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GNDU Question Paper-2023

Ba/Bsc 5

Semester

CHEMISTRY

(Inorganic Chemistry-IV)

Time Allowed: Three Hours Maximum Marks: 35

Note: Attempt Five questions in all, selecting at least One question from each section.

The Fifth question may be attempted from any section. All questions carry equal marks.

SECTION-A

1. (a) Discuss the Crystal field Splitting in tetrahedral complexes.

(b) Give reasons for the smaller value of Crystal field Splitting in tetrahedral than in

octahedral complexes. 2

2. (a) How does the Crystal field theory explain the colour of complexes?

(b) Define Crystal field stabilization energy. Calculate CFSE for d' in octahedral and

terahedral complexes. 2

SECTION-B

3. (a) What is magnetic susceptibility? How does it vary with temperature?

(b) Write a short note on ferromagnetic and anti ferromagnetic susbtances?

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4. (a) Write down the mechanism of nucleophilic substitution in square planar complexes.

(b) What is Trans Effect? Explain it by giving example.

SECTION-C

5. (a) What are Orgel diagrams? Give its limitation and also explain electronic transition

spectra of [Ti(H₂O)

]

(b) Write down the selection rules for d-d transition.

6. (a) What is Term symbol? How can you calculate Ground State term with spin

multiplicity for following octahedral ions V

,Ni

,Cu

(b) How do Russell Saunders states get splitted in octahedral fields? Explain with

diagram.

SECTION-D

7. (a) What is Wilkinson's catalyst? Give the mechanism of hydrogenation of alkene by

Wilkinson Catalyst.

(b) Discuss the bonding in metal-olefin complex.

8.(a) What is hapticity? Give an example of ligands of various hapticity.

(b) Give an account of preparation, properties and uses of organolithium compounds.

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GNDU Answer Paper-2023

Ba/Bsc 5

Semester

CHEMISTRY

(Inorganic Chemistry-IV)

Time Allowed: Three Hours Maximum Marks: 35

Note: Attempt Five questions in all, selecting at least One question from each section.

The Fifth question may be attempted from any section. All questions carry equal marks.

SECTION-A

1. (a) Discuss the Crystal field Splitting in tetrahedral complexes.

(b) Give reasons for the smaller value of Crystal field Splitting in tetrahedral than in

octahedral complexes. 2

Ans: Crystal Field Splitting in Tetrahedral Complexes

What is Crystal Field Theory?

Crystal Field Theory (CFT) is a model that describes the electronic structure of transition

metal complexes. It explains how the presence of surrounding ligands affects the energies of

the d-orbitals of the metal ion.

When a transition metal ion is surrounded by ligands (molecules or ions that donate

electron pairs to the metal), the degeneracy (equal energy) of its d-orbitals is disrupted,

leading to different energy levels. This is known as crystal field splitting.

Tetrahedral Complexes

Tetrahedral complexes are formed when a metal ion is surrounded by four ligands arranged

at the corners of a tetrahedron. For example, a common tetrahedral complex is

In a tetrahedral field, the d-orbitals split into two different sets:

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1. Higher Energy Set (e): Composed of d

d_{z^2and d

−y2d_

2. Lower Energy Set (t): Composed of d

d_.

The splitting occurs because the ligands approach the metal ion in a way that causes

different amounts of repulsion between the electrons in the d-orbitals and the electrons in

the ligands.

Visualization of Tetrahedral Splitting

To visualize this, imagine a sphere where the metal ion is at the center and the ligands are at

the corners of a tetrahedron. As the ligands come closer to the metal ion, the dz2d_{z^2}dz2

and dx2−y2d_{x^2-y^2}dx2−y2 orbitals experience more repulsion because they point

directly toward the ligands, raising their energy level compared to the dxyd_{xy}dxy,

dxzd_{xz}dxz, and dyzd_{yz}dyz orbitals, which are oriented between the ligands.

Crystal Field Splitting Energy (Δt\Delta_tΔt)

The energy difference between the two sets of orbitals in tetrahedral complexes is

represented by Δt\Delta_tΔt. This value is generally smaller than the splitting in octahedral

complexes, where the d-orbitals split into two sets (the t2gt_{2g}t2g and ege_geg levels)

with a larger energy difference (Δo\Delta_oΔo).

Electron Configuration

In tetrahedral complexes, the arrangement of electrons in the d-orbitals follows Hund's

Rule, where electrons occupy degenerate orbitals singly before pairing up. This influences

the magnetic properties of the complexes.

Reasons for Smaller Crystal Field Splitting in Tetrahedral Complexes Compared to

Octahedral Complexes

1. Ligand Arrangement:

o In tetrahedral complexes, the ligands are farther from the metal ion

compared to octahedral complexes. This greater distance reduces the

electrostatic repulsion between the d-orbitals and the ligands, resulting in a

smaller splitting energy Δt\Delta_tΔt.

2. Orientation of d-orbitals:

o The orientation of the d-orbitals also plays a crucial role. In octahedral

complexes, the ligands approach the metal ion along the axes (x, y, z), causing

significant repulsion for the t2gt_{2g}t2g orbitals (which point between the

ligands) and ege_geg orbitals (which point towards the ligands), resulting in a

larger energy difference. In contrast, in tetrahedral complexes, the ligands

are situated at the corners of the tetrahedron, leading to less effective

overlap and hence a smaller splitting.

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3. Crystal Field Splitting Energy Comparison:

o In general, the crystal field splitting energy for tetrahedral complexes is

approximately 4/9 that of octahedral complexes. This quantitative difference

is a result of both the arrangement of ligands and the spatial orientation of

the d-orbitals.

Spectrochemical Series

What is the Spectrochemical Series?

The spectrochemical series is an empirical list that ranks ligands based on their ability to

split the d-orbitals of transition metal complexes. Ligands that cause larger splitting (strong

field ligands) are placed higher in the series, while those that cause smaller splitting (weak

field ligands) are placed lower.

Importance of the Spectrochemical Series

1. Predicting Splitting:

o The position of a ligand in the spectrochemical series helps predict the extent

of crystal field splitting in a complex. For instance, ligands like CN−\text{CN}^-

CN− and CO\text{CO}CO are strong field ligands, leading to larger Δ\DeltaΔ

values, while ligands like I−\text{I}^-I− and Br−\text{Br}^-Br− are weak field

ligands that cause smaller splitting.

2. Influence on Color:

o The extent of splitting influences the wavelengths of light absorbed by the

complex, which in turn affects the color observed. Strong field ligands

typically cause greater splitting, resulting in absorption of higher energy light

(blue or violet), while weak field ligands lead to lower energy absorption (red

or yellow).

