remains as the intact [Fe(CN)6]4– ion and does not release Fe2+ or CN– ions into solution.

Coordination Entity

A coordination entity, also known as a coordination complex or simply a complex, is a molecular species or chemical compound consisting of a central metal atom or ion bonded to a fixed number of ions, molecules, or groups called ligands. The central metal atom or ion is typically a transition metal, and the ligands are responsible for coordinating to the metal through coordinate covalent bonds.

Here are some key points about coordination entities:

  1. Central Metal Atom/Ion: The central metal atom or ion is the focal point of the coordination entity. It serves as the anchor to which the ligands attach. Common examples of central metal ions include Fe²⁺, Cu²⁺, Co³⁺, and Pt⁴⁺, among others.
  2. Ligands: Ligands are molecules, ions, or groups of atoms that coordinate to the central metal through one or more donor atoms. Ligands provide pairs of electrons to form coordinate bonds with the metal. Common ligands include water (H₂O), ammonia (NH₃), cyanide (CN⁻), carbon monoxide (CO), and various organic molecules.
  3. Coordination Number: The coordination number is the total number of coordinate bonds formed between the central metal and the surrounding ligands. It represents how many positions around the central metal are occupied by ligands. For example, if a complex has a coordination number of 6, it means there are six ligands bonded to the central metal.
  4. Charge: Coordination entities can carry an overall charge, which is determined by the charge on the central metal ion and the charge on the ligands. The charges must balance to maintain the overall neutrality of the complex.

Examples of coordination entities include:

  • [CoCl₃(NH₃)₃]: In this complex, cobalt (Co) is the central metal ion, and it is coordinated to three ammonia (NH₃) molecules and three chloride (Cl⁻) ions. The coordination number is 6.
  • [Ni(CO)₄]: In this complex, nickel (Ni) is the central metal atom, and it is coordinated to four carbon monoxide (CO) ligands. The coordination number is 4.
  • [Fe(CN)₆]⁴⁻: In this complex, iron (Fe) is the central metal ion, and it is coordinated to six cyanide (CN⁻) ions. The coordination number is 6.
  • [Co(NH₃)₆]³⁺: In this complex, cobalt (Co) is the central metal ion, and it is coordinated to six ammonia (NH₃) molecules. The coordination number is 6.

Central atom/ion

  • Definition: In a coordination entity, the central atom/ion is the one to which a fixed number of ions/groups are bound in a definite geometrical arrangement.
  • Examples:
    1. [NiCl2(H2O)4]: Central atom/ion is Ni2+.
    2. [CoCl(NH3)5]2+: Central atom/ion is Co3+.
    3. [Fe(CN)6]3–: Central atom/ion is Fe3+.
  • Role: Central atoms/ions are also known as Lewis acids, meaning they can accept electron pairs in coordination bonds. They form the core of coordination complexes, surrounded by ligands.

Ligands

  • Definition: Ligands are ions or molecules that are bound to the central atom/ion in a coordination entity. They form coordination bonds with the central atom/ion and surround it in a specific arrangement.
  • Types of Ligands:
    1. Unidentate Ligands: These ligands bind to the central atom/ion through a single donor atom. Examples include Cl–, H2O, and NH3.
    2. Didentate Ligands: Ligands that can bind through two donor atoms are called didentate ligands. Examples include ethane-1,2-diamine (H2NCH2CH2NH2) and oxalate (C2O4^2–).
    3. Polydentate Ligands: Ligands with several donor atoms in a single molecule are called polydentate ligands. An important example is ethylenediaminetetraacetate ion (EDTA^4–), which can bind through two nitrogen and four oxygen atoms to a central metal ion.
    4. Chelate Ligands: When a di- or polydentate ligand uses its two or more donor atoms simultaneously to bind a single metal ion, it is termed a chelate ligand. The number of such ligating groups is called the denticity of the ligand. Chelate complexes are generally more stable than those containing unidentate ligands.
    5. Ambidentate Ligands: Some ligands have two different donor atoms and can bind to the central atom/ion through either of them. Examples include the NO2– ion (which can coordinate through nitrogen or oxygen) and the SCN– ion (which can coordinate through sulfur or nitrogen).
  • Chelate Ligands: These ligands form chelate complexes, which are known for their increased stability compared to complexes with unidentate ligands.
  • Ambidentate Ligands: Ligands like NO2– and SCN– can coordinate through different donor atoms, providing flexibility in bonding to the central metal atom/ion.

Coordination Number

  • Definition: The coordination number (CN) of a metal ion in a complex is defined as the number of ligand donor atoms directly bonded to the central metal ion.
  • Examples:
    1. In [PtCl6]2–, the coordination number of Pt is 6 because it is directly bonded to six Cl atoms.
    2. In [Ni(NH3)4]2+, the coordination number of Ni is 4 because it is directly bonded to four NH3 ligands.
    3. In [Fe(C2O4)3]3– and [Co(en)3]3+, both Fe and Co have a coordination number of 6 because they are bonded to six donor atoms. C2O4^2– and en (ethane-1,2-diamine) are didentate ligands, each contributing two donor atoms.
  • Considerations: The coordination number is determined solely by the number of sigma (σ) bonds formed between the ligands and the central metal atom/ion. Pi (π) bonds, if present, are not counted when determining the coordination number.

