Ch-9 Chemistry by pragya joshi

CHAPTER

Coordination Compounds In EDTA, a metal ion, two oxygen atoms and two nitrogen a toms comprise a square O

O C CH2

O CH2

C O Pb

CH2 CH2

N CH2

C O

Heavy metal toxicity in humans is the cause of many health disorders. The medical condition caused by increased levels of the heavy metal lead in the body is called lead poisoning and is known to interfere with a variety of body processes and is toxic to many organs and tissues. Chelation therapy using EDTA is the medically accepted treatment for lead poisoning.

CH2

O C

How does EDTA clear lead poisoning?

Conceptual Objectives • Explanation of Werner’s theory of coordination compounds. • Definitions of some important terms related to coordination compounds. • The nomenclature of coordination compounds. • The types of isomerism exhibited by coordination compounds.

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• The bonding in coordination compounds − Valence Bond Theory (VBT) and Crystal Field Theory (CFT). • The bonding in metal carbonyls. • The stability and important applications of coordination compounds.

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Chapter 9 Coordination Compounds Coordination chemistry is the branch of chemistry which distinctively deals with the study of coordination compounds for their chemical, structural, magnetic and spectral properties. Coordination compounds are a special class of compounds that consist of a central metal atom or ion, which is surrounded by oppositely charged ions or neutral molecules in more than its normal valence. They are also called complex compounds or simply complexes. There exist a large number of biologically important coordination compounds having complex organic species bound to a central metal ion. For example, 1. Chlorophyll: A green pigment vital for photosynthesis is a coordination complex of magnesium. 2. Hemoglobin: An important metalloprotein is a coordination complex of iron. 3. Vitamin B12: A coordination complex where central metal atom is cobalt. The use of coordination compounds extends to analytical chemistry, metal extraction and catalysis. They also find application in medicinal chemistry.

9.1 Werner’s Theory of Coordination Compounds The first systematic study on physical, chemical, structural and behavioral aspects of coordination compounds was done by Swiss chemist, Alfred Werner (1866−1919). He proposed his theory for coordination compounds in 1893 and after his painstaking work of 20 years won the Nobel Prize in Chemistry in 1913. The main postulate of the theory is that the metal ion encapsulated inside the coordination entity exhibits two types of valences: primary valence and secondary valence. 1. The primary valence is ionizable as well as non-directional. It corresponds to the oxidation state in the modern-day terminology. Normally, if a complex ion exhibits positive charge, the primary valence corresponds to the number of charges present on the complex ion and this charge is balanced by the same number of negative ions. Thus, primary valence can also be defined by the number of anions neutralizing the charge on the complex ion. The secondary valence is directional, non-ionizable and corresponds to coordination number in the modernday terminology. It equals the total number of ligands coordinately bonded to the central metal ion inside the coordination sphere. The metal complex first satisfies its secondary valence and then its primary valence. Every metal ion has fixed number of secondary valences, which are directed towards fixed positions in space around the central metal atom. This results in definite geometry and stereochemistry of the complex. 2. The spatial arrangement of coordination compounds is called coordination polyhedral. The common geometrical shapes for coordination compounds of transition metals are octahedral (for coordination number six,); tetrahedral and square planar (for coordination number four). 3. The coordination compounds are represented as [Ni(NH3)6]Cl2 where the entity in the square bracket is the coordination complex and the ions outside the bracket are called counter ions. Based on this theory, Werner deduced and explained the structure of various cobalt amines. He observed that when the compounds of cobalt (III) chloride reacted with ammonia, in the presence of excess silver nitrate, some chloride ions could be precipitated in cold, while some remained in solution. The difference in the color of solution and amount of AgCl produced was explained by the formation of following compounds: 1. CoCl3  6NH3: The central metal ion, cobalt, is in +3 oxidation state having primary valence of 3 and secondary valence (coordination number) of 6, which are satisfied, respectively, by three anionic Cl- ligands and six neutral NH3 ligands. This can be verified by the fact that on addition of AgNO3, three equivalents of AgCl are obtained. Primary valency ↑

Secondary valency ↑

[Co(NH3 )6 ] Cl3 + 3 Ag+ → [Co(NH3 )6 ]3+ + 3AgCl Coordination complex ( yellow )

Ions

Complex ion

3 equivalents

2. CoCl3  5NH3: Here again cobalt ion has secondary valence as 6 and primary valence as 3. The secondary valence is satisfied with five neutral NH3 ligands and one anionic Cl- ligand, which has dual character and satisfies both primary as well as secondary valences. The remaining primary valence of 2 is satisfied by two Cl- ions.

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9.1 Werner’s Theory of Coordination Compounds 3 Primary valency ↑

Secondary valency ↑

[CoCl(NH3 )5 ] Cl2 + 2 Ag+ → [CoCl(NH3 )5 ]2+ + 2AgCl Ions

Coordination complex (purple)

2 equivalents

Complex ion

On addition of AgNO3, only two equivalents of the AgCl are obtained. 3. CoCl3  4NH3: Here two of the Cl- ions exhibit dual character and satisfy both primary as well as secondary valences. In this case, only one Cl- will get precipitated out on the addition of AgNO3. Primary valency

Secondary valency ↑

↑

[CoCl2 (NH3 )4 ] Cl

Coordination complex (green)

Ag+ → [CoCl2 (NH3 )4 ]+ + AgCl Ions

Complexion

1 equivalent

4. CoCl3  3NH3: Here all the three chloride ions are satisfying primary as well as secondary valences, so none of the Cl- ions can be precipitated out by addition of AgNO3. Dual character ↑

[CoCl3 (NH3 )3 ] + Ag+ → No reaction Coordination complex (violet)

Ions

Difference Between Double Salt and a Complex Addition compounds are formed when two or more stable compounds join together in stoichiometric amounts. For example: KCl + MgCl2 + 6H2O → KCl  MgCl2  6H2O ( Carnallite )

K 2SO 4 + Al2 (SO 4 )3 + 24H2O → K 2SO 4  Al2 (SO 4 )3  24H2O (Potassium alum)

CuSO 4 + 4NH3 + H2O → CuSO 4  4NH3  H2O (Tetrammine copper (II) sulphate monohydrate)

Fe(CN)2 + 4KCN → Fe(CN)2  4KCN Potassium ferrocyanide

Addition compounds are of two types: 1. Double salts: Those which lose their identity in solution. 2. Complexes: Those which retain their identity in solution. When crystals of carnallite are dissolved in water, the solution shows the properties of K+, Mg2+ and Cl- ions. In the similar way, a solution of potassium alum shows the properties of K+, Al3+ and SO42- ions. These are both examples of double salts, which exist only in the crystalline state. When the other two compounds are dissolved in water, they do not form simple ions—Cu2+, or Fe2+ and CN-—but instead their complex ions remain intact. Thus, the cuproammonium ion [Cu(H2O)2(NH3)4]2+ and the ferrocyanide ion [Fe(CN)6]4- exist as distinct entities both in solid and in solution. These are complex ions and are represented by the use of square brackets. Compounds containing these ions are called coordination compounds.

Identify the primary and secondary valences in the following complexes: [Co(NH3)6]Cl3, [Co(NH3)4Cl2]Cl, [CoCl3(NH3)3]

Worked Problem 9.1

solution

[Co(NH3)6]Cl3: Primary valence is 3 and secondary valence is 6. [Co(NH3)4Cl2]Cl: Primary valence is 1 and secondary valence is 6. [CoCl3(NH3)3]: Primary valence is zero and secondary valence is 6.

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Chapter 9 Coordination Compounds Worked Problem 9.2

Explain why the complexes PtCl4  2KCl and PtCl2  2NH3 do not form precipitate of silver chloride when treated with AgNO3. solution

The complex PtCl4  2KCl has the structural formula K2[PtCl6] and the complex PtCl2  2NH3 has the structural formula [Pt(NH3)2Cl2]. In both the complexes, chloride ions are present as part of the coordination sphere and cannot be ionized. Hence, they do not form precipitate with AgNO3 solution.