Common Ligands in the Spectrochemical Series

Here is a simplified list of ligands arranged in the spectrochemical series from strong field to

weak field:

• Strong Field Ligands: CN−\text{CN}^-CN−, CO\text{CO}CO, NH3\text{NH}_3NH3,

en\text{en}en (ethylenediamine)

• Moderate Field Ligands: NO2−\text{NO}_2^-NO2−, Cl−\text{Cl}^-Cl−, F−\text{F}^-F−

• Weak Field Ligands: I−\text{I}^-I−, Br−\text{Br}^-Br−, H2O\text{H}_2\text{O}H2O

Conclusion

Understanding crystal field splitting, especially in tetrahedral complexes, is essential in

predicting the behavior of transition metal complexes, including their electronic

configurations, magnetic properties, and colors. The comparison of splitting energies

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between tetrahedral and octahedral complexes, along with the spectrochemical series,

provides valuable insights into the nature of bonding in these complexes.

This simplified explanation covers the fundamentals of crystal field splitting in tetrahedral

complexes, the reasons for their smaller splitting values compared to octahedral complexes,

and the significance of the spectrochemical series.

2. (a) How does the Crystal field theory explain the colour of complexes?

(b) Define Crystal field stabilization energy. Calculate CFSE for d' in octahedral and

terahedral complexes. 2

Ans: Crystal Field Theory and the Color of Complexes

1. Introduction to Crystal Field Theory (CFT)

Crystal Field Theory is a model that describes the electronic structure of transition metal

complexes. It explains how the arrangement of ligands around a central metal ion affects

the energy of the d-orbitals in the metal ion. According to CFT, ligands are thought of as

point charges that create an electric field, which interacts with the d-electrons of the metal

ion.

2. Explanation of Color in Complexes

The color of a complex is a result of the interaction of light with the d-electrons in the metal

ions. When light hits a complex, certain wavelengths (colors) are absorbed, while others are

transmitted or reflected. The color that we see is the complement of the color absorbed.

How CFT Explains the Color of Complexes:

• Ligand Field Splitting: In an octahedral complex, the five d-orbitals split into two

energy levels:

o t2g: Lower energy orbitals (dxy, dyz, dxz)

o eg: Higher energy orbitals (dx2-y2, dz2)

• Absorption of Light: When light hits the complex, some energy can excite an

electron from the lower-energy t2g orbitals to the higher-energy eg orbitals. The

energy difference between these two sets of orbitals corresponds to the energy of

specific wavelengths of light. For example, if a complex absorbs light in the red

region (around 700 nm), it may appear green to the human eye (complementary

color).

• Variation in Color: The color of the complex can change based on:

o The type of metal ion.

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o The oxidation state of the metal ion.

o The nature of the ligands (strong field vs. weak field ligands).

Examples:

• [Cu(H2O)6]²⁺: This complex absorbs light in the red region, appearing blue.

• [Ti(H2O)6]³⁺: This complex absorbs light in the blue-green region, appearing purple.

3. Crystal Field Stabilization Energy (CFSE)

Definition of CFSE: Crystal Field Stabilization Energy is the energy difference between the

stabilized state of the electrons in a crystal field and the energy of the same electrons in a

spherical field (where no splitting occurs). CFSE is important because it helps determine the

stability of a complex.

Calculation of CFSE:

• For Octahedral Complexes:

o The CFSE can be calculated using the formula:

o Where:

▪ nt

: Number of electrons in t2g orbitals

▪ n

: Number of electrons in eg orbitals

▪ Δ: Crystal field splitting energy (the energy difference between the t2g

and eg levels).

• For Tetrahedral Complexes:

o The CFSE is calculated differently since the splitting is inverted:

o Where:

▪ nt2n_: Number of electrons in t2 orbitals (equivalent to eg in

octahedral)

▪ nen_{e}ne: Number of electrons in e orbitals (equivalent to t2g in

octahedral).

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Example Calculations:

1. Octahedral Complex [Ti(H2O)6]³⁺:

o nt2g=6n_{t2g} = 6nt2g=6 (for Ti³⁺ with 3 d-electrons)

o neg=0n_{eg} = 0neg=0

o If Δ=10 Dq

2. Tetrahedral Complex [CuCl4]²⁻:

o nt2=2n_{t2} = 2nt2=2 (for Cu²⁺ with 9 d-electrons)

o ne=2n_{e} = 2ne=2

o If Δ=6 Dq\Delta = 6 \text{ Dq}Δ=6 Dq,

4. Factors Affecting the Magnitude of Crystal Field Splitting

Crystal field splitting is influenced by several factors:

a. Nature of the Metal Ion:

• Different transition metals have different numbers of d-electrons, which affects the

extent of d-orbital splitting.

• For example, heavier transition metals tend to have larger splitting compared to

lighter ones due to increased effective nuclear charge.

b. Oxidation State of the Metal:

• Higher oxidation states generally lead to greater splitting. This is because more

positive charge on the metal ion increases the attraction between the d-electrons

and the surrounding ligands, thus increasing the energy difference between the split

d-orbitals.

c. Type of Ligands:

• Ligands can be classified as strong field or weak field ligands based on their ability to

split d-orbitals:

o Strong Field Ligands: These create a large splitting (e.g., CN⁻, CO). They tend

to favor low-spin configurations.

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o Weak Field Ligands: These create a small splitting (e.g., H2O, Cl⁻). They favor

high-spin configurations.

d. Geometry of the Complex:

• The geometry (octahedral, tetrahedral, square planar, etc.) affects the extent of

splitting. For instance, octahedral complexes typically experience greater splitting

than tetrahedral complexes because of the arrangement of ligands around the metal

ion.

e. Ligand Field Theory Extensions:

• Beyond CFT, ligand field theory incorporates aspects of molecular orbital theory and

can explain the electronic transitions more accurately. It considers the covalent

character of metal-ligand bonds.

Conclusion

Crystal Field Theory is a valuable framework for understanding the color and stability of

transition metal complexes. By analyzing the interactions between d-orbitals and ligands,

CFT can explain the diverse colors observed in different metal complexes, the significance of

CFSE, and the various factors that influence crystal field splitting.

Understanding these concepts is essential for the study of coordination chemistry and the

application of these principles in fields such as material science, catalysis, and bioinorganic

chemistry.

This comprehensive overview should provide a solid foundation in CFT and its applications

regarding the color of complexes, CFSE calculations, and the factors affecting crystal field

splitting.

SECTION-B

3. (a) What is magnetic susceptibility? How does it vary with temperature?

(b) Write a short note on ferromagnetic and anti ferromagnetic susbtances?

Ans: 1. What is Magnetic Susceptibility?

Definition

Magnetic susceptibility (χ\chiχ) is a measure of how much a material will become

magnetized in an external magnetic field. It quantifies the degree to which a substance can

be magnetized when subjected to an external magnetic field. In simpler terms, it tells us

how a material responds to magnetism.

