COORDINATION COMPOUNDS: CLASS XII CHEMISTRY

 

COORDINATION COMPOUNDS:

Alfred Werner (1866-1919) was a Swiss chemist who made significant contributions to the field of coordination chemistry. Here are the key points of his work and theories:

1. Coordination Compound Pioneer: Werner was the first chemist to formulate ideas about the structures of coordination compounds. He extensively prepared and characterized various coordination compounds, studying their physical and chemical properties through simple experimental techniques.
2. Primary and Secondary Valence: Werner introduced the concept of primary and secondary valence for metal ions in coordination compounds. In binary compounds like CrCl₃, CoCl₂, or PdCl₂, the primary valences are 3, 2, and 2, respectively. 

   The concepts of primary and secondary valence introduced by Alfred Werner in the context of binary compounds like CrCl3, CoCl2, and PdCl2: Primary Valence (Oxidation State): The primary valence, also known as the oxidation state, refers to the charge that a metal ion would have if all its ligands in a coordination complex were removed, leaving behind only the metal and its associated counterions. It is essentially the charge of the metal ion in the absence of coordination with ligands. This valence can be determined using the usual rules for assigning oxidation states, considering the charge of the counterions. For example: In CrCl3, the primary valence of chromium (Cr) is +3. This is because if all chloride ions (Cl-) were removed, Cr would have a +3 charge to balance the three chloride ions. In CoCl2, the primary valence of cobalt (Co) is +2. Removing the two chloride ions would leave Co with a +2 charge. In PdCl2, the primary valence of palladium (Pd) is also +2 for the same reason. Secondary Valence (Coordination Number): The secondary valence, also known as the coordination number, refers to the number of ligands directly bonded to the central metal ion in a coordination complex. It represents the maximum number of bonds that the metal ion can form with ligands. The coordination number can vary depending on the metal and the specific complex being considered. For example: In CrCl3, the coordination number is 6. This means that chromium is coordinated to six chloride ions in a typical complex. In CoCl2, the coordination number is 4. Cobalt forms coordination bonds with four chloride ions or other ligands in various complexes. In PdCl2, the coordination number can also be 4, indicating that palladium can form coordination bonds with four chloride ions or other ligands. Relation Between Primary and Secondary Valence: The primary valence (oxidation state) and the secondary valence (coordination number) are related because the coordination of ligands to the metal ion often results in charge redistribution. The coordination complex's overall charge must balance the charge of the central metal ion and the ligands. This balance is achieved through the coordination number, which dictates how many ligands are needed to neutralize or stabilize the metal ion's charge. For example, in a +3 oxidation state (primary valence) of chromium (Cr), it typically forms complexes with a coordination number of 6, where six ligands (e.g., chloride ions) are required to balance the charge of the Cr3+ ion.

3. Precipitation Reactions: In a series of experiments with cobalt(III) chloride and ammonia, Werner observed that some chloride ions could be precipitated as AgCl when excess silver nitrate was added in cold conditions, while others remained in solution.
4. Stoichiometry and Colors: Werner's experiments led to stoichiometric relationships and color changes in the coordination compounds:
• 1 mol CoCl₃·6NH₃ (Yellow) produced 3 mol AgCl.
• 1 mol CoCl₃·5NH₃ (Purple) produced 2 mol AgCl.
• 1 mol CoCl₃·4NH₃ (Green) produced 1 mol AgCl.
• 1 mol CoCl₃·4NH₃ (Violet) also produced 1 mol AgCl.
5. Explanation of Observations: To explain these observations and conductivity measurements, Werner proposed that:
• Six groups, including chloride ions and ammonia molecules, remained bonded to the cobalt ion during the reaction.
• The compounds could be formulated as shown in a table, where entities within square brackets represented a single unit that did not dissociate under reaction conditions.
6. Secondary Valence: Werner introduced the term "secondary valence" to represent the number of groups directly bound to the metal ion. In these examples, the secondary valence was consistently six.
7. Isomerism: Werner's work revealed that compounds with identical empirical formulas, such as CoCl₃·4NH₃, could have distinct properties. Such compounds are called isomers. 

