Resonance booklet of Coordination Compounds

CHEMISTRY LECTURE NOTES COURSE : VIJAY (P) (LECTURE No. 1 TO 9)

TOPIC : COORDINATION COMPOUNDS (2012-13)

RESONANCE

COORDINATION COMPOUND - 1

LECTURE # 1 1.

GENERAL INTRODUCTION OF COMPLEX SALTS :

Salt : OR OR

The compound which contains acidic(anionic) or basic(cationic) radicals is known as salt. The compound which is formed by the neutralisation reaction of acid and base is known as salt. The compound which on hydrolysis gives acid and base is known as salt.

The most appropriate classification in context of present chapter will be as follows: (1)

Simple salt : The salt which contains one type of simple +ve ion and one type of simple –ve ion. e. g., NaCl, KCl, Na2CO3 etc.

(2)

Mixed salt : The salt which contains more than one type of +ve ions or more than one type of -ve ions or both the ions are more than one type. e.g., microcosmic salt NaNH4HPO4.4H2O, bleaching powder Ca(OCl)Cl.

(3)

Double salt : When two or more than two simple salts aqueous solutions are mixed in appropriate proportion and then are evaporated, the residue which will be left is known as double salt. when we prepare an aqueous solution of these double salts these will get ionised into all constituent ions. Hence the constituent ions do not loose there identity in double salts.(Generally in double salts one ion is common, either +ve ion or -ve ion in the two or more than two simple salts being mixed). Examples are potash alum - K2SO4 Al2(SO4)3 24H2O, mohr salt - (NH4)2SO4 FeSO4 6H2O and carnallite -KCl MgCl2 6H2O. K2SO4 + Al2(SO4)3 + 24H2O  K2SO4 . Al2(SO4)3 . 24H2O

2K+ (aq.) + 2Al+3 (aq.) + 4SO42– (aq.)

(4)

Complex salt/coordination compounds : A complex or a metal complex (sometimes even some electropositve non-metals like Se & Br also form complexes) is defined as a species in which metal atom or ion is bonded by donating some electrons pairs via coordinate bonds and these ions are found to retain their identity in aqueos solutions, hence the original constituent ions loose thier identity and forms a new group which acts as single unit in aqueos solutions (Do remember that this unit/group can be ungrouped or change its identity via chemical reactions). And the complex salts do atleast contain one complex ion. Do not that a complex ion or part is always written in square brackets. Fe(CN)2 + 4KCN  Fe(CN)2 . 4KCN or K4 [Fe(CN)6] (aq.) 4K+ (aq.) + [Fe(CN)6] (aq.) 4:1 comlex Other examples are, [Co(NH3)6]Cl3 (aq) [Co(NH3)6]3+(aq) + 3Cl– (aq) 1: 3 complex [Cu(NH3)4]SO4 (aq.) K2[Zn(CN)4] (aq.)

[Cu(NH3)4]2+ (aq.) + SO42– (aq.) 1 : 1 complex 2K+ (aq.) + [Zn(CN)4]2– (aq.) 2 : 1 complex

Other examples can be, K3[Fe(CN)6], [Cr(H2O)6]Cl3, [Pt(py)4][PtCl4] Coordination compounds are also acid-base adduct and are frequently called complexes or, if charged then complex ion.

(Also note the 1 : 4, 1 : 1, 2 : 1 or 3 : 1 terminology, often used in problems) 2.

DIFFERENT TERMS/DEFINITIONS TO BE USED Some important terms required for the description of coordination compounds are:

1.

COORDINATION ENTITY(COMPLEX) OR COORDINATION SPHERE : A coordination entity constitutes a central atom/ion, usually of a metal, to which are attached a fixed number of other atoms or groups each of which is called a ligand. It may be neutral or charged. In aqueous solutions this complex retain its identity. Examples being : [Co(NH3)6]3+, [PtCl4]2–, [Fe(CN)6]3–, [NiCl2(OH2)4]. The remaining ions apart from complex ions are called Counter Ions or Free Ions or Ionisation Sphere or

RESONANCE


Outer Sphere. Ex. Sol.

Designate the coordination entities and counter ions in the coordination compounds : [Cr(NH3)6]Cl3 ; K4[Fe(CN)6], K2[PtCl4] ; [Ni(CO)4]; K2[Ni(CN)4]. The respectively coordination entities are : [Cr(NH3)6]2– ; [Fe(CN)6]4– ; [PtCl4]2– ; [Ni(CO)4]; and the counter ions are Cl– , K+, K+, (no counter ion) and K+, respectively in the given coordination compounds.

2.

CENTRAL ATOM/ION : In a coordination entity – the atom/ion to which are bound a fixed number of ligands in a definite geometrical arrangement around it, is called the central atom or ion. For example, the central atom/ion in the coordination entities : [NiCl2(OH2)4], [CoCl(NH3)5]2+, [Fe(CN)6]3– are Ni2+, Co3+ and Fe3+, respectively.

3.

LIGANDS : The ligands are the ions or molecules bound to the central atom/ion in the coordination entity. This is better visualized as the combination of a Lewis acid (the central atom/ion) with a number of Lewis bases (ligands). The atom of Lewis base that forms the bond to the Lewis acid (central atom/ion) is called donor atom (because it donates the pair of electrons required for bond formation). The central atom/ ion is the acceptor atom/ion (because it receives the electron pairs from the ligands). Some of the common ligands in coordination compounds are : Br–, Cl–, CN–, OH–, O2–, CO32–, NO2–, C2O42–, NH3, CO, H2O, NH2CH2CH2NH2 (1,2-ethanediamine or enthane-1, 2 diamine).

CLASSIFICATION OF LIGANDS (ON BASIS OF DENTICITY) : DENTICITY AND CHELATION : When coordination of more than one sigma-electron pair donor group from the ligand to the same central atom/ion takes place, it is called chelation, and the ligand a chelating ligand. The number of such sigma bonds by such a ligand indicate the denticity of the ligand. For example : unidentate, didentate, terdentate or tridentate, tetradentate, etc. CHELATE LIGAND : Chelate ligand is a di or polydentate ligand which uses its two or more donor atoms to bind a single metal ion producing a ring. The complex formed is referred to as a chelate complex and the process of chelate formation is called chelation. Chelate rings may have any number of atoms; the most common contain five or six atoms, including the metal ion. Smaller rings have angles and distances that lead to strain; larger rings frequently result in crowding both within the ring and between adjoining ligands. Some ligands can form more than one ring; ethylene diaminetetracetate (EDTA) can form five by using the for carboxylate groups and the two amine nitrogens.

The chelate complexes are more stable than similar complexes containing unidentate ligands. The greater stability of the chelate complex in comparison to normal complex is called chelate effect. For example, Ni2+ (aq) + 6 NH3 (aq) [Ni(NH3)6]2+ (aq) Kformation = 108 2+ 2+ Ni (aq) + 3 NH2CH2CH2NH2(aq) [Ni(en)3] (aq) Kformation = 1018 The five and six membered rings are more stable. The following are type and examples if ligands on the basis of this classification. (a) Monodentate (one donor atom) :  : , N  H , CN– , NO – , OH–, CO, NO, C H N (py) etc. F–, Cl–, Bl–, H O 2

3

2

5

5

(b) Didentate (two donor atoms : Some most common bidentate ligands are

OR

NH2 – CH2 – CH2 – NH2 Ethylenediamine(en)

RESONANCE


NH2  CH  CH2  NH2 (1, 2 propanediamine) (pn) | CH3

O

O C

C C2O42–

OR

O O Oxalate ion (ox2-) O CO32–

OR

O

C || O Carbonate ion

2, 2 bipyridyl (bipy)

OR

H2N – CH2COO–

glycinate ion (gly)

OR acetylacetonate ion (acac) ion

CH3 – C  N – O – | CH3 – C  N – OH

OR

two nitrogen atoms as donor atoms

dimethyl glyoximate ion (dmg)

DIDENTATE CHELATION : In [PtCl2(en)], en represents the didentate ligand, NH2CH2CH2NH2 (1,2-ethanediamine or ethylenediamine) [Figure] H2 N

Cl Pt Cl

RESONANCE

CH2

CH2 N H2 [PtCl2(en)] (a)


(c) TRIDENTATE/TERDENTATE LIGAND/CHELATION :  H – CH – CH – N  – CH – CH – N  H H2N 2 2 2 2 2 Diethylenetriamine (dien)

OR

[N-(2-aminoethyl)-1, 2-ethanediamine]

TRIDENTATE/TERDENTATE : In the coordination entity (Pt Cl(dien)]+, dien, [N-(2-aminoethyl)-1, 2-ethanediamine] is a terdentate ligand [Figure] CH2

CH2

H2N Pt

NH

CH2

N H2

NH

Cl

+

[PtCl(dien)] (b)

(d) TETRADENTATE LIGAND/CHELATION :  H – CH – CH –  – CH – CH – N  – CH – CH – N  H H2N NH 2 2 2 2 2 2 2 Triethylenetetraamine (trien) OR [N, N+-bis-(2 aminoethyl)-1, 2-ethanediamine] In [Pt(trien)2+], trien, [N, N+-bis-(2 aminoethyl)-1, 2-ethanediamine represents a tetradentate ligand. H2C – CH2 H N

H2C

H N

CH2

Pt

H2C

N H2

CH2

N H2 [Pt (trien)] (c)

2+

(e) PENTADENTATE LIGAND/CHELATION

 H – CH – CH – N  H  H – CH – CH – N  H – CH – CH – N  – CH – CH – N H2N 2 2 2 2 2 2 2 2 2 tetraethylenepentaamine Ethelene diamine triacetate ion. O

3–

O

– O–C–CH – O–C–CH

2

N–(CH2 )2–N

CH2–C–O – H

2

or

O (Ethelenediaminetriacetato ion)

(f) HEXADENTATE LIGAND/CHELATION : EDTA is the most common hexadentate ligand and it is derived from ethylene diamine tetraacetic acid

This acid is directly not soluble in water but its sodium salt is soluble. O C

CH2

CH2 C :N

C

CH2

CH2

N:

O:

O CH2 C

: :

: :

:O

CH2

: :

: :

:O

4–

O

O:

or

O

O EDTA4–, Ethylenediaminetetraacetate ion

Therefore, acid is not the ligand but its tetra anion is a ligand and its usual denticity is 6 but it is sometimes also found to behave as tetradentate ligand and hence is a flexidentate ligand.

