CHE421 COORDINATION CHEMISTRY NOTES - Redeemer's ...

REDEEMER’S UNIVERSITY

COLLEGE OF NATURAL SCIENCES

DEPARTMENT OF CHEMICAL SCIENCES

INORGANIC CHEMISTRY IV

COURSE: CHE421

COORDINATION CHEMISTRY NOTES

2017/18 SESSION

SEMESTER 1

LECTURER

PROFESSOR G A KOLAWOLE

COMPILED BY PROFESSOR GA KOLAWOLE

1. Purpose of the Course

To extend the concept of periodicity to the f-block elements;

To deepen students‟ knowledge of coordination chemistry, introduced from first year;

To introduce students to the new area of Inorganic reaction mechanism

2. Course Outcomes

By the end of the course students should be able to:

• Account for the physical and chemical properties of the lanthanides and actinides in

relation to their electronic configurations highlighting differences and similarities

with the d-block transition metals.

• Identify the reasons for the similarities and differences in the properties of elements in

the two series, and with d-transition metals.

• Account for the prevalence of +3 oxidation state in the lanthanides

• Discuss the basic theories of bonding in Coordination compounds, particularly VBT,

CFT and LFT, bringing out the limitations and strengths of each theory.

• Extract structural information from the physico-chemical analyses of coordination

compounds especially electronic spectra and magnetic susceptibility measurements.

• To be able to extract reaction mechanisms from rate laws in inert octahedral, square

planar and redox reactions.

3. Recommended textbooks:

The under-listed books are some of the books consulted in the preparation of this note and

they are acknowledged.

N N Greenwood and A Earnshaw, Chemistry of the Elements, 2nd

Edition, Elsevier, 2009.

J D Lee, Concise Inorganic Chemistry, 5th Edition, Blackwell, Oxford, 1996.

P Atkins, T Overton, J Rouke, M Weller, F Armstrong, Shriever & Artkins Inorganic

Chemistry, 4th

Edition (or later), Oxford, 2006.

G L Miessler and D A Tarr, Inorganic Chemistry, 3rd

Edition, Pearson, 2004 or later.

4. Course Outline

UNIT Coverage Completion

Week (No.

of lectures)

Unit 1:

Coordination

Chemistry

Synthesis of coordination compounds;

nomenclature;

isomerism;

theories of structure and bonding

Physical methods of structural investigation;

electronic spectra; spectro-chemical series;

Jahn-Teller, tetragonal and trigonal

distortions,

Chelate effect, thermodynamic stability of

complexes,

Magnetochemistry

Weeks 1-5

(15 hours)

ASSESSMENT 1 (60 MINUTES)

Unit 2: The

Chemistry of f-

block elements

Physical and chemical properties related to

electronic structures.

Comparative discussion of the chemistry of

lanthanides and actinides, including

differences and similarities with the d-

transition metals: variable oxidation states,

reactivity, complex formation, electronic and

magnetic properties

Extraction from ores (lanthanides) and

preparation of trans-uranium metals

(actinides), including revision of relevant

areas of radiochemistry required to understand

the preparation of trans-uranium metals..

Electronic properties of lanthanides and

actinides

Magnetic properties of lanthanides and

actinides

6-9

(12 lectures)

MID-SEMESTER ASSESSMENT (90 MINUTES)

Unit 3: Inorganic

reaction

mechanisms

Substitution reactions in inert octahedral

complexes

Substitution reactions in square planar

complexes

Redox reactions: inner sphere and outer

sphere mechanisms

10-12

(9 lectures)

REVISION Week 13

EXAMINATION Week 14

5. Methods of Assessment to be used: (% weighting)

Activity %

Attendance at lectures 05

Other Assessments 20

Mid-semester assessment 15

Semester Examination 60

Total 100

NOTE

80% attendance is mandatory to qualify to write the final semester examinations,

for which a score of five is awarded. Any attendance less than 80% attracts zero

score and disqualification from writing the semester examination.

The other assessments will be in form of Tutorial Quizzes such that a set of

questions will be provided to guide your reading and periodically students are

given quizzes on the questions during one of the lectures to ascertain that each

student has read the note and has worked through the problems.

You are also given reading assignments that would warrant the use of the

library, for which a short typed report is submitted every two weeks.

