Coordination Theory

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CO-ORDINATION CHEMISTRY MEMORY MAP

Neutral

Bidentate

Monodentate

Polydentate

Bridging

Negative

Positive

Ambidentate

Ligands

Anionic Complexes

Werner theory VBT

Co-ordination compound

Nomenclature

Bonding

Neutral Complexes

Stability Constant

Isomerism

Stereooisomerism

Structural

Hydrate isomerism

Linkage isomerism

D U

C

A T

I O

N

S

CFT

Metal Carbonyls

Bridiging Complexes

E

Labile

Basic theory

Cationic Complexes

Ionisation isomerism

Chelating

Ligand isomerism

Coordination isomerism

Polymerisation isomerism

Geometrical

Optical

Coordination position isomerism

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DEFINITIONS (1) Simple Salts : When an acid reacts with an alkali, neutralisation takes place and a simple salt is produced, NaOH + HCl → NaCl + H2O When dissolved in water, these salts ionise and produce ions in solution. Mixed salts contain more than one acidic or basic radicals, e.g. NaKSO4. (2) Molecular or Addition Compounds : When solutions containing two or more salts in stoichiometric (i.e., simple molecular) proportions are allowed to evaporate, we get crystals of compounds known as molecular or addition compounds. These are of two types depending on their behaviour in aueous solution. (i) Double salts or Lattice Compounds : The addition compounds having the following characteristics are called double salts or lattice compounds (a)

They exist as such in crystalline state.

(b) When dissolved in water, these dissociate into ions in the same way in which the individual components of the double salts do. FeSO4.(NH4)2SO4.6H2O → Fe2+(aq) + 2NH4+(aq) + 2SO42–(aq) + 6H2O Mohr's salt K2SO4.Al2(SO4)3.24H2O → 2K+(aq) + 2Al3+(aq) + 4SO42–(aq) + 24H2O Potash alum In aqueous solution they give the test of all their constituent ions i.e. the individual components of a double salt do not lose their identity. (ii) Coordination (or complex) compounds : It has been observed that when solutions of Fe(CN)2 and KCN are mixed together and evaporated, potassium ferrocyanide, Fe(CN)24KCN is obtained which in aqueous solution does not give test for the Fe2+ and CN– ions, but gives the test for K+ ion and ferrocyanide ion, Fe(CN)64–. The ions [Fe(CN)6]4– is called complex ion & K+ is called counter ion. Fe(CN)2 + 4KCN → Fe(CN)2.4KCN l 4K+ + Fe(CN)64– Thus we see that in the molecular compound like Fe(CN)2. 4KCN, the individual compounds lose their identity. Such molecular compounds are called coordination (or complex) compounds. Complex ion : It may be defined as an electrically charged radical which consist of a central metal atom or ion surrounded by a group of ions or neutral molecules or both. A coordination compound consist of either (i)

A simple cation and a complex anion such as K4[Fe(CN)6], or

(ii)

A complex cation and a simple anion such as [Cu(NH3)4]SO4

(iii)

A complex cation and a complex anion such as [Co(NH3)6][Cr(C2O4)3] or

(iv)

A neutral molecule such as Cu(Gly)2, Ni(CO)4 etc.

Coordination sphere : The central metal ion and the ligands that are directly attached to it are enclosed in a square bracket which Werner has called coordination sphere or first sphere of attraction. The anions being outside the square bracket form the second sphere of attraction.

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WERNER'S THEORY Werner's coordination theory which could explain all the observed properties of complex compounds. More important postulates of this theory are : (i)

Most elements exhibit two types of valencies : Primary valency and (b) Secondary valency. (a) Primary Valency : This corresponds to oxidation state of the metal ion. This is also called principal, ionisable or ionic valency. It is satisfied by negative ions and its attachment with the central metal ion is shown by dotted lines. (b) Secondary or auxiliary valency : It is also termed as coordination number (usually abbreviated as CN) of the central metal ion. It is non-ionic or non-ionisable (i.e. coordinate covalent bond type). This is satisfied by either negative ions or neutral molecules.

(2)

Every element tends to satisfy both its primary and secondary valencies. In order to meet this requirement a negative ion may often show a dual behaviour, i.e. it may satisfy both primary and secondary valencies since in every case the fulfilment of coordination number of the central metal ion appears essential.

(3)

The ions attached to primary valencies possess ionizing nature whereas, the ions attached to secondary valencies do not ionise when the complex is dissolved in a solvent.

(4)

Every central ion tends to satisfy its primary as well as secondary valencies.

(5)

The secondary valencies are directional and are directed in space about the central metal ion. The primary valencies are non-directional. The presence of secondary valencies gives rise to stereo-isomerism in complexes.

(6)

The geometry of the complex ion depends on the coordination number.

(7)

Initially, Werner had pointed out co-ordination number of a metal atom to be four or six.

(8)

The six valencies were regarded to be directed to the corners of a regular octahedron circumscribed about the metal ion. For metals having four co-ordination number, the four valencies are either arranged in a square planar or tetrahedral nature.

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Sidgwick's concept of Effective Atomic Number - EAN concept (also called Noble Gas Rule) : Sidgwick suggested that after the ligands have donated a certain number of electrons to the central metal ion through bonding, the total number of electrons on the central atom, including those gained from ligands in the bonding is called the effective atomic number (EAN) of the central metal ion and in many cases this total number of electrons (i.e. EAN) surrounding the coordinated metal ion is equal to the atomic number of the inert gas. Example : EAN of Co(III) in [Co(NH3)6]3+ can be calculated as follows : Electrons in Coº atom = atomic number of Co

= 27 electrons

Electrons in Co3+ ion = 27 – 3

= 24 electrons

Electrons donated by 6 (: NH3) = 2 × 6

= 12 electrons –––––––––––––

EAN of Co(III) in [Co(NH3)6]

3+

= 24 + 12 = 36 –––––––––––––

EAN (= 36) of Co(III) is evidently equal to the atomic number of Kr. Exceptions of EAN Rule : EAN may be a few units more or less than the atomic number of the next inert gas. Complexes of Ni(II), Co(II), Ag(I) etc. which have more than one coordination number depending on the nature of the ligand, generally do not follow the EAN rule. LIGANDS The neutral molecules or ions which are directly linked with central metal ion or atom in a complex ion are called Ligands. e.g. in the complex ion, [Fe(CN)6]3– the six CN– ions are the ligands. The ligands are attached to the central metal ion or atom through coordinate bond or dative linkage. CLASSIFICATION OF LIGANDS : It is classified on the basis of (a) number of donor atoms present in the ligands (b) charge of ligands CLASSIFICATION BASED ON THE NUMBER OF DONOR ATOMS PRESENT IN THE LIGANDS The ligands of this class may be of the following types : (i)

Monodentate or unidentate ligands : The ligands which have only one donor atom and hence can coordinate to the central metal ion at one site only are called monodentate or unidentate ligands.

(ii)

Bidentate, tridentate....... polydentate ligands : The ligands having two, three, four, five or six donor atoms are called, bi, tri(or ter-), tetra- (or quadri-), penta-, and hexa- dentate ligands respectively. The bidentate, tridentate. etc. are called polydentate or multidentate ligands (literally dentate means toothed).