RESONANCE


(g) FLEXIDENTATE LIGAND : A polydentate ligand which is found to have different denticity in different coordination compounds is called a flexidentate ligand. Note that in a particular complex denticity of a particular ligand is fixed, it can not be flexible in the same compound. EDTA can act as hexa, penta as well as tetra dentate ligand. For example ; EDTA usually acts as hexadentate ligand but in [Cr(III)(OH)(EDTA)]2– and [Co(III)Br(EDTA)]2– as pentadentate and in [Pd(II)H2(EDTA)]0 as a tetradentate ligand. Sulphate ion, SO42– can also be mono or bi dentate ligand. For example ;

(h) AMBIDENTATE LIGAND : Ligands which can ligate through two different atoms present in it are called ambidentate ligands. Examples of such ligands are the NO2– and SCN¯ ions. NO2– ion can coordinate through either the nitrogen or the oxygen atoms to a central metal atom/ion. Similarly, SCN¯ ion can coordinate through the sulphur or nitrogen atom. Such possibilities give rise to linkage isomerism in coordination compounds. For examples : CN–

:

CNO–

:

– bonding may be through through C or N – CN , – NC – bonding may be through N or O.

nitrito-N

M  O—N=O M  SCN M  NCS

4.

nitrito-O thiocyanato or thiocyanato-S isothiocyanato or thiocyanato-N

OXIDATION NUMBER OF CENTRAL ATOM : The oxidation number of the central atom is defined as the charge it would carry if all the ligands are removed along with the electron pairs that are shared with the central atom. Oxidation number is represented by a Roman numeral in parenthesis following the name of the coordination entity. Some examplex are listed Table below.

Examples of important terms used in describing coordination entities : –––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––– Coordination entity (complex) Ligand list Central atom/oxidation number –––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––– [Co(NH3)6]3+ 6NH3 Co/ (III) [NiCl4]2– 4Cl– Ni / (II) [Co(CN)5F–]3– 5CN– + 1F– Co / (III) [Ni(CN)4]2– 4CN¯ Ni / (II) [Ni(CO)4 4CO Ni / (O) [Ni(H2O)6]2+ 6H2O Ni / (II) ––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––

RESONANCE


LECTURE # 2 3.

NOMENCLATURE OF COORDINATION COMPOUNDS : We first are writing down the rules to write down the names of two impotant parts of a complex, first of ligands and then that of central atom. Naming of Ligands: The names of ligands are written according to the charge (+ , – or nuetral) on them. Naming of Anionic ligands: the name of these ligands ends with –o. The name of anions can be of following three types 1. If name of the anion ends with –ide then generally –e is replaced with –o 2. If the name of anion ends with –ate then –e is replaced with –o 3. If the name of anion ends with –ite then –e is replaced with –o H– hydrido O2– oxido OH– hydroxido – X halido(chlorido, flourido, bromido, iodido) O22– peroxido O2– superoxido O3– ozonido N3– nitrido N3– azido NH2– amido NH2– imido CN– NC– [NCO]– [SCN]– NO2–

CH3COO– NO3– S2O32– SO4– SO3– HSO3– HO2– (NOS)– S2– HS– H2N – CH2COO–

cyanido or cyanido–C (second is preferred) isocyanido or cyanido–N cyanato –O or –N thiocyanato–N or –S nitro(is common when donation is through nitrogen) –NO2 nitrito–N –ONO nitrito–O acetato nitrato thiosulphato sulphato sulphito hydrogensulphito hydrogenperoxido thionitrito thio or sulphido hydrogensulphido(mercapto) RSH  mercatanes (greater affinity to mercury) glycinato (gly) acetylacetonato (acac)

CH3 – C = N – O– | CH3 – C = N – OH

dimethylglyoximato (dmg)

Naming of Neutral ligands : Generally names of neutral ligands are not modified except in some cases some special names are used for some common neutral ligands. NH3 ammine H2O aqua or aquo S8 octasulphur P4 tetraphosphorus CH3 NH2 methylamine

RESONANCE


C5 H5 N NH2 – CH2 – CH2 – NH2 O2 CO NO

pyridine (py) ethylenediamine (en) or ethane -1, 2-diamine dioxygen carbonyl nitrosyl

o - phenanthroline (o-phen) N..

..N

urea

2, 2 bipyridyl (bipy) C2H4 C6H6

ethylene or ethene benzene

NOTE : IN LAST TWO CASES THE PI ELECTRIN DENSITY IS GENERALLY DONATED Positive/Cationic ligands : The name of these ligands ends with –ium NO+ N2H5+ NO2+

: : :

nitrosonium or nitrosylium (old) hydrazinium nitronium

Naming of central atom/ion The name of central atom is not modified if the central atom is in the cationic part of the complex while its name is modified and it ends with –ate if it is present in the anionic part of the complex. Metal Name in Cationic part Name in Anionic part Cr chromium chromate Pt platinium platinate Co cobalt cobaltate Ni nickel nickelate Zn zinc zincate Pd palladium palladate Ti titanium titanate Va vanadium vanadate Mo molybdenum molybdate Hg mercury mercurate W tungstun tungstate Si silicon silicate Br bromine bromate B boron borate Os osmium osmate Cu copper cuprate Note : For some metals, the Latin names are used in the complex anions. Pb lead plumbate Ag silver argentate Au gold aurate Sn tin stannate Fe iron ferrate Note : The name of central atom is not modified if the central atom is in the neutral part of the complex.

RESONANCE


Steps for Writing the Name of a Complex :

(i) Calculate oxidation number of the central atom involved. (ii) The cation is named first as is the case with other ionic compounds. This rule applies to both positively and negatively charged coordination entities. For example in K4[Fe(CN)6] and [Co(NH3)6]Cl3, K+ and [Co(NH3)6]3+ cations are named first. Potassium hexacynidoferrate(II) Hexaaminecobalt(III) chloride (iii) Usually the number of free cation or anions is not given as can easily be calculated by charge balancing. K4 [Fe(CN)6] [Co(NH3)6] Cl3 Potassium but not tetrapotassium Chloride but not trichloride (iv) For writing the names of complex cation use the following rule. Names of ligand in english alphabetical order with their qunatity as prefix (but do not use mono for one and do not consider first letters of these number denoting prefixes in alphabetisation). This is followed by the name of metal followed by its oxidation number (in Roman numeral) in a brackette like (I), (II), (III), (O), (–I), (–II), etc. [Co(NH3)4Cl2]+, pentaamminechloridocobalt(III). (v) For writing the names of complex anion use the following rule. Names of ligand in english alphabetical order with their qunatity as prefix (but do not use mono for one and do not consider first letters of these number denoting prefixes in alphabetisation). This is followed by the name of metal(modified as per rule already specified) followed by its oxidation number (in Roman numeral) in a brackette. (I), (II), (III), (O), (–I), (–II), etc. (NH4)2 [Co(SCN)4], ammonium tetrathiocyanato-S-cobaltate(II). Note : When the names of the ligands include a numerical prefix or are complicated or whenever the use of normal prefixes creates some confusion, it is set off in parentheses and the second set of prefixes is used. 2 di bis 3 tri tris 4 tetra tetrakis 5 penta pentakis 6 hexa hexakis 7 hepta heptakis [CoCl2(NH2CH2CH2NH2)2]+, dichloridobis(ethane-1,2-diamine)cobalt(III). [NiCl2(PPh3)2], dichloridobis(triphenylphosphine)nickel(II). (vi) The neutral complex molecule is named similar to that of the complex cation. Examples ;

[Pt(NH3)BrCl(CH3NH2)], amminebromidochloridomethylamineplatinum(II). EXAMPLES ILLUSTRATING THE USE OF RULES OF NOMENCLATURE FOR COORDINATION COMPOUNDS: Formula of Coordination Compound K3[Fe(CN)6]

Name of Coordination Compound Potassium hexacyanidoferrate(III)

[Co(NH3)5Cl]Cl2

Pentaamminechloridocobalt(III) chloride

[Pt(NH3)2Cl(NH2CH3)]Cl

Diamminechlorido(methylamine)platinum(II) chloride

K2[PdCl4]

Potassium tetrachloridopalladate(II)

[Co(NH3)4(H2O)2]Cl3

Tetraamminediaquacobalt(III) chloride

[Pt[NH3)Cl2(C5H5N)]

Amminedichlorido(pyridine)platinum(II)

[Ni(NH3)4(H2O)2]SO4

Tetraamminediaquanickel(II) sulphate

K2[Ni(CN)4]

Potassium tetracyanidonickelate(II)

[Co(NH3)6]Cl(SO4)

Hexaamminecobalt(III) chloride sulpahte

Fe4[Fe(CN)6]3

Iron(III) hexacyanidoferrate(II)

RESONANCE


Na[Pt(NH3)BrCl(NO2)]

Sodium amminebromidochloridonitrito–N–platinate(II)

K3[Al(C2O4)3]

Potassium trioxalatoaluminate(III) or Potassium tris(oxalato)aluminate(III)

[Co(NH3)6][Co(NH3)2(NO2)4]3

Hexamminecobalt(III) diamminetetranitrito-N-cobaltate(III)

[Fe(NCCH3)6]Br2

Hexakis(methylcyanide)iron(II) Bromide

[CuCl2 {O = C (NH2)2}2]

Dichlorobis(urea)copper(III)

RULES FOR WRITING THE FORMULAE OF MONONUCLEAR COORDINATION COMPOUNDS : Mononuclear coordination entities contain a single central metal atom. The sequence of symbols within the formula of coordination entity is governed by the following rules : (i) The central atom is listed first ; (ii) The ligands are then placed in alphabetical order. The placement of a ligand in the list does not depend on its charge. (iii) Polydentate ligands are also placed alphabetically. In case of abbreviated ligand, the first letter of the abbreviation is used to determine the position of the ligand in the alphabetical order. When ligands are polyatomic, their formulas are enclosed in parentheses. Ligands abbreviations are also enclosed in parentheses. (iv) The formula of the coordination entity is enclosed in square brackets. (v) No space is kept between representations of ionic species within the formula ; (vi) When the formula of a charged coordination entity is written without the formula of the counter ion, the charge is indicated outside the square brackets as a right superscript with the number before the sign. (vii) The charge of the cation(s) is balanced by the charge of the anion(s). For example : [Co(CN)6]3–, [Cr(H2O)6]3+ etc. The following examples illustrate the above rules : [Co(NH3)6]Cl3 ; [Co(NH3)5Cl]Cl2 ; K2[PtCl4] [Co(NH3)4Cl(NO2)]Cl ; K3[Fe(CN)6] ; Na[AllII(H2O)2(OH)4] Ex.

Sol.