Practice Problems 1. How many ions per mole would be produced from the following complexes in solution: [Co(NH3)4Cl2]Cl, [Ni(H2O)6]Cl2, [Ni(CO)4], K4[Fe(CN)6] and [Pt(NH3)4][PtCl4],

2. How many chloride ions will be precipitated when the complex CoCl3  4NH3 is treated with AgNO3 solution?

9.2 Some Important Terms Pertaining to Coordination Compounds There are certain terms that are associated with coordination compounds and will be used throughout the chapter. The explanation of these terms is as follows: 1. Coordination entity: It is the central metal atom or ion which is bonded to a definite number of ions or molecules which is fixed. For example, in [Co(NH3)6]Cl3, a coordination entity, six ammonia molecules are surrounded by three chloride ions. 2. Central atom/ion: It is the central cation that is surrounded and coordinately bonded to one or more neutral molecules or negatively charged ions in a definite geometric arrangement. For example, in the complex [Co(NH3)6]Cl3, Co3+ represents the central metal ion which is positively charged and is coordinately bonded to six neutral NH3 molecules within the coordination sphere. The central metal/ion is also referred to as Lewis acid. 3. Ligands: It can be an atom, ion or molecule that binds to a central metal ion to form a coordination complex. The bonding between the metal and the ligand generally involves donation of one or more electrons from the ligand to the central metal atom. The metal−ligand bonding ranges from covalent to more ionic. Thus, ligands can be viewed as Lewis bases. The ligands bonding with the metal ion inside the coordination sphere can be neutral or anionic in nature; however, those bonding the metal ion outside the coordination sphere are bound to be anionic in nature, thereby satisfying the primary valence of the central metal ion. 4. Coordination number: The total number of ligands coordinately bonded to the central metal atom or ion in the primary coordination sphere represents the coordination number of that complex. The coordination number of the complex is fixed and represents the number of ligands that are arranged around the central metal cation in a particular spatial arrangement (Table 9.1). This implies that coordination number of the complex is the deciding factor for the geometry attained by the complex. 5. Coordination sphere: Primary coordination sphere of the coordination complex is represented by the central metal ion and the ligands coordinately bonded to it. It is that part of the complex which remains as the single entity, that is, it does not lose its identity and is non-ionizable. This part is written in square brackets to represent its autonomous identity. [Ni(NH3 )6 ] Cl2 

Coordination Counter sphere ion

6. Counter ion: The negative ions that are not part of the coordination sphere are called counter ions. These are the species that are written outside the square bracket, which defines the

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9.2 Some Important Terms Pertaining to Coordination Compounds 5

Table 9.1 Coordination number of various complexes

Complex

Coordination Number

[Ni(NH3)6]Cl2

[Co(NH3)6]Cl3

[Fe(CN)6]4-

[Ni(CO)4]

[Co(en)3]

coordination sphere of the coordination complex. For example, in the complex [Ni(NH3)6]Cl2, the 2 Cl- represent the counter ions that form the ionizable part of the complex. These also represent the primary valence of the coordination compound. When this complex is dissolved in solution of AgNO3, there is precipitation of AgCl which occurs due to reaction of Ag+ ions with the counter ions (Cl-). [Ni(NH3 )6 ]Cl2

Study tip

The central atom/ion’s coordination number is determined only by the number of s bonds (and not the p bonds) formed between the ligand and the central atom/ion.

Central Ligand Counter metal ion ion Counter ion

↑

2 Ag+

Coordination complex

Ions

↑

[Ni(NH3 )6 ] Cl2

[Ni(NH3 )6 ]2+ + 2AgCl Complex ion

Precipitate

7. Donor atom: The coordinating atom of the ligand which is actually donating electron pair to the central metal ion is called a donor atom. For example, in case of neutral ligands such as H2O or NH3, oxygen and nitrogen are donor atoms (Fig. 9.1). Donor atom

O H

Donor atom

Lone pair

Figure 9.1 Donor atoms and lone pairs in H2O and NH3. 8. Coordination polyhedron: A coordination polyhedron is the spatial arrangement of the ligand atoms that are directly attached to the central atom/ion. For example, [Co(NH3)6]3+ is octahedral, [Ni(CO)4] is tetrahedral) and [PtCl4]2 is square planar. Figure 9.2 shows some common coordination polyhedra.

Square planar

Tetrahedral

Trigonal bipyramidal

Square pyramidal

Octahedral

Figure 9.2 Shapes of different coordination polyhedra.

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Chapter 9 Coordination Compounds 9. O xidation number of central atom: When all the ligands are removed along with the electron pairs that are shared with the central atom, the charge that the central atom would carry is called the oxidation number. For example, the oxidation number of copper in [Cu(CN)4]3- is +1. It is represented in Roman numeral in parenthesis after the name of the central atom. For example, for the above mentioned complex, it is written as Cu(I). 10. Homoleptic and heteroleptic complexes: When the central metal ion is bound to only one kind of donor group, the complex is called homoleptic, for example, [Co(NH3)6]3+. The complex in which the metal atoms are bonded to more than one type of donor group is called heteroleptic species. For example, [Co(NH3)4Cl2]+ is called heteroleptic species.

Ligands Ligands are classified in many ways according to their charge, size (bulk), type of the coordinating atom(s) and the number of electrons donated to the metal (denticity or hapticity). In this subsection, we are going to differentiate between ligands on the basis of their denticity. Denticity refers to the number of times a ligand bonds to a central metal ion through noncontiguous donor sites. There are many ligands capable of binding metal ions via multiple sites, since the ligands have lone pair of electrons on more than one atom. A ligand that binds through two sites is classified as didentate; that through three sites as tridentate and so on. In general, the ligands binding via more than one site are referred to as polydentate ligands. 1. Unidentate ligands: These ligands are capable of binding to the central metal ion via single donating site or atom. They contain only one atom capable of donating the lone pair of electrons to the metal center. The number of such known ligands is high, but their stability is usually less than polydentate ligands. These include Cl-, Br-, I-, F-, NO3-, N3-, OH-, H2O, C5H5N, NH3, NO2-, etc. 2. Didentate ligands: These ligands are capable of binding to the central metal ion through two binding sites. The complexes formed from these ligands are higher in stability than those formed by unidentate ligands, keeping the central metal ion same. The most common example of didentate ligand is ethylenediamine, which can coordinate with the metal ion via two of its nitrogen atoms, each having a lone pair of electrons. This is a neutral ligand, which forms a chelate-like structure when bonding in didentate fashion. For example, nickel(II) ion can form six such bonds, so a maximum of three ethylenediamine molecules can be attached to one Ni2+ ion. ii

Study tip

Complexes of polydentate ligands are called chelate complexes. They tend to be more stable than complexes derived from monodentate ligands.

 Key Point

Effective atomic number (EAN) is the total number of electrons neighboring the nucleus of a metal atom in a coordination complex. It is composed of the metal atom’s electrons and the bonding electrons from the surrounding ligands. For example, the EAN of the cobalt atom in the complex [Co(NH3)6]3+ is 36, which is obtained by calculating the sum of the number of electrons in the trivalent cobalt ion (24) and the number of bonding electrons from six surrounding ammonia molecules, each of which contributes an electron pair (2 × 6 =12).

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NH2 − CH2 − CH2 − NH2 Ethylenediamine

CH2 CH2

NH2

CH2

NH2

OH2 Ni OH2

OH2

CH2

NH2

OH2

CH2

NH2

OH2 Ni OH2

NH2

CH2

NH2

CH2

NH2

NH2 Ni NH2

CH2 NH 2 NH2 CH2

CH2 Chelate with one ethylenediamine ligands

Chelate with two ethylenediamine ligands

Chelate with three ethylenediamine ligands

3. Polydentate ligands: These include tridentate, tetradentate, pentadentate, hexadentate and all types of ligands that bind to the metal ion through multiple bonding sites. These ligands form chelate-like ring structures which are kinetically as well as thermodynamically very stable. The most important example of this type includes the hexadentate ligand, ethylenediamine tetraacetic acid (EDTA). The bonding in this ligand occurs through two nitrogen atoms and four oxygen atoms as shown in Fig. 9.3: 4. Ambidentate ligands: These ligands are capable of ligating through two different atoms. For example, (a) NO2-: Nitro or Nitro-N − Bonding occurs through N. (ONO-)/NO2-: Nitrito or Nitro-O − Bonding occurs through O. (b) CN-: Cyano − Bonding occurs through C. NC-: Isocyano − Bonding occurs through N.

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9.3 Nomenclature of Coordination Compounds 7 O O

O C

CH2

HO HO

CH2

O−

C OH

N N

M O−

Ethylenediamine ligand with various binding sites

O O− O− O

EDTA forming chelate-like structure when bonded to the metal ion in hexadentate fashion

Figure 9.3 Bonding in EDTA.

Practice Problems 3. In the following complex, indicate the ligands and the central metal ion, its coordination number and oxidation state: [Co(NO2)2(C5H5N)2(NH3)2]NO2.

4. Indicate the coordinate number and the oxidation state of iron in the following complexes: (a) K4[Fe(CN)6] (b) Fe(CO)6 (c) Fe(EDTA)]- (d) [Fe(CN)6]3-

9.3 Nomenclature of Coordination Compounds The general rules for naming the coordination complexes and writing their formulae have been set by IUPAC to avoid any ambiguity.