Coordination Sphere

  • Coordination Sphere: The coordination sphere in a coordination complex consists of the central metal atom/ion and the ligands attached to it. This entire unit is enclosed within square brackets. For example, in the complex K4[Fe(CN)6], the coordination sphere is [Fe(CN)6]4–.
  • Counter Ions: Counter ions, also known as counterions or countercharges, are ionizable groups located outside the square brackets that balance the charge of the coordination sphere. These counter ions ensure that the complex as a whole is electrically neutral. In the given example, K+ is the counter ion that balances the charge of [Fe(CN)6]4–, resulting in an overall neutral complex, K4[Fe(CN)6]. The counter ions are necessary to maintain charge neutrality in the compound.

Coordination Polyhedron

  • Definition: The spatial arrangement of ligand atoms directly attached to the central atom/ion defines a coordination polyhedron around the central atom. This polyhedron represents the three-dimensional shape of the complex.
  • Common Coordination Polyhedra:
    1. Octahedral: In an octahedral coordination polyhedron, the central atom/ion is surrounded by six ligand atoms or groups arranged at the corners of an octahedron. An example is [Co(NH3)6]3+.
    2. Tetrahedral: In a tetrahedral coordination polyhedron, the central atom/ion is surrounded by four ligand atoms or groups arranged at the corners of a tetrahedron. An example is [Ni(CO)4].
    3. Square Planar: In a square planar coordination polyhedron, the central atom/ion is surrounded by four ligand atoms or groups arranged in a flat, square plane. An example is [PtCl4]2–.
  • Visualization: Fig. 5.1 illustrates the shapes of these different coordination polyhedra, providing a visual representation of their spatial arrangements.

Oxidation Number of Central Atom

  • Definition: The oxidation number of the central atom in a complex is defined as the charge it would carry if all the ligands are removed, along with the electron pairs that are shared with the central atom.
  • Representation: The oxidation number of the central atom is typically represented by a Roman numeral in parentheses following the name of the coordination entity.
  • Example: For instance, consider the complex [Cu(CN)4]3–. To determine the oxidation number of copper (Cu), you would imagine removing all the ligands (CN–) along with the electron pairs that are shared with copper. In this case, copper would carry a charge of +1, so it is written as Cu(I).

The oxidation number of the central atom is a crucial concept in coordination chemistry, as it helps in understanding the electron transfer and redox reactions involving coordination complexes.

Homoleptic and Heteroleptic Complexes

  • Homoleptic Complexes: In homoleptic complexes, the central metal atom/ion is bonded to only one type of donor group or ligand. These complexes have a uniform ligand environment around the central atom. An example is [Co(NH3)6]3+, where cobalt (Co) is bonded to six ammonia (NH3) ligands.
  • Heteroleptic Complexes: Heteroleptic complexes are characterized by the central metal atom/ion being bonded to more than one type of donor group or ligand. These complexes have a diverse ligand environment around the central atom, with different types of ligands. An example is [Co(NH3)4Cl2]+, where cobalt (Co) is bonded to both ammonia (NH3) and chloride (Cl–) ligands.

The distinction between homoleptic and heteroleptic complexes is essential in understanding the structural diversity and properties of coordination compounds. Heteroleptic complexes often exhibit more complex geometries and reactivity due to the presence of different ligand types.

Nomenclature of Coordination Compound

Nomenclature in coordination chemistry is crucial for accurately describing the composition and structure of coordination compounds, especially when dealing with various isomers. The International Union of Pure and Applied Chemistry (IUPAC) provides a standardized set of rules and recommendations for naming coordination entities. Here are some key points about the nomenclature of coordination compounds:

  1. Central Atom/Ion: The name of the central atom/ion is usually placed first in the formula. For example, in [Co(NH3)6]3+, “Co” represents the central cobalt atom.
  2. Ligands: Ligands are named before the central atom/ion. The names of ligands end with “o” for anions and “amine” for neutral ligands. For example, “NH3” is named as “ammine,” and “Cl–” is named as “chloro.”
  3. Ligand Multiplicity: The number of each type of ligand is indicated by prefixes such as “di-” for two, “tri-” for three, and so on. For example, [Co(NH3)4Cl2]+ is named “tetrachloridocobaltate(III) ammine.”
  4. Charged Coordination Entities: The charge of the coordination entity is indicated in Roman numerals in parentheses following the name of the central atom/ion. For example, [Fe(CN)6]4– is named “hexacyanidoferrate(II).”
  5. Neutral Coordination Entities: In cases where the coordination entity is neutral, the name of the central atom/ion is followed by the names of the ligands, listed in alphabetical order. For example, [PtCl2(NH3)2] is named “diamminedichloridoplatinum(II).”
  6. Isomerism: Nomenclature helps distinguish between various isomers. Isomers are given different names to reflect their distinct structures.
  7. Use of Prefixes: When there are multiple ligands of the same type, numerical prefixes such as “bis-” for two, “tris-” for three, and so on, are used. For example, [Co(NH3)3Cl3] is named “triamminetrichloridocobalt(III).”