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Table : Some Common Polydentate Ligands. (Donor atoms having the lone pairs are indicated by asterisks) Type of Ligand

Name of Ligand

Structure of Ligand

2–

O Carbonato

`

Acetylacetonato

CO 3 2–

`

O—C—O * *

LM — O OP O MMH C — C| = C —||C — CH PP | MM PP H N Q *

Bi-dentate

Abbreviation

*

3

3

2 : 2'-Dipyridyl or

−

(acac) –

dipy

2, 2'-dipyridine

N *

Oxalato

N *

LM O O MMO — ||C —||C — O N −

*

*

OP P —P Q

(ox)2–

LMH C —||C — —||C — CH OP MM *N N* PP | | MN PQ O OH 3

Dimethyl glyoximato

2−

−

3

(dmg)– or (DMG)–

−

D U

C

A T

I O

N

S

*

H2N —CH2—CH2— NH2

en

Propylene diamine or

1 2 3 CH2 — CH — CH3 | | NH2 NH2 * *

pn

1, 2-diamino propane

E

*

Ethylene diamine

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Name of Ligand

Structure of Ligand

Abbreviation

Isobutylene diamine

* * H2N —C(CH3)2— NH2

i-bn

butylene diamine

* * H2N —CH(CH3)—CH(CH3)— NH2

bn

Tetra methyl-ethylene

* * H2N —C(CH3)2—C(CH3)2— NH2

tetrameen

diamine

– (oxin)– or (oxinate)–

8-hydroxyquinolinato

O *

N

1, 10 phenanthroline

o-phen

or o-phenathroline

*

As(CH3)2

o-phenylene bis

diars or D

dimethyl arsine

As(CH3)2 *

Glycinato

Biguanido

Diethylene triamine

E

D U

C

A T

I O

N

S

LM MMH N — CH N* 2

2

LM MMNH * MN

2

O || —C—O *

OP PP Q

−

(gly)–

OP NH N || || P — C — NH — C — NH P PQ *−

−

2

* * * H2N — (CH2 )2 — NH — (CH2 )2 — NH2

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(Big)–

dien

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Type of Ligand

Tridentate

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Name of Ligand

Structure of Ligand

Imino-di-acetato

LMOOC − MN* — H C — *NH — CH 2

Abbreviation

2

* N

OP PQ

− — COO *

2−

(IDA)2–

* N * N

2, 2', 2"-terpyridine

Tetradentate

Triethylene tetramine

Nitrilo triacetato

Pentadentate

terpy

H H H2N — (CH2 )2 — N — (CH2 )2 — N * * * —(CH2 )2 — NH2 *

*N

Ethylene diamine triacetato

Hexadentate

Ethylene diamine tetraacetato

E

D U

C

A T

I O

N

S

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- 3CH2.COO * CH2.COO * CH2.COO *

-* OOC—CH2 -* OOC—CH2 * —N

*

trien

(NTA)3–

3-

*CH2—COO

—

H

OOC—H2C * OOC—H2C * * —N

*

4-

CH2—COO * CH2—COO *

CO-ORDINATION CHEMISTRY

(EDTA)4– Y4–

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CLASSIFICATION BASED ON THE CHARGE These ligands may be neutral molecules, negatively charged ions (anions) or positively charged ions (cations), e.g., (a)

Neutral ligands :

(i)

Ligands which are named as molecule

(ii)

The names of neutral ligands are not systematic.

Ligands

Name

C5H5N

Pyridine

(C6H5)3 P

Triphenylphosphine

NH2CSNH2

Thiourea

H2NCH2CH2NH2

Ethylenediammine

Ligands which are not named as molecules

(iii)

Ligands

Name

Ligands

Name

H2O

aquo

NH3

ammine

CO

carbonyl

CS

thiocarbonyl

NO

nitrosyl

NS

thionitrosyl

Organic free radicals when act as ligands are named as usual radical •

Methyl

CH3 •

Ethyl

C 2H5 •

Phenyl

C 6H5

(b)

Negative ligands : If the names of the anions end in -ide, -ite or -ate, the endings of the names of the ligands used are -ido, -ito, and -ato respectively. Some examples of negative ligands are : Ligands

Name

CH3COO–

acetato

F–

fluoro

Cl–

Chloro

Br–

Bromo

I–

iodo

CN–

cyano

OCN–

cyanato

SCN–

thiocyanato when S-atom of the ligand coordinates with the central metal ion

NO2

E

D U

C

A T

I O

N

S

–

nitro, coordinates with the metal ion through N-atom

NO2–

nitrito, when the ligand ion coordinates with the metal through O-atom

OH–

hydroxo or sometimes hydroxy

H–

hydrido etc.

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Positive ligands : Examples of positively charged monodentate ligands are NO+

nitrosylium

NH2NH3+

hydrazium

Quite obviously the names of these ligands have the suffix-ium. OTHERS TYPE OF LIGANDS 1.

Bridging ligands : It is also possible that a monodenatre ligand may have more than one free electron pairs and thus may simultaneously coordinate with two or more atoms, i.e. the ligand forms two σ-bonds with two metal atoms & thus acts a bridge between the metal atoms. Such a ligand is called a bridging ligand and the resulting complex is known as bridged complex. Examples of bridging ligands are : OH–, F–, Cl–, NH2–, CO, O2–, SO42– etc.

2.

Ambidentate ligands : These are the ligands which have two or more donor atoms but in forming complexes only one donor atom is attached to the metal ion at a given time. Such ligands are called ambidentate ligands. Some examples of such type of ligands are given below :

O

ion

Nitro (M—NO2)

N

Nitrito (M—ON = O)

O CN–

Cyano (MCN)

ion

NCS–

Isocyano (MNC)

ion

Thiocyanate (MSCN) Isothinocyanate (MNCS)

3.

Chelating ligands : Polydentate ligands whose structures permit the attachment of their two or more donor atoms (or sites) to the same metal ion simultaneously and thus produce one or more rings are called chelate or chelating ligands (from the Greek for claw) or chelating groups. Example : NH2—CH2—CH2—NH2

(4) Coordination Number (C.N.) or Ligancy : It is the total number of the atoms of the ligands that can coordinate to the central metal ion. Numerically coordination number represents the total number of the chemical bonds formed between the central metal ion and the donor atoms of the ligands. Table : Fimiliar C.N.'s of some common metal ions Univalent

C.N.

Divalent

C.N.

ion

ions

ions

ions

Ag

2

V

6

Sc 3+ 6

Pt 4+ 6

2, 4

Fe2+ 6

Cr3+ 6

Pd4+ 6

Tl +

2

Co2+ 4, 6

Fe 3+ 6

Cu

2, 4

Ni

4, 6

Co3+ 6

Cu2+ 4, 6

Os3+ 6

Zn2+ 4

Ir3+

6

Au

4

+

Au +

2+

2+

2+

4

2+

4

Pd Pt

Ag

E

D U

C

A T

I O

N

S

2+

3+

Trivalent C.N. Tetravelent

C.N.

4

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CO-ORDINATION CHEMISTRY

NOMEN CLATURE OF COORDINATION COMPOUNDS The following rules are adopted for naming all types of coordintion compounds (1)

If a coordination compound is ionic, the name of cation is given first whether or not it is the complex ion followed by the name of the anion just like naming a simple salt.

(2)

The name of cation and anion are separated by a space.

(3)

With in a complex ion, the ligands are name first followed by the metal ion.

(4)

The ligands are named according to the following rules :

(1) Naming of Ligands : If the coordination sphere of a given complex compound contains various types of ligands. The ligands are named in alphabetical order. The prefixes di, tri etc. are not to be considered while determining this alphabetical order. For example [Co(NH 3) 4 (NO 2)Cl] + ion is named as tetramminechloronitro cobalt (III) ion. (2) Naming of the Negative Ligands : In general, if the anion name ends in -ide, -ite or -ate, the final -e is replaced by o, giving -ido, -ito and -ato respectively, e.g. SO32– (sulphito), SO42 (sulphato), CH3COO– (acetato), S2– (thio or sulphido), NO3– (nitrato), NH2– (imido), NH2– (amino or amine), N3– (azido), NHOH– (hydroxylamido), HON = C(CH3)C(CH3) = NO– (dimethylglyoximato). Some exceptions to this rule are : F– (fluoro), Cl– (chloro), CN– (cyano), O2– (oxo), OH– (hydroxo), O22– (peroxo), O2H– (perhydroxo). (3) Naming of the Neutral Ligands : For neutral ligands, the names are not systematic. For less common neutral ligands (e.g. PH3), the names of free molecules is used as such. For some of the more common neutral ligands, special names are used e.g., H 2O (aquo), NH3(ammine), CO(carbonyl), NO(nitrosyl), CS(thiocarbonyl), NS(thionitrosyl). (4) Naming of the Positive Ligands : Positively charged ligands have suffix ium, e.g. NH 2 NH3+ is called hydrazinium, and NO+ is nitrosylium. (5) Indication of the Number of Ligands : (a)

If the number of a particular ligands is more than one in the complex ion, the number is indicated by using Greek prefixes; di-, tri-, tetra-, penta-, hexa- etc.