Formulas of Mononuclear Coordination Entities Mononuclear coordination entities contain a single central metal atom. The formula of a mononuclear coordination compound is a concise representation of the constituent central atom and all elements/groups coordinated with it in primary and secondary valences. The general rules for writing the formula areas : 1. In a coordination formula, the central atom is listed first. 2. The formally anionic ligands appear next, listed in alphabetical order according to the first symbols of their formulae. The neutral ligands follow, also in alphabetical order, according to the same principle. 3. The formula of the entire coordination entity, whether charged or not, is enclosed in square brackets. 4. If the coordination entity is negatively charged, the formula is preceded by the cationic formula. 5. When ligands are polyatomic, their formulae are enclosed in parentheses. Ligand abbreviations are also enclosed in parentheses. Polydentate ligands are also listed alphabetically. When the ligand is abbreviated, the first letter of the abbreviation is used to determine the position of the ligand in alphabetical order. 6. The ligands and the metal within a coordination sphere should be written without a space. 7. The charge in a coordination entity without a counter ion is written outside the square brackets as a right superscript with the number before the sign. For example, [Co(CN)6]3-. 8. The cation charges are balanced by the anion charges.

Naming of Mononuclear Coordination Compounds Naming of coordination compounds follows additive nomenclature, that is, the groups surrounding the metal atom are identified in the name and are given as prefixes.

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Chapter 9 Coordination Compounds

1. The name of the ligands is listed before the name(s) of the central atom(s). 2. The name of the cationic part is written first followed by the anionic part. 3. No space is left between names that refer to the same coordination entity. 4. Ligand names are listed in the alphabetical order (multiplicative prefixes indicating the number of ligands are not considered in determining that order). This is opposite to that of writing a formula. 5. Prefixes are used to designate the number of each type of ligand in the complex ion; for example, di-, tri- and tetra-. If the ligand already contains a prefix (e.g., ethylenediamine) or if it is a polydentate ligand then Greek prefixes bis-, tris-, tetrakis-, pentakis-, etc. are used instead. Number Prefix

mono

di (bis)

tri (tris)

tetra penta hexa (tetrakis) (pentakis) (hexakis)

hepta

octa

nona (ennea)

deca

undeca

dodeca

6. For naming the central metal if the complex ion is a cation, the metal is named same as the element. For example, Co in a complex cation is named cobalt and Pt is named platinum. If the complex ion is an anion, the name of the metal ends with the suffix -ate. For example, Co in a complex anion is called cobaltate and Pt is called platinate. For some metals, the Latin names are used in the complex anions, for example, Fe is called ferrate (not ironate). 7. Anionic ligands should end in “-o”. Thus, anions that end in “-ide” (e.g., chloride), “-ate” (e.g., sulphate, nitrate), and “-ite” (e.g., nitrite) should be changed into -ido, -ato and -ito, respectively. 8. For neutral ligands, the common name of the molecule is used, for example, H2NCH2CH2NH2 (ethylenediamine). Important exceptions include: water called as “aqua”, ammonia called as “ammine”, carbon monoxide called as “carbonyl”, N2 called as “dinitrogen” and O2 called as “dioxygen”. Table 9.2 shows the names of anionic and neutral ligands. 9. The oxidation state of the metal is written as the Roman numeral in parenthesis. Table 9.2 The naming convention of anionic and neutral ligands

Anionic Ligands

Naming Convention

Neutral Ligands

Naming Convention

Br-

Bromido

NH3

Ammine

F-

Fluorido

H2O

Aqua

Oxo

Nitrosyl

Hydroxo

Carbonyl

CN-

Cyanido

Dioxygen

Oxalato

Dinitrogen

Carbonato

C5H5N

Pyridine

Acetato

H2NCH2CH2NH2

Ethylenediamine

C2O

CH3COO

Some examples of naming coordination compounds are as follows: 1. [Cr(NH3 )3 (H2O )3 ]Cl3 IUPAC name: triamminetriaquachromium(III) chloride Since there are three chlorides binding with the complex ion, the charge on the complex ion must be +3. From the charge on the complex ion and the charge on the ligands, we can calculate the oxidation number of the metal. As all the ligands are neutral molecules, the oxidation state of chromium must be +3. 2. [Pt(NH3 )5 Cl]Br3 IUPAC name: pentaamminechloridoplatinum(IV) bromide The charge of the complex ion must be +3 since it bonds with 3 bromides. The five NH3 are neutral molecules while the chloride carries -1 charge. Therefore, the oxidation state of platinum must be +4.

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9.3 Nomenclature of Coordination Compounds 9

3. [Pt(H2NCH2CH2NH2 )2 Cl2 ]Cl2 IUPAC name: dichlorobis(ethylenediamine)platinum(IV) chloride Ethylenediamine is a didentate ligand, the bis- prefix is used instead of di-. 4. K 4 [Fe(CN)6 ] IUPAC name: potassium hexacyanidoferrate(II) Potassium is the cation and the complex ion is the anion. Since there are 4 K+ ions binding with a complex ion, the charge on the complex ion must be -4. Since each ligand carries -1 charge, the oxidation state of Fe must be +2. 5. Pt(NH3 )2 Cl4 IUPAC name: diamminetetrachloridoplatinum(IV) This is a neutral molecule because the charge on Pt (+4) equals the negative charges on the 4 Clligands. If the compound is [Pt(NH3)2Cl2]Cl2, even though the number of ions and atoms in the molecule are identical to the above example, it should be named as diamminedichloroplatinum(II) chloride. 6. Fe(CO )5 IUPAC name: pentacarbonyliron(0) Since it is a neutral complex, it is named in the same way as a complex cation. The common name of this compound, iron carbonyl, is used more often. 7. (NH4 )2 [Ni( C2O 4 )2 (H2O )2 ] IUPAC name: ammonium diaquabis(oxalato)nickelate(II) The oxalate ion is a didentate ligand so -bis is used. 8. [Ag(NH3 )2 ][Ag(CN)2 ] IUPAC name: diamminesilver(I) dicyanidoargentate(I) The first complex is a cationic complex, so Ag is named as silver, whereas the second complex is an anionic complex, so Ag is named as argentate. Worked Problem 9.3

Write the formula for each of the following complexes: (a) hexamminecobalt(III) chloride (b) potassium iron(III) hexacyanidoferrate (c) diamminedichloridoplatinum(II) (d) tetracarbonylnickel(0) (e) triamminechloroidocyanonitrocobalt(III) (f) sodium bis(thiosulphato)argentate(I) (g) nickel hexachloroidoplatinate(IV) (h) tetraammineplatinum(II) amminetrichloridoplatinate(II) solution

(a) [Co(NH)6]3+ (b) K3[Fe(CN)6] (c) [Pt(NH3)2Cl2]

(d) [Ni(CO)4] (e) [Co(NH3)3ClCN(NO2)] (f) Na3[Ag(S2O3)2] (g) [NiCl4][PtCl2] (h) [Pt(NH3)4][Pt(NH3)Cl3]

Practice Problems 5. Give the IUPAC names of the following compounds: (a) [Co(NH3)5ONO]Cl3 (b) Ca2[Fe(CN)6] (e) Mg[Cr(NH3)2(NO2)5] (f) [Pt(NH3)Br(NO2)Cl]Cl 6. Write the formula for each of the following complexes: (a) tetraamminecopper(II) sulphate (c) bis(cyclopentadienyl)iron(II) (e) diamminebis(ethylenediamine)cobalt(III) chloride (g) tetrapyridineplatinum(II) tetrachloroplatinate(II)

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(d) Na[Au(CN)2]

(g) [Cr(SCN)2(NH3)4]

(h) [Fe(C2O4)3]3-

(b) potassium tetracyanidonickelate(0) (d) tetrathiocyanato-N-zinc(II) (f) potassium trioxalatoaluminate(III)

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Chapter 9 Coordination Compounds

9.4 Isomerism in Coordination Compounds Isomers are compounds with same chemical formula but different structural arrangement of atoms. Some of their physical and chemical properties are different owing to this difference in spatial arrangement. The isomerism in coordination compounds can be broadly classified into two categories: 1. Stereoisomerism: These are further of two types, geometrical isomerism and optical isomerism. 2. Structural isomerism: These are further of several types, ionization isomerism, coordination isomerism, hydrate isomerism and linkage isomerism. Stereoisomerism arises in coordination compounds due to difference in arrangement of ligands in the space around the central metal atom. Structural isomerism arises due to the difference in arrangement of atoms/ions surrounding the central metal ion. These isomers are discussed in detail in the following subsections.