Formulas of Mononuclear Coordination Entities

In coordination chemistry, writing the formula of a mononuclear coordination entity (a single central metal atom/ion bonded to ligands) follows specific rules to provide a clear and concise representation. These rules help in accurately conveying the composition of the complex. Here are the guidelines for writing such formulas:

  1. Central Atom/Ion: The central metal atom/ion is listed first in the formula.
  2. Ligands: Ligands are listed next, in alphabetical order. The charge of the ligand does not affect its alphabetical placement.
  3. Polydentate Ligands: When polydentate ligands are present, they are also listed alphabetically. If ligands are abbreviated, the first letter of the abbreviation is used to determine their position in the alphabetical order.
  4. Parentheses and Square Brackets: The formula for the entire coordination entity, whether charged or not, is enclosed in square brackets “[ ]”. Polyatomic ligands are enclosed in parentheses “()”. Ligand abbreviations, if used, are also enclosed in parentheses.
  5. Spacing: There should be no space between the ligands and the central metal atom/ion within the coordination sphere.
  6. Charge Indication: If the coordination entity is charged, the charge is indicated outside the square brackets as a right superscript with the number placed before the sign. For example, [Co(CN)6]3– or [Cr(H2O)6]3+.
  7. Balancing Charges: The charge of the cation(s) (if present) is balanced by the charge of the anion(s) (if present) to ensure that the overall complex is electrically neutral.

Naming of Mononuclear Coordination Compounds

Nomenclature in coordination chemistry follows systematic rules to describe the composition and structure of coordination compounds accurately. Here are the key rules along with examples:

  1. Order of Naming:
    • The cation is named first in both positively and negatively charged coordination entities.
    • Ligands are named in alphabetical order before the name of the central atom/ion (opposite of writing formulas).
  2. Ligand Names:
    • Anionic ligands end in “-o.”
    • Neutral and cationic ligands have the same names, except for specific cases: “aqua” for H2O, “ammine” for NH3, “carbonyl” for CO, and “nitrosyl” for NO.
  3. Numerical Prefixes:
    • Prefixes “mono,” “di,” “tri,” etc., indicate the number of individual ligands in the coordination entity.
    • When ligand names include numerical prefixes, terms like “bis,” “tris,” “tetrakis,” etc., are used. The ligand to which they refer is placed in parentheses.
  4. Oxidation State:
    • The oxidation state of the metal in a cation, anion, or neutral coordination entity is indicated by a Roman numeral in parentheses.
  5. Naming of Metal:
    • If the complex ion is a cation, the metal is named the same as the element (e.g., “cobalt” for Co).
    • If the complex ion is an anion, the metal name ends with “-ate” (e.g., “cobaltate” for Co). Latin names are used for some metals (e.g., “ferrate” for Fe).
  6. Neutral Complex Molecule:
    • The name of a neutral complex molecule is similar to that of the complex cation.

Examples of Nomenclature:

  1. [Cr(NH3)3(H2O)3]Cl3 is named as:
    • Triamminetriaquachromium(III) chloride
    • Explanation: The cation is inside the square brackets, and it’s a cation. Ligands are named alphabetically. The charge on the complex ion is +3, and this determines the oxidation number of chromium, which is also +3.
  2. [Co(H2NCH2CH2NH2)3]2(SO4)3 is named as:
    • Tris(ethane-1,2-diamine)cobalt(III) sulfate
    • Explanation: Sulphate is the counter anion. The complex cation has a charge of +3, and the oxidation number of cobalt is +3.
  3. [Ag(NH3)2][Ag(CN)2] is named as:
    • Diamminesilver(I) dicyanidoargentate(I)
    • Explanation: Two different cations are present, and they both contain silver (Ag) ions with different ligands. The names reflect this differentiation.

Isomerism in Coordination Compounds

Isomerism in coordination compounds is a phenomenon where two or more compounds have the same chemical formula but differ in the arrangement of atoms or ligands, resulting in distinct physical or chemical properties. There are two principal types of isomerism in coordination compounds:

(a) Stereoisomerism
(i) Geometrical isomerism
(ii) Optical isomerism

(b) Structural isomerism
(i) Linkage isomerism
(ii) Coordination isomerism
(iii) Ionisation isomerism
(iv) Solvate isomerism

Geometric Isomerism

Geometrical isomerism is a type of stereoisomerism that arises due to different possible spatial arrangements of ligands in coordination compounds. It is particularly significant in complexes with coordination numbers of 4 and 6. Let’s explore some examples and characteristics of geometrical isomerism:

  1. Square Planar Complexes:
    • In square planar complexes of the form [MX2L2], where X and L are unidentate ligands, geometrical isomerism can occur.
    • For example, in a square planar complex, the two ligands X can be arranged adjacent to each other in a cis isomer or opposite to each other in a trans isomer.
  2. Other Square Planar Complexes:
    • Complexes of the form MABXL (where A, B, X, and L are unidentate ligands) can exhibit three isomers: two cis isomers and one trans isomer.
    • These isomers have distinct arrangements of the ligands around the central metal atom.
  3. Tetrahedral Complexes:
    • Geometrical isomerism is not possible for complexes with tetrahedral geometry because the arrangement of ligands in a tetrahedron does not lead to different spatial isomers.
  4. Octahedral Complexes:
    • In octahedral complexes of the form [MX2L4], where X are unidentate ligands, geometrical isomerism can occur.
    • The two ligands X can be oriented cis or trans to each other, leading to distinct isomeric forms.
  5. Didentate Ligands:
    • Geometrical isomerism can also arise when didentate ligands (e.g., ethylenediamine, en) are present in complexes of the form [MX2(L-L)2].
    • The arrangement of the didentate ligands can lead to different geometric isomers.
  6. Octahedral Coordination Entities:
    • In octahedral coordination entities of the type [Ma3b3], such as [Co(NH3)3(NO2)3], geometrical isomerism can occur.
    • Two specific types of isomers are observed:
      • Facial (fac) Isomer: Three donor atoms of the same ligands occupy adjacent positions at the corners of an octahedral face.
      • Meridional (mer) Isomer: The positions of the three donor atoms are around the meridian of the octahedron.

Optical Isomerism

Optical isomerism, also known as enantiomerism, is a type of stereoisomerism observed in coordination compounds where two isomers are non-superimposable mirror images of each other. These isomers are called enantiomers. Here are the key characteristics of optical isomerism:

  1. Chiral Molecules: Optical isomers are non-superimposable mirror images, meaning they cannot be perfectly overlapped or aligned. Molecules or ions that exhibit this property are referred to as chiral.
  2. Enantiomers: The two forms of optical isomers are called dextro (d) and laevo (l). They are labeled based on their ability to rotate the plane of polarized light in a polarimeter.
    • Dextro (d): Dextro isomers rotate the plane of polarized light to the right (clockwise).
    • Laevo (l): Laevo isomers rotate the plane of polarized light to the left (counterclockwise).
  3. Common in Octahedral Complexes with Didentate Ligands: Optical isomerism is commonly observed in octahedral coordination complexes that involve didentate ligands. Didentate ligands have two donor atoms, which create chiral environments around the central metal atom/ion.
  4. Example: In a coordination entity of the type [PtCl2(en)2]2+, where en represents ethylenediamine (a didentate ligand), only the cis-isomer exhibits optical activity. This means that it can rotate the plane of polarized light either to the right (dextro) or to the left (laevo) depending on its specific spatial arrangement.

Linkage Isomerism

Linkage isomerism is a type of structural isomerism that arises in coordination compounds containing ambidentate ligands. Ambidentate ligands are ligands that have multiple donor atoms and can bind to the central metal atom/ion through different atoms. A common example of linkage isomerism involves the thiocyanate ligand, NCS–, which can bind through either the nitrogen or sulfur atom.

Here’s how linkage isomerism works:

  1. Ambidentate Ligands: Ambidentate ligands have more than one potential binding site. In the case of NCS–, it can bind to the central metal through either the nitrogen atom (–NCS) or the sulfur atom (–SCN).
  2. Examples of Linkage Isomerism: The phenomenon of linkage isomerism can be illustrated with examples. One such example is the complex [Co(NH3)5(NO2)]Cl2. In this complex, the nitrite (NO2–) ligand is ambidentate and can bind in two different ways:
    • In the red form, the nitrite ligand is bound to the central cobalt atom through oxygen (–ONO).
    • In the yellow form, the nitrite ligand is bound to the central cobalt atom through nitrogen (–NO2).

These two forms of the complex are linkage isomers because they have the same chemical formula but different arrangements of ligands due to the variation in how the ambidentate ligand binds.

Coordination Isomerism

Coordination isomerism is a type of structural isomerism that arises from the interchange of ligands between cationic and anionic entities of different metal ions present in a complex. This phenomenon results in two different coordination isomers of the same compound. Here’s how coordination isomerism works:

  1. Different Metal Ions: Coordination isomerism occurs when there are two different metal ions present in a complex. Each metal ion binds to a different set of ligands.
  2. Interchange of Ligands: In coordination isomerism, the ligands attached to one metal ion are exchanged with the ligands attached to the other metal ion, creating two distinct isomeric forms.
  3. Example: An example of coordination isomerism can be seen in the complex [Co(NH3)6][Cr(CN)6], where NH3 ligands are bound to the Co3+ ion, and CN– ligands are bound to the Cr3+ ion. In its coordination isomer, [Cr(NH3)6][Co(CN)6], the NH3 ligands are bound to the Cr3+ ion, and the CN– ligands are bound to the Co3+ ion.

In these two coordination isomers, the arrangement of ligands around the central metal ions is different, leading to distinct chemical and physical properties. Coordination isomerism is one of the ways in which coordination compounds with multiple metal ions can exhibit structural diversity.