6. Plagiarism

Plagiarism involves copying printed work or idea developed by other people without

acknowledging the source of your information. This amounts to stealing the intellectual

property of other people and is punishable.

It is easily detected when you copy other peoples‟ work verbatim (word for word) or even

when you reframe it without acknowledging the origin of the information you use in your

work. You need to keep this in mind when you are given an assignment to do that involves

consulting books, scientific journals or even newspapers or even use information contained in

the printed lecture notes given to you. When detected, you can easily lose critical marks due

to you in an assignment or even face disciplinary action.

Like in all human endeavours, intellectual honesty is a virtue and it pays to cultivate the habit

right now and keep it up for the rest of your life.

To acknowledge the source of information, you give a reference number after the statement

and at the end of the work you provide a list of references corresponding to the numbered

references.

References normally should include: names of authors (Initials and surname); title of article;

name of journal, book, newspaper, etc; volume of the journal, book, newspaper, etc; pages of

the article in the source and, if a book, and the publisher, the edition and year of publication.

If the information is from a website, quote the website and the date you download the

information in addition to the above.

Coordination Chemistry

1.0 Preparation of coordination compounds

Preparation of compounds is the trade mark of every chemist. Research in chemical

industries is largely oriented toward the synthesis of new and useful materials. The chemist

is interested in preparing new compounds because it is an invaluable way of expanding our

knowledge of chemistry.

There are many routes, but related experimental methods to preparing metal complexes. The

method chosen depends upon the metal, the oxidation state of the metal, the ligand and the

electron configuration of the ion. Not all methods can therefore be employed to the synthesis

of a particular compound. Having found a suitable method for making the compound in good

yield one needs to find a suitable way to isolate the product from its reaction mixture.

Some of the commonly used techniques to obtain products from reaction mixtures are:

Evaporation of the solvent to concentrate (could be under reduced pressure usin a

rotary evaporator) and the cooling in an ice-bath (or a refrigerator). Adding a seed

crystal of the desired compound (if available, and often it is not available) or

scratching the inside of the beaker below the liquid surface helps to induce

crystallization.

A slow addition of a solvent that is miscible (but less polar) with the solvent of the

reaction mixture but which does not dissolve the desired product followed by cooling,

(seeding), and scratching.

For a cationic complex an appropriate anion with which it forms an insoluble salt can

be added. A suitable cation may be added to the reaction mixture containing an

anionic complex. E.g. to precipitate [Ni(CN)5]3-

from a solution add a large trivalent

cation like [Cr(en)3]3+

to give [Cr(en)3][Ni(CN)5].

Chromatography can be used to separate and purify complexes.

Other techniques are distillation (could be under reduced pressure if the compound

decomposes before its boiling point) and sublimation (for volatile complexes), and

Soxhlet extraction (of either the complex, if soluble in extractor solvent or of the

impurity if the complex is insoluble).

1.1 General principles of synthesis coordination compounds

There are two important variables that occur in reactions involving transition metals:

Coordination number

Oxidation state.

Either may increase, decrease, or remain unchanged in a reaction. It is, in practice, not

possible to predict either of these variables in a reaction. This is because ligands behave in

peculiar way depending on a number of constraints. E.g. a tetradentate ligand may behave as

a bidentate ligand. It is also possible for a ligand, which should be anionic, to coordinate

neutral or as a radical. Whether a reaction results in a change of oxidation state of the central

metal or not would depend on the mode of coordination of the resulting complex.

In general the following classifications hold:

Addition reaction: Coordination number of an electron acceptor (the metal/metal ion)

increases.

Substitution reaction: Coordination number is unchanged.

Dissociation reaction: Coordination number decreases.

Oxidation – reduction reaction: There is a change in oxidation state.

Coordination compounds are also classified according to the speed at which they undergo

substitution reaction:

Complexes that undergo substitution reaction at room temperature spontaneously are

said to be kinetically labile.

Those where substitution takes hours/days are said to be kinetically inert.

1.3 General rules guiding lability/inertness

1.3.1 Labile complexes

Complexes with central metal atom having d electrons in the eg orbitals, e.g.

[Ga(C2O4)3]3-

, d10

; [Co(NH3)6]2+

, d7+

; [Cu(H2O)6]2+

, d9; [Ni(H2O)6]

2+, d

8 and

[Fe(H2O)6]3+

, d5.