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Types of Magnetic Susceptibility

Magnetic susceptibility can be classified into different types based on how the material

behaves in a magnetic field:

1. Diamagnetic (χ<0\chi < 0χ<0):

o Behavior: Materials that show negative susceptibility are called diamagnetic.

They are weakly repelled by a magnetic field.

o Examples: Bismuth, copper, and lead.

2. Paramagnetic (χ>0\chi > 0χ>0):

o Behavior: Paramagnetic materials have a positive susceptibility. They are

weakly attracted by a magnetic field due to the presence of unpaired

electrons.

o Examples: Aluminum, platinum, and manganese.

3. Ferromagnetic:

o Behavior: Ferromagnetic materials have a large positive susceptibility. They

can retain magnetization even after the external magnetic field is removed.

o Examples: Iron, nickel, and cobalt.

4. Antiferromagnetic:

o Behavior: These materials have a unique property where the magnetic

moments of adjacent atoms or ions are aligned in opposite directions,

resulting in no net magnetization.

o Examples: Manganese oxide (MnO) and iron oxide (FeO).

How Magnetic Susceptibility Varies with Temperature

The variation of magnetic susceptibility with temperature is significant for understanding

the behavior of magnetic materials:

1. Diamagnetic Materials:

o The magnetic susceptibility remains relatively constant with temperature

since diamagnetism arises from the paired electrons in atoms.

o As temperature changes, the thermal motion of atoms increases, but this has

little effect on their magnetic properties.

2. Paramagnetic Materials:

o The susceptibility decreases with increasing temperature. This is because

thermal agitation at higher temperatures disrupts the alignment of unpaired

electrons, reducing magnetization.

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o This relationship can be described by Curie’s Law, which states: χ=CT\chi =

\frac{C}{T}χ=TC Where CCC is the Curie constant, and TTT is the absolute

temperature. As TTT increases, χ\chiχ decreases.

3. Ferromagnetic Materials:

o At low temperatures, ferromagnetic materials exhibit strong magnetization.

As the temperature increases, there is a critical point known as the Curie

temperature, above which the material loses its ferromagnetic properties

and becomes paramagnetic.

o This change occurs due to increased thermal agitation overcoming the

exchange energy that maintains the alignment of magnetic moments.

4. Antiferromagnetic Materials:

o These substances also display a variation in susceptibility with temperature.

At absolute zero, they exhibit complete antiferromagnetic order. As

temperature rises, thermal motion can disrupt this order.

o There is a specific temperature, called the Néel temperature, above which

antiferromagnetic materials become paramagnetic.

2. Ferromagnetic and Antiferromagnetic Substances

Ferromagnetic Substances

Ferromagnetism is a phenomenon where certain materials can become magnetized; they

have strong, permanent magnetic moments.

Characteristics:

• Alignment: In ferromagnetic materials, the magnetic moments of atoms align

parallel to each other in the presence of an external magnetic field, resulting in a net

magnetic moment.

• Hysteresis: These materials exhibit hysteresis, meaning their magnetization depends

on their previous magnetic history. They retain magnetization even after the

external field is removed.

• Curie Temperature: The temperature above which a ferromagnetic material loses its

permanent magnetism is called the Curie temperature (T_C).

Examples:

• Iron: Iron is one of the most well-known ferromagnetic materials. It is used

extensively in magnets and magnetic devices.

• Nickel: Nickel also exhibits ferromagnetic properties and is used in various

applications, including batteries and coins.

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• Cobalt: Cobalt has ferromagnetic properties and is often alloyed with other metals

to enhance magnetic properties.

Applications:

• Magnets: Ferromagnetic materials are widely used in the production of permanent

magnets.

• Electric Motors: They are essential components in electric motors and generators

due to their magnetic properties.

Antiferromagnetic Substances

Antiferromagnetism is a type of magnetism that occurs in certain materials where adjacent

magnetic moments align in opposite directions.

Characteristics:

• Opposite Alignment: In antiferromagnetic materials, the magnetic moments of

neighboring atoms or ions are aligned antiparallel to each other, leading to a

cancellation of magnetic effects and a net magnetic moment of zero.

• Néel Temperature: The temperature below which antiferromagnetic ordering occurs

is called the Néel temperature (T_N). Above this temperature, the material behaves

like a paramagnet.

Examples:

• Manganese Oxide (MnO): This compound is a classic example of an

antiferromagnetic material, showing strong antiferromagnetic ordering at low

temperatures.

• Iron Oxide (FeO): FeO exhibits antiferromagnetic properties and is often studied for

its interesting magnetic behavior.

• Nickel Oxide (NiO): This material also displays antiferromagnetic behavior at low

temperatures.

Applications:

• Magnetic Storage: Antiferromagnetic materials are used in advanced magnetic

storage devices due to their ability to maintain a stable magnetic state.

• Spintronics: Antiferromagnetic materials are being investigated for use in

spintronics, a technology that exploits the intrinsic spin of electrons.

3. Thermodynamic and Kinetic Stability of Complexes

Thermodynamic Stability

Thermodynamic stability refers to the stability of a complex in terms of its energy state. A

thermodynamically stable complex is one that is lower in energy compared to its reactants

and is less likely to undergo spontaneous reactions under given conditions.

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Factors Influencing Thermodynamic Stability:

• Formation Energy: If the formation of a complex releases a significant amount of

energy, it tends to be more thermodynamically stable.

• Ligand Field Strength: The nature of the ligands (molecules or ions that bind to the

central metal ion) can significantly influence stability. Strong field ligands lead to

greater stability.

• Entropy: Complexes with higher entropy (greater disorder) may have improved

thermodynamic stability at high temperatures.

Example:

• [Cu(NH₃)₄]²⁺: This complex is thermodynamically stable due to strong interactions

between copper and ammonia ligands, resulting in lower energy.

Kinetic Stability

Kinetic stability refers to the rate at which a complex undergoes reactions. A kinetically

stable complex may have a high activation energy barrier, making it slow to react, even if it

is thermodynamically unstable.

Factors Influencing Kinetic Stability:

• Activation Energy: A high activation energy means that a complex is less likely to

react quickly, leading to kinetic stability.

• Ligand Arrangement: The spatial arrangement of ligands around the metal center

can impact the rate of reaction. A complex with a stable arrangement may resist

change.

• Temperature: Increased temperature can provide the energy necessary for a

reaction to occur, potentially reducing kinetic stability.

Example:

• [Co(NH₃)₆]³⁺: This complex may be thermodynamically unstable but can exhibit

kinetic stability due to its high activation energy barrier for dissociation.