Ionisation Isomerism

Ionisation isomerism is a type of structural isomerism that occurs in coordination compounds when the counterion in a complex salt can also act as a potential ligand. In ionisation isomerism, the counterion and one of the ligands can exchange positions, resulting in two distinct isomeric forms of the same compound. Here’s how ionisation isomerism works:

  1. Counterion as a Potential Ligand: In some coordination compounds, the counterion (anion or cation) has the ability to act as a ligand and bind to the central metal ion. This is possible because the counterion may have donor atoms capable of forming coordinate bonds.
  2. Exchange of Ligands: Ionisation isomerism arises when the counterion displaces one of the ligands in the complex, and the displaced ligand becomes the new counterion. This interchange of ligands and counterions leads to two different coordination isomers.
  3. Example: An example of ionisation isomerism can be observed in the isomers [Co(NH3)5(SO4)]Br and [Co(NH3)5Br]SO4. In the first isomer, the SO4 ion acts as a counterion, and in the second isomer, the Br ion acts as a counterion. The exchange between SO4 and Br results in two distinct coordination isomers.

These two coordination isomers have different arrangements of ligands and counterions, leading to variations in their chemical properties. Ionisation isomerism is an intriguing aspect of coordination chemistry that highlights the versatility of ligands and counterions in complex coordination compounds.

Solvate Isomerism

Solvate isomerism, also known as hydrate isomerism when water is involved as the solvent, is a type of structural isomerism observed in coordination compounds. Solvate isomerism is similar to ionisation isomerism in that it involves variations in the presence or absence of solvent molecules as ligands. In solvate isomerism, the coordination compound can exist in two forms that differ in whether solvent molecules are directly bonded to the central metal ion or merely present as free solvent molecules in the crystal lattice. Here’s how solvate isomerism works:

  1. Involvement of Solvent Molecules: Solvate isomerism typically arises when the coordination compound involves solvent molecules (commonly water, in the case of hydrate isomerism) in its structure.
  2. Direct Bonding vs. Presence in the Lattice: In one form of the isomer, the solvent molecules are directly bonded to the central metal ion, forming coordinate bonds. In the other form, the solvent molecules are not directly bonded but are present as free solvent molecules within the crystal lattice.
  3. Example: An example of solvate isomerism can be illustrated with the aqua complex [Cr(H2O)6]Cl3 (violet) and its solvate isomer [Cr(H2O)5Cl]Cl2·H2O (grey-green). In the first complex, all six water molecules are directly bonded to the chromium ion, while in the solvate isomer, one water molecule is not directly bonded but is present as a free solvent molecule in the crystal lattice.

These two coordination isomers exhibit differences in their coordination environments and properties due to the presence or absence of direct bonds between the central metal ion and solvent molecules. Solvate isomerism provides insights into how solvent molecules can interact with coordination compounds and impact their behavior.

Bonding in Coordination Compounds

Coordination compounds exhibit unique bonding features that were first described by Alfred Werner. While Werner’s theory was groundbreaking, it had limitations and couldn’t answer some fundamental questions about coordination compounds. Over time, various theories and approaches have been developed to explain the nature of bonding in coordination compounds. Two important theories are the Valence Bond Theory (VBT) and the Crystal Field Theory (CFT).

Valence Bond Theory

VBT is used to explain the bonding in coordination compounds by considering the hybridization of metal atomic orbitals and their overlap with ligand orbitals. This leads to the formation of coordination bonds and the prediction of complex geometry and magnetic properties.

  1. Hybridization of Metal Orbitals:
    • According to VBT, the central metal atom or ion in a coordination compound can undergo hybridization of its atomic orbitals.
    • These hybridization possibilities include (n-1)d, ns, np, or ns, np, nd orbitals.
    • The goal of hybridization is to create a set of equivalent hybrid orbitals with specific geometries, such as octahedral, tetrahedral, or square planar.
  2. Overlap with Ligand Orbitals:
    • Once hybridization occurs, these hybrid orbitals are allowed to overlap with the orbitals of ligands.
    • The overlap leads to the formation of coordination bonds between the metal and the ligands.
  3. Predicting Geometry:
    • VBT can often predict the geometry of a coordination complex based on the complex’s magnetic behavior.
    • Diamagnetic complexes, lacking unpaired electrons, tend to have octahedral or tetrahedral geometries.
    • Paramagnetic complexes, with unpaired electrons, can have various geometries.
  4. Example: Diamagnetic Octahedral Complex [Co(NH3)6]3+:
    • In this complex, cobalt (Co) is in a +3 oxidation state with the electronic configuration 3d6.
    • The hybridization scheme involves inner orbital hybridization, utilizing the inner d orbital (3d).
    • Six pairs of electrons, contributed by each NH3 ligand, occupy the resulting six hybrid orbitals.
    • The complex exhibits an octahedral geometry due to the arrangement of these hybrid orbitals.
    • Importantly, it is diamagnetic because it lacks unpaired electrons.
    • Such complexes are called “inner orbital” or “low spin” or “spin-paired” complexes.
  5. Paramagnetic Octahedral Complex [CoF6]3–:
    • In contrast, the complex [CoF6]3– is paramagnetic, indicating the presence of unpaired electrons.
    • This complex utilizes outer orbital hybridization, specifically the 4d orbitals.
    • The use of outer orbitals results in an octahedral geometry.
    • Complexes like this one are referred to as “outer orbital” or “high spin” or “spin-free” complexes.
  6. Hybrid Orbitals as Theoretical Concepts:
    • It’s important to note that hybrid orbitals are theoretical constructs used for mathematical calculations.
    • Hybridization is a mathematical manipulation of wave equations for atomic orbitals.
  7. Tetrahedral Complexes:
    • In tetrahedral complexes, one s orbital and three p orbitals from the central metal atom are hybridized to create four equivalent hybrid orbitals.
    • These hybrid orbitals are oriented tetrahedrally around the metal atom.
    • For example [NiCl4]2–, where nickel (Ni) is in a +2 oxidation state with an electronic configuration of 3d8.
    • Each Cl– ion donates a pair of electrons to form coordination bonds.
    • The complex is paramagnetic, indicating the presence of two unpaired electrons.
  8. Example of Diamagnetic Tetrahedral Complex [Ni(CO)4]:
    • Another example given is [Ni(CO)4], which has a tetrahedral geometry.
    • In this case, nickel is in a zero oxidation state (Ni2+) and has the electronic configuration 3d8.
    • Despite having the same tetrahedral geometry as [NiCl4]2–, [Ni(CO)4] is diamagnetic because nickel has no unpaired electrons.
  9. Square Planar Complexes:
    • In square planar complexes, the hybridization involved is dsp2.
    • For example [Ni(CN)4]2–, where nickel is in a +2 oxidation state with an electronic configuration of 3d8.
    • The hybridization scheme results in the creation of hybrid orbitals that can accept pairs of electrons from ligands.
    • The complex is diamagnetic, indicating the absence of unpaired electrons.
  10. Nature of Hybrid Orbitals:
    • It’s emphasized that hybrid orbitals are theoretical constructs used for mathematical calculations.
    • Hybridization is a mathematical manipulation of wave equations for the atomic orbitals involved in the formation of coordination bonds.
    • Hybrid orbitals are not physically existing entities but are useful tools for explaining bonding in coordination compounds.