Complexes containing less than 3 electrons in the d orbitals , e.g. [Ti(H2O)6]3+

, d1;

V(phen)3]

3+, d

2 and [Ca(EDTA)]

2+, d

0+.

1.3.2 Inert complexes

Octahedral low-spin d4, d

5 and d

6 complexes, e.g. [Fe(CN)6]

3-; [Co(NH3)6]

3+ and

[PtCl6]2-

, d6.

Octahedral d3 complexes, e.g. [Cr(H2O)6]

3+, d

3.

Crystal field approach helps to see the picture clearly.

1.3.3 Addition reactions

Addition reactions lead to increase in coordination number, usually accompanied by colour

changes.

[ML4] + Y → [ML4Y]

[ML4Y] + Y → [ML4Y2]

Y is an adduct and can be the solvent molecule or another molecule, e.g.

[Cu(acac)2] + py → [Cu(acac)2py]

The product may or may not be isolable but the formation of the product can be detected

because of the change in coordination number.

1.3.4 Substitution reactions

Majority of complexes can be prepared by substitution reactions, in a number of cases,

displacing water. However, the method employed depends on whether the complex being

substituted is labile or inert.

1.3.5 Preparation of labile complexes

Formation of labile complexes is virtually instantaneous upon mixing of the reactants hence

there are few practical difficulties in their preparation, but three points must be remembered:

It is, in practice, difficult to prepare such complexes with several non-ionic ligands

bonded to the same metal atom, although anionic species may be coordinated together

with a neutral ligand.

Although it may be possible to isolate and characterize a solid complex quite a

different complex may be the predominant species in solution.

Some complex ions display incongruent solubility (arising from the second point

above), e. g. if an aqueous solution containing iron(II) sulphate and ammonium

sulphate in 1:1 ratio is allowed to crystallize then [Fe(H2O)6]SO4.(NH4)2SO4 is

formed. The iron(II) ammonium sulphate is said to show congruent solubility.

However, if solutions containing KCl and CuCl2 at ratio 2:1 are allowed to crystallize,

crystals of KCl are obtained first and only later does the complex K2[Cu(H2O)2Cl4]

crystallize. If the complex is re-crystallized there is an initial deposition of KCl again.

The complex is said to display incongruent solubility.

1.3.6 Basic principles for the preparation of metal complexes

Labile complexes are prepared in aqueous medium from hydrated salts.

Inert complexes are prepared from anhydrous complexes if non-aqueous medium is to be

used. Where only hydrated salts are available, salts have to be dehydrated first before use. If

preparation is to be done in aqueous medium then a labile complex of a lower oxidation state

is oxidized or a salt at a higher oxidation state is reduced.

Some ions are unstable to oxygen. Complexes of such ions are prepared in an inert

atmosphere, e.g. under N2 gas.

Examples

Substitution reaction in aqueous solution is the most common method for labile complexes.

The method involves a reaction between a metal salt in water and a coordinating agent.

1. Action of excess ammonia on aqueous solution of copper(II) salts:

[Cu(H2O)6]2+

+ 4 NH3(l) → [Cu(NH3)4]2+

+ 6 H2O

or

[Cu(H2O)6]2+

+ 4 NH3(aq) → [Cu(H2O)2(NH3)4]2+

+ 4 H2O

The instant replacement of water by ammonia at room temperature is shown by a change in

colour. Other species corresponding to stepwise substitution of the water by NH3 exist in

solution, e.g. [Cu(H2O)5(NH3)]2+

, [Cu(H2O)4(NH3)2]2+

, etc.

There are many labile complexes which may be studied readily in solution but which are very

difficult to obtain in the solid state.

1.3.7 Preparation of uncharged complexes

A neutral complex is usually precipitated from either aqueous solution or aqueous alcohol

and, unless highly polymeric, may be re-crystallized from organic solvents. For example,

[Cu(H2O)6]2+

+ 2 Hacac [Cu(acac)2] + 4 H2O + 2 H3O+

For this equilibrium to be displaced to the right (i.e. in favour of the formation of the

complex) the system is usually buffered to about pH 6. Sodium actate is commonly used.

1.3.8 Preparation of inert complexes

As mentioned earlier, substitution is usually slow and preparations can be done in three ways:

1. If water is not being displaced and reaction is being done in water more drastic

experimental conditions are imposed.