Relationship Between Thermodynamic and Kinetic Stability

• Thermodynamic Stability: A complex can be thermodynamically stable, meaning it is

low in energy and less likely to decompose or react. However, this does not

guarantee that it will react quickly.

• Kinetic Stability: A complex can be kinetically stable, meaning it does not react

quickly even if it is not the lowest energy form (thermodynamically unstable).

In summary, thermodynamic stability relates to the energy state of the complex, while

kinetic stability pertains to the rate at which reactions occur. Both types of stability are

crucial in understanding the behavior of coordination complexes.

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Conclusion

Understanding magnetic susceptibility, ferromagnetic and antiferromagnetic substances,

and the concepts of thermodynamic and kinetic stability in complexes is fundamental in the

study of inorganic chemistry. These concepts help explain the behavior of materials in

magnetic fields and the stability of complex compounds, which are essential for various

applications in science and technology.

4. (a) Write down the mechanism of nucleophilic substitution in square planar complexes.

(b) What is Trans Effect? Explain it by giving example.

Ans: Inorganic chemistry, particularly the study of coordination complexes, introduces a fascinating

area involving reactions, stability, and behavior of metal complexes. In the fifth semester of a BA/BSc

course in chemistry, understanding nucleophilic substitution in square planar complexes, the Trans

effect, and the factors that affect the stability of complexes is crucial. Let's break these concepts

down step by step in a simple, easy-to-understand language while maintaining a detailed

explanation.

(a) Mechanism of Nucleophilic Substitution in Square Planar Complexes

Nucleophilic substitution reactions involve replacing one group (called a leaving group)

attached to the central metal atom in a complex with another group (called a nucleophile).

These reactions are common in square planar complexes, where the metal center is

coordinated to four ligands arranged in a square plane.

Understanding Square Planar Complexes:

• Structure: A square planar complex has a central metal atom or ion that is

surrounded by four ligands in the same plane, forming a square-like arrangement.

• Examples: Some examples of metals that commonly form square planar complexes

include platinum (Pt), palladium (Pd), nickel (Ni), and gold (Au). The most famous

example is [Pt(NH₃)₂Cl₂], where platinum is at the center, two ammonia (NH₃)

molecules, and two chloride (Cl) ions occupy the corners of the square plane.

Now, in nucleophilic substitution reactions in these complexes, one of the ligands gets

replaced by a nucleophile. This process can follow two different mechanisms: Associative (A)

and Dissociative (D) mechanisms.

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Associative Mechanism (A Mechanism):

In this mechanism, the incoming nucleophile attacks the metal complex before the leaving

group departs, leading to a temporary intermediate where the metal is bonded to more

than four ligands (five ligands in total).

1. Step 1: Nucleophile Attacks – The nucleophile approaches the metal center and

bonds with it, forming a five-coordinate intermediate.

2. Step 2: Leaving Group Departs – After forming the intermediate, the leaving group is

expelled from the complex, resulting in a new square planar complex with the

nucleophile in place of the leaving group.

This type of mechanism is favored when the incoming nucleophile is strong and can easily

attach to the metal center.

Example: The reaction of [PtCl₄]²⁻ with an incoming nucleophile like ammonia (NH₃) can

proceed via an associative mechanism, where ammonia temporarily forms a five-

coordinated intermediate before one of the chlorides leaves.

Dissociative Mechanism (D Mechanism):

In the dissociative mechanism, the leaving group departs first, creating a vacancy at the

metal center (three-coordinate intermediate), and then the nucleophile comes in to fill the

empty position.

1. Step 1: Leaving Group Departs – The leaving group breaks away from the complex,

leaving behind a three-coordinate intermediate.

2. Step 2: Nucleophile Attacks – The nucleophile then attacks the metal and occupies

the vacant position.

This mechanism is more likely when the leaving group is weakly bound to the metal center,

making it easier to leave. It is typically seen in complexes where the metal has a strong

preference for square planar geometry, even in the intermediate stage.

Summary of Mechanisms:

• Associative (A Mechanism): Nucleophile comes in first → Five-coordinate

intermediate forms → Leaving group leaves.

• Dissociative (D Mechanism): Leaving group leaves first → Three-coordinate

intermediate forms → Nucleophile comes in later.

Both mechanisms depend on the specific nature of the metal, the nucleophile, and the

leaving group involved.

(b) What is the Trans Effect?

The Trans Effect is a phenomenon observed in square planar complexes, particularly those

of transition metals like platinum, palladium, and gold. It refers to the ability of a ligand to

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influence the substitution rate of another ligand located opposite to it (the trans position) in

the square planar structure.

How the Trans Effect Works:

In a square planar complex, one ligand can exert an influence on how easily the ligand

directly opposite (trans to it) is replaced during a nucleophilic substitution reaction. Some

ligands have a stronger trans effect than others, meaning they can make the trans ligand

more easily replaced by a nucleophile.

Examples of Trans-Directing Ligands:

• Strong Trans Effect Ligands: Ligands like CN⁻ (cyanide), CO (carbonyl), and PR₃

(phosphines) have a strong trans effect. When one of these ligands is present in the

complex, the ligand trans to it is more easily substituted.

• Weak Trans Effect Ligands: Ligands like NH₃ (ammonia) and Cl⁻ (chloride) have a

weaker trans effect, meaning they do not accelerate the replacement of the trans

ligand as strongly.

Example of the Trans Effect:

A classic example is the substitution reaction of [Pt(NH₃)₂Cl₂], also known as cisplatin (an

anti-cancer drug). In this complex:

• NH₃ has a weaker trans effect.

• Cl⁻ has a stronger trans effect.

When cisplatin undergoes substitution, the chloride ion that is trans to the stronger trans-

effect ligand is replaced faster by a nucleophile (like water or ammonia) compared to the

ligand trans to a weaker trans-effect ligand.

How the Trans Effect Occurs:

The trans effect can arise due to two main reasons:

1. Electronic Effects: Some ligands donate electrons to the metal, weakening the bond

between the metal and the trans ligand, making it easier to replace.

2. Steric Effects: Some ligands are bulky and can push away or destabilize the ligand

trans to them, allowing for easier substitution.

Applications of the Trans Effect:

The trans effect is particularly useful in the synthesis of coordination complexes. Chemists

can use ligands with a strong trans effect to direct which ligands will be substituted, allowing

for selective and controlled synthesis of metal complexes.

The stability of a coordination complex refers to how tightly the ligands are bound to the

central metal ion. Some complexes are very stable and do not easily dissociate or undergo

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substitution reactions, while others are more reactive. Several factors influence the stability

of these complexes.

1. Nature of the Metal Ion:

• Charge of the Metal Ion: A metal ion with a higher positive charge attracts ligands

more strongly, leading to a more stable complex. For example, Fe³⁺ (ferric ion) forms

more stable complexes than Fe²⁺ (ferrous ion) because the higher charge on Fe³⁺

attracts ligands more strongly.