Write the formulae for the following coordination compounds : (i) tetrahydroxidozincate (II) (ii) pentaaquachloridochromium (III) chloride (iii) tetrabromidocuparate (II) (iv) pentacarbonyliron (O) (v) potassium tetracyanidocuprate (II) (i) Zn(OH)4]2– (ii) [Cr(OH2)5Cl]Cl2 (iii) [CuBr4]2– (iv) [Fe(CO)5]

(v) K2[Cu(CN)4]

Naming of Some Special Complexes and of Complexes with Bridging ligands : The name of a bridge complex is prefixed by – If the situation on both sides of the bridge is symmetrical then we can write the name of remainig complex at one place like (i)

 3 1  3   (NH3 )5 Cr  OH – Cr (NH3 )5   

5

Cl5

–hydroxidobis(pentaamminechromium(III)) chloride OR we could also have written the name of one side of the bridge then the name of bridging complex and then the name of the other side of the bridge, like  3 1  3   (NH3 )5 Cr  OH – Cr (NH3 )5   

5

Cl5

pentaamminechromium(III)––hydroxidopentaamminechromium(III) chloride But if the compound is unsymmetrical on both sides of the bridge then we have to follow the second rule , i.e. write the name of one side then that of the bridge and then that of the second side, like

Tetraaquacobalt(III)––amido––hydroxidotetramminechromium(III) sulphate Some other examples are

RESONANCE


(ii) –amido––hydroxidobis(tetramminecobalt (III)) ion

(iii)

(SO4)2 di––hydroxidobis(tetraaquairon(III)) sulphate

(iv)

(SO4)2 –amido––hydroxidobis(bis(ethylenediammine)cobalt(III)) sulphate

OR

bis(ethylenediammine)cobalt(III)––amido––hydroxidobis(ethylenediammine)cobalt(III) sulphate.

Naming of special compounds :

Fe

The name of the above compund is bis(cyclopentadienyl)iron(II). This is commonly known as Ferrocene. In the above compound the aromatic pi electron density of the ring is being donated to the Fe2+ ion, hence in this way iron is bonded to all the five carbons of the upper flat ring and also to all five carbons of below flat ring. This is also called a sandwich compound because of its shape. Its not a true complex this is an organometallic compound which we study in detail in the last part of the chapter. To write down the formula of such compound a special symbol  is written before the name of such ligand and on superscript to this  the number of atoms is written between which the electron density is being shared. So the formula of this compound will be Fe[  5–C5H5]2 . The other such example is bis(benzene)chromium which will be represented by the formula Cr[  6–C6H6]2 – H

H C

Cl

C

Ti+4

Pt H Cl

H Cl

The last compound is commonly known as Zeisse’s salt. The name of the compound will be trichloridoethyleneplatinate(II) ion and the formula will be written as [PtCl3(  2–C2H4)]

RESONANCE


LECTURE # 3 (i) (a)

BONDING IN COORDINATION COMPOUNDS : Initial bonding theories and EAN rule WERNER’S COORDINATION THEORY The systematic study of coordination compounds was started by Alfred Werner whose pioneering work opened an entirely new field of investigation in inorganic chemistry. He prepared and characterized a large number of coordination compounds and studied their physical, chemical and isomeric behaviour by simple experimental techniques. On the basis of these studies. Werner, in 1898, propounded his theory of coordination compounds.

The main postulates of Werner’s theory are : 

Metals exert two types of linkages/valencies ; (i) the primary or ionizable links/valencies which are satisfied by negative ions and equal the oxidation state of the metal and (ii) the Secondary or nonionizable links/ valencies which can be satisfied by neutral or negative ions/groups(or sometimes by cationic species discovered later). The secondary linkages equal the coordination number of central metal atom/ion. This number is fixed for a metal.



The ions/groups bound by the secondary linkages have characteristic spatial arrangements corresponding to different co-ordination numbers. In the modern terminology, such spatial arrangements are called coordination polyhedra. The various possibilities are CN = 2 Linear CN = 3 Triangular CN = 4 Tetrahedral or sq. planar



CN = 5

square pyramidal or TBP

CN = 6

Octahedral

The secondary valencies are generally represented by solid lines while the primary valencies are represented by dashed lines and the ions which satisfy both primary and secondary valencies will be drawn with both solid and dashed lines. For example the complex [CoCl (H2O)5] Cl2 is represented as

On the basis of the above postulates, Werner formulated the coordination compounds, CoCl3 . 6NH3, CoCl3 . 5NH3 and CoCl3 . 5NH3 and CoCl3 . 4NH3 as : [Co(NH3)6]Cl3, [CoCl(NH3)5]Cl2 and [CoCl2(NH3)4]Cl. respectively - the species within the square brackets being the coordination entitles (complexes) and the ions outside the square brackets the counter ions. He further postulated that octahedral, square, planar and tetrahedral geometrical shapes are more common in coordination compounds of transition metals. Thus, [Co(NH3)6]3+, [CoCl(NH3)5]2+, [CoCl2(NH3)4]+ are octahedral entities, while [Ni(CO)4] and [PtCl4]2– are tetrahedral and square-planar, respectively. Series of coloured compounds obtained by the interaction of aqueous CoCl3 and NH3 Compound Colour Name according to colour CoCl3 . 6NH3 Yellow Luteo Complex CoCl3 . 5NH3 Purple Purpureo Complex CoCl3 . 4NH3 Green Praseo Complex CoCl3 . 4NH3 Violet Violeo Complex

RESONANCE


Werner used some experimental methods to find out the number of free ions and coordination number in many different complexes. Such methods involved freezing point dpression measurement or boiling point elevation measurement of aqueous solution of these complexes( we will study later about these methods). Other methods included conductivity measurement of the aqueous solutions of these complexes. Assume here that the conductivity of a solution only depends on the number of ions present per unit volume or on concentration of the ions( which actually is not the case, we will study later that it also depends on the nature of the on), try and arrange the solution of the following complexes in increasing order of conductivity. S.No.

Werner complex

Modern notation

Ionisation

1

CoCl3.6NH3

[Co(NH3)6]Cl3

[Co(NH3)6] +3Cl

2

CoCl3.5NH3

[Co(NH3)5Cl]Cl2

[Co(NH3)5Cl] +2Cl

3

CoCl3.4NH3

[Co(NH3)4Cl2]Cl

4

CoCl3.3NH3

[Co(NH3)3Cl3]

3+

Secondary valency satisfied by

Primary valency satisfied by

six (NH3)

three (Cl )

–

–

–

2+

–

five (NH3) and one (Cl )

[Co(NH3)4Cl2] + Cl

+

–

four (NH3) and two (Cl )

[Co(NH3)3Cl3]

three (NH3) and three (Cl )

–

–

–

–

three (Cl ) including one (Cl ) with dual nature – three (Cl ) including – two (Cl ) with dual nature –

three (Cl ) all with dual nature

From the above table it is clear that (i) the solution conductivity of complexes 1,2 and 3 corresponds to 1:3, 1:2 and 1:1 electrolyte respectively and thus the increasing order of the conductivity can be represented as CoCl33NH3 < CoCl3. 4NH3 < CoCl3.5NH3 < CoCl3.6NH3 and (ii) the complexes 1,2 and 3 will react with silver nitrate and give 3, 2 and 1 mole of the white precipitate of silver chloride. Ex. Sol.

Ex.

Ans. Sol.

For the compound represented by the formula CoCl3.5H2O write down the different complexes possible and repesent them as per Werners notation and also write down their names. CN = 6 octahedral [CoCl(H2O)5]Cl2; [CoCl2(H2O)4]Cl.H2O; [CoCl3(H2O)3].2H2O; generally water of crystallization is found with complexes which have bogth cationic and anionic parts, it is less common with species which are neutral.

Match the pairs of complexes listed in column-I with the method(s) used for their differentiation listed in column-II. Column – I Column–II (A) [Cr(H2O)6]Cl3 and Cr(H2O)5Cl] Cl2.H2O

(p) Can be differentiated by amount, nature or colour of precipitate formed.

(B) [Co(NH3)5Br]SO4 and [Co(NH3)5SO4]Br

(q) Can be differentiated by electrical conduction measurement method (appreciable difference)

(C) [Co(NH3)5Cl]Cl2 and [Co(NH3)6]Cl3

(r) Can be differentiated using cryoscopic measurement method.

(D) [Cu(H2O)4]SO4.H2O and [Cu(H2O)6](NO3)2

(s) Can be differentiated by heating with concentrated H2SO4

(A) – p,q,r,s ; (B) – p,q ; (C) – p,q,r ; (D) – p,r,s (A) [Cr(H2O)6]Cl3 (aq) [Cr(H2O)6]3+ (aq) + 3Cl– (aq) [Cr(H2O)5Cl] Cl2.H2O (aq) (B)

(C)

[Co(NH3)5Br]SO4 (aq)

[Co(NH3)5Br]2+ (aq) + SO42– (aq)

[Co(NH3)5SO4]Br (aq)

[Co(NH3)5SO4]+ (aq) + Br– (aq)

[Co(NH3)5Cl]Cl2 (aq) [Co(NH3)6]Cl3 (aq)

(D)

[Cr(H2O)5Cl]2+ (aq) + 2Cl– (aq)

[Cu(H2O)4]SO4.H2O (aq) [Cu(H2O)6](NO3)2 (aq)

RESONANCE

[Co(NH3)5Cl]2+ (aq) + 2Cl– (aq) [Co(NH3)6]3+ (aq) + 3Cl– (aq) [Cu(H2O)4]2+ (aq) + SO42– (aq) [Cu(H2O)6]2+ (aq) + 2NO3– (aq) COORDINATION COMPOUND - 13

Note : (1) (2) (3) (4) (b)

Ag+ + Cl–  AgCl  (white)

; Ag+ + Br–  AgBr  (pale yellow)

Ba2+ + SO42–  BaSO4  (white) Conductivity depends on ; (a) the charge on cation and anion and (b) the number of ions in aqueous solution. Cryoscopic measurement method depends on the number of ions in aqueous solution. H2SO4 acts as dehydrating agent and thus removes the water molecules present in the ionisation sphere.