(b)

In case of chelating ligands like ethylenediamine, trialkyl phosphine which contain the prefixes di-, tri- etc. in their ligand names, the prefixes bis (for two), tris-, (for three), tetrakis- (for four), pentakis- (for five), liexakis- (for six) etc. are used before their names.

(c)

For example : [CoIII (NH3)2 (en)2] Cl3 - diammine -bis (ethylene diamine) cobalt (III) chloride, [Co III (en)3]2 (SO4)3 - tris (ethylene diamine) cobalt (III) sulphate, [FeII (CN)2 (CH3NC)4] - dicyano tetrakis (methyl-lisocyanide) iron (II).

(6) Naming of the bridging ligands of the bridged polynuclear complexes : The complexes having two or more metal atoms are called polynuclear complexes. In these complexes the bridging group is indicated in the ormula of the complex by separating it from the rest of the complex by hyphens and by adding the prefix µ before its name. The Greek letter µ should be repeated before the name of each different bridging group.

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Examples :

OH (N H 3 ) 5 C r

(N H 3 ) 5

4+

NH2 (N H 3 ) 4 C o

µ-hydroxo-bis {penta-ammine chromium (III)} chloride

C l5

C o(N H 3 ) 4

µ-amido-µ-nitro octa-ammine dicobalt (III) ion

NO2 [(CO)3 Fe(CO)3Fe(CO)3]

Tri-µ-carbonyl-bis (tricarbonyl iron)

(7) Naming of ambidentate ligands : The ligands which can be coordinated to the central metal ion through either of the two donor atoms are called ambidentate ligands. Such ligands are either named by special names such as thiocyanate for –SCN and isothiocyanate for –NCS, nitro for –NO2 and nitrito for –ONO or the symbol of the element coordinated with the metal ion is written after the name of the ligand, e.g. thiooxalate-S for a thio-oxalate group coordinated to the metal ion through sulphur atom and thio-oxalato-O for a thio-oxalate group coordinated through oxygen atom. [Co(NH3)5 – ONO] Cl2

→ Nitro-O or Nitrito complex

[Co(NH3)5 – NO2] Cl2

→ Nitro-N or Nitrato complex

[Pd (diph) (SCN)2]

→ Thiocyanate-S complex

[Pd (diph) (SCN)2]

→ Thiocynate-N complex

NAMING OF THE CENTRAL METAL ION AND MONONUCLEAR COMPLEX : (1) Anionic complexes : (a)

In naming anionic complexes like [Pt(NH3)Cl5]–, [Cr(en)I4]– etc. ligands are named first and then the central metal ion.

(b)

To name the central metal ion the suffi "ate" is attached to its name and in order to indicate the oxidation state of the metal ion this suffix is followed by Roman numeral (such as I, II, III etc.) in the parentheses at the end of the name of the complex without a space between the two. (o) is used for an oxidation state of zero.

(c)

For negative oxidation state the negative sign is placed before the Roman numeral. Examples : [CrIII (en) I4]–

Tetra iodo (ethylene diamine) chromate (III) ion.

[NI0 (CN)4]4–

Tetracyanonicklate (0)

[Co (CO)4]

Tetracarbonylcobaltate (–I)

–

–

(2) Cationic and neutral complexes : In case of cationic and neutral complexes like [Cr(H2O)4Cl2]+, [Cu(NH3)4]2+, [Ni(CO)4], central metal ion followed by a Roman numeral in parentheses to indicate its oxidation state. The suffix ate is not attached to the name of the metal in case of these complexes. Examples : [CrIII (H2O)4 Cl2]+

Dichloro-tetra-aquo chromium (III) ion

Cationic complexes

[Al (OH) (H2O)5]

Hodroxo pentaaquo aluminium (III) ion

Cationic complexes

[Ni0(CO)4]0

Tetra carbonyl Nickel (0)

Neutral

III

E

D U

C

A T

I O

N

S

3+

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ISOMERISM IN CO-ORDINATION COMPOUNDS Similar to the isomerism exhibited by organic compounds, specific orientation of ligands in space around the central atom gives rise to isomerism in co-ordination compounds. The isomerism in co-ordination compounds can be classified as : 1. Structural isomerism 1.

2. Stereoisomerism

STRUCTURAL ISOMERISM : (i)

Ionization isomerism (a)

Complexes having same empirical formula but producing different ions in solution state, e.g. [Co(NH3)5Br] SO4

(b)

(ii)

and

[Co(NH3)5SO4]Br

Pentaamminebromocobalt (III) sulphate

Pentaamminesulphatocobalt (III) bromide

(Violet)

(Red)

Scuh an isomerism is due to the interchange of groups between the co-ordination sphere of the metal ion and the ions outside the sphere.

Hydration isomerism : This isomerism arises when different number of water molecules are present within and outside the co-ordination sphere. For example, 3 hydration isomers of CrCl3.6H2O are ; [Cr(H2O)6]Cl3 Violet

[Cr(H2O)5Cl]Cl2.H2O

[Cr(H2O)4Cl2]Cl.2H2O

Green

Green

(iii) Linkage isomerism : This isomerism arises when ligand has two possibilities in its mode of attachment to the metal atom. For example, the ligand – NO2 binds with metal atom through N or O in the following example. [Co(NH3)5ONO]Cl2

[Co(NH3)5NO2]Cl2

Pentaamminenitritocobalt(III) chloride

Pentaamminenitrocobalt(III) chloride

(Red)

(Yellow)

(iv) Ligand Isomerism : As the name imlies, ligands isomerism arises from the presence of ligands which can adopt different isomeric forms An example is pronded by diaminopropane which have the amine groups in the terminal or adjacent position. e.g., [Co(pn)2Cl2]+, [Co(tn)2 Cl2]+, [Co(pn)(tn)Cl2]+ (iv) Co-ordination isomerism : This isomerism exist in polynuclear complex (where both complex anions and complex cations are present) by partial or complete exchange of ligands. For example, [Cr(NH3)6].[Cr(SCN)6] and [Cr(NH3)4(SCN)2].[Cr(NH3)2.(SCN)4] [Co(NH3)6].[Cr(C2O4)3] and [Cr(NH3)6].[Co(C2O4)3] [Co(en)3].[Cr(CN)6] and [Cr(en)3].[Co(CN)6] (v) Co-ordination position Isomerism : This type of isomerism is exhibited by polynuclear complexes by changing the position of ligands with respect to different metal atoms present in the complex e.g.,

OH (NH3)4 Co

OH Co (NH3)2Cl2

SO4

OH

E

D U

C

A T

I O

N

S

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Cl(NH3)3 Co

Co (NH3)3Cl SO4 OH

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(vi) Polymerisation Isomerism : This isomerism arises when compounds have the same stoichiometric composition but different molecular composition (which is an integer multiple of stoichiometric composition), (a)

It is special case of coordination isomerism.

(b)

It can generate a series of compound having same empirical formula but different molecular masses. This referred as polymerisation isomerism.

(c)

In polymerisation isomerism no polymerisation takes place.