• Size of the Metal Ion: Smaller metal ions tend to form more stable complexes

because the ligands can get closer to the metal and form stronger bonds. For

instance, Cu²⁺ forms more stable complexes than larger metal ions like Ba²⁺.

• Electronic Configuration: The electron arrangement of the metal ion plays a role in

stability. Metal ions with completely filled or half-filled d orbitals (like d⁵ in Mn²⁺ or

d¹⁰ in Zn²⁺) often form more stable complexes.

2. Nature of the Ligands:

• Charge on the Ligand: Ligands that are negatively charged (like CN⁻, OH⁻) tend to

form more stable complexes because they are more strongly attracted to the

positively charged metal ion.

• Size of the Ligand: Smaller ligands can approach the metal ion more closely, forming

stronger bonds, which increases stability. Large ligands may face steric hindrance,

reducing stability.

• Electron Donating Ability: Ligands that can donate electrons effectively (like PR₃ and

CO) stabilize the metal center and form stronger complexes.

• Chelation Effect: When ligands form more than one bond with the metal ion (like

EDTA or en), they wrap around the metal, forming a more stable complex due to the

chelation effect. The chelation effect significantly increases the stability of the

complex.

3. The Chelate Effect:

The chelate effect refers to the increased stability of complexes where the ligands form

rings with the metal ion by binding through multiple points. These ligands are called

chelating ligands.

For example, a bidentate ligand like ethylenediamine (en), which can attach to the metal ion

at two points, forms a more stable complex than two separate monodentate ligands. The

chelate effect occurs because forming a ring structure reduces the entropy (disorder) of the

system and creates a more rigid, stable structure.

4. The Nature of the Solvent:

The solvent in which the complex is dissolved also plays a role in its stability. In some cases,

polar solvents like water can stabilize charged complexes through solvation (interaction with

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the solvent). In other cases, non-polar solvents may stabilize complexes by reducing the

solubility of the ligands, encouraging them to stay bound to the metal.

5. Steric Effects:

If the ligands or the metal center are too large, steric hindrance (physical blocking) can

reduce the stability of the complex. Large ligands may not be able to approach the metal ion

closely enough to form strong bonds, leading to less stable complexes.

6. Entropy and Enthalpy Considerations:

• Enthalpy: Complexes that release a large amount of energy when forming bonds

(highly exothermic reactions) are more stable.

• Entropy: Complexes that lead to a significant increase in disorder (entropy) are often

more stable. For example, the formation of a chelate complex releases several water

molecules, increasing entropy and stabilizing the complex.

Conclusion:

In summary:

• Nucleophilic substitution in square planar complexes can follow either the

associative or dissociative mechanism, depending on the strength of the nucleophile

and the leaving group.

• The Trans effect is a fascinating phenomenon where one ligand can influence the

substitution rate of another ligand located trans to it in a square planar complex.

• The stability of coordination complexes is influenced by several factors, including the

nature of the metal ion, the ligands, the chelate effect, solvent, steric effects, and

thermodynamic considerations like entropy and enthalpy.

These fundamental concepts are critical for understanding the reactivity and behavior of

coordination complexes in inorganic chemistry.

SECTION-C

5. (a) What are Orgel diagrams? Give its limitation and also explain electronic transition

spectra of [Ti(H₂O)

]

(b) Write down the selection rules for d-d transition.

Ans: (a) Orgel Diagrams

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Orgel diagrams are graphical representations that help in understanding the electronic

transitions of transition metal complexes, specifically for high-spin octahedral and

tetrahedral systems. These diagrams are particularly useful for explaining the energy levels

of the d-orbital electrons in such complexes. Orgel diagrams focus on transitions that are

spin-allowed and are most commonly used for d^1, d^4, d^6, and d^9 configurations for

tetrahedral complexes, and d^2, d^3, high-spin d^7, and d^8 configurations for octahedral

complexes.

Orgel diagrams represent the different electronic states of these metal complexes and show

how they split in an octahedral or tetrahedral field. The diagrams are divided into two types

based on the ground terms of the metal ions:

1. D Orgel Diagrams: These are used for systems where the ground term is D (like d^1

and d^9 configurations).

2. F Orgel Diagrams: These apply to complexes with F ground terms (like d^2, d^3, and

d^8 configurations).

Limitations of Orgel Diagrams:

1. Only for Spin-Allowed Transitions: Orgel diagrams only show spin-allowed

transitions, meaning they ignore the possibility of spin-forbidden transitions. This

makes them less versatile for certain complexes where spin-forbidden transitions

may occur.

2. Limited to High-Spin Complexes: Orgel diagrams are specifically designed for high-

spin complexes. They do not account for low-spin complexes, which are more

common in strong-field environments.

3. No Quantitative Information: While they provide qualitative information about the

relative positions of energy levels, Orgel diagrams do not offer exact quantitative

data regarding the energy differences between levels.

4. Not Effective for All Configurations: Orgel diagrams are best suited for simple d^1,

d^2, d^4, d^6, and d^9 systems. They are less useful for more complex

configurations or for those systems with multiple transitions.

Electronic Transition Spectrum of [Ti(H₂O)₆]³⁺:

The complex [Ti(H₂O)₆]³⁺ is an octahedral transition metal complex, where titanium exists in

the +3 oxidation state with an electron configuration of d1d^1d1. This means that the

electron occupies one of the t2gt_{2g}t2g orbitals in an octahedral crystal field.

• In an octahedral field, the five degenerate d-orbitals split into two sets: three

t2gt_{2g}t2g orbitals (lower in energy) and two ege_geg orbitals (higher in energy).

• The electronic transition observed in the spectrum of this complex occurs when the

electron is excited from the t2gt_{2g}t2g set to the ege_geg set, resulting in an

absorption of energy that corresponds to this transition.

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• For [Ti(H₂O)₆]³⁺, this corresponds to a d-d transition (from t2gt_{2g}t2g to ege_geg).

Since there is only one electron in the system, only one transition is possible, which

is represented by a single peak in the UV-visible spectrum.

The absorption band in the spectrum for this complex occurs due to this electron excitation

from a lower-energy t2gt_{2g}t2g orbital to a higher-energy ege_geg orbital, resulting in a

transition that can be detected using spectroscopy techniques.

(b) Selection Rules for d-d Transitions

Selection rules govern the probability of electronic transitions between different energy

levels. For d-d transitions in transition metal complexes, two important selection rules must

be considered:

1. Laporte Selection Rule:

o This rule states that transitions between states of the same parity (symmetry)

are forbidden. In simpler terms, transitions are not allowed between orbitals

that both have the same type of symmetry.

o For d-d transitions, since both the initial and final states are d-orbitals (which

are of the same parity), these transitions are Laporte-forbidden. However, in

real complexes, distortions or mixing of orbitals due to interactions with

ligands can lead to a partial lifting of this restriction, making weak d-d

transitions observable.