For example the preparation of K3[Rh(C2O4)3] is done in boiling concentrated aqueous

solutions of K3[RhCl6] and K2C2O4 for 4 h and then evaporated until product crystallizes

from the solution.

H2O, 4 h

100 oC

K3[RhCl6] + 3 K2C2O4 K3[Rh(C2O4)3] + 6 KClwine red yellow

To prepare [Co(en)3]Cl3 from [CoCl(NH3)5]Cl2, the rection is carried out on a steam bath

because the reaction is slow at room temperature;

[CoCl(NH3)5]Cl2 + 3 en [Co(en)3]Cl3 + 5 NH3

2. If water is being displaced then the water has to be replaced first before he correct

product can be obtained.

For example, Cr(III) complexes cannot be made from aqueous solvents if water is undesirable

in the complexes (alternative routes are available).

For example potassium thiocyanate (m.p. 173o) can be used as a solvent at elevated

temperatures above 173o. Under this condition water is readily displaced from [Cr(H2O)6]

3+.

[Cr(H2O)6]3+

+ 6 NSC-

180o

molten KCNS[Cr(NSC)6]

3-+ 6 H2O

In certain instances the salt is first dehydrated before the product can be obtained.

Dehydration can be effected by use of thionyl chloride or 2,2-dimethoxypropane. Preparation

of anhydrous complexes are best done from organic solvents and starting with, preferably,

anhydrous salts, where they are available.

Examples

1. Consider the preparation of (NEt4)2[NiCl4] from [Ni(H2O)6]Cl2.

[Ni(H2O)6]Cl2 + 6 SOCl2 NiCl2 + 6 SO2 + 12 HCl

NiCl2 + 2 NEt4ClSOCl2

Reflux (NEt4)2[NiCl4] Alternatively,

[Ni(H2O)6]Cl2 + 6 (MeO)2CMe2NiCl2 + 6 Me2CO + 12 MeOH

NiCl2 + 2 NEt4Cl (NEt4)2[NiCl4]

2. Preparation of [Cr(en)3]Cl3 from CrCl3.6H2O

[Cr(H2O)6]Cl3 + 6 (MeO)2CMe6 CrCl3 + 6 Me2CO + 12 MeOH CrCl3 + 3 en Cr(en)3]Cl3

1.3.9 Oxidation-reduction reaction in the preparation of inert complexes

Oxidation

Co(II) salts are usually used as starting materials for the preparation of Co(III) complexes;

CoCl2 + 2NH4Cl + 10 NH3 + H2O2 2 [Co(NH3)6]Cl3 + 2 H2OCharcoal

Charcoal acts as a catalyst. In its absence the product is mostly [Co(NH3)5X] complexes, X =

H2O or Cl.

If it is only hydrated salts of an inert complex that is available it is best to dehydrate the salt

first before use.

Reduction

The preparation of K3[Cr(C2O4)3] from K2Cr2O7

Here the dichromate is reduced by an aqueous solution of oxalic acid and potassium oxalate:

K2Cr2O7 + 7 H2C2O4 + 2 K2C2O42 K3[Cr(C2O4)3] + 6 CO2 + 7H2O

Note that in the instances where oxidation-reduction reactions are used the starting

compounds are labile. Cr(III) complexes can also be prepared from Cr(II). However, Cr(II)

compounds are rather unstable and can only be stored under inert atmosphere.

Thermal dissociation reactions

In certain instances a new complex can form by controlled heating of another complex,

usually with evolution of a volatile compound.

For example the preparation of anhydrous copper(II) sulphate from the hydrated salt is done

by controlled heating of the latter.

CuSO4.5H2O CuSO4.4H2O CuSO4.H2OCuSO4.3H2O

CuSO4

96.5o 102

o115

o

220o

The blue hydrated copper(II) sulphate loses water stepwise until all the water of

crystallization is lost to give the off-white anhydrous salt. Controlled heating under vacuum

is therefore a viable method for making a number of complexes.

Examples

The conversion of [Cr(en)3]Cl3 to cis-[Cr(en)2Cl2]Cl is done by controlled heating:

[Cr(en)3]Cl3 → cis-[Cr(en)3Cl2]Cl

1.3.10 Some methods employed in characterizing coordination compounds

Before embarking on characterization, compounds must be ascertained to be pure. To

establish purity the following could be done:

Determine the melting point. The melting point of a pure compound is expected to be

sharp. However, a sharp melting point does not necessarily refer to the melting point

of the compound anticipated. Where the melting point of the compound has been

reported, the melting point of the compound could be compared with the literature

value.