(d)

For e.g., with empirical formula [Co(NH3)3(NO)3] x, can show many coordination compound with following composition. (i)

x = 2,

[Co(NH3)6] [Co(NO2)6]

(ii)

x = 2,

[Co(NH3)4 (NO2)2] [Co(NH3)2(NO2)4]

(iii)

x = 3,

[Co(NH3)5(NO2)] [CO(NH3)2(NO2)4]2

(iv)

x = 4,

[Co(NH3)6] [Co(NH3)2(NO2)4]3 e.g., [Pt(NH3)2Cl2] and [Pt(NH3)4] [PtCl4]

2.

STEREOISOMERISM OR SPACE ISOMERISM : (a)

When two compounds contain the same ligands coordinated to the same central ion, but the arrangement of ligands in space is different, the two compounds are said to be stereoisomers and the type of isomerism is called steroisomerism.

(b)

It is also called space isomerism.

(c)

Stereoisomerism is of two types : (i) Geometrical or cis-trans isomerism, and (ii) optical or mirror-image isomerism.

(i)

Geometrical Isomerism : This type of isomerism arises due to different positions occupied by ligands around the metal atom. When two similar ligands are adjacent to each other, then isomer is referred as cis-isomer and when they are at opposite positions then trans-isomer. Note : 1.

Geometrical isomers is not observed in complex of coordination number 2 and 3.

2.

Geometrical isomerism is not observed in complexes of coordination number 4 of tetrahedral geometry.

3.

Geometrical isomerism cannot arise in a tetrahedral complex (C.N. = 4) because this geometry contains all the ligands in cis (i.e., adjacent from the other three ligands and all bond angles are the same (= 109.50) This isomerism is, however, found in many square planar (C.N. = 4) and octahedral (C.N. = 6) complexes.

4.

The complexes of general formula, Ma3b or Mab3, or Ma4, of square planar geometry do not show geometrical isomerism.

5.

The complexes of general formula, Ma6 and Ma5b, of octahedral geometry do not show geometrical isomerism.

Geometrical isomerism in Square planar complex : Square planar complexes having formula M A2B2 , MA 2BC , M ABCD show geometrical isomerism

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EXAMPLES : (i)

[Pt(NH3)2Cl2] resembling to formula MA 2B2 exists in cis and trans forms

(ii)

[Pt(py)2NH3Cl]+ resembling to formula MA 2BC exists as.

(iii)

Complexes of MABCD type exist in 3-geometrical isomeric forms e.g., [Pt(NH3) (NH2OH) (py) (NO2)]+

(iv) Complexes of M A2 type may also show geometrical isomerism in square planar complexes where A is unsymmetrical bidentate ligand. e.g., [Pt(NH2CH2COO)2]

(v)

Bridged complex of type M2 A2B 4 also exists as cis and trans-isomers. e.g., [Pt.Cl2.P(C2H5)3]2

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MEMORY CHART : GEMETRICAL ISOMERISM IN SQUARE PLANAR COMPLEXES

Geometry of molecule having Co-ordination No. 4

Tetrahedral (Will not show G.I.)

Square Planar

Bidentate Ligand

Mondentate Ligand

MA BC (Will show G.I.)

MA" (Will not show G.I.)

MA!B (Will not show G.I.)

MA B (Will show G.I.)

MABCD (Will show G.I.)

MAB!

Symmetrical bidentate ligand

Unsymmetrical bidentate ligand

M(AA)

M(AA’)

(Will not show G.I.)

(Will show G.I.)

(Will not show G.I.)

(b) Geometrical isomerism in Octahedral Complexes : Positions A–B, B–C, C–D, A–D etc. which are at 90º to each other are cis with respect to each other while A–F, B–D, C–E etc. (at 180º to each other) are trans positions.

Complexes of type M A 4B2 , MA 4BC will show geometrical isomerism.

(i)

e.g., [Fe(CN)4 (NH3)2]2– resembling with M A 4B2 type.

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Facial (fac) and meridional (mer) isomers : Such type of geometrical isomers are in compounds of type M A3B3 type. If three ligands of same type are at the corners of a triangular face then it will be fac-isomer. If they are at the corners of square plane then it will be mer-isomer.

e.g., [RhCl3(py)3]

(iii)

Complexes of type M(A)3 also show geometrical isomerism where A is bidentate unsymmetrical ligand. e.g., [Cr(NH2CH2COO)3]

(iv) Complexes of type M(A)2 B2 or M(A)2 BC also show geometrical isomerism, here A is symmetrical bidentate ligand.

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Complex having Co-ordination No. 6

Number of Geometrical isomers

1.

MA6

0

2.

MA5B

0

3.

MA4B2

2

4.

MA3B3

2

5.

MA 2B2C2

5

6.

MABCDEF

15

OPTICAL ISOMERISM : (a)

A coordination compound which can rotate the plane of polarised light is said to be optically active.

(b)

When the coordination compounds have same formula but differ in their abilities to rotate directions of the plane of polarised light are said to exhibit optical isomerism and the molecules are optical isomers.

(c)

The optical isomers are pair of molecules which are non-superimposable mirror images of each other.

(d)

The isomer which rotates the plane of polarised light to right direction is termed dextro (d-form) while the isomer which rotates plane of polarised light to left direction is termed laevo (l-form).

(e)

The two optically active isomers are collectively called enantiomers. Enantiomers are mirror imgae to each other and their physical properties are different.

OPTICAL ISOMERISM IN TETRAHEDRAL COMPLEXES (a)

Optical isomerism is expected in tetrahedral complexes of the type Mabcd but no optical isomer has been isolated until now.

(b)

However, compounds containing two unsymmetrical bidentate ligands have been resolved into optical isomers and are known for Be(II), Zn(II) and B(III). For example, Bis-benzoylacetonatoberyllium (II) exhibits optical isomerism.

C

C

Be C

O

C C

H—

H—C H#C$

O

O

O

CH! O

C C

H— C

C

C C$H#

CH!

O

O

O

H—

CH!

H!C

C H#C$

H#C$

Mirror plane OPTICAL ISOMERISM IN SQUARE PLANAR COMPLEXES Optical isomers rarely occur in square planare complexes on account of the presence of axis of symmetry. OPTICAL ISOMERISM IN OCTAHEDRAL COMPLEXES Optical isomerism is very common in octahedral complex.

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Octahedral complexes of general formulae show optical isomerism. (i)

[Ma2 b2c2 ]n±

(ii)

[Mabcdef]

(iii)

[M(AA)3]n±

(iv) [M(AA)2a2]n± (v)

(where AA = symmetrical bidentate ligands)

[M(AA)2ab]n±

(vi) [M(AB)3]n±

(where AB = unsymmetrical bidentate ligands)

Examples : [Ma2 b2 c2 ]n±,

(i)

e.g., [Pt(py)2(NH3)2Cl2]2+

py Cl

py

py

2+

Cl

Pt Cl

(ii)

Pt

NH3 NH3 Cis-d-isomer

[Mabcdef];

NH3 Mirror

Cl NH3

Cis-l-isomer

e.g., [Pt(py)NH3NO2ClBrI]

Br

Br

py

O2N

NO2 Pt

py Pt

NH3

Cl

Cl

H3N

I d-isomer (iii)

Mirror

[M(AA)3]n±;

I I-isomer

e.g., [Co(en)3]3+

en

en

3+

en

Co

d-form

D U

C

A T

I O

N

S

3+

Co en

E

2+

py

en Mirror

l-form

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en

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Co

en en

'Meso' or optically inactive form

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(iv) [M(AA)2a2]n±,

e.g., [Co(en)2Cl2]+

en

Cl

en

+

Cl

Co

Co en

Cl

Cl

en

+

Cis-d-isomer

Cis-l-isomer

Mirror

Trans form of [M(AA)2a2]n± does not show optical isomerism (v)

[M(AA)2ab]n±;

e.g., [Co(en)2 NH3Cl]2+

en

Cl

Cl Co

Co en

H3N

2+

en

2+

Cis-d-isomer (vi) [M(AB)3];