   Werner's coordination theory, which you've described here, provides an excellent framework for explaining the observed phenomena in the series of cobalt(III) chloride compounds with ammonia. Let's break down the observations and explanations step by step: Yellow [Co(NH3)6]3+3Cl– (1:3 electrolyte): In this compound, all six ammonia (NH3) molecules remain bonded to the cobalt(III) ion, and no chloride ions dissociate. It forms a 1:3 electrolyte, which means for every cobalt ion (Co3+) that dissolves, three chloride ions (Cl-) accompany it. When you add silver nitrate (AgNO3), Ag+ ions react with the three Cl- ions to form AgCl (silver chloride) precipitate: Co(NH3)6]3+ + 3AgNO3 → [Co(NH3)6]3+  3Cl– + 3AgCl (s) Purple [CoCl(NH3)5]2+2Cl– (1:2 electrolyte): In this compound, five ammonia molecules remain bonded to the cobalt(III) ion, and two chloride ions dissociate. It forms a 1:2 electrolyte, which means for every cobalt ion that dissolves, two chloride ions accompany it. When you add silver nitrate, Ag+ ions react with the two Cl- ions to form AgCl precipitate: [CoCl(NH3)5]2+ + 2AgNO3 → [CoCl(NH3)5]2+ 2Cl– + 2AgCl (s) Green [CoCl2(NH3)4]+Cl– (1:1 electrolyte): In this compound, four ammonia molecules remain bonded to the cobalt(III) ion, and one chloride ion dissociates. It forms a 1:1 electrolyte, which means for every cobalt ion that dissolves, one chloride ion accompanies it. When you add silver nitrate, Ag+ ions react with the one Cl- ion to form AgCl precipitate: [CoCl2(NH3)4]+ + AgNO3 → [CoCl2(NH3)4]+ Cl– + AgCl (s) Violet [CoCl2(NH3)4]+Cl– (1:1 electrolyte): This is the same as the green compound, where four ammonia molecules remain bonded to the cobalt(III) ion, and one chloride ion dissociates. It also forms a 1:1 electrolyte, and when you add silver nitrate, Ag+ ions react with the one Cl- ion to form AgCl precipitate. Overall, the key points to understand are: Werner's theory introduced the concept of primary and secondary valences. Primary valences refer to the number of ions directly coordinated to the central metal ion, and secondary valences refer to the number of groups that remain attached to the metal ion during reactions. In each compound, the cobalt(III) ion has six secondary valences because it is surrounded by six ligands (either ammonia or chloride ions). The compounds exhibit different electrolytic behaviors and precipitation reactions based on the number of chloride ions that can be replaced by silver ions due to the coordination of ammonia molecules around the cobalt ion. The observed behaviors are consistent with the theory of coordination compounds and the concept of secondary valences proposed by Werner.

8. Werner's Coordination Theory (1898): In 1898, Werner formulated his theory of coordination compounds, with the following main postulates:
• Metals in coordination compounds exhibit two types of linkages: primary and secondary.
• Primary valences are typically ionizable and are satisfied by negative ions.
• Secondary valences are non-ionizable and are satisfied by neutral molecules or negative ions. The secondary valence is equal to the coordination number and is fixed for a metal.
• The ions or groups bound by secondary linkages to the metal have specific spatial arrangements corresponding to different coordination numbers.
9. Coordination Polyhedra: In modern terms, these spatial arrangements are referred to as coordination polyhedra.
10. Common Geometrical Shapes: Werner proposed that octahedral, tetrahedral, and square planar geometrical shapes are more common in coordination compounds of transition metals. For example:
• [Co(NH₃)₆]³⁺, [CoCl(NH₃)₅]²⁺, and [CoCl₂(NH₃)₄]⁺ are octahedral entities.
• [Ni(CO)₄] and [PtCl₄]^2– are tetrahedral and square planar, respectively.

Werner's groundbreaking work laid the foundation for our understanding of coordination chemistry and the complex structures of coordination compounds. His concepts and theories remain fundamental in modern chemistry.

 

CONCEPT OF DOUBLE SALT:

Double Salts:

1. Definition: Double salts are compounds formed by the combination of two or more stable compounds in a stoichiometric ratio.
2. Dissociation: Double salts, when dissolved in water, completely dissociate into simple ions. For example, KCl.MgCl₂.6H₂O dissociates into K⁺, Mg²⁺, Cl^-, and H₂O ions.
3. Examples: Examples of double salts include carnallite (KCl.MgCl₂.6H₂O), Mohr’s salt (FeSO₄.(NH₄)2SO₄.6H₂O), and potash alum (KAl(SO₄)₂.12H₂O).

Complexes:

1. Definition: Complexes are compounds where a central metal atom or ion is bonded to a fixed number of ions or molecules, called ligands.
2. Dissociation: Complex ions do not dissociate into individual metal and ligand ions when dissolved in water. For example, [Fe(CN)₆]^4- in K₄[Fe(CN)₆] remains intact. 