Microanalysis. The percentages of all the elements present in the compound are

determined and compared with the values calculated theoretically.

When the purity of the compound has been established the compound could then be

characterized using some of the following techniques:

Conductivity measurement. The molar ionic conductance of a compound (the

conductance of 1 mole of ions from the compound at infinite dilution) is determined.

The contribution to the molar ionic conductance of an ion Ix±

is about 60 ohm-1

cm2.

For ions Mm+

and Xn-

in the salt MnXm the contribution will be 60m ohm-1

(from Mm+

)

and 60n ohm-1

(from Xn-

). Multiplying by the number of ions of each sort and adding

leads to the conclusion that a salt MnXm will have a molar conductivity of about

120nm ohm-1

at 20 oC .

Example: The molar conductance of CoCl3.5NH3 is 261 ohm-1

cm2, hence

120nm = 261; nm = 2;

n = 1 and m = 2; n + m = 3.

That is, the number of ions is 3, hence the structure is [Co(NH3)5Cl]Cl2.

Generally m 120nm

Chemical reactions (already mentioned under introduction)

Infrared spectroscopy (i.r.): Normally the infrared spectrum of the ligands and the

complexes are required for meaningful comparison. The differences between the

spectra fall into four categories:

1. Band positions may change

2. Relative band intensities may change. Usually new, often weak, bands may

appear.

3. Single peaks in the free ligands may split into several, closely spaced, bands in

the complex.

4. Some peaks in the ligand may disappear while new ones, particularly in the region

due to Metal – Ligand bond (below 600 cm-1

may appear. Evidence for coordination and

atoms in coordination can be established from i.r. spectra. The mode of coordination of some

anionic ligands can also be detected.

Examples

CO32-

can be monodentate or bidentate. Free CO32-

absorbs at 890 cm-1

; coordinated

CO32-

absorbs at ~850 cm-1

(if monodentate) and at ~830 cm-1

(if bidentate).

SCN- can be S-bonded, where (C-S) is at ~700 cm

-1 or N-bonded where (C-S) is at

~820 cm-1

.

NO2- could be N-bonded [(N-O) is at ~1310 cm

-1] or O-bonded [(N-O) is at ~1065 cm

-

1].

Stretching frequency in metal-oxygen double bond, as in V=O occurs at 960±50 cm-1

.

Metal-ligand absorptions are generally weak and occur at 600 – 100 cm-1

region, which

may present instrumental problems. Usually any i.r. spectrometer that does not record to

200 cm-1

is of limited use in coordination chemistry. This is the region to concentrate on

for M-N, M-O, and M-X frequencies.

Infrared spectra for coordination compounds are usually recorded preferably in the solid

(KBr or CsBr pellets). There are overlaps when run in Nujol and one cannot go below

600 cm-1

in Nujol.

Other techniques that can be used in the characterization of coordination compounds include:

UV-Visible spectroscopy

Photoluminescence

Magnetochemistry

Thermogravimetric analysis

Differential thermal analysis

Cyclic Voltammetry (for oxidation-reduction properties of complexes)

Mass spectrometry

Nuclear magnetic resonance spectrometry

Mössbauer Spectrometry

Optical rotatory dispersion and circular dichroism (for optically active complexes)

X-ray diffraction (which gives the ultimate structure of the compound unequivocally)

2. Nomenclature

The International Union of Pure and Applied Chemistry (IUPAC) system will be discussed.

2.1 Naming ligands

Ligands can be anionic or neutral. Both anionic and neutral ligands can

be monodentate or polydentate.