NH3

en Mirror

Cis-l-isomer

e.g., [Cr(gly)3]

gly

gly Cr

gly

gly

Cr

gly Cis or trans-d-isomer

gly Mirror

Cis or trans-l-isomer

Some more examples are :

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A T

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N

S

(a)

[Cr(ox)3]3–

(b)

[Fe(dipy)3]2+

(c)

[Cr(ox)2(H2O)2]–

(d)

[Pt(en)3]4+

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Possible number of stereoisomers for complex Formula

Number of sterioisomers

Pairs of enantiomers

Ma6

1

0

Ma5b

1

0

Ma4 b2

2

0

Ma3 b3

2

0

Ma4 bc

2

0

Ma3bcd

5

1

Ma2bcde

15

6

Mabcdef

30

15

Ma 2 b 2 c 2

6

1

Ma2 b2 cd

8

2

Ma3 b 2 c

3

0

M(AA)(BC)de

10

5

M(AB)(AB)cd

11

5

M(AB)(CD)ef

20

10

M(AB)3

4

2

M(ABA)cde

9

3

M(ABC)2

11

5

M(ABBA)cd

7

3

M(ABCBA)d

7

3

M(AA)3

2

1

Capital letters represent chelating ligands, lower case represent monodenatate ligands.

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MODERN THEORIES OF BONDING CO-ORDINATION COMPOUNDS Four distinct approaches have been given to explain properties of complex compounds such as colour, geometry and magnetic properties. These are : 1. Valence bond theory

2. Crystal field theory

3. Molecular orbital theory

4. Ligand field theory

VALENCY BOND THEORY : Valence bond theory makes the following assumptions : (i)

The central metal atom provides a number of empty orbitals equal to its co-ordination number for the formation of covalent bonds (or more accurately co-ordinate bonds) with ligand orbitals.

(ii)

The empty orbitals of the metal ion hybridize to give an equal no. of hybrid orbitals of equivalent energy.

(iii)

The metal atom or ion can use (n – 1)d, ns, np, nd orbitals for hybridization to yield square planar, tetrahedral or octahedral geometry.

(iv) These hybridized orbitals then overlap with ligand orbitals which can donate an electron pair for bonding. (v)

If the complex contains unpaired electrons, the complex is paramagnetic in nature whereas, if it does not contain unpaired electrons, the complex is diamagnetic in nature.

(vi) Under the influence of a strong ligand (like, NH3, CN–, CO, etc.), the electrons can be forced to pair up against the Hund's rule of maximum multiplicity. Step-1 : Central metal ion should be identified. Step-2 : Determine its oxidation state. Step-3 : Write electronic configuration of central metal atom Step-4 : Using oxidation state write electronic configuration of central ion. Step-5 : Now look the number of ligands and decide hybridisation 4 – sp3

6 – d2sp3

4 – dsp2

6 – sp3d2

Step-6 : Now vacant appropriate number of 'd' orbital to have an empty set of obitals for hybridisation. Step-7 : Name hybrid orbitals. Step-8 : Donate electron pair from the ligands.

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Table : Co-ordination numbers, hybridized orbitals and geometry of some co-ordination compounds Co-ordination

Hybridized orbital

Examples of complexes

2

sp

[Ag(NH3)2]+, [Ag(CN)2]–

3

sp2

[HgI3]–

4

sp3

[FeCl4]–, [Ni(CO)4], [Zn(NH3)4]2+, [ZnCl4]2–

Number

[CuX4]2–, X = CN–, Cl–, Br–, I–, CNS– 4

dsp2 the d-orbital involved is dx2 − y2 orbital of the inner

shell, i.e., it is (n – 1) dx2 − y2 orbital. 5

dsp3, the d-orbital is (n – 1)

[CuCl5]3–, [MoCl5]0, [Fe(CO)5]0

dz2 orbital. 5

sp3d, the d-orbital is ndx2 − y2

[SbF5]2–, IF5

orbital. 6.

d2sp3, when d-orbitals are

[Cr(NH3)6]3+, [Ti(H2O)6]3+, [Fe(CN)6]2–

(n–1) d-orbitals (Inner orbital

[Co(NH3)6]3+, [PtCl6]2–, [CoF6]3–

complexes) or sp3d2, when d-orbitals are nd-orbital (Quter orbital complexes), in both cases d-orbitals are dz2 and dx2 − y2 orbitals.

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Table : Geometry (shape) and magnetic nature of some of the complexes (Application of valence bond theory) Complexes

Oxidation

Type of

Geometry

No. of unpaied

Magnetic

state

hybrid.

Shape

electrons

[NiCl4]2–

+2

sp3

Tetrahedral

2

Paramagnetic

[Ni(CN)4]2–

+2

dsp2

Square planar

0

Diamagnetic

[Ni(NH3)6]2+

+2

sp3d2 (Outer)

Octahedral

2

Paramagnetic

[Mn(CN)6]4–

+2

d2sp3 (Inner)

Octahedral

1

Paramagnetic

[CuCl4]2–

+2

sp3

Tetrahedral

1

Paramagnetic

[Cu(NH3)4]2+

+2

dsp2

Square planar

1

Paramagnetic

[Cr(H2O)6]3+

+3

sp3d2(Outer)

Octahedral

3

Paramagnetic

[CoF6]3–

+3

sp3d2 (Outer)

Octahedral

4

Paramagnetic

[Co(NH3)6]3+

+3

d2sp3 (Inner)

Octahedral

0

Diamagnetic

[Fe(CN)6]4–

+2

d2sp3(Inner)

Octahedral

0

Diamagnetic

[Fe(H2O)6]2+

+2

sp3d2 (Outer)

Octahedral

4

Paramagnetic

[Fe(NH3)6]2+

+2

sp3d2 (Outer)

Octahedral

4

Paramagnetic

[Fe(CN6)]3–

+3

d2sp3 (Inner)

Octahedral

1

Paramagnetic

[Fe(CO)5]

0

dsp3 (Inner)

Trigonal

0

Diamagnetic

bipyramidal

Limitations of valence bond theory : The valence bond theory was fairly successful in explaining qualitatively the geometry and magnetic properties of complexs. However, it has a number of limitations. (i)

The theory does not offer any explanation why most of complexes are coloured

(ii)

The theory does not offer any explanation for the existence of inner-orbital and outer-orbital complexes.

(ii)

The theory does not explain why certain complexes are labile while others are inert.

(iii)

In the formation of [Cu(NH3)4]2+, one electron is shifted from 3d to 4p-orbital. The theory

is silent about the energy availability for shifting such as electron. Such an electron can be easily lost, then why [Cu(NH3)4]2+ complex does not show reducing properties.

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CRYSTAL FIELD THEORY (CFT) (i)

CFT regards a complex as a combination of a central ion surrounded by other ions or molecules with electrical dipoles (i.e., ligands). It considers these ligands as point charges or as point dipoles.

(ii)

In its simplest treatment CFT does not consider covalent bonding in complexes, but the bonding between the metal cation and ligands arises from the electrostatic attraction between the nucleus of the metal cation and the partial negative charge invariably present on the ligands. The interaction between the electrons of cation and those of the ligands is entirely repulsive. It is these repulsive forces that are responsible for causing the splitting of the d-orbitals of the metal cation. The bonds between the metal and the surrounding ligands are thus purely ionic.

(iii)

The CFT does not provide for electrons to enter the metal orbitals. Thus the metal ion and the ligands do not mix their orbitals or share electrons i.e., it does not consider any orbital overlap.

(iv)

CFT gives a representation of bonding that is purely an electrostatic or coulombic interaction between positively charged (i.e., cation) and negatively charged (i.e., anions or dipole molecules which act as ligands) species.