Magnetic Properties of Coordination Compounds

  1. Measuring Magnetic Properties:
    • The magnetic moment of coordination compounds can be determined through magnetic susceptibility experiments.
    • Magnetic data analysis is a useful tool for understanding the electronic configurations and structures of metal complexes.
  2. Magnetic Behavior of First Transition Series:
    • Coordination compounds of first-row transition metals (e.g., Ti³⁺, V³⁺, Cr³⁺) with up to three electrons in the d orbitals have two vacant d orbitals available for octahedral hybridization with 4s and 4p orbitals.
    • The magnetic behavior of these metal ions and their coordination entities is similar.
  3. Complexities with More Than Three 3d Electrons:
    • When more than three 3d electrons are present (e.g., d⁴, d⁵, d⁶), obtaining the required pair of 3d orbitals for octahedral hybridization becomes complicated due to Hund’s rule.
    • For d⁴ (Cr²⁺, Mn³⁺), d⁵ (Mn²⁺, Fe³⁺), and d⁶ (Fe²⁺, Co³⁺) cases, the number of unpaired electrons varies (two, one, and zero unpaired electrons, respectively) due to pairing of 3d electrons.
  4. Anomalies in Magnetic Data:
    • An apparent anomaly exists in the magnetic data of certain coordination compounds.
    • For example, [Mn(CN)₆]³⁻ has a magnetic moment of two unpaired electrons, while [MnCl₆]³⁻ has a paramagnetic moment of four unpaired electrons.
    • Similarly, [Fe(CN)₆]³⁻ has a magnetic moment of one unpaired electron, while [FeF₆]³⁻ has a paramagnetic moment of five unpaired electrons.
    • [CoF₆]³⁻ is paramagnetic with four unpaired electrons, while [Co(C₂O₄)₃]³⁻ is diamagnetic.
  5. Explanation using Valence Bond Theory:
    • The anomalies are explained by valence bond theory in terms of the formation of inner orbital and outer orbital coordination entities.
    • [Mn(CN)₆]³⁻, [Fe(CN)₆]³⁻, and [Co(C₂O₄)₃]³⁻ are inner orbital complexes involving d²sp³ hybridization, resulting in paramagnetic and diamagnetic properties.
    • [MnCl₆]³⁻, [FeF₆]³⁻, and [CoF₆]³⁻ are outer orbital complexes involving sp³d² hybridization, resulting in paramagnetic properties corresponding to four, five, and four unpaired electrons.

Limitations of Valence Bond Theory

  1. Assumptions: VBT relies on a number of assumptions about the nature of bonding in coordination compounds. While these assumptions work well for many complexes, they are not always applicable to all situations.
  2. Quantitative Interpretation of Magnetic Data: VBT can provide qualitative explanations for the magnetic behavior of coordination compounds (e.g., the presence of unpaired electrons), but it does not offer a quantitative interpretation of magnetic data. It cannot predict the exact magnetic moments or behaviors in a precise manner.
  3. Explanation of Color: VBT does not provide a satisfactory explanation for the color exhibited by coordination compounds. The theory cannot account for the relationship between the color of a complex and its electronic structure.
  4. Quantitative Interpretation of Thermodynamic and Kinetic Stabilities: VBT does not offer a quantitative framework for predicting or interpreting the thermodynamic or kinetic stabilities of coordination compounds. It does not provide detailed insights into factors affecting stability.
  5. Prediction of Tetrahedral and Square Planar Structures: VBT cannot make precise predictions about whether a 4-coordinate complex will adopt a tetrahedral or square planar geometry. It lacks the predictive power to determine the exact structure in such cases.
  6. Distinction Between Weak and Strong Ligands: VBT does not effectively distinguish between weak and strong ligands. It cannot provide a quantitative ranking of ligand strengths or explain why certain ligands form more stable complexes than others.