2. Spin Selection Rule:

o This rule states that transitions are allowed only if there is no change in the

spin state of the electron during the transition. In other words, transitions

where the total spin quantum number remains the same are spin-allowed.

o For high-spin complexes, d-d transitions are more likely to be spin-allowed.

However, if a transition involves a change in the electron spin state, it

becomes spin-forbidden and is less likely to occur.

Summary of the Spectral Behavior and Rules:

• In complexes like [Ti(H₂O)₆]³⁺, only one spin-allowed d-d transition is observed

because of the d^1 configuration.

• According to the Laporte rule, the transition is forbidden in the strictest sense, but in

practice, octahedral complexes experience distortions (like Jahn-Teller distortions),

which relax the Laporte rule, allowing weak transitions.

• The spin selection rule is satisfied in this case since the transition does not involve a

change in the spin state of the electron.

Therefore, the electronic spectra of [Ti(H₂O)₆]³⁺ feature a single broad absorption band

corresponding to this allowed d-d transition, which is weak due to the Laporte-forbidden

nature but still detectable.

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Orgel diagrams and selection rules together help predict the nature of such electronic

transitions in transition metal complexes and explain the observable features in their

spectra

6. (a) What is Term symbol? How can you calculate Ground State term with spin

multiplicity for following octahedral ions V

,Ni

,Cu

(b) How do Russell Saunders states get splitted in octahedral fields? Explain with

diagram.

Ans: (a) Term Symbol and Calculating Ground State Terms

A term symbol represents the quantum states of atoms and ions, particularly their total

angular momentum. It is crucial in spectroscopic notation and helps describe the state of an

electron configuration based on quantum mechanics. The term symbol is written as:

(2S+1)

Here:

• S is the total spin quantum number, which represents the overall spin of the

electrons.

• L is the total orbital angular momentum, which is derived from the orbital angular

momentum of each electron.

• J is the total angular momentum (a combination of SSS and LLL).

Steps to calculate the ground state term symbol:

1. Find the electron configuration: Identify the distribution of electrons in orbitals (s, p,

d, f).

2. Determine SSS: Calculate the total spin quantum number by considering the spins of

individual electrons.

3. Determine LLL: Sum the orbital angular momenta of the electrons.

4. Calculate the spin multiplicity: It is given by 2S+12S + 12S+1.

5. Calculate JJJ: It can range from ∣L−S∣|L-S|∣L−S∣ to ∣L+S∣|L+S|∣L+S∣, depending on

whether the subshell is more or less than half-filled.

Let’s calculate the ground state terms for some specific ions:

1. V³⁺ (Vanadium ion):

o Configuration: 3d23d^23d2

o S=1S (since two unpaired electrons contribute to spin).

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o L=3 (as both electrons are in the d-orbital, their orbital angular momenta add

up).

o Ground state term: 3F

2. Ni²⁺ (Nickel ion):

o Configuration: 3d

o S=1S (as there are two unpaired electrons in the d-orbital).

o L=3L (for a d⁸ configuration, the possible orbital angular momentum leads to

L=3L ).

o Ground state term: 3F

3. Cu²⁺ (Copper ion):

o Configuration: 3d93d^93d9

o S=1/2S = 1/2S=1/2 (since one unpaired electron contributes to spin).

o L=2L = 2L=2.

o Ground state term: 2D^{2}D2D.

(b) Splitting of Russell-Saunders States in Octahedral Fields

In coordination complexes, especially in octahedral geometries, the electronic states of

transition metal ions are split due to the interaction of the d-orbitals with the surrounding

ligands. This phenomenon is known as crystal field splitting. In an octahedral field, the five

degenerate d-orbitals split into two groups:

• ege_ (higher energy, consisting of dz2d_{z^2}dz2 and dx2−y2d_{x^2 - y^2}dx2−y2).

• t2gt_{2g}t2g (lower energy, consisting of dxyd_{xy}dxy, dxzd_{xz}dxz, and

dyzd_{yz}dyz).

The splitting of these orbitals affects the term symbols of the atom or ion. For example, for a

d1d^1d1 configuration in an octahedral field:

• The t2gt_{2g}t2g set of orbitals (lower energy) corresponds to a triply degenerate

state.

• The ege_geg set (higher energy) corresponds to a doubly degenerate state.

The Russell-Saunders term symbols (which describe the overall spin and orbital state of the

atom) are also affected by this splitting, and the degenerate states may break into different

components based on symmetry considerations.

For example:

• The 2D^{2}D2D state can split into 2T2g^{2}T_{2g}2T2g and 2Eg^{2}E_g2Eg in an

octahedral field.

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• The 3F^{3}F3F state can split into multiple terms depending on how the orbitals

interact with the ligand field.

A diagram can visually represent the splitting, with the d-orbital levels (five-fold degenerate)

splitting into the t2gt_{2g}t2g and ege_geg orbitals under the influence of the octahedral

field.

L-S coupling, or Russell-Saunders coupling, describes the coupling between the orbital

angular momentum (L) and the spin angular momentum (S) in atoms. This type of coupling

is dominant when the interaction between the electrons' spins and their mutual interactions

are stronger than the spin-orbit interaction.

In L-S coupling:

• First, the total orbital angular momentum (LLL) is obtained by summing the

individual orbital angular momenta of the electrons.

• Similarly, the total spin angular momentum (SSS) is obtained by summing the spins

of all the electrons.

• The total angular momentum (JJJ) is then calculated by combining LLL and SSS.

This form of coupling works well for light atoms (low atomic number), where the spin-orbit

interaction is relatively weak. For heavier atoms, another type of coupling called j-j coupling

becomes more relevant.

In L-S coupling:

• If the subshell is less than half-filled, J=L−SJ = L - SJ=L−S.

• If the subshell is more than half-filled, J=L+SJ = L + SJ=L+S.

• For half-filled subshells, J=SJ = SJ=S.

For example, in a p2p^2p2 configuration:

• S=1S = 1S=1 (since both electrons are unpaired).

• L=1L = 1L=1 (since the p-orbital has l=1l = 1l=1 for each electron).

• The term symbol is 3PJ^{3}P_J3PJ, where JJJ can take values 2,1,02, 1, 02,1,0,

depending on how LLL and SSS combine.

In an octahedral field, the L-S coupled states can further split due to the crystal field effects,

leading to complex splitting patterns. The combination of these factors gives rise to the

observed spectra of transition metal complexes.