Geometric Isomerism This type of isomerism occurs in di-substituted (heteroleptic) complexes with coordination numbers 4 and 6 having square planar and octahedral geometries, respectively. Geometrical isomerism is not possible for coordination numbers 2 and 3 and in tetrahedral complexes with coordination number 4. This isomerism can be of two types: 1. When the two identical ligands are adjacent to each other, the isomer is known as cis. 2. When the two identical ligands are on the opposite side, the isomer is known as trans. Some examples of geometrical isomerism are discussed as follows: 1. Geometrical isomerism in complexes with coordination number 4: Square planar complexes of the type MA2X2, MA2X4, MABX2, MABX4 can exist as geometrical isomers. Here A and B are neutral ligands and X and Y are anionic ligands. Consider the complex [Pt(NH3)2Cl2], it may exist as two geometrical isomers: NH3

NH3

Pt NH3

NH3

Cis

Trans

Similarly, a complex with an unsymmetrical didentate ligand of the type [M(AB)2] may show geometrical isomerism. For example, [Pt(gly)2] where gly is H2NCH2COO-; the two geometrical isomers are

NH2

CH2

NH2

C NH2

CH2

Cis

Trans

2. Geometrical isomerism in complexes with coordination number 6: Consider an octahedral complex: 1 5

2 M

3 6

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9.4 Isomerism in Coordination Compounds 11

Here positions 1−6, 2−4 and 3−5 are trans, while positions 1−2, 2−3, 2−5, 6−3, 6−4, 1−3, 1−4 and 1−5 are cis to each other. The complexes of the type MA4X2, MA2X4, MA4X4 may exhibit geometrical isomerism. For example, consider the complex [Co(NH3)4Cl2]-. CI

CI CI

H3N

Co H3N

NH3

NH3 CI

NH3

Cis (Violet)

Trans (Green)

Similarly, for the complex [Fe(CN)4(NH3)2]-, the two geometrical isomers are NH3

NH3 NC

NH3 Fe

CN Fe

NH3

Cis

Trans

(a) C omplexes of type M(ABCDEF) can also exhibit geometrical isomerism. These isomers may be written by fixing a ligand at one position and then placing the other ligands trans to it. An example of this type is [Pt Br Cl I NO2 py NH3]. (b) Complexes of the type M(AA)2X2 and M(AA)2XY, where (AA) is a symmetrical didentate ligand, also exhibit geometric isomerism. Examples of this type include [Co(en)2Cl2]+. CI

CI CI

en en

Cis

Trans

(c) C omplexes having unsymmetrical didentate ligands also exhibit geometrical isomerism. Example of this type includes [Cr(gly)3] where gly is H2NCH2COO- (Fig. 9.4). O

C CH2

C O

CH2

H2N

NH2

CH2

C O

O C

CH2 NH2

CH2

Cr O

C H2N

O CH2 C O O

O Cis

Trans

Figure 9.4 Geometrical isomers of [Cr(gly)3].

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Chapter 9 Coordination Compounds X X MX6

Only one form

X Y X

MX5Y

MX4Y2

Y X

MX3Y3

X Fac-

M Y

M Y Trans Y

Y X

Only one form as all six corners are equivalent

M X Cis

M X

Y X

M Y Mer-

Two isomers cis and trans

Two isomers and facial (fac) and meridional (mar)

Figure 9.5 Isomers in octahedral complexes. In octahedral entities such as [MA3B3] for example [Co(NH3)3(NO2)3] another type of isomerism exists. When there are three donor atoms of the same ligand occupying adjacent positions at the corners, it is called a facial (fac) isomer. When the positions are around the meridian of the octahedron, it is called a meridional (mer) isomer (Fig. 9.5).

Optical Isomerism There are certain compounds which have the ability to rotate the plane of polarized light. The compounds are said to be optically active and the phenomenon is called optical isomerism. The isomers that rotate the plane of polarized light equally, but in opposite direction are called enantiomers or enantiomorphs. The isomers that rotate the plane to the right are called dextrorotatory (d) complexes, and the isomers that rotate the plane of light to the left are called laevorotatory (l) complexes. The 1:1 mixture of both the isomers will yield a racemic mixture. These isomers are non-superimposable mirror images of each other and are called chiral. These have identical chemical as well as physical properties. Some examples of optical isomerism in coordination complexes are discussed as follows: 1. Octahedral complexes of the type M(AA)3 where (AA) is a symmetrical didentate ligand like ethylenediamine show optical isomerism. For example, the complexes [Co(en)3]3+ and [Cr(ox)3]3- can exist as optical isomers. 3+

en en

en Co

Mirror images of each other 3−

ox ox

3−

ox Co

Co ox

Mirror images of each other

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9.4 Isomerism in Coordination Compounds 13

2. Also complexes of the type M(AA)2X2 and M(AA)2XY, where AA is again a symmetrical didentate ligand show optical isomerism. Examples include [Co(en)2Cl2]+, [RhCl2(en)2]+. Consider the complex [Co(en)2Cl2], we have already seen that it exists as geometrical isomers. The trans isomer has superimposable mirror image as represented below and is thus optically inactive. CI

Superimposable mirror images, optically inactive isomers

The cis form can exist as non-superimposable images and is thus optically active. en

en Co CI

Co CI

Non-superimposable mirror images, optically active isomers

Other examples of this type include [PtCl2(en)2]2+ where again trans form is optically inactive and cis form is optically active. 3. Complexes of the type [M(AA)X2Y2] having only one symmetrical didentate ligand also show optical isomerism. Examples include [CoCl2(en)(NH3)2]. NH3

NH3

Non-superimposable mirror images and optically active isomers

Linkage Isomerism This type of isomerism arises due to presence of ligands with two different donor atoms, which may thus attach to the central metal atom through either of the two atoms. This is exhibited by the complexes having ambidentate ligands. For example, NO2-: nitro-N, where bonding occurs through N and (ONO-)/NO2-: nitro-O, where bonding occurs through O and CN-: cyano, where bonding occurs through C and NC-: isocyano where bonding occurs through N. An example of linkage isomerism is [Co(ONO)(NH3)5]Cl2 (red form) and [Co(NO2)(NH3)5]Cl2 (yellow form) which JØrgensen discovered.

Coordination Isomerism This type of isomerism occurs in those complexes which have both anionic as well as cationic entities and there is a difference in the distribution of ligands within these entities. For example, [Cu(NH3)4]2+[PtCl4]2- and [Pt(NH3)4]2+[CuCl4]2-. Both these complexes share the same molecular formula, but they differ in the cationic and anionic entities. Other examples of this type include [Pt(NH3)4]2+[PtCl4]2-; [PtCl(NH3)3]1+[PtCl3(NH3)]1- and [Cr(NH3)6]3+[Cr(CN)6]3-.

Ionization Isomerism This type of isomerism occurs in complexes which have the same molecular formula but give different ions in solution on ionization. It arises due to difference in the position of the atoms/ions inside

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Chapter 9 Coordination Compounds or outside the coordination sphere. It occurs when counter ions and ligands inside the sphere exchange positions with each other. Consider the example of a coordination complex having the formula Co(NH3)5BrSO4. There are two possible structures for this complex: [Co(NH3 )5 Br ]SO 4

Pentamminebromocobalt(III) sulphate (Violet −rred)

[Co(NH3 )5 SO 4 ]Br

Pentamminesulphatecobalt(III) bromide (Red)

Both these complexes share the same molecular formula, but the difference arises in the mode of their ionization. This can be explained as follows: 2 [Co(Br)(NH3 )5 ]SO 4 BaCl  → [CoBr(NH3 )5 ]Cl2 + BaSO 4 3 [CoSO 4 (NH3 )5 ]Br AgNO  → [CoSO 4 (NH3 )5 ]NO3 + AgBr

Some other complexes that show this type of isomerism are [CoCl(NO2)(NH3)4]Cl and [CoCl2(NH3)4], [Co(NO3)(NH3)5]SO4 and [Co(NH3)5SO4]NO3 and [PtCl2(NH3)]Br2 and [PtBr2(NH3)]Cl2.

Hydration Isomerism This isomerism is very similar to ionization isomerism, but the difference is that in this case the solvent molecule can act as the ligand. The solvate isomers are different from each other on the basis of whether or not a solvent molecule is directly bonded to the metal ion or is present as free solvent molecule in the crystal lattice. This isomerism is known as hydrate isomerism if we are considering is water molecules as solvent. These complexes have the same molecular formula, but differ in the number of water molecules of hydration, that is, water molecules outside the coordination sphere. Consider the example of CrCl3  6H2O. This can have any of the following structures: [Cr(H2O)6 ]Cl3 , [CrCl(H2O)5 ]Cl2 ⋅ H2O, [CrCl2 (H2O)4 ]Cl ⋅ 2H2O, [CrCl3 (H2O)3 ] ⋅ 3H2O All these are totally different complexes, but are hydrate isomers of each other. Other complexes of this type include [CoCl(en)2(H2O)]Cl2 and [CoCl2(en)2]Cl  H2O, [CoCl(H2O)(NH3)4]Cl2 and [CoCl2(NH3)4]Cl  H2O and [CrCl2(py)2(H2O)2]Cl and [CrCl3(Py)2H2O]H2O.