Crystal Field Theory

Crystal Field Theory (CFT) is an important model in coordination chemistry that provides insights into the electronic structure and properties of coordination compounds. Here are some key points about CFT:

  1. Electrostatic Model: CFT is an electrostatic model that considers the metal-ligand bond as primarily ionic in nature. It focuses on the electrostatic interactions between the metal ion and the ligands.
  2. Treatment of Ligands: In CFT, ligands are treated as point charges (for anions) or point dipoles (for neutral molecules). This simplification allows for a mathematical description of the interactions between the metal and ligands.
  3. Degeneracy of d Orbitals: In an isolated gaseous metal atom/ion, the five d orbitals have the same energy and are degenerate, meaning they have identical energy levels.
  4. Effect of Ligands: When ligands surround the metal atom/ion in a complex, they create an asymmetrical negative field. This field affects the energy levels of the d orbitals, leading to the splitting or lifting of degeneracy.
  5. Pattern of Splitting: The pattern of splitting of the d orbitals depends on the nature of the crystal field generated by the ligands. There are two common types of crystal field splitting: octahedral and tetrahedral.
    • Octahedral Splitting: In octahedral complexes, the d orbitals split into two sets. Three of the d orbitals (dxy, dxz, and dyz) are raised in energy (t2g set), while the other two (dx2-y2 and dz2) are lowered in energy (eg set). This splitting leads to a d-orbital energy diagram with a gap between the t2g and eg sets.
    • Tetrahedral Splitting: In tetrahedral complexes, the d orbitals also split into two sets. However, the eg set is raised in energy, while the t2g set is lowered in energy. The splitting pattern is reversed compared to octahedral complexes.
  6. Explanation of Color: CFT provides a qualitative explanation for the color exhibited by transition metal complexes. The energy difference between the split d orbitals corresponds to specific wavelengths of light, resulting in the absorption of certain colors.
  7. Magnetic Properties: CFT helps explain the magnetic properties of coordination compounds. The number of unpaired electrons in the partially filled d orbitals influences the magnetic behavior, with paramagnetic complexes having unpaired electrons and diamagnetic complexes having no unpaired electrons.

Crystal Field Theory is a valuable tool for understanding the electronic structure, colors, and magnetic properties of coordination compounds. It provides a qualitative framework for explaining these phenomena based on the interactions between metal ions and ligands in a crystal field.

Crystal Field Splitting in Octahedral Coordination Entities

In octahedral coordination entities, the arrangement of six ligands around the central metal atom/ion leads to a splitting of the metal’s d orbitals due to repulsion between the metal’s d electrons and the electrons or negative charges on the ligands. This splitting is known as crystal field splitting, denoted as Δ₀ (Delta naught), and it results in the d orbitals adopting different energy levels.

Here are the key points regarding crystal field splitting in octahedral coordination entities:

  1. Repulsion Effect: Ligands surrounding the metal atom/ion create repulsion between their electrons and the metal’s d electrons. The extent of repulsion depends on the orientation of the metal d orbitals relative to the ligands.
  2. Energetic Changes: In an octahedral complex, the d orbitals that point directly toward the ligands (dx²-y² and dz²) experience greater repulsion and are raised in energy. These orbitals form the higher energy set, denoted as the eg set.
    • dx²-y² and dz² orbitals are directed along the axes toward the ligands.
  3. Lower Energy Orbitals: The d orbitals that point between the axes (dxy, dyz, and dxz) experience less repulsion and are lowered in energy. These orbitals form the lower energy set, denoted as the t₂g set.
    • dxy, dyz, and dxz orbitals are directed between the axes, away from the ligands.
  4. Energy Separation: The energy separation, Δ₀ (Delta naught), represents the difference in energy between the eg and t₂g sets of d orbitals.
  5. Factors Affecting Crystal Field Splitting: The magnitude of crystal field splitting, Δ₀, depends on the nature of the ligands and the charge on the metal ion.
  6. Spectrochemical Series: Ligands can be ranked in a spectrochemical series based on their ability to create strong or weak crystal field splitting. Ligands at the high end of the series, such as CN⁻ and CO, are strong field ligands, leading to larger Δ₀ values. Ligands at the low end, such as I⁻ and Br⁻, are weak field ligands, resulting in smaller Δ₀ values.
  7. Electron Distribution: The crystal field splitting affects the distribution of d electrons in the metal ion. For example, in d² and d³ configurations, the Hund’s rule dictates that electrons occupy the t₂g orbitals singly before pairing up.
  8. High Spin vs. Low Spin: The relative magnitude of Δ₀ and the pairing energy (P) determines whether a complex is in a high spin or low spin state:
    • If Δ₀ < P, the complex forms a high spin configuration, with unpaired electrons in the eg orbitals.
    • If Δ₀ > P, the complex forms a low spin configuration, with all d electrons paired in the t₂g orbitals.
  9. Stability of Complexes: Calculations show that d⁴ to d⁷ coordination entities are more stable for strong field ligands, as they prefer low spin configurations. Conversely, weak field ligands favor high spin configurations.