SIDGWICK’S EFFECTIVE ATOMIC NUMBER RULE (EAN) : On the basis of some limited number of complexes available at that time, Sidgwick observed some stable complexes and formulated the following rule, which is found to be incorrect when many more omplexes have been prepared.( So the rule is not scientific, but sometimes questions are directly asked on calculation of EAN, so we should be knowing it.) EAN Rule: The complexes in which the EAN of the central atom equals the atomic number of the next noble gas, are found to be extra stable. EAN(Effective Atomic Number) = No. of e– present on the metal atom/ion + no. of e–s donated by ligands to it. OR

EAN = (Atomic no. of C. M. – O. S. of CM) + [No. of electron donated by ligands] For fast calcultions remember the atomic number of the noble gases which are He 2 Ne 10 Ar 18 Kr 36 for 3d series metals Xe 54 for 4d series metals Rn 86 for 5d series metals Ex.(a) [Co3+(NH3)6]Cl3 Number of electrons in Co3+ = 27 – 3 = 24 6NH3 = 6 × 2 = 12 So, EAN = 24 + 12 = 36 (krypton) (b)

[Fe2+(CN)6]4– Number of electrons in Fe2+ = 26 – 2 = 24 6CN– = 6 × 2 = 12 So, EAN = 24 + 12 = 36 (krypton)

(c)

[Fe3+(CN)6]3– Number of electrons in Fe3+ = 26 – 3 = 23 6CN– = 6 × 2 = 12 So, EAN = 23 + 12 = 35

NOTE : (a) The EAN rule is generally found to be not valid in case of most of the complexes but in case of metal carbonyls( which is an important class of comlexes, we will be studying later) this rule is found to be valid in all cases except one or two exceptions, so do remember it for metal carbonyls and do know how to calculate the EAN for any metal. (b)

(i) (ii) (iii)

NO ligand is found to act as three electron donar, as indicated by the following reactions in which when the carbonyl compound is heated in atmosphere of excess of NO. [Fe(CO)5] + 2NO  [Fe(CO)2(NO)2] + 3CO [Cr(CO)6] + 4 NO  [Cr(NO)4] + 6 CO But also remember that NO+ is only two electron donar The Werner theory and some other initial bonding theories cannot answer basic questions like. Why only certain elements possess the remarkable property of forming coordination compounds ? Why the bonds in coordination compounds have directional properties ? Why coordination compounds have characteristic magnetic and optical properties ?

RESONANCE


LECTURE # 4 VALENCE BOND THEORY : The valence bond theory, VBT, was extended to coordination compounds by Linus Pauling in 1931. The basic principles involved in this treatment are : (a)

Metal will provide vacant orbitals depending on its CN. These vacant orbitals will be hybridised. Thus hybrid vacant orbital of metal will overlap with the filled orbital of the ligand to form a coordinate bond.

(b)

There is also found some kind of back bonding from metals to the ligands to maintain the charge neutrality. The extent of this back bonding also depends on the nature of the ligand. We will study about this a bit in metal carbonyls and all other complexes generally neglect this.

 bond  It is a partial bond between filled orbitals of metal & vacant of ligand. The ligands which allow back bonding to sufficient extent are also called  acid acceptors. For example CO, CN–, NO etc. (c)

The type of hybridisation of metal & shape of complex involved can be decided conveniently if some chracteristics of the complex like magnetic nature, geometry or whether exhibits isomerism or not etc. be known. Hence along with the coordination number some other property must also be given to analyse the complex using VBT. sp3, dsp2, dsp3 and d2sp3 yield tetrahedral, square planar, trigonal-bipyramidal or square-pyramidal and octahedral spatial arrangements commonly encountered in coordination compounds. It is found that [Ni(CN)4]2–, is diamagnetic in nature. But we know that Ni2+ is paramagnetic with two unpaired electrons, hence to be diamagnetic the electrons must have got paired up and one d orbital will be available for the dsp2 hybridisation. For one d-orbital lobe available for hybridisation, pairing of electrons takes place in the remaining d-orbitals, as shown below in the case of [Ni(CN)4]2–, 3d 4s 4p Ni       Ni2+      [Ni(CN)4]2–         CN¯ CN¯ CN¯ CN¯ dsp2 hybrids used [Ni(CN)4]2– is square planar and diamagnetic. [NiCl4]2–, on the other hand, is paramagnetic and has tetrahedral geometry. In this case, the VB treatment assumes that (i) the d-orbital occupancy remains the same as in the free Ni2+ ion and (ii) the metal uses sp3 hybrids (involving 4s and 4p orbitals) for bonding with the ligands as shown below : [NiCl4]2– is paramagnetic as there are two unpaired electrons. 3d 4s 4p Ni       Ni2+      [Ni(CN)4]2–          Cl¯ Cl¯ Cl¯ Cl¯ 3 sp hybrids used The VB theory explains the formation of 6– coordinate, octahdral coordination entities by invoking the use of (n-1)d2nsnp3 or nsnp3nd2 hybrid orbitals by the central metal ion in forming bonds with the ligands. Whether the hybridisation is dsp3 or sp3d, it can be decided if number of unpaired e– or magnetic moment is known to us else not. An example involving use of (n-1)d2nsnp3 hybrids is the coordination entity, [Fe(CN)6]4–, which is founf to be diamagnetic The double occupancy of (n-1)d orbitals orbitals confers extra stability and the absence of unpaired electrons renders this entity diamagnetic.

RESONANCE


ns np            L L L L L L (d2sp3 hybrid orbitals filled by electron pairs donated by the respectively ligands. L = CN¯).

Fe Fe2+ [Fe(CN)6]4–

(n-1)d   

  

  

Ex. Sol.

Predict the number of unpaired electrons in the square planar [Pt(CN)4]2– ion. Pt2+ is a 5d8 ion. For square-planar geometry, dsp2 hybrids are required. For the availability of one d orbital, pairing of electrons takes place in the remaining d-orbitals. Hence there are no unpaired electrons in [Pt(CN)4]2– ion.

Ex. Sol.

For the comlex [Ma2 b2], It is given that it does not show geometrical isomerism. Predict the hybridisation. CN = 4 Hence either square planar( dsp2 ) or tetrahedral( sp3 ) but since no geometrical isomerism, hence must be tetrahedral( sp3 ).

Some More Examples With Different Coordination number (CN) : CN = 2 The coordination two is relatively rare, occurring mainly with the +1 cations of Cu, Ag, Au and with Hg2+. The coordination geometry is linear. For example Ex.

[Ag(CN)2]–

Ag()  4d10 5s0

Ag – C bond in [Ag(CN)2]– is obtained by overlapping of sp hybridised orbitals of Ag() and sp hybridised orbitals of C of CN– Such complexes are typically unstable towards the addition of further ligands as [Cu(CN)2]– + 2CN–  [Cu(CN)4]3– CN = 3 The most important geometries for the complexes with CN = 3 are trigonal planar and the trigonal pyramidal. Examples are the planar HgI3– and [Cu(CN)3]2– and the pyramidal SnCl3–. Ex.

[Hg2+3–3]–1

Hg()  5d10 6s0 p0

triiodomercurrate (II) ion

C.N. = 4 This is one of the most important coordination number and results in tetrahedral and square planar geometry. Ex.

Ni (CO)4 (liquid), is found to be diamagnetic, explain the shape and hybridisation.

Sol.

Ex.

[NiCl4]2–, is found to be paramagnetic with two unpaired electrons, explain the shape and the hybridisation.

RESONANCE


Sol.

Tetrahedral

Cl– are weak ligand( will study later in CFT) So, they cannot force 3d8 electrons to pair up in the 4 orbitals. So, it is paramagentic complex. Ex. Sol.

[Ni2+(CN)4]2– , the complex is found to be diamagnetic, explain the shape and hybridisation of the complex. 3d8 4s0 p0 d0

square planar

diamagnetic

net = 0

CN– are strong field ligands (will study in CFT) and therefore they will force the 3d8 electrons to pair up in 4 orbitals. CN = 5 Coordination number five is less common than coordination number four or six, the possible geometries are trigonal bipyramidal and the square pyramidal. In most of the complexes these geometries are also found to interchange into each other. For both these geometries sp3d and dsp3 are two possible hybridisations. Ex.

[Fe(CO)5], this complex is founf to have its total dipole moment to be equal to zero. Explain the hybridisation.

Sol.

Fe  3d6 4s2 p0 d0

Trigonal bipyramidal and diamagnetic

net = 0

Ex.

[Ni(CN)5]3– is found to be dimagnetic with 2 types of Ni – C bond lengths out of which four bond lengths are found to be equal and the fifth one is different. Explain the hybridisation, shape and thus the observed properties.

Sol.

CN = 5

RESONANCE


CN = 6 This is the most common coordination number and the geometry associated is octahedron. For hybridisation there are two possibilities sp3d2 or d2sp3

Outer d complexes Inner d complexes (Octahedral) (Octahedral) High spin = no. of unpaired e–  generally Outer d complexes Low spin = no. of unpaired e–  generally Inner d complexes Ex. Sol.

[Co (NH3)6]3+, this complex is found to be diamagnetic, explain shape and hybridisation. Co3+  3d6 4s0 p0 d0 Therefore, pairing of electrons in 3d sub shell of Co(III) is forced for the complex to be diamagnetic.

Spin = 0 Ex. Sol.

Low spin  NA

[Fe(CN)6]4– complex is found to be diamagnetic while the [Fe(CN)6]3– complex is found to be paramagnetic with one unpaired electron. Explain the shape and hybridisations of both. [Fe(CN)6]4– Fe2+  3d6

diamagentic and low spin complex

[Fe3+(CN)6]3– Fe3+  3d5 Ex. Sol.

paramagentic (I) low spin complex

[FeF6]3– complex is found to be highly paramagnetic with five unpaired electrons, explain the shape and hybridisation. Fe3+  3d5

paramagentic (5) high spin complex/outer -d- complex While the VB theory, to a large extent, explains the formation, structures and magnetic behaviour of coordination compounds, it suffers from the following shortcomings : 

A number of assumptions are involved.



There is no quantitative interpretation of magnetic data.



It has nothing to say about the spectral(color) properties of coordination compounds.



It does not give a quantitative interpretation of the thermodynamic or kinetic stabilities of coordination compounds.



It does not make exact predictions regarding the tetrahedral and square-planar structures of 4-coordinate complexes.



It does not distinguish between strong and weak ligands. The drawbacks of VBT of coordination compounds are, to a considerable extent, removed by the Crystal Field Theory.