(v)

Crystal field theory assumes ligands as point charges and the attraction between the central metal and the ligands in a complex is purely electrostatic. In an isolated gaseous metal ion, the five d orbitals do all have the same energy (degenerate). But in octahedral & tetrahedral complexes, the energy of the dorbitals is raised because of repulsion between the field produced by the ligand and the electrons on the metal.

d-orbitals e.g., t2g degeneracy

Low spin & high spin complexes

Explanation of Colour & Magnetic Properties

Assumptions in CFT

Crystal field Theory

Relation between ∆O & ∆t

Octahedral splitting of d-orbitals

Tetrahedral spiliting of ‘d’ orbitals

Factors affecting ∆O & ∆t Nature of ligand

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SPLITTING OF D-ORBITALS IN OCTAHEDRAL COMPLEXES : (i)

Let us consider an octahedral complex, [ML6]n+, in which the central metal ion is placed at the centre of the octahedron and is surrounded by six ligands which reside at the six corners of the octahedron as shown ahead.

(ii)

The electrons in d-orbitals of the metal cation are repelled by negative point charge or by the negative end of the dipole of the ligands.

(iii)

This repulsion increase the energy of all the five d-orbitals.

(iv) If all the ligands approaching the metal cation are at an equal distance from each of d-orbitals (i.e., the ligand field is spherically symmetrical), the energy of each of the five d-orbitals increases by the same amount, i.e., all the d-orbitals remains degenerate, although they possess higher energy. (v)

However, this is not possible as it is only hypothetical situation.

(vi) Since the lobes of two eg orbitals is lie in the path of approaching ligands see, fig. The electrons in these orbitals experience greater force of repulsion than those in t2g orbitals whose lobes are directed in space between the path of the ligands see fig. i.e, energy of eg orbitals is increased while that of t2g orbitals is decreased.

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(vii) Thus, an energy difference exists between two sets of orbitals. (viii) This energy difference is called crystal field splitting energy and is represented by ∆O (the subscript O stands for octahedral). It measures the crystal field strength of the ligands. (ix)

The crystal field splitting occurs in such a way that average energy of the d-orbitals does not change.

Fig : Crystal field splitting of d-orbitals

(x)

in an octahedral complex

(a)

Five degenerate d-orbitals of free metal cation

(b)

Hypothetical degenerate d-orbitals at higher energy level under spherically symmetrical ligand field

(c)

Spilitting of d-orbitals under the influence of ligands.

Thus, three orbitals lie at an energy that is lie at an energy

2 ∆ below the average orbital energy and two d-orbitals 5 0

3 ∆0 above the energy. 5

Fig : (xi)

The energy gap between t2g and eg sets is also denoted by 10 Dq.

(xii) Energy of t2g orbitals is 4Dq less than that of hypothetical degenerate d-orbitals and that of eg orbitals is 6Dq above that of the hypothetical degenerate d-orbitals. Thus, t 2g set loses an energy equal to 0.4 ∆0 or Dq while eg set gains energy equal to 0.6 ∆0 or 6 Dq.

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CRYSTAL FIELD SPLITTING IN TETRAHEDRAL COMPLEXS : (i)

The tetrahedral arrangment of four ligands surrounding the metal ion Mm+ may be depicted as shown in fig.

(ii)

The three d-orbitals, i.e., t 2g orbitals are close to the approaching ligands. As a result of this t2g electrons suffer more repulsion than eg electrons.

(iii)

The energy of f2g orbitals increases more than eg orbitals.

(iv) The splitting is shown in fig.

Fig : Crystal field splitting of d-orbitals in a tetrahedral complex (v)

The energy gap between two sets of orbitals is designated as tetrahedral complex).

∆t (The subscript t indicates

(vi) It is observed that ∆t is considerably less than ∆O. (vii) It has been found that ∆t =

4 ∆O 9

DISTRIBUTION OF d-ELECTRONS IN t2g AND eg ORBITALS IN OCTAHEDRAL COMPLEX The distribution of d-electron in t2g and eg orbitals takes place on the basis of the nature of ligands. Two cases may arise. (a) When the ligands are weak : (i)

Under the influence of weak ligands, the energy difference, ∆O, between t2g and eg sets is relatively small and hence all the five d-orbitals may be supposed to be degenerate, i.e., all the d-orbitals have nearly the same energy and the distribution of electrons in t2g and eg sets occurs according to Hund's rule, i.e., electron will pair up only when each of the five d-orbitals is at least singly occupied.

(ii)

When the ligands are weak, the first three electrons numbered 1, 2, 3 go to t 2g set, those numbered 4, 5, go to eg set, those numbered 6, 7, 8, go to t2g set and the remaining two electron numbered 9, 10 wil occupy eg set. This can be shown as : t2g1,

(iii)

E

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C

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N

S

2, 3

4,  → eg

5

 → t2g

6, 7, 8

9,  → eg

10

In complexes of weak ligands, ∆O is less than P(P is called average pairing energy which is the energy required to pair two electrons in the same orbital). ∆O, the octahedral crystal field splitting energy, tends to force as many electrons to t2g set while P tends to prevent the electrons to pair in the t2g level.

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(b) When the ligands are strong : (i)

Under the influence of strong ligands, the energy difference between t 2g and eg sets is relatively high and thus the distribution of d-electrons in t2g and eg sets does not obey Hund's rule.

(ii)

The first electrons numbered 1, 2, 3, 4, 5, 6 will go to t 2g set and remaining four electrons numbered 7, 8, 9 and 10 will go to eg set. t2g1,

2, 3, 4, 5, 6

 → eg7,

8, 9, 10

For these complexes, ∆O is higher than P.

(ii)

The above points are summarised in the following table :

Strong field (Low spin complexes) Weak field (High spin complexes) dxions

∆O > P Configuration

∆O < P

No. of unpaired

spin

Configuration

electrons

No. of unpaired

Common spin

examples

electrons

d1

t 2g 1 eg o

1

1/2

t 2g1 ego

1

1/2

Ti3+

d2

t 2g 2 eg o

2

1

t 2g2 ego

2

1

V3+

d3

t 2g 3 eg o

3

3/2

t 2g3 ego

3

3/2

Cr3+

d4

t 2g 4 eg o

2

1

t 2g3 e g1

4

2

Mn3+

d5

t 2g 5 eg o

1

1/2

t 2g3 e g2

5

5/2

Mn2+,Fe3+

d6

t 2g 6 eg o

0

0

t 2g4 e g2

4

2

Fe2+,Co3+

d7

t 2g6 eg 1

1

1/2

t 2g5 e g2

3

3/2

Co2+

d8

t 2g6 eg 2

2

1

t 2g6 e g2

2

1

Ni2+

d9

t 2g6 eg 3

1

1/2

t 2g6 e g3

1

1/2

Cu2+

d 10

t 2g6 eg 4

0

0

t 2g6 e g4

0

0

Zn2+

The following conclusion are derived from the table :

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(i)

The distribution of electrons of d1, d2, d3, d8, d9 and d10 in t2g and eg sets for both strong and weak octahedral ligands field is the same.

(ii)

For each of d4, d5, d6 and d7 there is difference in the arrangement of electrons in weak and strong ligands field.

(iii)

Weak field complex of d4, d5, d6 and d7 ions have greater number of unpaired electrons than those of strong field complexes and thus, the resultant spins of weak field complexes have higher value than strong field complexes.

(iv)

Hence, the complexes of weak field ligands are called high spin complexes and the complexes of strong field ligands are called low spin complexes. Calculations show that coordination entities with four to seven d-electrons are more stable for strong field as compared to weak field cases.

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Distribution of d-electrons in t2g and eg orbitals in tetrahedral complexes The distribution of d-electrons in the t2g and eg sets in a tetrahedral complex in presence of weak and strong ligands field has been shown in the following table.