Conclusion

In summary, the term symbols describe the quantum states of atoms and ions, providing

information about their angular momentum. The ground state term is determined by

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maximizing the total spin and orbital angular momentum. In an octahedral field, Russell-

Saunders states undergo splitting due to the interaction of the d-orbitals with the

surrounding ligands. The splitting pattern depends on the symmetry of the field, and L-S

coupling further describes how orbital and spin angular momenta combine in light atoms.

SECTION-D

7. (a) What is Wilkinson's catalyst? Give the mechanism of hydrogenation of alkene by

Wilkinson Catalyst.

(b) Discuss the bonding in metal-olefin complex.

Ans: Wilkinson's Catalyst

Wilkinson's catalyst, named after Sir Geoffrey Wilkinson, is a well-known coordination

complex used primarily for the hydrogenation of alkenes. The chemical formula for

Wilkinson’s catalyst is [RhCl(PPh₃)₃], where Rh stands for rhodium, Cl for chlorine, and PPh₃

for triphenylphosphine. It is an organometallic compound, meaning that it contains a metal-

carbon bond, which plays a crucial role in its catalytic activity.

The catalyst is a square planar complex, which makes it highly reactive in the coordination

of hydrogen and alkenes, and is effective in homogeneous catalysis. In this process, both the

catalyst and the reactants are in the same phase (usually dissolved in the same solvent),

which provides excellent control over the reaction conditions and often leads to high

selectivity.

Applications of Wilkinson's Catalyst

One of the primary uses of Wilkinson’s catalyst is in the hydrogenation of alkenes, which is

the process of adding hydrogen (H₂) across the double bond of an alkene to form an alkane.

This reaction is significant in both organic synthesis and industrial applications where

selective hydrogenation is required. The catalyst is known for its efficiency and ability to

selectively hydrogenate alkenes without affecting other functional groups like carbonyls

(C=O), nitriles (CN), or aromatic rings (aryl groups)

Other applications include:

1. Selective Hydrogenation: It can selectively hydrogenate certain compounds without

affecting sensitive groups like carbonyls or aromatic systems.

2. Decarbonylation of Aldehydes: This reaction is used to remove a carbonyl group

after activation in a synthetic sequence.

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3. Formation of Carbon-Carbon Bonds: The catalyst is employed in reactions where

bonds between carbon atoms are formed, which is useful in constructing complex

organic molecules.

Mechanism of Hydrogenation of Alkenes by Wilkinson’s Catalyst

The mechanism of hydrogenation by Wilkinson's catalyst involves several steps, including

oxidative addition, migratory insertion, and reductive elimination. Here’s a step-by-step

explanation of the catalytic cycle:

1. Formation of the Active Catalyst: Wilkinson’s catalyst, as originally formulated, is a

precatalyst. Upon exposure to hydrogen (H₂), one of the triphenylphosphine (PPh₃)

ligands dissociates, forming a more reactive species [RhCl(PPh₃)₂], which is the active

catalyst.

2. Oxidative Addition: The active catalyst binds to molecular hydrogen (H₂) in a process

called oxidative addition. This increases the oxidation state of the rhodium from +1

to +3 and forms a rhodium-dihydride complex (Rh-H₂), which now contains two

hydrogen atoms attached to the rhodium metal center.

3. Coordination of the Alkene: The alkene (C=C bond) approaches and coordinates

with the rhodium complex. In this step, the alkene binds to the rhodium center,

positioning itself in a way that allows the hydrogen atoms to transfer from the metal

to the alkene.

4. Migratory Insertion: One of the hydrogen atoms from the rhodium-hydride complex

migrates to the carbon of the alkene, forming a new C-H bond. This process results in

the formation of a new alkyl-rhodium complex.

5. Reductive Elimination: In the final step, the remaining hydrogen atom on the

rhodium migrates to the adjacent carbon of the alkyl group, completing the

hydrogenation and releasing the saturated alkane (an alkane is a hydrocarbon with

single bonds). This process regenerates the active form of the catalyst, ready to

begin another cycle of hydrogenation

The catalytic cycle is highly efficient and selective, making it ideal for use in both small-scale

organic synthesis and large-scale industrial processes.

Metal-Olefin Bonding

In metal-olefin (alkene) complexes, the bonding involves a synergistic interaction between

the metal and the olefin. The metal donates electron density to the π* orbitals of the olefin,

while the olefin donates electron density back to the metal through its π-bonding electrons.

1. π-Backbonding: This is one of the key interactions in metal-olefin complexes. The

metal center, particularly a transition metal, donates electron density from its d-

orbitals into the empty π* anti-bonding orbitals of the alkene. This stabilizes the

complex and weakens the double bond of the olefin. It’s a synergistic process

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because while the metal donates electron density to the olefin, the olefin also

donates back to the metal.

2. σ-Bonding: The olefin’s π-electrons interact with the metal’s vacant orbitals, forming

a sigma (σ) bond. This interaction allows the olefin to bind to the metal in what is

called a "side-on" fashion, where both carbons of the alkene interact with the metal

center simultaneously.

This bonding interaction is highly useful in catalysis, as it allows the metal to activate the

alkene toward further reactions, such as hydrogenation or polymerization. Metal-olefin

complexes are key intermediates in many catalytic cycles, especially in processes like

hydrogenation (as seen with Wilkinson's catalyst) and polymerization reactions (as in

Ziegler-Natta catalysis)(

Conclusion

Wilkinson's catalyst is a versatile and highly effective tool for the hydrogenation of alkenes,

offering both high selectivity and efficiency. Its mechanism, involving oxidative addition,

migratory insertion, and reductive elimination, highlights its importance in organic and

industrial chemistry. The bonding between metals and olefins plays a crucial role in many

catalytic processes, facilitating the activation of olefins toward reactions that would

otherwise be difficult to achieve. The synergistic nature of metal-olefin bonding, through π-

backbonding and σ-interactions, makes these complexes essential for various catalytic

applications.

8.(a) What is hapticity? Give an example of ligands of various hapticity.

(b) Give an account of preparation, properties and uses of organolithium compounds

Ans: Hapticity in Chemistry

Hapticity is a term used in coordination chemistry to describe how a ligand attaches to a

metal atom or ion. It refers to the number of atoms in a ligand that are directly bonded to

the metal center. The concept of hapticity is especially important in organometallic

chemistry, where ligands often bind through multiple atoms.

In simple terms, if a molecule (ligand) attaches to a metal ion through one atom, it has a

hapticity of 1 (monohapto, denoted as η¹), and if it attaches through more than one atom, it

has a higher hapticity (like η² for two atoms, η³ for three atoms, and so on).

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Different Types of Hapticity

The hapticity of a ligand is represented by the Greek letter η (eta) followed by a superscript

number indicating how many atoms of the ligand are attached to the metal center. Here are

some common hapticities with examples:

1. Monohapto (η¹): The ligand binds to the metal through one atom.

o Example: In many metal-alkyl compounds, such as methyl lithium (CH₃Li), the

carbon atom in the CH₃ group is attached to the lithium atom, making it an η¹

ligand.