   Composition and Formation: Double Salt: Double salts are formed by the combination of two or more distinct salts or ionic compounds in a fixed stoichiometric ratio. They retain their identity even in the solid state and have a well-defined chemical formula. For example, carnallite (KCl.MgCl2.6H2O) is a double salt formed by combining potassium chloride (KCl) and magnesium chloride (MgCl2) with water molecules. Complex: Complexes are formed by the coordination of a central metal ion with surrounding ligands (molecules or ions). The ligands form coordinate bonds with the metal ion. Complexes often have a specific structure, but their composition may not always be as well-defined as that of double salts. For example, [Fe(CN)6]4– is a complex ion formed by the coordination of six cyanide (CN–) ligands with an iron (Fe) ion. Dissociation Behavior: Double Salt: When double salts are dissolved in water, they readily dissociate into their constituent ions. The ions are fully separated in solution, and they can conduct electricity. For example, when carnallite (KCl.MgCl2.6H2O) dissolves in water, it dissociates into K+, Mg2+, Cl– ions, and water molecules. Complex: Complex ions do not dissociate into their constituent ions when dissolved in water. Instead, they remain intact as a single species in solution. The coordination bonds between the central metal ion and the ligands are quite strong, preventing dissociation. For example, [Fe(CN)6]4– does not dissociate into Fe2+ and CN– ions in solution; it exists as a stable complex. Electrolytic Behavior: Double Salt: Due to their complete dissociation into ions, double salts exhibit strong electrolytic behavior in aqueous solutions. They conduct electricity efficiently because of the presence of free ions. Complex: Complexes do not exhibit significant electrolytic behavior in solution because they do not dissociate into free ions. Consequently, they do not conduct electricity to the same extent as double salts. Examples: Double Salt Examples: Carnallite (KCl.MgCl2.6H2O), Mohr's salt (FeSO4.(NH4)2SO4.6H2O), potash alum (KAl(SO4)2.12H2O), etc. Complex Examples: [Fe(CN)6]4– of K4[Fe(CN)6], [Cu(NH3)4]2+ of Cu(NH3)4SO4, [PtCl4]2– of K2[PtCl4], etc. In summary, the primary distinction between double salts and complexes lies in their behavior in solution. Double salts dissociate into their constituent ions upon dissolution, while complexes remain as intact entities. This behavior is due to the difference in the strength of the chemical bonds involved. Double salts involve ionic bonds, while complexes involve coordinate covalent bonds between the central metal ion and the ligands.

3. Coordination Entity: A coordination entity in a complex consists of the central metal atom or ion surrounded by ligands. Examples include [CoCl₃(NH₃)₃] and [NiCl₂(H₂O)₄]. 

   Definition of a Coordination Entity: A coordination entity is a molecular species consisting of a central metal atom or ion bonded to a fixed number of ions or molecules. This grouping is often referred to as a coordination complex or coordination compound. Central Metal Atom or Ion: At the core of a coordination entity, there is a central metal atom or ion. This metal atom or ion is typically a transition metal due to its ability to form various coordination bonds with other molecules or ions. Ligands: Surrounding the central metal atom or ion are ligands. Ligands are molecules or ions that coordinate with the central metal through coordinate covalent bonds. These bonds involve the donation of electron pairs from the ligands to the metal atom or ion. Ligands can be neutral molecules (such as ammonia, NH3) or negatively charged ions (such as chloride, Cl-). Coordination Number: The fixed number of ligands directly bonded to the central metal atom or ion is known as the coordination number. It represents the maximum number of bonds the metal can form with ligands. In the example you provided, [CoCl3(NH3)3], the coordination number of cobalt (Co) is 6 because it is surrounded by three ammonia (NH3) molecules and three chloride (Cl-) ions. Examples of Coordination Entities: Coordination entities can vary in composition and coordination number. Some examples include: [Ni(CO)4]: In this coordination entity, the central metal is nickel (Ni), and it is coordinated to four carbon monoxide (CO) ligands. The coordination number of nickel in this case is 4. [PtCl2(NH3)2]: Here, platinum (Pt) serves as the central metal, and it is coordinated to two ammonia (NH3) molecules and two chloride (Cl-) ions, resulting in a coordination number of 4. [Fe(CN)6]4-: In this coordination entity, the central metal is iron (Fe), and it is coordinated to six cyanide (CN-) ions. The coordination number of iron is 6. [Co(NH3)6]3+: In this example, cobalt (Co) is the central metal, and it is coordinated to six ammonia (NH3) molecules, giving it a coordination number of 6.