RESONANCE


LECTURE # 5 (III)

CRYSTAL FIELD THEORY : The Crystal Field Theory (CFT) was originally proposed for explaining the optical properties of crystalline solids.It was applied to the study of coordination compounds in the 1950s. CFT assumes the ligands to be point charges and the interaction between them and the electrons of the central metal to be electrostatic in nature. The five d-orbitals in an isolated gaseous metal atom/ion have same energy, i.e., they are degenerate. This degeneracy is maintained if a spherically symmetrical field of negative charges surrounds the metal atom/ion. However, when this negative field is due to ligands (either anions or the negative ends of dipolar molecules like NH3 and H2O) in a complex, it becomes asymmetrical and the degeneracy of the d-orbitals is lifted. It results in splitting of the d-orbital energies. The pattern of splitting depends upon the nature of the crystal field. We will first consider :

(A)

CRYSTAL FIELD EFFECTS IN OCTAHEDRAL COORDINATION ENTITIES :

dxy dyz dzx planar d orbitals

d

,d

2 y 2 z2 x  

axial d orbitals

or t2g

or

eg

For convenience, let us assume that the six ligands are positioned symmetrically along the cartesian axes, with the metal atom at the origin. As the ligands approach, first there is an increase in the energy of d orbitals to that of the free ion just as would be the case in a spherical filed. Next, the orbitals lying along the axes (dz2 and dx 2  y 2 ) get repelled more strongly than dxy, dyz and dxz orbitals, which have lobes directed between the axes. The d xy , d yz , dxz orbitals are lowered in enrgy relative to the average energy in the spherical crystal filed. Thus, the degenerate set of d orbitals get split into two sets : the lower energy orbitals set, t2g and the higher enrgy, eg set. The energy seperation is denoted by 0 (the subscript o is for octahedral (Fig.)

RESONANCE


Let us now learn about the significance od 0 by considering first the d1 coordination entity, [Ti(H2O)6]3+, formed in aqueous solutions of Ti3+ (d1) ion. Obviously, the single d electron occupies one of the lower energy t2g orbitals. In d2 and d3 coordiantion entities, the d electrons occupy the t2g orbitals singly in accordance with the Hund’s rule. For d4 ions, two possible patterns of electron distribution arise : i) the fourth electron may enter an eg orbital (of higher energy) or ii) it may pair an electron in the t2g level. The actual configuraion adopted is decided by the relative values of 0 and P. P represents the enrgy required for electron pairing in a single orbital. If is less than P (0 < P), we have the so called weak field, high spin situaiton, and the fourth electron 3 1 entres one of the eg orbitals giving the configuraiton t 2g eg . If now a fifth electron is added to a weak field 3

2

coordination entity, the configuraiton becoms t 2g eg . When 0 > P, we have the strong field, low spin situation, and pairing will occur in the t2g level with the eg level remaining unoccuped in tentities of d1 to d6 ions. Calculations show that coordination entities with four to seven d electrons are more stable for strong field as compared to weak field cases. This extra stabilization due to ligand field in comparison to normal situation( if there would have been no spilitting or if there would have been spherically symmetrical ligand field) is called CFSE(Crystal field stabilization energy). This CFSE can be calculated as follows

General formula : CFSE = [– 0.4 (n) t2g + 0.6 (n) eg] o + nP. where n & n are number of electron(s) in t2g & eg orbitals respectively and o crystal field splitting energy for octahedral complex. n represents the number of extra electron pairs formed because of the ligands in comparison to normal degenerate configuration. Metal ion Configuration

Example

Configuration in Ligand field

CFSE

d0

Sc 3+

t2g0,0,0, eg0

0

d1

Ti3+

t2g1,0,0, eg0

–0.4 0

2

3+

d

3

d

8

V

3+

2+

Cr , V 2+

t2g

1,1,0

t2g

1,1,1 2,2,2

, eg

0

, eg

0

eg

–0.8 0 –1.2 0

1,1

–2.4 0 + 1.2 0 = –1.2 0

d

Ni

t2g

d9

Cu2+

t2g2,2,2 eg2,1

–2.4 0 + 1.8 0 = –0.6 0

d10

Zn2+

t2g2,2,2 eg2,2

–2.4 0 + 2.4 0 = 0

Therefore, for the above configurations, there is no effect of the nature of ligand. They may be strong or weak, the formula for CFSE will remain the same. For configuration from d4 to d7 : Metal ion Configuration 4

d

Example 2+

Cr S.L. W.L.

` 5

d

2+

Mn S.L.

Conf. in Ligand field

t2g

2,1,1

t2g

1,1,1

t2g

2,2,1

0,0

, eg

1,0

, eg

0,0

, eg

CFSE

–1.6 0 + P –0.6 0 –2.0 0 + 2P

3+

d6 7

d

Fe W.L.

t2g1,1,1, eg1,1

Co3+,Fe2+ S.L.

t2g2,2,2, eg0,0

–2.4 0 + 2P

W.L.

t2g2,1,1, eg1,1

– 0.4 0

2+

Co S.L. W.L.

t2g

2,2,2

t2g

2,2,1

1,0

, eg

1,1

, eg

0

–1.8 0 + P – 0.8 0

Note : S.L. = Strong field ligands W.L. = Weak field ligands Ex.

For the complex[Cr(H2O)6]2+ , P = 23500 cm–1 and its hybridisation and net CFSE in kJ/mol.

Sol.

Cr2+  3d4  t2g1, 1, 1 eg1

P(4)

P > 0 weak ligand

sp3d2

0 = 12600 cm–1, then for this complex calculate

outerd complex.

CFSE = (– 3 × 0.4 + 1 × 0.6) 0 = – 0.6 0 net stabilisation energy = CFSE + P (wherever there is pairing)

RESONANCE

= CFSE + 0

= – 0.6 0


SPECTROCHEMICAL SERIES On the basis of the behaviour (the splitting produced with different metal ions) of ligands these are arranged in a series according to their strentgh which is called Spectrochemical Series. weak Ligand

I–<Br– < S2– < SCN– < Cl– < NO3 < N3 < F– < OH– < C2O24 < H2O < NCS– < CH3CN < Py < NH3 < en < bpy < NO 2 < PPh3 < CN– < CO Strong Ligand Do note one thing in the series that in all strong field ligands the donor atom is either Nitrogen atom or Carbon atom while in all weak field ligands the donar atoms are either Oxygen atoms or Sulphur or halogen atoms.( Remember) The intermediate ligands can in some cases act as strong field ligands or can also act as weak field ligands depending on the nature of the metal involved or some other factors. Generally for kexamples in our course we can consider ligands upto water as weak and beyond water as strong. But do remember the following exception. Ex.

[Co(C2O4)3] Co3-  3d6

t2g2, 1, 1 eg 1,1

If C2O24 is considered W.L.

sp3d2 and paramagnetic with four unpaired electrons. high spin /outer orbital complex.

If C2O24 is considered S.L

d2sp3 and diamagnetic

t2g2,2,2 eg 0,0

low spin/inner orbital complex.

C2O24 is a S.L.

Experimentally, it is diamagnetic Ex. Sol.

What is electronic configuration of Fe2+ in complex [Fe(CN)6]4– and in complex [FeF6]3– [Fe(CN)6]4– Fe+2  3d6 t2g2,2,2 eg0,0 d2sp3 [Fe+3 F6]3–

Fe+3  3d5

t2g1,1,1

eg1,1

sp3d2

Ex.

In the complex [Fe(Cl)(CN)4(O2)]4– , the electronic configuration of metal ion in this ligand field is found to be t2g2,2,2, eg0,0. Then find out whether the complex is paramagnetic or diamagnetic and whether the O–O bond lentgh in the complex is more than that in dioxygen O2. Also write the IUPAC name of the complex.

Sol.

3d6

Fe2+

O21 superoxido

 Bond order = 1.5 Paramagnetic (1) B.L. more than O2 Chlorotetracyano superoxidoferrate (II) ion.  Diamagnetic

RESONANCE


LECTURE # 6 (B) CRYSTAL FIELD EFFECTS IN TETRAHEDRAL COORDINATION ENTITIES : Tetahedral complex (sp3) In a tetrahedral field

ˆ e  55º < M g

and

ˆ t  35º < M 2g

Consider a cube such that a metal atom or ion is situated at its centre of symmetry through which the axis of geometry are passing and joining the face centres of this cube. Therefore, lobes of eg orbitals will be directed towards the face centres but those of t2g orbitals will be pointing towards edge centres. Now consider 4 monodentate ligands approaching the metal, the 4 alternate corners of this cube so as to make a tetrahedron. ˆ e  55º ˆ t  35º  110º  (i) < M ˆ L = 109º28’ (ii) < M (iii) < M g 2g Thus it is clear that t2g orbitals are nearer to the ligands than the eg orbitals. Hence t2g orbitals will experience more repulsion than eg orbitals. Therefore, crystal field splitting will be reversed of octahedral field which can be shown as below.

In tetrahedral complexes none of the ligand is directly facing any orbital so the splitting is found to be small in comparison to octahedral complexes. For the same metal, the same ligands and metal-ligand distances, it can be shown that t = (4/9)o. This may attributes to the following two reasons. (i) there are only four ligands instead of six, so the ligand field is only two thirds the size ; as the ligand field spliting is also the two thirds the size and (ii) the direction of the orbitals does not concide with the direction of the ligands. This reduces the crystal field spliting by roughly further two third. So t =

2 2 4 × = o. 3 3 9

Consequently, the orbital splitting energies are not sufficiently large for forcing pairing and, therefore, low spin configurations are rarely observed. Ex. Sol.

 2  Ni Cl4   

2

calculate the CFSE of this tetrahedral complex.

Ni2+ 3d8 4s0p0 Weak filed ligand Configuration of Ni(II) in tetrahedral field P = eg2, 2, t2g 2,1,1 t P(2) For square planar geometry one d-orbital must be vacant. So hybridisation  sp3 3  2  t  + 4  t  CFSE = 4 ×   5  5 

= – 2.4 t + 1.6 t = – 0.8 t.

In tetrahedral complexe, all ligands are considered to be weaker ligands and, therefore, electrons filling takes place continuously from equal to t2g. [MnO4]–  permanate ion

d3s / sp3 tetrahedral

RESONANCE


FACTORS FAVOURING TETRAHEDRAL COMPLEXES : Tetrahedral complexes are favoured by steric requirements, either simple electrostatic repulsion of charge ligands or vander wall's repulsions of large one. A valence bond (VB) point of view ascribed tetrahedral structure to sp3 hybridisation. Tetrahedral complexes are thus generally favoured by large ligands like Cl–, Br–, – and PPh3 and metal ions of six types ; (1) those with a noble gas configuration such as Be2+ (nsº) ; (2) those with pseudo noble gas configuration (n–1)d10 ns0np0, such as Zn2+, Cu+ and Ga3+, and (3) those transition metal ions which do not strongly favour other structure by virtue of the CFSE, such as Co2+, d7. (4) those transtion metal which have lower oxidation state. (5) those metals generally with electronic configuration dº, d5 and d10 perfer to form such complexes. (6) It is abserved that

Sc Ti V Cr Mn Fe usually form tetrahedral complexes with oxo ligands eg. : [CrO4]2– , [MnO4]2– [FeO4]2–

Co Ni Cu Zn form tetrahedral complexes with hallide ions eg. : [NiCl4]2– , [CuCl4]2–

Examples : 2

2

2

2

3d10 , [ Z n(NH3 )4 ] ; 3d10 , [ Zn(CN)4 ] CFSE = O (diamagnetics) [Ni(CO)4] diamagnetics [NiCl4]2– paramagnetics with two unpaired electrons. [NiCl2 (pph3)2] paramagnetics with two unpaired electrons. [MnCl4]2– [FeCl4]2– [Cu(py)4]+ Cs2 [CuCl4] (NH4)2 [CuCl4]

paramagnetics with five unpaired electrons. paramagnetics with four unpaired electrons. diamagnetics. orange tetrahedral paramagnetics with one unpaired electrons. yellow square planaer paramagnetics with one unpaired electrons.