Weak field (high spin complexes) Configuration

No. of

spin

unpaired

Strong field (Low spin complexes) Config-

No. of

uration

unpaired

electrons

E

D U

C

A T

spin

electrons

d1

t2g0 e g1

1

1/2

t 2g0 e g1

1

1/2

d2

t2g0 e g2

2

1

t 2g0 e g2

2

1

d3

t2g1 e g2

3

1

t 2g0 e g3

1

1/2

d4

t2g2 e g2

4

2

t 2g0 e g4

0

0

d5

t2g2 e g2

5

2

t 2g1 e g4

1

1/2

d6

t2g3 e g3

4

2

t 2g2 e g4

2

1

d7

t2g3 e g4

3

1

t 2g3 e g4

3

1

d8

t2g4 e g4

2

1

t 2g4 e g4

2

1

d9

t2g5 e g4

1

1/2

t 2g5 e g4

1

1/2

d 10

t2g6 e g4

0

0

t 2g6 e g4

0

0

I O

N

S

1 2

1 2

1 2

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1 2

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SPECTROCHEMICAL SERIES (i)

For any given metal cation, the magnitude of crystal field splitting energy depends on the nature of ligands.

(ii)

The greater the ease with which the ligand can approach the metal ion, the greater wil be the crystal field splitting caused by it.

(iii)

The ligands which affect only a small degree of crystal field splitting are termed weak field ligands.

(iv) When the ligands are arranged in order of the magnitude of crystal field splitting, the arrangement, thus, obtained is called spectrochemical series. I– > Br– < Cl– < NO3– < F– < OH– < OX2– < H2O < py = NH3 < en < dipy < o-Phen < NO2– < CN– < CO Weak field ligands

Increasing crystal field

Strong field ligands

From the above arrangement it is clear that ligands before H2O such as I–, NO3–, OH–, etc. are weak field ligands while the ligands after H2O such as NO2–, CN–, CO etc., are strong field ligands.

(v)

(vi) Stronger field ligands cause greater crystal splitting i.e., ∆O value for octahedral compelx is high. (vii) Besides the nature of ligands, there are some other factors which affects the crystal field splitting energy. These factors are : (a)

The identity of the metal : The crystal field splitting, ∆, is about 50 percent higher for the second transition series compared to the first whereas the third series is about 25 percent higher than second. There is a small increase in the crystal field splitting along each series.

(b)

The oxidation state of the metal : Generally, the higher the oxidation state of the metal, the greater the crystal field splitting. For example, most of the cobalt (II) complexes have low values of ∆ whereas all cobalt (III) complexes have high values of ∆. 4 The number of ligands : The crystal field splitting for a tetrahedral environment is about 9 that for an octahedral environment.

(c)

Application of CFT Crystal field theory have the following applications

E

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A T

(1)

It can provide an explanation for the observed colours of transition metal complexes.

(2)

It is helpful in determining the number of unpaired electrons in a high spin (HS) and low spin (LS) octahedral complexes.

(3)

It is used in determining magnetic moment of complexes.

(4)

It is used in calculating crystal field stabilisation energy (CFSE).

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STABILITY OF COORDINATION COMPOUNDS IN SOLUTIONS (i)

A coordination compound is formed in solution by the stepwise addition of ligands to a metal ion.

(ii)

Thus, the formation of the complex, MLn (M = central metal cation, L = monodentate ligand and n = coordination number of metal ion) may be supposed to take place by the following n consecutive steps. ML

M + L

K1 =

[ML] [M][L]

K2 =

[ML 2 ] [ML][L]

ML2 + L

ML + L

ML2

ML3

[ML 3 ] K3 = [ML ][L] 2 MLn–1 + L

MLn

[ML n] Kn = [ML ][L] n−1 K1, K2, K3,......, Kn are called stepwise stability constants. (iii)

With a few exceptions, the values of successive stability constants decrease regularly from K1 to Kn.

(iv) The overall stability constant, K is given as : M + nL

Mn

K = K1 K2 K3,......Kn = (v)

[ML n ] [M][L]n

The higher the overall stability constant value of the complex, the more stable it is.

(vi) Alternatively 1/k value called instability constant explain the dissociation of the complex into metal ion and ligands in the solution.

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(vii) The values of stability constants for some of the complex are given below : Complex

Stability constant

[Cu(NH3)4]2+

4.5 × 1011

[Ag(NH3)2]+

1.6 × 107

[Co(NH3)6]2+

1.12 × 106

[Co(NH3)6]3+

5.0 × 1033

[AgCl2]–

1.11 × 105

[AgBr2]–

1.28 × 107

[Ag(CN)2]–

1.0 × 1022

[Cu(CN)4]2–

2.0 × 1027

[Fe(CN)6]3–

7.69 × 1043

From the above values, some conclusion are drawn : (a)

The values of stability constants differ widely depending on the nature of the metal ion and the ligand. In general, higher the charge density on the central ion, the greater is the stability of its complexes, i.e., the higher value of

ch arg e , the greater is the stability of its radius of the ion

complexes. (b)

Electronegativity of the central ion influences the stability. The higher the electronegativity of the central ion, the greater is the stability of its complexes.

(c)

The more basic a ligand, the greater is the ease with which it can donate its lone pairs of electrons and therefore, the greater is the stability of the complexes formed by it.

(d)

The higher the oxidtation state of the metal, the more stable is the complex. The charge density of Co3+ ion is more than Co2+ ion and thus, [Co(NH3)6]3+ is more stable than [Co(NH3)6]2+. Similarly, [Fe(CN)6]3– is more stable than [Fe(CN)6]4–.

(e)

The cyano and ammine complexes are far more stable than those formed by halide ions. This is due to the fact that NH3 and CN– are strong Lewis bases. The complexes of bivalent cations (M2+) of 3d-series show the following order of stability :

(f)

Cation

Mn2+

Fe2+

Co2+

Ni2+

Cu2+

Ionic size

0.91

0.83

0.82

0.78

0.69

decrease Stability of the complex (e)

E

D U

C

A T

I O

N

S

increase

Chelating ligands form more stable complexes as compared to monodentate ligands.

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METAL CARBONYLS (1)

Metal carbonyl show exceptionally high stability than the expected stability by their normal bonding This can be explained by concept of back bonding.

(2)

Metal have 2 types of orbital (a)

Vacant valence orbitals

(b)

Some partilly d-orbitals.

In metal carbonyls metals forms a σ bond with ligand where ligand donates its e– pair into vacant

(3)

orbitals of metals. The excess e– deposited on metal by the ligand is relieved by back bonding where metal donates back from its filled orbitals into vacant anti-bonding orbitals of CO. This extra back-bonding is called

(4)

Synergic back-bonding which imparts extra bond strength and metal carbonyls turns more stable

π molecular orbital on CO molecule Vacant * Ni-Co π bond

– 2

O:

Ni +

+

:C

O:

+

:C

Ni

–

–

–

+

+

–

2

Filleed dx -y orbital on Ni Atom

ORGANOMETALS Organometal compounds are divided into two class : (1)

These are sigma complexes (σ-organometals)

(2)

π-complexes. In σ– complexes metal and ligands are bonded by direct σ-bond whereas π- complexes has ligands with molecular π-electron which are donated into metal vacant orbital. In π complexes ligands molecular donate π-electrons of their molecular orbital which are denoted η3, η4 ..... etc depending on number of atom which constitute the π molecular orbital.

η2,

by

Note : (a)

Ferrocene and dibenzne chromium are examples of sandwitch compounds where the ligands ring

covers the metal ion from both top & bottom.

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A T

(b)

Zieses salt is used as antitumour drug in which ethene ligand act as π-ligand.

(c)

K4[Fe(CN)6] is not called an organometal although there is direct sigma bond between metal and carbon.

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σ-Complexes

Examples of

π-Complexes

(i)

CH3MgX

(i)

Ziesel salt K[Pt(η2CH2=CH2)]

(ii)

R2Zn

(ii)

Dibenzene cromium [Cr (C6H6)2]

(iii)

Et4Pt (TEL)

(ii)

Ferrocene [Fe(η5 – C5H5)2]

(iv)

R2CuLi

(v)

R4Sn

π -Complexes

(a)

Zeise's salt

(b)

(c)

[Cr(η6–(C6H6)2]

[Fe(n5 – C5H5)2]

Note : (a)

Ferrocene is a type of metallocene it can exist in both eclipsed and staggered form, in solid state it is in straggered form and in gaseous form it is in eclipsed from.