Worked Problem 9.4

Name the type of isomerism depicted by the following: (a) [Pt(NH3)2Cl2]; (b) CrCl3  6H2O; (c) [Co(NH3)5NO2](NO3)2; (d) [Co(NH3)6][Cr(CN)6]. solution

(a) Ionization isomerism (b) hydration isomerism (c) linkage isomerism (d) coordination isomerism

Practice Problems 7. Give formula for the following: (a) Linkage isomer of [CoCl3(SCN)Cl2 (b) Ionization isomer of [Co(NH3)5Br]SO4 (c) Coordination isomer of [Pt(NH3)[PtCl6] 8. Which of the following complexes can exhibit geometrical isomerism: (a) [CoCl2(en)2] (b) [Co(NH3)Br]SO4 (c) [CoCl2(NH3)2] (tetrahedral) (d) [CrCl3(NH3)3] (e) [Pt(NH2OH)(NO2)(NH3)py]+ (square planar)

9. Octahedral complexes of unidentate ligands exhibit optical isomerism only when they have at least three different ligands. Explain.

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9.5 Bonding in Coordination Compounds 15

9.5 Bonding in Coordination Compounds Although Werner described bonding in coordination compounds, it could not explain some observations related to them, such as: 1. Ability of only certain elements to form coordination compounds. 2. Directional nature of bonds in coordination compounds. 3. Magnetic and optical properties of coordination compounds. Various theories have been put forward to explain the bonding of metal ion with ligands in complexes. These include Valence Bond Theory (VBT), Crystal Field Theory (CFT), Ligand Field Theory and Molecular Orbital Theory. The application of the first two theories to explain bonding in coordination compounds is discussed in the following subsections.

Valence Bond Theory This theory was developed by Pauling and gave valence bond treatment of bonding in complexes. It could satisfactorily explain the structure and magnetic properties of a large number of coordination compounds. The main postulates of the theory are as follows: 1. For the coordination bond formation between central metal ion and the ligand, the central metal ion makes available a number of empty orbitals; and the number of these empty orbitals is equal to the coordination number of the metal ion. For example, if the coordination number is 6, the central metal ion will make available six empty orbitals and if it is 4, the metal ion will make available four empty orbitals. 2. The atomic orbitals made available by the central metal ion will be a mixture of s, p, d orbitals, all of which are of different energies and orientation. All these orbitals hybridize to give orbitals that are equivalent in energy and symmetry, which then form bonds with the ligands. These equivalent orbitals result in some particular geometry of the complex such as tetrahedral, square planar, octahedral, etc. (Table 9.3).

Table 9.3 Hybridization involved in complexes of different geometries

Coordination Number 4

Hybridization Involved

Geometry Tetrahedral

Square planar

dsp

sp d

sp d , d sp

Square bipyramidal

3 2

Octahedral

3. The d orbitals that are involved in the process of hybridization can be either from the inner d orbital (n − 1) or the outer d orbital (n). The complexes thus formed are referred to as inner orbital and outer orbital complexes, respectively. 4. Each ligand will have atleast one lone pair of electrons. It can have more than one pair of electrons to donate, but should have atleast one. 5. The coordinate bond is formed when there is an overlap between hybridized empty orbitals of the central metal ion and filled orbitals of the ligand. The greater the extent of overlap, the stronger will be the bond and more stable will be the complex. 6. The number of unpaired electrons in the complex indicates the geometry of the complex and vice versa. Under the influence of a strong ligand, the electrons can be forced to pair up against the Hund’s rule of maximum multiplicity. If the complex contains unpaired electrons, it is paramagnetic in nature; whereas, if it does not contain unpaired electrons, it is diamagnetic in nature. With the help of these postulates, we can predict the geometry of various coordination complexes. The two examples in the following subsections illustrate the application of valence bond theory in the formation of octahedral and tetrahedral complexes.

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Chapter 9 Coordination Compounds Coordination Number 6 1. [Co(NH3)6]3+: Cobalt atom (Z = 27) has electronic configuration of 3d74s2. To achieve the oxidation state of +3 to form the complex, it loses two electrons from 4s and one electron from 3d, thereby generating Co3+ ion. This leads to four vacant orbitals but there are six electrons to be accommodated. In the presence of strong field ligand NH3, rearrangement or pairing of electrons in 3d orbitals takes place in Co3+. This rearrangement results in two 3d orbitals becoming vacant. These vacant orbitals combine with one vacant 4s and three vacant 4p orbitals, and hybridize to give six equivalent d2sp3 hybridized orbitals directed towards the corner of an octahedral. Six pairs of electrons from six NH3 molecules occupy these vacant orbitals and give rise to coordination number 6, and thus account for octahedral geometry for this complex. 3d Co atom

↑↓ ↑↓ ↑

↑

3d Co (lll)

↑↓ ↑

↑↓ 4s

↑

↑↓ ↑↓ ↑↓

Co3

d2sp3 hybrid ↑↓ ↑↓ ↑↓

[Co(NH3)6]3

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ Six pairs of elelectrons from six NH3 molecules

In absence of any unpaired electron, the complex is diamagnetic. Also since inner d (n – 1) orbitals are used for bond formation, the complex is known as inner orbital or low spin or spinpaired complex 2. [CoF6]3+: In this complex cobalt is in oxidation state +3. On formation of Co(III), the electronic configuration changes from 3d74s2 (in ground state) to 3d64s0 as shown

3d Co atom

↑↓ ↑↓ ↑

↑

↑↓ 4s

Co (lll)

↑↓ ↑

↑

4p 4p

↑

In the presence of F-, which is a weak-field ligand, the 3d electrons do not show any tendency to pair up. Therefore 4s, three 4p and two 4d orbitals hybridize to form six equivalent sp3d2 orbitals. The lone pair of electrons from six fluoride ligands occupy these orbitals to form an octahedral, paramagnetic complex. 3d Co3

↑↓ ↑

↑

4s ↑

↑ sp3d2 hybridization 4s

3d [CoF6]3

↑↓ ↑

↑

4p xx xx xx

4d xx xx

The complex is paramagnetic due to presence of unpaired electrons. Since outer nd orbitals are used in hybridization, the complex is called outer orbital or high spin or spin free complex.

Coordination Number 4 For coordination number 4 we can have two types of hybridization: sp3 (tetrahedral geometry) and dsp2 (square planar geometry). 1. [NiCl4]2-: In this complex ion Ni (Z = 28) is in +2 oxidation state. The configuration of Ni atom is 3d8 4s2 which changes to 3d 84s0 on the formation of Ni(II).

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9.5 Bonding in Coordination Compounds 17 3d Ni atom

↑↓ ↑

↑

↑↓

↑↓ ↑

↑

3d Ni(II)

↑

Since Cl is a weak-field ligand, it does not have the tendency to pair up the electrons in the 3d level. The empty 4s and 4p orbitals hybridize to form four equivalent sp3 orbitals. The lone pair of electrons from four chloride ions occupy these hybridized orbitals, giving rise to a tetrahedral and paramagnetic complex. 3d Ni2

↑↓ ↑↓ ↑↓

4s ↑

↑ sp3 hybridization

3d [NiCl4]2 ↑ ↓ ↑ ↓ ↑ ↓

4s ↑

↑

4p xx

2. [Ni(CN)4]2-: In this case, again Ni is in (+2) oxidation state. The configuration of Ni atom is 3d84s2 which changes to 3d84s0 on the formation of Ni(II) with the loss of two electrons. 3d Ni atom ↑ ↓ ↑ ↓ ↑ ↓

4s ↑

↑

↑↓

3d Ni(II) ↑ ↓ ↑ ↓ ↑ ↓

4s ↑

4p 4p

↑

Since CN- is a very strong-field ligand, it has the tendency to pair up electrons in the 3d level so as to set up space for its own electrons to be accommodated. Now, the configuration changes as shown 3d

NI2 ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓

↑ dsp2 hybridization

 Key Point

Thus empty 3d orbitals, 4s and two 4p orbitals hybridize to form four equivalent dsp orbitals. The four lone pairs from cyanide ligands occupy these orbitals to form square planar and diamagnetic complex. 2

3d [Ni(CN)4]2

↑↓ ↑↓ ↑↓ ↑↓

xx xx

The hybrid orbitals actually do not exist. They are theoretical (mathematical) manipulations of the wave function of the participating atoms.