Cu (II)  3d9 4s0p0

sp3

3d8 4s0 p1 or 5s1 dsp2

unpaired electron is present in p or s orbitals. Ist complex is more paramagnetic than IInd. (C) CRYSTAL FIELD EFFECTS IN SQUARE PLANAR COMPLEXES : The square planar arrangement of ligands may be considered to be one derived from the octahedral field by removing two trans-ligands located along the Z-axis. In the process, the eg and t2g sets of orbitals is lifted i.e., these orbitals will no longer be degenerate. The four ligands in square planar arrangement around the central metal ion are shown in Fig. As the ligands approach through the axes, they would have greatest influence on d 2 2 orbital, so the energy of this x y orbital, will be raised most. The dxy orbital, lying in the same plane, but between the ligands will also have a greater energy though the effect will be less than that on the d 2 2 orbitals. On the other hand, due to x y absence of ligands along Z-axis, the d 2 orbital becomes stable and has energy lower than that of dxy orbital. z Similarly dyz and dxz become more stable. The energy level diagram may be represented as Fig. along with tetrahedral and octahedral fields.

RESONANCE


The value of sp has been found larger than o because of the reason that dxz and dyz orbitals interact with only two ligands in the square planar complexes, while in octahedral complexes the interaction takes place only with four ligands. sp has been found equal to 1.3o. 1 > 2 > 3 ; sometimes it can be 1 > 3 > 2 But, and

sp = (1 + 2 + 3) > o sp = 1.3 o.

[Ni(CN)4]2– Ni(II) 3d8 Electron distrubution in case of square planar complex :

d2yz, d2zx, d2z2, d2xy , d0 x 2  y 2

FACTORS FAVOURING SQUARE PLANAR (i) Metals (atom/ion) with configuration 4d8 or 58 always form square planar complexes irrespective of nature of the liqand. For Example : 2

[Pt Cl4 ] 2 * , although Cl– are W.L. yet is is square planar complex 5d8 [Pt(CN)4]2– [Pt(NH3)4]2+

square planar square planar

Such metal atom or ions are Rh+1 r+1 Pt+2 Ag3+ Pd2+ Au3+

Z = 45 Z = 77 Z = 78 Z = 47 Z = 46 Z = 79

Rh(I) Pd(II) Ag(III) 45 46 47

r(I) Pt(II) 77 78

Au(III) 79

(ii) But with the metal atom or the ion with 3d8 configuration, for example Ni(II)) complex will be square planar only with the strong field ligands. (tetrahedral with weak ligand).

RESONANCE


Examples : (i) [Ni(CN)4]2– (ii) [Ni(en)2]2+ (iii) [Ni(dmg)2]0 (iv) [PtCl4 ]2– (v) [Pt(NH3)4]2+ (vi) [AgF4]– (vii) [RhCl(PPh3)2 ]0 (viii) [Ni(CR)2 (PMe3)2] Note : D = diamagnetic ;

D D D D D D D D

Wilkinson catalyst. (hydrogenation of Alkenes)

PMe3 = trimethyl phosphine.

Important factors which determine the magnitude of  the orbital splitting energy, are : (i) Oxidation state of the metal ion : Generally, the higher the ionic charge on the central metal ion, the greater the value of . (ii) Nature of the metal ion : For analogous entities within a group, the values differ. The general trend being 3d < 4d < 5d. Thus while going from Cr to Mo or Co to Rh, the 0 value increases by ~ 50%. As a consequence of this, coordination entities have a greater tendency to be low spin as compared to the first transition series. (iii) Geometry of the coordination entity : The t value is only ~ 50% as large as that of 0. sp > 0 > t approximately, for the same set of ligands sp  1.3 0

4  9 0 (iv) Nature of the ligand : The ligands can be arranged in the order of increasing field strength. With strong field ligands, increases With weak field ligand, decreases. t 

RESONANCE


LECTURE # 7 5.

SOME PROPERTIES OF COORDINATION COMPOUNDS :

(a)

COLOUR IN COORDINATION COMPOUNDS : Coordination compounds of transition metals have fascinating colours. According to the crystal field theory the colour is due to the d-d transition of electron under the influence of ligands. We know that the colour of a substance is due to the absorption of light at a specific wavelength in the visible part of the electromagnetic spectrum (400 to 700 nm) and transmission or reflection of the rest of the wavelengths. An object that absorbs all visible light appears black. The mechanism of light absorption in coordination compounds is that photons of appropriate energy can excite the coordination entity from its ground state to an excited state. Consider the Ti(III) ion in solution, that is [Ti(H2O)6]3+. This is a violet colour octahedral complex, where in the ground state of the complex a single electron is present in t2g level. The next higher state available for the transition is the empty eg level. If the light corresponding to the energy of yellow-green is absorbed by the complex, it would excite the electron from t2g level to eg level. Consequently the complex appears violet in colour. In case of copper (II) ions in solution, for example, it can be imagined that one of the d-electrons from the t2g set (dxy, dyz, dxz) gets excited to the eg set ( d x 2  y 2 ,d z 2 orbitals). In this case since high energy is transmitted it means that low energy light (red region) is absorbed. For copper (II) ions in aqueous solution, the energy gap 0 is relatively small. Table below gives the relationship of the wavelength absorbed and the colour observed. By using spectroscopic data for a number of coordiantion compounds having the same metal ion but different ligands, the crystal field splitting for each ligand has been calculated and the spectrochmeical series has been prepared.

Relationship between the wavelength of light absorbed and the colour observed In some coordination entitles ––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––– Coordination entity Wavelength of light Colour of light Colour of coordination absorbed (nm) absorbed entity –––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––– [CoCl(HN3)5]2+ 535 Yellow Violet [Co(NH3)5(H2O)]3+ 500 Blue Green Red [Co(NH3)6]3+ 475 Blue Yellow Orange [Co(CN)6]3310 Ultraviolet Pale Yellow 2+ [Cu(H2O)4] 600 Red Blue [Ti(H2O6)3+ 498 Blue Green Purple ––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––– You need not remember the above table, you will be provided with it whenever the question is asked and also you need not remember the color of the complexes except some common complexes which will be used in qualitative analysis and hence in detection of different ions. Note : (a) In absence of ligand, crystal field splitting does not occur and as a consequence the substance appears colourless. For example (i) removal of water from violet coloured complex [Ti(H2O)6]Cl3 on heating makes it colourless. (ii) Similarly anhydrous copper sulphate (CuSO4) is white, but hydrated copper sulphate (CuSO4.5H2O) is blue coloured. (b) The nature of the ligand and the molar ratio of metal : ligand also influence the colour of the complex for example in the pale green complex of [Ni(H2O)6], the colour change is absorbed when ethylenediamine is progressively added to it. Molar ratio of en : Ni

Coloured observed

1 :1

Pale blue

2 :1

Blue/Purple

3 :1

Violet

Note : Ruby is Al2O3 in which 0.5–1% Cr3+ ions (d3 electron system) are randomly distributed in the positions normally occupied by Al3+. We may consider Cr(III) species as octahedral Cr(III) complexes incorporated into

RESONANCE


the alumina lattice ; d-d transition of electron at these centres give rise to the colour (red). Emerland is the mineral beryl (Be3Al2Si6O18) in which Cr3+ ions occupy octahedral sites, but in this case low energy corresponding to yellow red and blue is absorbed and light corresponding to green region is transmitted. (b)

MAGNETIC PROPERTIES OF COORDINATION COMPOUNDS : Additional information for understanding the nature of coordination entities Is provided by magnetic susceptibility measurements. We have noted that coordination compounds enerally have partially fllled d orbitals and as such they are expected to show characteristic magnetic properties depending upon the oxidation state, electron configuration, coordination number of the central metal and the nature of the ligand field. It is experimentally possible to determine the magnetic moments coordination compounds which can be utilized for understanding the structures of these compounds. The number of unpaired electrons in any complex can be easily calculated from the configuration of the metal ion, its coordination number and the nature of the ligands involved( strong or weak from the spectrochemical series) and after that the magnetic moment of the complexes can be easily calculated using Magnetic Moment = n(n  2) Bohr Mangeton Though experimentally it is the magnetic moment calculation which gives us which one is a weak ligand and which one is a strong ligand. For example, [Fe(CN)6]3– has magnetic moment of a single unpaired electron while [FeF6]3– has a paramagnetic moment of five unpaired electrons. [CoF6]3– is paramagnetic with four unpaired electrons while [Co(C2O4)3]3– is dimagnetic. This apparent anomaly is explained by VB theory in terms of formation of inner orbital and outer orbital orbital coordination entities as shown below and also gives us the hint that which one is a weaker ligand and which one is stronger.