4. Central Atom/Ion: The central atom or ion in a coordination entity is the one to which ligands are bonded. Examples include Ni²⁺ (in [NiCl₂(H₂O)₄]), Co³⁺ (in [CoCl(NH₃)₅]²⁺), and Fe³⁺ (in [Fe(CN)₆]^3-).
5. Ligands: Ligands are ions or molecules that are bound to the central atom/ion in a coordination entity. They can be unidentate (e.g., Cl-, H₂O, NH₃), didentate (e.g., ethane-1,2-diamine, oxalate), polydentate (e.g., EDTA^4-), or chelate ligands (ligands that use two or more donor atoms simultaneously to bind a metal ion).
6. Coordination Number: The coordination number (CN) of a metal ion in a complex is the number of ligand donor atoms directly bonded to it. For example, in [PtCl₆]^2-, Pt has a coordination number of 6, and in [Ni(NH₃)₄]²⁺, Ni has a coordination number of 4.
7. Ambidentate Ligands: Some ligands, like NO^2- and SCN^-, have multiple donor atoms and can coordinate through different atoms to the central metal ion.

Double salts are compounds that fully dissociate into simple ions in water, while complexes are formed by the coordination of a central metal atom/ion with ligands and do not dissociate into their constituent ions when dissolved. Coordination entities, central atoms/ions, ligands, coordination numbers, and the concept of ambidentate ligands are important aspects of complex chemistry.

 

(8) Coordination Number Determination:

• The coordination number of the central atom/ion is determined solely by the number of sigma (σ) bonds formed by the ligands with the central atom/ion.
• Pi (π) bonds, if present between the ligand and the central atom/ion, are not considered when determining the coordination number.

(9) Coordination Sphere:

• The coordination sphere comprises the central atom/ion and the ligands attached to it.
• It is collectively enclosed in square brackets in chemical notation.
• The ionizable groups are written outside the bracket and are referred to as counter ions.
• For example, in the complex K₄[Fe(CN)₆], the coordination sphere is [Fe(CN)₆]^4- and the counter ion is K⁺.

(10) Coordination Polyhedron:

• The spatial arrangement of the ligand atoms directly attached to the central atom/ion defines a coordination polyhedron around the central atom/ion.
• Common coordination polyhedra include octahedral, square planar, and tetrahedral shapes.
• Examples:
• [Co(NH₃)₆]³⁺ exhibits an octahedral coordination polyhedron.
• [Ni(CO)₄] has a tetrahedral coordination polyhedron.
• [PtCl₄]^2- displays a square planar coordination polyhedron.
• The shapes of different coordination polyhedra are illustrated in Figure 5.1.

In summary, the coordination number is determined by the number of sigma bonds formed between the central atom/ion and the ligands. The coordination sphere encompasses the central atom/ion and the ligands enclosed in square brackets, while the coordination polyhedron represents the spatial arrangement of ligand atoms directly attached to the central atom/ion and can take various shapes such as octahedral, square planar, or tetrahedral.

(11) Oxidation Number of Central Atom:

• The oxidation number of the central atom in a complex is determined by considering the charge it would carry if all the ligands were removed, along with the electron pairs shared with the central atom.
• This oxidation number is indicated by a Roman numeral in parentheses following the name of the coordination entity.
• For instance, the oxidation number of copper in [Cu(CN)₄]^3- is +1, and it is represented as Cu(I).

(12) Homoleptic and Heteroleptic Complexes:

• Homoleptic Complexes: These are complexes in which a metal is bound to only one kind of donor group. For example, [Co(NH₃)₆]³⁺ is a homoleptic complex because it contains only ammonia (NH₃) ligands binding to the central cobalt ion.
• Heteroleptic Complexes: These are complexes in which a metal is bound to more than one kind of donor group. For example, [Co(NH₃)₄Cl₂]⁺ is a heteroleptic complex because it includes both ammonia (NH₃) and chloride (Cl^-) ligands binding to the central cobalt ion.

The oxidation number of the central atom in a complex reflects the charge it would carry after the removal of ligands and shared electron pairs. Roman numerals denote this oxidation number in the complex's name. Additionally, complexes are categorized as homoleptic when they involve only one type of ligand and heteroleptic when they contain multiple types of ligands bound to the central metal ion.

 

 

Importance of nomenclature in coordination chemistry and the role of IUPAC recommendations:

1. Importance of Nomenclature:
• Nomenclature is crucial in coordination chemistry to provide a systematic and unambiguous way of describing the formulas and names of coordination entities.
• It ensures clarity and consistency in communication among chemists, researchers, and educators.
2. Dealing with Isomers:
• Coordination compounds often exhibit various isomers, which are compounds with the same molecular formula but different structural arrangements or spatial orientations.
• Nomenclature is essential for distinguishing between these isomers and accurately representing their structures.
3. IUPAC Recommendations:
• The International Union of Pure and Applied Chemistry (IUPAC) plays a central role in establishing guidelines and recommendations for naming coordination compounds.
• These recommendations are widely accepted and adopted by the global scientific community, providing a standardized nomenclature system.
4. Systematic Names:
• IUPAC guidelines ensure that coordination compounds are given systematic names that reflect their composition and structure.
• This systematic approach allows chemists to deduce the nature of the compound from its name and vice versa.