Magnetic Properties of Coordination Compounds The magnetic properties of coordination compounds are important because when properly interpreted, these are useful for identifying and characterizing the compounds. The magnetic moment of the coordination compounds can be obtained by measuring the magnetic susceptibility. The magnetic nature of coordination compounds is determined by the number of unpaired electrons present. If we consider the first three metal ions of the first transition series(Ti3+, V3+ and Cr3+), each d electron singly occupies a d orbital and the remaining d orbitals are available for forming d2sp3 hybrid orbitals with 4s and 4p orbitals. Hence, these compounds show similar magnetic behavior. When more than three d electrons are present, that is for d4 (Cr2+, Mn3+), d5 (Mn2+, Fe3+), d6 (Fe2+, Co3+), these d orbitals are not available for octahedral hybridization. These can be made vacant only if the electrons pair up in the 3d orbitals and this leaves lesser number of unpaired electrons, thus affecting the magnetic behavior. The experimental magnetic moment data is in agreement with spin pairing in some cases, particularly in compounds with electronic configuration d6. However, for variation in magnetic properties is observed due to different number of unpaired electrons. For example: 1. [Mn(CN)6]3- has magnetic moment corresponding to two unpaired electrons while [MnCl6]3shows paramagnetic character corresponding to four unpaired electrons. 2. [Fe(CN)6]3- has magnetic moment corresponding to single unpaired electron while [FeF6]3- shows paramagnetic character corresponding to five unpaired electrons. 3. [CoF6]3- is paramagnetic with four unpaired electrons while [Co(C2O4)3]3- is diamagnetic.

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Chapter 9 Coordination Compounds This variation in magnetic properties can be explained in terms of use of inner (n − 1) d orbitals and outer (n) d orbitals for hybridization. [Mn(CN)6]3-, [Fe(CN)6]3- and [Co(C2O4)3]3- are inner orbital complexes with d2sp3 hybridization. On the other hand, [MnCl6]3-, [FeF6]3- and [CoF6]3- are outer orbital complexes with sp3d2 hybridization. The inner d orbitals remain unpaired and so these complexes are paramagnetic.

More to Explore

The magnetic moment (µ) can be measured from spin-only formula, which assumes that the magnetic moment arises entirely from the unpaired electron spin.

µs = [n(n + 2)]1/2 where n is the number of unpaired electrons. This formula works reasonably well with the metal ions of the first transition series. However, in case of metals of second and third transition series, this formula is not applicable since there we also have large contribution from the angular orbital moment along with the spin motion. For the complex [CoF6]3-, spin-only magnetic moment (µs) can be calculated as follows:

µs = [n(n + 2)]1/2 = [ 4 × ( 4 + 2)]1/2

(as n = 4 in [CoF6 ]3− )

= (24 )1/2 = 4.90 BM

Limitations of Valence Bond Theory There are certain limitations of the valence bond theory: 1. It is only a qualitative explanation for bonding in coordination compounds. 2. It does not explain why some complexes are colored, whereas others are not. Complexes of the same metal in the same oxidation state may have different colors with different ligands. 3. The magnetic and spectral properties of some of the complexes are not explained. 4. It does not explain the kinetic and thermodynamic stabilities of the coordination complexes. 5. The predictions given regarding the complexes having square planar and tetrahedral geometries are not exact. 6. The demarcation between strong-field and weak-field ligands is not proper.

Crystal Field Theory In this theory, the interaction between the ligand and the central metal ion is treated as a purely electrostatic interaction, as opposed to the valence bond theory where the interaction is treated as covalent in nature. According to this theory proposed by H. Bethe and Van Vleck in 1930s, the transition metal which forms the central atom in the complex is regarded as a positive ion of charge equal to the oxidation state. This is surrounded by negative ligands or neutral molecules which have a lone pair of electrons. If the ligand is a neutral molecule such as NH3, the negative end of the dipole in the molecule is directed towards the metal ion. The electrons on the central metal are under repulsive forces from those on the ligands. Thus, the electrons occupy the d orbitals furthest away from the direction of approach of ligands. In the crystal field theory, the following assumptions are made: 1. Ligands are treated as point charges. 2. There is no interaction between metal orbitals and ligand orbitals. 3. The d orbitals on the metal all have the same energy (i.e., degenerate) in the free atom. However, when a complex is formed the ligands destroy the degeneracy of these orbitals, that is,

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9.5 Bonding in Coordination Compounds 19 the orbitals now have different energies. In an isolated gaseous metal ion, all the five d orbitals have the same energy, and are termed degenerate. If a spherically symmetrical field of negative charges surrounds the metal ion, the d orbitals remain degenerate. However, the energy of the orbitals is raised because of repulsion between the field and the electrons on the metal. In most transition metal complexes, either six or four ligands surround the metal, giving octahedral or tetrahedral structures. In both of these cases, the field produced by the ligands is not spherically symmetrical. Thus the d orbitals are not all affected equally by the ligand field.

Thus, under the influence of the ligands, the five degenerate d orbitals of the metal ion will split into two groups of orbitals of different energies. This effect is known as crystal field splitting or energy level splitting.

Crystal Field Splitting in Octahedral Coordination Entities In an octahedral complex, the coordination number is 6. The central metal ion is at the center and the six ligands occupy the six corners of the octahedron. The phenomenon of splitting of d orbitals in an octahedral complex is illustrated in Fig. 9.6. dx 2 − y 2 dz 2 (eg)

0.6∆ο ∆

Energy

State II 0.4∆ο State I

dxy dyz dzx (t2g ) State III

Figure 9.6 Splitting of d orbitals in an octahedral complex. In the free metal ion, all the five d orbitals are degenerate (State I, Fig. 9.6). When the six ligands approach the central metal cation along the axes, they exert an electrostatic force of repulsion on the outer d electrons; that is, the d electrons are repelled by the lone pairs of the ligands. This repulsion raises the energy of the degenerate d orbitals to give five excited degenerate orbitals (State II, Fig. 9.6). Since the lobes of d 2 and d 2 2 orbitals (eg set of orbitals) lie directly in the path of the z x −y approaching ligands, the electrons in these orbitals experience greater repulsion exerted by the electron clouds of the ligands than that experienced by the electrons in the dxy, dyz and dzx orbitals (t2g set of orbitals), which are directed in space between the x, y and z axes. Hence, under the influence of the approaching ligands, the orbitals d 2 2 and d 2 are raised in energy, whereas the x −y z orbitals dxy, dyz and dzx are lowered in energy relative to the excited d levels (State III, Fig. 9.6). The separation of five d orbitals into t2g and eg sets of different energies is known as crystal field splitting. The eg set, which is of higher energy, is doubly degenerate; whereas the t2g set, which is lower in energy, is triply degenerate. The energy difference between eg and t2g sets is denoted by ∆o (the subscript o stands for octahedral). The magnitude of ∆o depends on the field strength of the ligand as well as on the metal ion. ∆o is called the crystal field stabilization energy (CFSE).The energy of the two eg orbitals increases by 0.6∆o (3/5∆o) and that of t2g orbitals is lowered by 0.4∆o (2/5∆o).The magnitude of crystal field splitting depends on a number of factors, the most important amongst which is the nature of the ligand. If it is easier for a ligand to approach metal atom and interact with it, then the extent of crystal field splitting is high. The ligands that cause only a small degree of crystal field splitting are called weak ligands, whereas the ligands that cause a large degree of crystal field splitting are called strong ligands. The strong ligands give a higher value of ∆o, and weak ones give a lower value. Some common ligands arranged in the increasing order of their splitting power are as follows: I− < Br − < S2− < Cl− < NO3− < F − < OH− < EtOH < (COO − )2 < H2 < EDTA < NH3 < ethylenediamine < diphenyl < o-phenanthroline < NO2− < CN− < CO

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Chapter 9 Coordination Compounds This arrangement of ligands in the order of their abilities to split the d orbitals energies is called the spectrochemical series. When the first, second and third electrons enter the d orbital, they occupy the three t2g (lower energy) orbitals one by one. However, when the fourth and the next electron enter the d orbitals, the following two alternative options are available for the electron: 1. The fourth and the next electron may pair up with the electrons in the t2g orbitals (contrary to the Hund’s rule of maximum multiplicity). 2. The fourth and the next electron may enter the higher eg orbitals in accordance with Hund’s rule of maximum multiplicity. The exact path chosen by the electron depends upon the strength of the ligand: Case I: In the presence of a strong-field ligand, such as CO, CN-, etc., the electrons pair up because ∆o is large enough to overcome the energy required to pair up the electron in the same orbital, that is pairing energy (P). This results in filling of the lower energy t2g d orbitals. Thus, in the presence of a strong-field ligand, low-spin (diamagnetic) complexes are formed, for example, [Fe(CN)6]4-. Case II: In the presence of a weak-field ligand, such as F-, H2O, etc., since ∆o is not large enough to overcome the pairing energy of the electrons, the next (fourth and fifth) electrons will enter the eg orbitals. Thus, in the presence of a weak-field ligand, high-spin complexes (paramagnetic) are formed, for example, [Fe(H2O)6]2+.