2. Dihapto (η²): The ligand binds through two atoms.

o Example: Ethylene (C₂H₄) is an example of a dihapto ligand when it binds to

metals in compounds like Zeise's salt (K[PtCl₃(C₂H₄)]). Here, the two carbon

atoms of the ethylene molecule are attached to the metal, giving it an η²

hapticity.

3. Trihapto (η³): The ligand binds through three atoms.

o Example: In allyl complexes like [Cr(C₃H₅)(CO)₅], the three carbon atoms of

the allyl group are attached to the chromium atom, giving the ligand an η³

hapticity.

4. Pentahapto (η⁵): The ligand binds through five atoms.

o Example: Cyclopentadienyl (C₅H₅) is a common pentahapto ligand in

organometallic chemistry. In compounds like ferrocene (Fe(C₅H₅)₂), each

cyclopentadienyl ring binds to the iron atom through all five carbon atoms,

giving it an η⁵ hapticity.

5. Hexahapto (η⁶): The ligand binds through six atoms.

o Example: Benzene (C₆H₆) is an example of a hexahapto ligand when it forms a

complex with metals. In bis(benzene)chromium (Cr(C₆H₆)₂), each benzene

ring binds to the chromium atom through all six carbon atoms, making it an

η⁶ ligand.

Organolithium Compounds: Preparation, Properties, and Uses

1. Introduction to Organolithium Compounds

Organolithium compounds are organometallic compounds that contain a carbon-lithium (C–

Li) bond. These compounds are highly reactive and serve as important reagents in organic

synthesis. They are widely used in the formation of carbon-carbon bonds, which is

fundamental in the creation of complex molecules.

2. Preparation of Organolithium Compounds

Organolithium compounds are typically prepared by one of the following methods:

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(a) Reaction of Alkyl Halides with Lithium Metal

This is the most common method for preparing organolithium compounds. The reaction

involves treating an alkyl halide (R–X, where R is an alkyl group and X is a halogen like

bromine or chlorine) with metallic lithium in an organic solvent like hexane or ether.

• Example: Methyl lithium (CH₃Li) is prepared by reacting methyl bromide (CH₃Br) with

lithium metal:

(b) Metal-Halogen Exchange

In this method, a more reactive organolithium compound is used to replace a halide in an

alkyl halide, forming a new organolithium compound.

• Example: Phenyl lithium (C₆H₅Li) can be prepared by reacting n-butyllithium (C₄H₉Li)

with bromobenzene (C₆H₅Br):

In this method, a proton from an organic molecule (such as an aromatic compound) is

directly replaced by lithium. This process typically requires a strong base, like n-butyllithium,

to deprotonate the organic compound.

• Example: Phenyl lithium (C₆H₅Li) can be formed by treating benzene with n-

butyllithium:

3. Properties of Organolithium Compounds

Organolithium compounds have several unique properties due to the polar nature of the

carbon-lithium bond. These properties include:

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(a) High Reactivity

The C–Li bond is highly polar, making the carbon atom strongly nucleophilic (electron-rich)

and the lithium atom strongly electrophilic (electron-poor). This makes organolithium

compounds very reactive and useful for forming carbon-carbon bonds in organic synthesis.

(b) Aggregation

In the solid state and in non-polar solvents, organolithium compounds tend to form clusters

or aggregates. These aggregates can contain several lithium atoms and organic groups

bonded together. For example, methyl lithium (CH₃Li) often exists as a tetramer (four CH₃Li

units bonded together) in the solid state and in solution.

Organolithium compounds are generally soluble in non-polar solvents like hexane, ether, or

benzene. This solubility is important because many reactions involving organolithium

compounds take place in these solvents.

(d) Air and Moisture Sensitivity

Organolithium compounds are highly reactive toward air and moisture. In the presence of

moisture, they can react violently to form lithium hydroxide and the corresponding

hydrocarbon:

Because of this sensitivity, organolithium compounds are usually handled under an inert

atmosphere of nitrogen or argon and in anhydrous (water-free) solvents.

4. Uses of Organolithium Compounds

Organolithium compounds are highly versatile reagents with many important applications in

organic synthesis and industrial chemistry. Some key uses include:

(a) Formation of Carbon-Carbon Bonds

One of the most important uses of organolithium compounds is their ability to form carbon-

carbon bonds. They act as strong nucleophiles, attacking electrophilic carbon atoms (such as

those in carbonyl groups) to create new C–C bonds.

• Example: Organolithium compounds can react with aldehydes or ketones to form

alcohols. For instance, methyl lithium (CH₃Li) reacts with acetone (CH₃COCH₃) to

produce tert-butyl alcohol (C₄H₉OH):

(b) Synthesis of Complex Organic Molecules

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Organolithium reagents are widely used in the pharmaceutical and agrochemical industries

to synthesize complex organic molecules. They are key reagents in the preparation of drug

intermediates, natural products, and other biologically active compounds.

• Example: Organolithium compounds are used in the synthesis of anti-inflammatory

drugs and antibiotics.

Organolithium compounds are also used as initiators in the polymerization of certain

monomers, especially in the production of synthetic rubbers like polybutadiene and styrene-

butadiene rubber (SBR). These materials are widely used in the manufacture of tires and

other rubber products.

• Example: n-butyllithium is used as a catalyst in the polymerization of butadiene to

produce synthetic rubber:

(d) Preparation of Other Organometallic Compounds

Organolithium compounds are often used as starting materials to prepare other

organometallic reagents, such as Grignard reagents (RMgX) and organocopper compounds

(R₂CuLi). These reagents are widely used in organic synthesis for forming C–C bonds.

• Example: n-butyllithium (C₄H₉Li) can react with copper(I) iodide (CuI) to form lithium

di-n-butylcuprate (C₄H₉₂CuLi), which is used in nucleophilic substitution reactions.

5. Safety and Handling

Organolithium compounds are highly reactive and can be dangerous if not handled properly.

They are pyrophoric, meaning they can spontaneously ignite when exposed to air. Special

care is required when working with these compounds, including using inert gas atmospheres

(such as nitrogen or argon) and ensuring all equipment is completely dry.

Conclusion

Organolithium compounds are powerful tools in the field of organic and organometallic

chemistry. Their ability to form carbon-carbon bonds makes them invaluable reagents in the

synthesis of a wide variety of compounds, including pharmaceuticals, polymers, and fine

chemicals. However, due to their high reactivity, they must be handled with caution and

under controlled conditions. Understanding their preparation, properties, and uses provides

insight into their pivotal role in modern chemistry.

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