Nomenclature in coordination chemistry is essential for maintaining clarity, consistency, and accuracy in describing coordination compounds, especially when dealing with isomers. The recommendations provided by IUPAC serve as a globally accepted standard for naming these compounds systematically.

 

Mononuclear coordination entities:

 (i) Central Atom First:

• The formula begins with the central atom, naming it first.

(ii) Ligands in Alphabetical Order:

• Next, list the ligands in alphabetical order.
• The charge of the ligand does not influence its position in the alphabetical list.

(iii) Handling Polydentate Ligands:

• For polydentate ligands, follow alphabetical order as well.
• In the case of abbreviated ligands, use the first letter of the abbreviation to determine their alphabetical position.

(iv) Use of Square Brackets and Parentheses:

• Enclose the formula for the entire coordination entity, whether it is charged or not, within square brackets.
• If the ligands are polyatomic, enclose their formulas in parentheses.
• Ligand abbreviations should also be enclosed in parentheses.

(v) No Space Between Ligands and Metal:

• Ensure there is no space between the ligands and the central metal atom/ion within the coordination sphere.

(vi) Indicating Charge:

• If the coordination entity is charged, indicate the charge outside the square brackets as a right superscript with the number appearing before the sign.
• For example, [Co(CN)₆]^3-, [Cr(H₂O)₆]³⁺, etc.

(vii) Balancing Charges:

• In a complex containing both cations and anions, the charge of the cation(s) should be balanced by the charge of the anion(s).

These rules ensure a systematic and consistent method for representing the formulas of mononuclear coordination entities, making it easier to understand their composition and charge.

 

 

Rules and examples for naming coordination compounds:

 (i) Cation First:

• Start by naming the cation, whether it is positively or negatively charged, before naming the anion.

(ii) Ligand Alphabetical Order:

• List the ligands in alphabetical order, ahead of the central atom/ion's name.
• This order is the reverse of writing the formula.

(iii) Ligand Names:

• Anionic ligands have names ending in "-o."
• Neutral and cationic ligands retain the same name as the ligand, except for specific ligands like "aqua" for H₂O, "ammine" for NH₃, "carbonyl" for CO, and "nitrosyl" for NO.
• When writing the formula, these names are enclosed in brackets.

(iv) Use of Prefixes:

• Utilize prefixes "mono," "di," "tri," etc., to indicate the number of individual ligands in the coordination entity.
• If ligand names include numerical prefixes, use terms like "bis," "tris," "tetrakis," etc., placing the ligand in parentheses. For example, [NiCl₂(PPh₃)₂] is named as dichloridobis(triphenylphosphine)nickel(II).

(v) Oxidation State:

• Indicate the oxidation state of the metal in the coordination entity with a Roman numeral in parentheses.

(vi) Metal Naming Conventions:

• For cationic complex ions, name the metal the same as the element (e.g., Co for cobalt, Pt for platinum).
• For anionic complex ions, add the suffix "-ate" to the metal name (e.g., cobaltate for Co).
• For certain metals, use Latin names in complex anions (e.g., ferrate for Fe).

(vii) Neutral Complex Molecules:

• Name neutral complex molecules similarly to complex cations.

Examples:

1. [Cr(NH₃)₃(H₂O)₃]Cl₃ is named as:
• Triamminetriaquachromium(III) chloride
• Explanation: The cationic complex ion is named first. Ligands (ammine and aqua) are listed alphabetically. The oxidation state of chromium (III) is indicated in parentheses.
2. [Co(H₂NCH₂CH₂NH₂)₃]2(SO₄)₃ is named as:
• Tris(ethane-1,2–diamine)cobalt(III) sulphate
• Explanation: The anion is sulphate. The charge on the complex ion is determined by the charge balance with the anion.
3. [Ag(NH₃)₂][Ag(CN)₂] is named as:
• Diamminesilver(I) dicyanidoargentate(I)
• Explanation: Two different complex ions are present. The first ion is named as a cation, while the second is named as an anion.

These rules for naming coordination compounds provide a systematic and consistent way to describe their composition, charge, and structural components in a clear and organized manner.