Crystal Field Splitting in Tetrahedral Complexes The coordination number in tetrahedral complexes is 4. In a tetrahedral field, the four ligands approach in a different fashion. Figure 9.7 suggests that the ligands interact more with the t2g (dxy, dyz and dzx) orbitals pointing close to the direction of the approaching ligands than the eg ( and ) orbitals lying between the ligands. The three t2g orbitals are of higher energy and two eg orbitals are of lower energy. Here, the d orbitals are split into two groups but in reverse order. This is represented in Fig. 9.8.

y z

L M

L L

Figure 9.7 Approach of the four ligands in tetrahedral field.

dxy dyz dzx (t2g )

0.4 ∆t ∆

Energy

State II 0.6 ∆t State I

dx 2 − y 2 dz 2 (eg ) State III

Figure 9.8 Splitting of d orbitals in tetrahedral complex.

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9.5 Bonding in Coordination Compounds 21

The crystal field splitting in this case is denoted by ∆t (the subscript “t” indicating tetrahedral complex). In the tetrahedral field, since the d orbitals are not interacting directly with the ligand field, the splitting of d orbitals is less than that in the octahedral complexes. t < o  t = 49  o The lower value of crystal field splitting (∆t) in tetrahedral complexes may also be attributed to the lesser number of ligands (4).

Color in Coordination Compounds Most of the transition metal complexes are colored in their solution or solid state. The transition metals have the ability to absorb some of the radiations from the visible spectrum of light and transmit other radiations. When a transition complex is placed in white light, it absorbs certain portion of this white light and the portion which is not absorbed is reflected back from the complex. It is this reflected portion that imparts color to the complex. Thus, the actual color of the complex is depicted by the reflected light and not the light it absorbs. Crystal field theory explains the color in coordination compounds and attributes it to the d−d transition of the electron between the split t2g and eg levels. The energy gap between t2g and eg levels is very low in case of transition metal complexes and when light falls on them, the electrons in the lower energy level jump to the higher energy level. The absorbed light is actually that portion of light which is sufficient to excite the electrons from lower energy level to the higher energy level. The portion of light reflected back is responsible for the color of the complex. Table 9.4 shows the color of the light absorbed by a coordination entity and the color of the coordination entity. Consider the example of the complex [Ti(H2O)6]3+, in which Ti (III) has a single d electron and hence d1 configuration. The d electron will occupy t2g orbital. On irradiation with visible light, it should be possible for the ion to capture a quantum of radiation with frequency ∆o/h, (where h is Planck’s constant) and get excited from t2g orbital to eg orbital. Solutions containing the hydrated Ti3+ ion are reddish−violet in color because yellow and green light are absorbed to excite the electron and the transmitted light is the complimentary color.

 Key Point

In absence of a ligand, the crystal field splitting does not occur, so d−d transitions are not possible and the substance is colorless. For example, hydrated copper sulphate (CuSO4∆  5H2O) is blue in color whereas anhydrous CuSO4 is white.

Limitations of Crystal Field theory Crystal field theory successfully explains the formation, structures, color and magnetic properties of coordination compounds. However, the theory has certain limitations: 1. According to crystal field theory, anionic ligands are considered as point charges and hence should exert the greatest splitting effect. However, they occur at the low end of spectrochemical series.

Table 9.4 Relationship between the colors of the coordination entities and the colors absorbed by them.

Chap_09.indd 21

Coordination Entities

Color of Light Absorbed

Color of Coordination Entity

[CoCl(NH3 )5 ]2+

Yellow

Violet

[Co(NH3 )5 (H2O )]3+

Blue Green

Red

[Co(CN)6 ]3−

Ultraviolet

Pale Yellow

[Cu(H2O )4 ]2+

Red

Blue

[ Ti(H2O )6 ]3+

Blue Green

Violet

[Co(NH3 )6 ]3+

Blue

Yellow Orange

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Chapter 9 Coordination Compounds 2. Covalent bonding between the ligand and the central atom is not considered in this theory. 3. This theory ignores the attractive influence of the nuclear charge in the ligand atom on the metal ion d electrons.

Practice Problems 10. Ni(CO)4 is a diamagnetic complex, whereas NiCl4 is paramagnetic. Explain. 11. Both the complexes Ni(CO)4 and [Ni(CN)4]2- have coordination number 4; however, Ni(CO)4 is tetrahedral in shape whereas [Ni(CN)4]2- is square planar. Explain the difference in geometry in terms of the difference in hybridization.

12. According to crystal field theory, how do five degenerate d orbitals split in the presence of ligands in octahedral complexes?

9.6 Bonding in Metal Carbonyls The first metal carbonyls, Ni(CO)4 and Fe(CO)5 were discovered in 1890 and 1891, respectively, and now carbonyl derivatives of at least one type are known for all the transition metals. The simplest carbonyls are of the type M(CO)x which have simple well-defined structures. For example, V(CO)6 and Cr(CO)6 are octahedral, Fe(CO)5 is trigonal bipyramidal and Ni(CO)4 is tetrahedral. Of the d-block elements, the ones that form stable mononuclear carbonyls are those that require an integral number of carbonyl ligands to attain the number of electrons of the succeeding noble gas atom. OC Ni OC

Fe CO

OC OC

There are numerous polynuclear carbonyls that are homonuclear, such as Fe3(CO)12 or heteronuclear, such as MnRe(CO)10. In these compounds, there are not only linear M−C−M groups but also either M−M bonds or M−M bonds with bridging carbonyls. For example, Mn2(CO)10 is made up of two square pyramidal Mn(CO)5 units joined by an Mn−Mn bond and Co2(CO)8 has a Co−Co bond bridged by two carbonyl groups. OC

OC OC OC

OC CO Fe

CO CO

Fe2(CO)8

OC CO OC OC Mn OC CO

CO Mn CO CO

O C

Mn2(CO)10

C O

Co C O

CO CO

Co2 (CO)8

The bonding in linear M−C−O groups is depicted in Fig. 9.9

C O σ bond

C O π back-bond

Figure 9.9 Bonding in linear M−C−O groups. First, there is a dative overlap of the filled carbonyl carbon s-orbital with the vacant metal orbital. The second dative overlap is of the filled dp metal orbital with an empty antibonding pp orbital of the carbonyl group. This bonding mechanism is synergic, that is, the effects of s-bond formation strengthen the p bonding and vice versa.

Chap_09.indd 22

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9.6 Bonding in Metal Carbonyls 23

9.7 Stability of Coordination Compounds In general, the coordination complexes are highly stable complexes, where the stability arises from the interaction between the metal ion and the ligands. The stronger the interaction, higher will be the stability of the complex. Now considering the general reaction for the formation of the coordination complex, Ma+ + nLb− → [ML n ]c + where a+, b- and c+ are the respective charges on the metal, ligand and complex. The stability constant is represented as K =

[(ML n )c + ] [Ma+ ][Lb− ]n

The reaction may proceed through various steps: M + L → ML ML + L → ML 2

[ML] [M][L] [ML 2 ] K2 = [ML][L] K1 =

  ML n−1 + L → ML n K n =

[ML n ] [ML n − 1][L]

In these steps of the reaction, K1, K2, … are referred to as “stepwise stability constants” and overall stability constant is given by K = K1 × K 2 × K 3  K n Now, let us again consider the formation of the complex MLn by the following pathway [M(H2O)n ] + nL → [ML n ] + H2O Here, the overall stability constant is written as ML  H O  β n =  n   2 n M(H2O )n ][L  Since concentration of water is high, it can be considered constant and the stability constant can be written as

βn =

ML n  n M(H2O )n  L 

This reaction actually proceeds in steps and can be written as [M(H2O)n ] + L → [M(H2O)n−1L] + H2O

K1 =

[M(H2O)n−1L] [M(H2O)n ][L]

[M(H2O)n−1L] + L → [M(H2O)n−2L2 ] + H2O

K2 =

[M(H2O)n−2L 2 ]2 [M(H2O)n−1L][L]

Kn =

[ML n ] [M(H2O)L n−1][L]

  [M(H2O)L n−1] + L → [ML n ] + H2O

The overall stability constant can be related to the stepwise stability constants by the relation:

βn = K1 × K 2 × K 3  K n

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Chapter 9 Coordination Compounds Most of the practical calculations and measurements are actually made from aqueous mediums and solutions; and thus, the complex is formed by the displacement of H2O molecules by the ligand in consideration. The value of K for some systems is listed as follows: Ag+ + 2NH3 → [Ag(NH3 )2 ]+

K = 1.65 × 107

Ag+ + 2CN− → [Ag(CN)2 ]−

K = 5.48 × 1018

Cu2+ + 4NH3 → [Cu(NH3 )4 ]2+

K = 4.65 × 1023

Cu2+ + 4CN− → [Cu(CN)4 ]2−

K = 2.48 × 1027

From the above data, it can be seen that CN- behaves as a better ligand in complex formation, since the stability constant values for CN- are much higher as compared to that of NH3 as ligand. The reciprocal of the formation constant is called the stability or dissociation constant of coordination compounds.