(b)

Ferrocene can be prepared by rxn of ferrous sulphate by cyclopentadienyl magnesium bromide.

Applications of Organometallics : Organometallics are of immense importance. (i) Homogeneous catalysis : A large number of reactions in solutions are catalysed by organometallic compounds or intermediates derived from transition metal complexes. For example, Wilkinson's catalyst, (Ph3P)3RhCl. (ii) Heterogeneous catalysis : Organometallic compounds can also be used as heterogeneous catalysts, e.g., trialkyl aluminium. (iii) Organic synthesis : They ar widely used in synthesis of various types of organic compounds. e.g., organolithium and organomagnesium compounds are commonly used for the synthesis of different types of organic compounds. (iv) Medicine : A number of organo arsenic compounds are used as main remedy for syphilis. (v) Agriculture : To prevent the infection of young plants, the seeds are treated with organometallics such as ethyl mercury chloride.

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APPLICATIONS OF COORDINATION COMPOUNDS : The complexes are of immense importance on account of their applications in various fields. During complex formation there are drastic changes in the properties of metal atom or ion and these changes in properties are made use of in the applications of metal complexes. 1.

Analytical chemistry : Coordination compounds find their applications in both separation and detection of cation and anion. (A) Separation : (i)

The separation of Ag+ from Hg22+ in the first group of analysis is based on the fact that white silver chloride is soluble in aqueous ammonia and Hg2Cl2 forms a black insoluble material. AgCl + 2NH4OH → [Ag(NH3)2]Cl + 2H2O Soluble

(NH2 )Cl + Hg + HCl + H O Hg2Cl2 + NH4OH → Hg " """ ! "" 2 Black inso lub le

(ii)

The separation of II B group sulphides from IIA group sulphides is based on the fact that sulphides of IIB group form complex sulphides with yellow ammonium sulphide which are soluble while sulphides of IIA group do not react. Sb2S5 + 3(NH4)2S → 2(NH4)3 [SbS4] As2S5 + 3(NH4)2S → 2(NH4)3 [AsS4] SnS2 + (NH4)2S → (NH4)2 [SnS3]

(B) Detection : Detection of cation / anion is based on the formation of coloured complexes. S.No.

Cation / Anion to

Complex formed

Colour of complex

be detected

E

D U

C

A T

1.

Cu2+

[Cu(NH3)4]SO4

Deep blue solution

2.

Cu2+

Cu2[Fe(CN)6]

Chocolate ppt.

3.

Fe 3+

[Fe(SCN)]2+

Blood red colour

4.

Fe 3+

KFeIII[FeII(CN)6]

Prussian blue

5.

Fe 2+

KFeII[FeIII(CN)6]

Turnbull's blue

6.

Co2+

K3[Co(NO2)6]

Yellow ppt.

7.

Co2+

Na3[Co(CO3)3]

Apple green

8.

Cd2+

[Cd(NH3)4](NO3)2

Colourless

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CO-ORDINATION CHEMISTRY S.No.

Cation / Anion to

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Complex formed

Colour of complex

be detected

9.

OH | CH3 − C = N | CH3 − C = N

Ni2+

O

A

Ni

B

O

N = C − CH3 | N = C − CH3 | OH

Scarlet red complex

10.

NH4+

(NH4)2 [HgI4]

Brown complex.

11.

CH3COO–

(CH3COO)3Fe

Blood red

[Fe(H2O)5NO]SO4

Brown ring

(acetate) 12.

Nitrate NO3–

(Ring Test) 13.

Sulphide S2–

Na4[Fe(CN)5NOS]

Violet

14.

Nitrite NO2–

K3[Co(NO2)6]

Yellow

(C) Many ligands (organic reagents) are used for the gravimetric estimation of a number of metal ions. Metal ion to be estimated

Organic reagent used

Cu2+

Benzoin oxime

Ni2+

Dimethyl glyoxime

Fe 3+

1,10-phenanthroline

Al3+

8-hydroxyquinoline

Co2+

α-nitroso β-naphthol

(D) EDTA is used as a complexing agent in volumetric analysis of metal ions like Ca2+, Mg2+ and Zn2+. (E) The coordination compounds of the transition metals exhibit a variety of colours. The property is utilised in colorimetric analysis for the estimation of many metals. (F) Other analytical applications of coordination compounds are oxidation-reudction indicators, estimation of hardness in water, sequestering reagents and solvent extraction. The complex 'Ferroin' [FeII(Phen)3]2+ (Bright red colour) is a useful redox indicator.

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Metallurgical operations : Silver and gold are extracted by the use of complex formation. Silver ore is treated with sodium cyanide solution with continuous passing of air through solution. Silver dissolves as a cyanide complex and pure silver is precipitated by the addition of scrap zinc. Ag2S + 4NaCN

2Na[Ag(CN)2] + Na2S

4Na2S + 5O2 + 2H2O → 2Na2SO4 + 4NaOH + 2S 2Na[Ag(CN)2] + Zn → Na2[Zn(CN)4] + 2Ag Native Ag and Au also dissolve in NaCN solution in presence of oxygen of the air. 4Ag + 8NaCN + O2 + 2H2O → 4Na[Ag(CN)2] + 4NaOH Ag and Au are precipitated by addition of scrap zinc. Nickel is extracted by converting it into a volatile complex, nickel carbonyl by use of carbon monoxide (Mond's process). The complex decomposes on heating again into nickel and carbon monoxide. Heating Ni + 4CO → Ni(CO)4   → Ni + 4CO

(Associated with other metals) 3.

Photography : In photography, the image on the negative if fixed by dissolving all the remaining silver halides with hypo solution in the form of a soluble complex. AgBr + 2Na2S2O3 → Na3[Ag(S2O3)2] + NaBr (Soluble)

(Soluble)

4.

Electroplating : Metal complexes release metal slowly and thus give a uniform coating in electroplating of the metal on the desired object. Cyano complexes of silver, gold, copper, etc., are used for the electrodeposition of these metals.

5.

Biological processes : (i)

Haemoglobin : Haemoglobin is a protein which is present in blood. The quadridentate macrocyclic ligand, the porphine molecule is an important part of haemoglobin structure. The two H+ ions bonded to nitrogen atoms are displaced and the metal ion coordiante simultaneously with all four nitrogen atoms. The complex formed from porphine is called porphyrin. The Fecomplex is called heme.

The iron-heme complex is present in another class of proteins, called cytochromes.

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The water ligand can be replaced eadily by molecular oxygen to form red coloured oxyhaemoglobin.

Structure of oxy-haemoglobin (ii)

Chlorophyll : The chlorophyll molecule, which plays an essential role in photosynthesis also contains the porphyrin ring but the metal ion there is Mg2+ rather than Fe2+.

(iii) Vitamin B12 : It is a complex of cobalt with a quadridentate ligand which is similar to porphyrin ligand of haemoglobin. 6.

Plant growth : Iron in +3 state present in the soil is mostly hydrolysed to form insoluble iron hydroxide Fe(OH)3, which cannot be taken up by plants. To overcome iron deficiency, the complex Fe(III)-EDTA is added to the soil. This complex is soluble in water and readily enters the roots of trees and reach to various parts of the plants where it is converted into useful compounds.

7.

In medicinal field : (i) The complex of calcium with EDTA is used for the treatment of lead poisoning. Lead readily replaces calcium in the complex and lead-EDTA complex is finally eliminated from the body in urine. (ii) The platinum complex cis [Pt(NH3)2Cl2] known as cis-platin is used as an antitumor agent in the treatment of cancer.

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