Crystal Field Splitting in Tetrahedral Coordination Entities

In tetrahedral coordination entities, the splitting of d orbitals differs from that in octahedral complexes. Here are the key points regarding crystal field splitting in tetrahedral coordination entities:

  1. Inverted Splitting: In tetrahedral coordination, the splitting of the metal’s d orbitals is inverted compared to octahedral complexes. This means that the d orbitals that were higher in energy (eg set) in octahedral complexes become lower in energy in tetrahedral complexes, and vice versa.
  2. Smaller Splitting: The magnitude of the splitting energy, denoted as Δt (Delta sub t), is smaller in tetrahedral complexes compared to octahedral complexes. Mathematically, it can be shown that Δt is approximately equal to (4/9)Δ₀, where Δ₀ represents the crystal field splitting energy in octahedral complexes.
  3. Ineffectiveness of Pairing: Due to the smaller splitting energy in tetrahedral complexes, it is generally insufficient to force electron pairing within the d orbitals. As a result, low spin configurations are rarely observed in tetrahedral complexes. Most tetrahedral complexes exhibit high spin configurations.
  4. Absence of ‘g’ Subscript: In octahedral and square planar complexes, which possess a center of symmetry, the energy levels are labeled with ‘g’ subscripts to denote this symmetry. However, tetrahedral complexes lack this center of symmetry, so the ‘g’ subscript is not used to label the energy levels.

Colour in Coordination Compounds

The color of coordination compounds is a fascinating property that can be explained by the Crystal Field Theory. Here’s how the theory elucidates the colors of coordination compounds:

  1. Absorption of Light: When white light passes through a coordination compound, some of the visible spectrum is absorbed by the compound. The color we perceive is the complementary color to that which is absorbed. For example, if green light is absorbed, the compound will appear red.
  2. Crystal Field Theory Explanation: Crystal Field Theory (CFT) provides a molecular-level explanation for these colors. It considers the metal-ligand bond as ionic, arising from electrostatic interactions between the metal ion and the ligands.
  3. Splitting of d Orbitals: In coordination compounds, especially those with transition metals, the d orbitals of the metal ion split in energy due to the electrostatic interactions with the ligands. This splitting depends on the nature of the ligands and the metal-ligand distances.
  4. d-d Transitions: The color arises from d-d electron transitions within the metal’s d orbitals. In the ground state of a complex, electrons occupy certain d orbitals. When light of the appropriate energy (wavelength) is absorbed, it can promote an electron from a lower energy d orbital to a higher energy d orbital.
  5. Complementary Color: The observed color of the complex is complementary to the absorbed wavelength. For example, if a complex absorbs light in the blue-green region, it will appear violet because violet is complementary to blue-green.
  6. Effect of Ligands: The choice of ligands can significantly influence the color of a complex. Different ligands will cause different splitting of the d orbitals, leading to variations in absorbed wavelengths and hence, colors. Ligands can be arranged in a spectrochemical series based on their ability to cause d-d transitions.
  7. Colorless Complexes: In the absence of ligands, crystal field splitting does not occur, and the substance remains colorless. For example, removing water from a hydrated complex can render it colorless.
  8. Influence of Ligands: The influence of ligands on the color of a complex can be demonstrated by gradually adding a different ligand to a metal complex. As the ligand changes, the color of the complex can shift accordingly.

Limitations of Crystal Field Theory

The Crystal Field Theory (CFT) is a useful model for explaining many properties of coordination compounds. However, it does have some limitations, including:

  1. Assumption of Point Charges: CFT assumes that ligands are point charges. In reality, ligands have molecular structures and possess covalent characteristics in their bonding to the central metal ion. This oversimplification can lead to inaccuracies in predicting the electronic and magnetic properties of certain coordination compounds.
  2. Spectrochemical Series Anomalies: CFT predicts that anionic ligands should have the greatest splitting effect due to their negative charge. However, the experimental spectrochemical series places anionic ligands like fluoride (F-) and oxide (O2-) at the weaker field end of the series, which contradicts CFT predictions.
  3. Neglect of Covalency: CFT treats metal-ligand bonding as purely ionic, disregarding the covalent character of the bonds. In reality, metal-ligand interactions often have both ionic and covalent components, especially for transition metals. CFT does not account for this covalency, which can be significant in some cases.
  4. Ligand Field Theory and Molecular Orbital Theory: While CFT provides valuable insights, it has limitations in explaining certain aspects of coordination compounds. To address these limitations, more advanced theories like Ligand Field Theory (LFT) and Molecular Orbital Theory (MOT) have been developed. These theories offer a more comprehensive understanding of the electronic structure, bonding, and properties of coordination compounds.

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