9.8 Importance and Applications of Coordination Compounds The applications of coordination compounds in a number of general fields are summarized as follows: 1. Analytical chemistry: The analytical applications of coordination chemistry are in (a) Qualitative and Quantitative analysis: Metals ions form colored coordination compounds on reaction with a number of ligands. These reactions are used for detection of these metal ions. The colored complexes formed can be used for the estimation of metals by classical or instrumental methods such as gravimetry or colorimetry. Some examples are given as follows:

•

Nickel ions (Ni2+) can be detected by the addition of dimethylglyoxime (DMG) which results in the formation of a brilliant coordination complex. The complex formed can be used for gravimetric estimation of nickel ions present. −

N Ni2+ + 2 N

OH OH

− +2OH −2H2O

N Ni

N O

•

N O−

The presence of iron ions (Fe3+) can be detected by the addition potassium ferrocyanide solution, which results in formation of Prussian blue complex. Fe2+ + K 3Fe(CN)6 → KFe[Fe(CN)6 ] + 2K +

•

Cobalt ions (Co2+) can be detected by the addition of amonium thiocyante solution, which results in formation of blue color solution due to the formation of following complex. N

NH4+ S S

Co−2 N

S S NH4+

Ammonium tetrathiocyanatocobaltate (II) (blue)

(b) Volumetric analysis: Hardness of water can be estimated by titration with EDTA. The metal ions causing hardness, that is, Ca2+ and Mg2+ form stable complexes with EDTA.

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Concept Review 25

2. Metal extraction and purification: Extraction of metals, such as silver and gold is carried out by forming their water soluble cyanide complexes with the ore. Pure gold can then be obtained from the solution by addition of zinc. Similarly, metals can be purified by formation and then decomposition of their coordination compounds. For example, impure nickel obtained after extraction may be converted into pure nickel by first converting it to nickel carbonyl and then decomposing. 3. Catalysis: Coordination compounds are used as catalysts in important commercial processes. For example, (a) The Zeigler−Natta catalyst (TiCl4 and trialkyl aluminium) is used as a catalyst in the formation of polyethylene. (b) The Wilkinson catalyst RhCl(PPh3)3 is used in the hydrogenation of alkenes. (c) In the Monsanto acetic acid process, various rhodium complexes, such as [Rh(CO)2I2], [Rh (Cl)(CO)(PPh3)2] or [Rh(Cl)(CO)2]2 are used as a catalyst in the presence of CH3I, I2 or HI as activator. 4. Electroplating: Coordination compounds of gold, silver and copper are used as components in the baths used for electroplating articles of other metals with these metals. For example, in silver plating, K[Ag(CN)2] is used as an electrolyte, in gold plating K[Au(CN)2] is used as an electrolyte and in copper plating K3[Cu(CN)4] is used as an electrolyte. 5. Biological importance: Some important biological compounds are coordination complexes. For example, chlorophyll is a complex of Mg2+. This green pigment plays a vital role in photosynthesis in plants. Similarly, haemoglobin, the red pigment present in blood, is a coordination complex of Fe2+ and vitamin B12, an essential nutrient, is complex compound of Co3+. 6. Medicinal uses: Complexing or chelating agents are used in treating metal poisoning, wherein, the coordination complex is formed between toxic metal in excess metal and the complexing agent. For example, EDTA is used in lead poisoning. EDTA, when injected intravenously into the bloodstream, traps lead forming a compound that is flushed out of the body with the urine. Other heavy metal poisonings that can be treated similarly with chelation therapy are mercury, arsenic, aluminum, chromium, cobalt, manganese, nickel, selenium, zinc, tin and thallium. Similarly chelating ligands D-penicillamine and desferrioxime B are used for removal of excess copper and iron, respectively. New potent drugs are being created using various derivatives of metallocene. A platinum complex [PtCl2(NH)32] called cis-platin is used in treatment of cancer.

Concept Review 1. Coordination compounds are not only important in 2. 3.

4. 5.

Chap_09.indd 25

biological systems but these are also useful in chemical industries. Werner’s theory was the first attempt to explain bonding in coordination complexes. He concluded that in entities, the metal shows two types of valences: primary valences (non-directional) and secondary valences (directional). Because of secondary valences, a coordination entity has a particular shape. The Valence Bond Theory (VBT) developed by Pauling explains which atomic orbitals on the metal were used for bonding. The main limitations of VBT are (a) it does not provide explanation for the electronic spectra of the colored compounds and (b) it does not explain why magnetic properties vary with temperature. The Crystal Field Theory (CFT) proposed by Bethe and van Vleck explained that the attraction between the metal

and the ligands in the complex is considered to be purely electrostatic, and the bonding could be due to ion−ion interaction. 7. CFT is simple and explains the electronic spectra and magnetism of transition metal compounds. 8. Most of the transition compounds are colored in their solution or solid state. 9. In metal carbonyls, the metal−carbon bond has both s and p character. The ligand to metal is s bond and the metal to ligand is p bond and this bonding stabilizes the metal carbonyls. 10. The stability of coordination compounds is given in terms of stepwise stability constant. 11. Coordination compounds find applications in metallurgy, analytical chemistry and as catalysts.

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Chapter 9 Coordination Compounds

Exercises 1. Discuss the concept of primary and secondary valences

in coordination complexes as postulated by Werner’s coordination theory. 2. How does the formation of a coordination compound take place according to valence bond theory? What are the limitations of the theory? 3. Using valence bond theory explain why [Ni(H2O)6]2+ is outer orbital complex and [Cr(H2O)6]3+ is inner orbital complex. 4. Explain why [Ni(CO)4] is tetrahedral, [Ni(CN)6]2- is square planar and [Ni(NH3)6]2+ is octahedral using valence bond theory. 5. What are the salient features of crystal field theory? 6. Explain the crystal field splitting of d orbitals by the strong ligands in octahedral complexes with coordination number six. 7. Both [Co(NH3)6]3+ and [CoF6]3- are octahedral complexes. Use crystal field theory to explain the magnetic nature of the two complexes. 8. Apply crystal field theory to explain why: (a) [Ti(H2O)6]3+ is colored and [Sc (H2O)6]3+is colorless. (b) [Co(NH3)6]3+ is orange yellow in color and [CoF6]3- is blue in color. 9. Discuss the factors that affect the stability of coordination compounds. 10. Write the IUPAC names of the following: K4[Fe(CN)6]; K3[Al(C2O4)3]; [Au(CN)4]-; Li[AlH4]; Ca2[Fe(CN)6]; Na2[ZnCl4]; K2[NiF6]. 11. Define and give one example of each of the following: (a) linkage isomerism; (b) optical isomerism; (c) geometric isomerism.

Chap_09.indd 26

12. What is the type of isomerism shown by the tris(ethylenediamine)cobalt(III) ion? Draw the structures.

13. Why are d−d electronic transitions forbidden? Why are they weakly absorbing and why do they occur at all?

14. Why are compounds of Ti4+ and Zn2+ typically white? 15. 16. 17. 18. 19. 20. 21. 22.

Why are Mn2+ compounds very pale in color? What d−d transitions are spin allowed for a d5 ion? What is the spectrochemical series, and what is its importance? Which of the following complexes will show optical isomerism: MABXY (tetrahedral); MABXY(square planar); MA2B2X2 (octahedral)? trans-Dichlorobis(ethylenediammine)cobalt(III) ion does not exhibit optical isomerism while the cis isomer does. Explain. The compound CrCl3  6H2O exists in three distinct forms which are violet, green and dark green in color. Explain suggesting possible structures for the compounds. Why are most of the transition metal complexes colored in their solution or solid state? Which among CN- and CO will form a more stable complex and why? Give an example each role of coordination compounds in: (a) metallurgical process and (b) catalysis. What is the significance of coordination compounds in analytical chemistry?

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