(c)

STABILITY OF COORDINATION COMPOUNDS : The stability of a coordination compound [MLn] is measured in terms of the stability constant (equilibrium constant) given by the expression,  n = [MLn]/[M(H2O)n][L]n for the overall reaction : M(H2O)n + nL

MLn + nH2O

By convention, the water displaced is ignored, as its concentration remains essentially constant. The above overall reaction takes place in steps, with a stability (formation) constant, K1, K2, K3, ...... Kn for each step as represented below : M(H2O)n + L

RESONANCE

ML(H2O)n–1 + H2O

K1 = [ML(H2O)n–1] / {[M(H2O)n][L]} COORDINATION COMPOUND - 27

MLn–1 (H2O) + L M(H2O)n + nL

MLn + H2O MLn + nH2O

Kn = [MLn] / {[MLn–1 (H2O)] [L]}  n = K1 x K2 x K3 x ........ x Kn

 n, the stability constant, is related to thermodynamic stability when the system has reached equilibrium. Most of the measurements have been made from aqueous solutions, which implies that the complex is formed by the ligand displacing water from the aqua complex of the metal ion. Ignoring the charge and taking L as un unidentate ligand, the stepwise formation of the complex is represented as shown above. K1, K2, K3 ..... Kn representing the stepwise stability (or formation) constants. The above is thermodynamic stability criteria, there can be another kind of stability called Kinetic Stability, which measures the rate of ligand replacement. The value of stability constant for some of the complexes are given below: Complex

Stability constant

[Cu(NH3)4]2+ [Ag(NH3)2]2+ [Co(NH3)6]2+ [Co(NH3)6]2+ [AgCl2]– [AgBr2]– [Ag(CN)2]– [Cu(CN)4]2– [Fe(CN)6]3–

4.5 × 1011 1.6 × 107 1.12 × 1011 5.0 × 1033 1.11 × 105 1.28 × 107 1.0 × 1022 2.0 × 1027 7.69 × 1043

The above table is only for teachers information, we dont expect the students to remember the constants or the order of stability of the complexes. Important generalizations derived from the vast data on stability constants are : 

For a given metal and ligand the stability is generally greater, the greater the charge on the metal ion. Thus, stability of coordination entities of ions of charge 3+ is greater than the entities of 2+ ions. Further, for the divalent ions of the first row transition elements, irrespective of the ligand involved, the stabilities vary in the Irving-Williams order : MnII < FeII < CoII < NiII < CuII > ZnII This order is according to the size of the ions, smaller the size of the ion or greater the charge density on the metal greater is the stability of the complex.



The stability also depends on the formation of chelate rings. If L is an unidentate ligand and L–L, a didentate ligand and if the donor atoms of L and L–L are the same element, then L–L will replace L. The stabilisation due to chelation is called the chelate effect. It is of great importance in biological systems and analystical chemistry. The chelate effect is maximum for the 5- and 6-membered rings. In general, rings provide greater stability to the complex.



If a multidentate ligand happens to be cyclic and there are no unfavourable steric effects, a further increase in stability occurs. This is termed the macrocyclic effect.

RESONANCE


LECTURE # 9 6. (i)

ISOMERISM IN COORDINATION COMPOUNDS : Structural Isomerism :

(A)

IONISATION ISOMERISM : This type of isomerism occurs when the counter ion in a coordination compound is itself a potential ligand. Thus, the pairs of the following compounds : [Co(NO3)(NH3)5]SO2, and [Co(SO4)(NH3)5](NO3) [PtCl2(NH3)4]Br2, and [PtBr2(NH3)4]Cl2 produce different ions in solution. [Co(en)2(NO2)Cl]SCN, [Co(en)2(NO2)SCN]Cl and [Co(en)2(SCN)Cl]NO2

(i) (ii) (iii)

A special form of ionisation isomerism is the so called Hydrate isomerism. It occurs when water forms a part of the coordination entity or is outside it. For example, CrCl3 . 6H2O exists in three distinct isomeric forms : [Cr(H2O)6]Cl3, violet ; [CrCl(H2O)5]Cl2.H2O, pale green : [CrCl2(H2O)4]Cl.2H2O, dark green. Apart from their distinctive colours, the three isomers can be identified by the addition of excess of aqueous silver nitrate to their aqueous solutions, which precipitates chloride in the molar ratio of 3 : 2 : 1 respectively. Complex Reaction with AgNO3 Reaction with conc. H2SO4(dehydrating agent) [Cr(H2O)6]Cl3 in the molar ratio of 3:1 No water molecule is lost or no reaction [CrCl(H2O)5]Cl2.H2O in the molar ratio of 2:1 one mole of water is lost per mole of complex [CrCl2(H2O)4]Cl.2H2O in the molar ratio of 1:1 two mole of water is lost per mole of complex (B)

LINKAGE ISOMERISM The type of isomerism is theoretically possible in any coordination compound containing an ambidentate ligand like NO2– or SCN–, [Co(NO2)(NH3)5]2+ , for example, has two isomeric forms having red and yellow colours. The red coloured isomer is [Co(ONO)(NH3)5]2+, pentaamminenitrito-O-cobalt(III) cation while the yellow isomer is formulated as [Co(NO2)(NH3)5]2+. It is named as : pentaamminenitrito-N-cobalt(III) cation. In the first cation we have Co-(ONO) link, while in the second case, NO2– coordinates to the metal ion through the nitrogen atom, Co-(NO2).

(C)

COORDINATION ISOMERISM Coordination compounds made up of cationic and anionic coordination entities show this type of isomerism due to the interchange of ligands between the cation and anion entities. Some of the examples are : [Co(NH3)6][Cr(CN)6] and [Cr(NH3)6](Co(CN)6] [Cu(NH3)4](PtCl4] and [Pt(NH3)4][CuCl4] [Co(NH3)6][Cr(SCN)6] and [Cr(NH3)4(SCN)2](Co(NH3)2(SCN)4] [Pt(NH3)4][PtCl6] and [Pt(NH3)4Cl2][PtCl4] Such isomers are expected to have significant differences in their physical and chemical properties.

(i) (ii) (iii) (iv)

(D)

LIGAND ISOMERISM If there is possibility of isomerism in the ligands themselves for example many organic ligands have such possibilities, if complexes shows isomerism from this source then it is said to be ligand isomerism. For example ligands, 1,2-diaminopropane(propylenediamine or pn) and 1,3-diaminopropane(trimethylenediamine or tn) are such pairs. Similarly ortho-, meta- and para-toluidine (CH3C6H4NH2).

(E)

POLYMERISATION ISOMERISM Considered to be a special case of coordination isomerism, in this the various isomers differ in formula weight from one another, so not true isomers in real sense. For example [Co(NH3)4(NO2)2][Co(NH3)2(NO2)4], [Co(NH 3 ) 6 ][Co(NO 2 ) 6 ], [Co(NH 3 ) 5 (NO 2 )][Co(NH 3 ) 2 (NO 2 ) 4 ] 2 , [Co(NH 3 ) 6 ][Co(NH 3 ) 2 (NO 2 ) 4 ] 3 , [Co(NH3)4(NO2)2]3[Co(NO2)6] and [Co(NH3)5(NO2)2]3[Co(NO2)6]2. These all have the empirical formula Co(NH3)3(NO2)3, but they have formula weights that are 2,2,3,4,4 and 5 times this, respectively.

RESONANCE


(ii)

Stereoisomerism : Geometrical Isomerism :

GEOMETRIC ISOMERISM : First we consider the cases of coordination number 4. The tetrahedral compounds can not show geometrical isomerism as we all know that all four positions are equivalent in tetrahedral goemetry. So moving on to square planar complexes there are different possibilities depending upon the formula of the compound. So we are taking the different cases in sequence (a) Square planar complexes with all monodentate ligands In all following examples, we will be representing metal atom by symbol M, monodentate ligands by small letters a, b, c, d, e, f ... etc and the symmetrical bidentate ligands are by AA, BB, CC etc while the umsymmetrical bidentate ligands by AB, CD.. etc. Formula No. of GI’s Trans pairs Ma4 No GI (aa)(aa) Ma3b No GI (aa)(ab) Ma2b2

2

(aa)(bb) (ab)(ab)

b a

a Ma2 b2

b

M a Cis

trans cis

M b b

trans

a

The above compound is cisplatin used in cancer therapy. Ma2bc

2

(aa)(bc) (ab)(ac)

b a

a Ma2bc

Cis Mabcd

c

M a

M c b

3

a trans (ab)(cd) (ac)(bd) (ad)(bc)

b a

a

d aTc

aTb aTc aTd

d

M

Mabcd

trans cis

M c c

aTb

b

b

a M c aTd

RESONANCE

d


(b) Square planar complexes with one bidentate ligand and other two monodentate ligands Complex No.of geometrical isomers Trans pairs M(AA)a2 M(AA)ab M(AB)a2

0 0 0

(Aa)(Aa) (Aa)((Ab) (Aa)(Ba)

M(AB)ab

2

(Aa)(Bb) (Ab)(Ba)

b

A

a

A

M

M

B

b

B

a

(c) Square planar complexes with two bidentate ligands Complex

No.of geometrical isomers

Trans pairs

M(AA)2

0

(AA)(AA)

M(AB)2

2

(AB)(AB) (AA)(BB)

A

A

M(AB)(AB)

A

B M

M B

B

B

A

M(AA)(BB) M(AA)(CD)

0 0

(AB)(AB) (AC)(AD)

M(AB)(CD)

2

(AC)(BD) (AD)(BC) C

A M(AB)(CD)

ATD

RESONANCE

M

M B

D

A

D

B ATC

C


LECTURE # 8 CN = 6 (Octahedral Complexes) Again the different cases possible are (I)

Complexes containing only mondentate ligands Total possible formula can be

i  Ma6 ii  Ma5b

  no.G.I only one form is possible  

(iii) Ma2b4

2

–

(aa)(bb)(bb) (ab)(ab)(bb)

b a

a

b

M

a

a

a

b

a

a

|||

M

a

b

a

a

a

Cis

Trans

2

(aa)(ab)(bb) (ab)(ab)(ab)

b

a

b a

facial or fac2 triangular faces such that each corner is occupied by similar ligands

RESONANCE

b M

M a

(optically inactive) (optically inactive)

b

b a

a c

a

–

a M

M a

c


b

c b

a

a

both a are in trans (aa)(aa)(bc) (aa)(ab)(ac)

a

Ma3b3

b

b

both a are in cis (iv) Ma4bc – 2

(v) Ma3b3

M b

b

b

b

|||

M

b

Ma4bc


b

b

Ma2b4

( both will be optically inactive)

a

a b

Meridional or Mer-


Like [Co(NO2)3(NH3)3]. If each trio of donor atoms occupy adjacent positions at the corners of an octahedral face, we have the facial (fac) isomer. When the positions are around the meridian of the octahedron we, get the meridional (mer) isomer

b a

b b

b

M

M b

a a Fac-

(a)

NO2

H2N

Co

Co NH3

NO2 Fac-

b MerNH3

NH3

O2N

a

a

NH3 O2N

a

O2N

(b)

NO2 NO2 Mer-

(vi) Ma3b2c – 3 The trans pairs for the three isomers are (aa)(ab)(bc) (aa)(bb)(ac) (ab)(ab)(ac)

(optically inactive) (optically inactive) (optically inactive)

(vii) Ma3bcd – 4 The trans pais for the four possible isomers are (aa)(ab)(cd) (aa)(ac)(bd) (aa)(ad)(bc) (ab)(ac)(ad)

(optically inactive) (optically inactive) (optically inactive) (optically active)

(viii) Ma2b2c2 – 5 The trans pairs for the possible isomers are (aa)(bb)(cc) (aa)(bc)(bc) (bb)(ac)(ac) (cc)(ab)(ab) (ab)(ac)(bc)

(optically inactive) (optically inactive) (optically inactive) (optically inactive) (optically active)

RESONANCE


(II)

(ix) Ma2b2cd

–

6

(x) Ma2bcde

–

9

(xi) Mabcdef

–

15

Compounds containing bidentate ligand a a

[M(AA)a4] – 1 form possible (No G.I.)

A M

a

A a

a

b A

A

a

a M

M

M(AA)a3b – 2 A

M(AA)a2b2 – 3

a

bTa

aTa b

a b

A

M A

b

b

a A

A

a

a

A M

M b

A

b

A

a

b

b

aTa

aTb

bTb

M(AA)abcd – 6 M(AA)3 – 1 form (No G.I.) MI(AB)3 – 2 form A A

A

M

Fac. B

B M

Mer. B

B

RESONANCE

A

A

B

B B COORDINATION COMPOUND - 34

Optical isomerism (is in the NCERT text so will be covered) : A Coordination compound which can rotate the plane of polarised light is said to be optically active. When the coordination compounds have same formula but differ in their ability to rotate directions of the plane of polarised light are said to exhibit optical isomerism and the molecules are optical isomers. The optical isomers are pair of molecules which are non-super impossible mirror images of each other. This is due to the absence of elements of symmetry in the complex. (A) Optical isomerism is very common in octahedral complexes. (a) [Ma2B2C2]n± Out of five different geometrical isomers, only one is optically active [Pt(py)2(NH3)2Cl2]2+

(b)

[Mabcdef]

All 15 different geometrical isomers are optically active

[Pt(py)NH3NO2ClBrI] py

py

py

NO2

py

O2 N

Pt Cl I

Pt H3N

NH3 Mirror

d-isomer

(c)

Cl I l-isomer

[M(AA)3]n±

[Co(en)3]3+ en

3+

3+

en

en Co

Co en

en d-form

(d)

[M(AA)2a2]n±

en Mirror

l-form

[Co(en)2Cl2]+

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

RESONANCE


(e)

[M(AA)2ab]n±

(Co(en)2NH3Cl]2+

(f)

[M(AA)a2b2]

[CrCl2(en)(NH3)2]+

H3N H3N

(g)

N en N Cr Cl Cl

+

[M(AB)3]

en N

N NH3

+

Cr NH3

Cl Cl

(Cr(gly)3]

METAL CARBONYLS : A widely studied and important class of organometallic compounds is that of metal carbonyls. Binary compounds of metals with CO are called metal carbonyls. The homoleptic carbonyls (compounds containing carbonyl ligands only) are formed by most of the transition (d metals). The metals constituting the central part of the d block form stable, neutral binary carbonyls like : [V(CO)6], [Cr(CO6], [Mo(CO)6], [W(CO6], [Mn2(CO)10], [Fe(CO)5], [Fe2(CO)9], [Co2(CO)8], [Co4(CO)12], [Ni(CO4)], etc. Outside the central part of d-block, the metal carbonyls are usually unstable. METAL CARBONYLS–STRUCTURE AND BONDING : Homoleptic binary metal carbonyls have simple, welldefined structures. Tetracarbonylnickel(o) is tetrahedral, pentacarbonyliron(o) is trigonal bipyramidal while hexacarbonylchromoum (o) is octahedral. Decacarbonyldimanganese(o) is made up of two square pyramidal Mn(CO)5 units joined by a Mn–Mn bond. Octacarbonyldicobalt(o) has a Co–Co bond bridged

RESONANCE


Structures of some representative homoleptic metal carbonyls. Metal carbonyls are mostly solids at room temperature and pressure. Exceptions being iron and nickel carbonyls which are liquids. The mononuclear carbonyls are vaolatile and toxic. With the exception of Fe2(CO)9, enneacarbonyldiiron(o), carbonyls are soluble in hydrocarbon solvents. Mononuclear carbonyls are either colourless or light-coloured. For example, Fe(CO)5 is light straw-coloured liquid. Polynuclear carbonyls are more deeply coloured. Fe3(CO)12, dodecacarbonyltriiron(o), for example, is a deep grass green solid. The reactivity of metal carbonyls is due to (a) the metal centre of the carbonyl, and (b) the CO ligands. Metal carbonyls find use as industrial catalysts and as precursors in organic synthesis. The bonding in organometallic compounds in general and in carbonyls in particular is best explained by the molecular orbital theory. Carbon monoxide, CO, as a ligand binds itself to metal atoms through the carbon atom. It is a weak donor (a weak base). It forms a weak  bond to the central atom. CO is also an acceptor ligand and froms a  bond to the metal. This characteristic property of back bonding stabilised the metal-ligand interaction. For a better understanding of the nature of bonding in metal carbonyls, we have to consider first the molecular orbital energy level diagram of carbon monoxide given in the figure. 3, the highest occupied molecular orbital in CO is essentially a lobe projecting away from the carbon atom. When CO acts as a lgand. This orbital serves as a weak donor to the metal atom, and forms a bond as represented in Figure. The lowest unoccupied molecular orbltals of CO are the * (2) orbitals. These play an important part in bonding as they can overlap metal d orbltals having  symmetry. The resultant Interaction leads to the delocalisation of electrons from, filled d orbltals on the metal into the empty orbitals on the CO ligands.This in fact, is back bonding from metal to CO. The metal to ligand bonding creates a synergic effect which strengthens the bond between CO and the metal (Figure).

The MO energy level diagram for CO The filled 3 and the vacant 2 oribttols are important for bonding in the formation of metal carbonyls.

C

O









M









RESONANCE


Synergic bonding Example of synergic bonding interactions in a carbonyl complex. (i) As M — C  bonding increases, the C — O bond becomes weaken. The greater the positive charge on the central metal atom, the less readily the metal can donate electron density into the * orbitals of the carbon monoxide ligands to weaken the C — O bond. (ii) In contrast, in the anionic complex (i.e. carbonylate anion) the metal has a greater electron density to be dispersed, with the result that M — C  bonding is enhanced and the C — O bond is diminished in strength. For example in isoelectronic complexes the strength of metal-ligand bond increases and strength of C — O bond in CO decreases (because bond order decreases) as the negative charge on the complexes increases. Thus order of CO bond strengths ; (a) [M(CO)6]+ > [Cr(CO)6] > [V(CO)6]– > [Ti(CO)6]2–. (b) [Ni(CO)4] > [Co(CO)4]– > [Fe(CO)4]2–.

Classification of metal carbonyls

Monomeric (metal atom = 1)

Bridged or binuclear (metal atom = 2)

 Ni(CO) 4 Fe(CO)5 3d  Cr(CO)6   V(CO)6

Mn2CO10

Polynuclear (metal atom > 2)

Fe 3 (CO)12 CO 4 (CO)12

Mo(CO)6 4d  Ru(CO)5

 W (CO)5 5d  Os(CO)6 Note : Besides the metal carbonyls, there are some metal carbonyl anions which generally form salts with alkali metal ions. Ex : Na[V– (CO)6]–1 Sodium hexacarbonyl vanadate (–I) K2[Ni2(CO)6] STRUCTURES OF MONOMERIC CARBONYLS :









In all Carbonyls, the oxidation state of the metal is O and they all are found to be diamagnetic except one, V(CO)6  P(l) Ni(CO)4 Ni(O)  3d84s2  3d10 sp3 tetrahedral D.M. = 0 Fe(CO)5 Fe(O)  3d64s2  3d8 dsp3 TBP D.M. = 0 (dz2) Cr(CO)6 Cr(O)  3d54s1  3d6 4s0p0 d2sp3 D.M. = 0 Octahedral dx2–y2 dz2 V(CO)6 V(O)  3d34s2  3d5 4s0p0

Octahedral

Out of the 4, last one is most unstable

RESONANCE


18e– Rule The relative stabilites of metal carbonyls and their ions can be explained on the basis of a rule which is known 18 e– rule. If no. of valence shell e–s of the metal + no. of e–s donated by bonded CO grps. If this comes to be 18 e– then the carbonyl is said to be stable Ni(CO)4  Ni(O)  3d8 4s2 10 + 8 = 18 Fe(CO)5  8 + 10 = 18 Cr(CO)6  6 + 12 = 18 V(CO)6  5 + 12 = 17. Unstable [V(CO)6]–  17 + 1 = 18. More stable than V(CO)6

Sigma () bonded complexes : In these complexes, the metal atom and carbon atom of the ligand are joined together with a sigma bond in which ligand contributes one electron and is therefore called one electron donor. For example: (a) Grignard’s Reagent R – Mg – X where R is a alkyl or aryl group and X is halogen. (b) (CH3)4 Sn, (C2H5)4 Pb, Al2 (CH3)6, Al2 (C2H5)6 etc. Bridging groups

CH3

CH3

CH3

Al

Al CH3

CH3

CH3

Pie ()-bonded organometallic compounds : These are the compounds of metal with alkenes, alkynes, benzene and other ring compounds. Zeise's salt : In Zeise’s salt structure, the ethylene occupies the fourth coordination site of the square planar complex with the C — C axis perpendicular to the platinum ligand plane. Relative to free ethylene the C — C bond is lengthened (from 133.77 pm to 137.5 pm), and the hydrogen are slightly tilted back from a planar arrangement. The bond between the ethylene molecule and the metal ion may be considered as a dative  bond to an available orbital on the metal. The bonding scheme is analogous to that in carbon monoxide complexes in which there is a ligand metal  donation and a reciprocal metal to ligand  bonding. The extent of back bonding varies depending on the metal, the substituents on ethylene, and the other ligands on the metal, H –

H

K [PtCl3 (2 – C2H4 )]

C || C

Cl Pt Cl

H

K+ H

Cl

Ferrocene and bis(benzene)chromium : Ferrocene obey 18-electron rule. Depending on the electron counting method adopted, the cyclopentadienyl ligand may be viewed as either a five electron donor (neutral atom) or a six electron donor (oxidation state). Similarly, the benzene ligand may be viewed as a six electron donor in neutral atom as well as in the oxidation state. Ferrocene show thermal stability and is not oxidised by air.

Fe (5 – C5H5)2

RESONANCE

and

Cr (6 – C6 H6)2

Cr


Resonance booklet of Coordination Compounds

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