2.1.1 Anionic ligands

Anionic ligands end in “-o”

Monodentate anionic ligands:

Ligand Name Ligand Name

Cl-

Chloro O2-

Oxo

Br-

I-

OH-

NH2-

NO2-; M-NO2

ONO-; M-ONO

CN-; M-CN

NC-; M-NC

Bromo

Iodo

Hydroxo

Amido

Nitro

Nitrito

Cyano

Isocyano

H-

O22-

CH3COO-

SCN-; M-SCN

NCS-; M-NCS

Hydrido

Peroxo

Acetato

S-thiocyanato

or thiocyanato

N-thiocyanato

or isothiocyanato

Bidentate anionic ligands

Ligand Name

acac-; [CH3COCHCOCH3]

-

glyox-; [HONC2H2NO]

-

ox2-

; [C2O4]2-

Hdmg-, [CH3CNCNCH3]

-

Acetylacetonato or Pentane-2,4-dionato

Glyoximato

Oxalato

Dimethylglyoximato

Polydentate ligand: EDTA4-

, [(O2CCH2)2N(CH2)2N(CH2COO)2]4-

whose IUPAC name is

ethylenediaminetetraacetato is one of the most commonly used polydentate ligand. The

neutral tetrabasic acid is represented as H4EDTA.

2.1.2 Neutral ligands

Monodentate

Ligand Name Ligand Name

H2O

NH3

CO

NO

Aqua or aquo

Ammine

Carbonyl

Nitrosyl

RNH2

py

(C6H5)3P

Alkylamine

Pyridine

Triphenylphosphine

Bidentate

Lgand Name Structure

en Ethylenediamine NH2CH2CH2NH2

Bipy

Bipyridine

N N

1,10-phen or phen Phenathroline

N N

2.2 Naming complexes

2.2.1 If it is a salt:

Name cation first and then the anions, like in all salts. There is a space between the

cation and the anion.

Within a complex:

Name negative ligands.

Name the neutral ligands.

Name the metal (with oxidation state, in Roman numeral, in brackets).

Where there is more than one type of ligand in a complex, name in alphabetical order or in

order of complexity where they start with the same alphabet.

For the number of ligand (of the same type) use di-, tri-, tetra-, penta- and hexa- for 2, 3, 4,

5, and 6 respectively.

If the ligands are multi-syllabic put the name of the ligand in parenthesis. The numerical

prefixes are changed to bis-, tris-, tetrakis-, pentakis, and hexakis- for 2, 3… 6.

In anionic complexes the name of the metal ends “-ate”. In some cases the Latin name is

used; e.g. iron becomes ferrate.

In cationic complexes the metal retains its English name.

Examples

[Cr(NH3)6](NO3)3 hexaamminechromium(III) nitrate

K2[PtCl6] potassium hexachloroplatinate(IV)

K3[Fe(ox)3].3H2O potassium trioxalatoferrate(III), trihydrate or 3-water

Na[Co(CO)4] sodium tetracarbonylcobaltate(-I)

K4(Ni(CN)4] potassium tetracyanonickelate(0)

[Co(en)2Cl2]Cldichlorobis(ethylenediamine)cobalt(III) chloride

[Co(NO2)3(NH3)3] triamminetrinitrocobalt(III)

In a neutral complex, the ligands are named first adopting the rules above followed by the

metal.

Example

[Ni(Hdmg)2] bis(dimethylglyoximato)nickel(II)

2.2.2 Bridging complexes

Ligands that bridge two centres of coordination (polynuclear) are preceded by the Greek

letter, , which is repeated before the name of each different kind of bridging group.

[(H2O)4Fe

HO

OH

Fe(H2O)4](SO4)2

[(en)2Co

-dihyroxobis[tetraaquairon(III)] sulphate or tetraaquairon(III)--dihydroxotetraaquairon(III) sulphate

HN

OH

Co(en)2]Cl4

bis(ethylenediamine)--imido-m-hydroxo-bis(ethylenediamine)cobalt(III) chlorideor hydroxo--imidobis[bis(ethylenediamine)cobalt(III)] chloride

[(NH3)4Co

NO2

HN

Co(NH3)4] (NO3)4

-amido--nitrobis[tetraamminecobalt(III)] nitrate

2.2.3 Point of attachment

Whenever necessary the point of attachment of a ligand is designated by placing the symbol

(in italics) of the element attached after the name of the group is separated by hyphen.

(NH3)3[Cr(NCS)6] ammonium hexathiocyanato-N –chromate(III) or

hexaisothiocynatochromate(III)

(NH3)2[Pt(SCN)6] ammonium hexathiocyanato-S-platinate(IV)

2.2.4 Naming geometric isomers

Geometric isomers are generally named by the use of the terms cis to designate adjacent (90o

apart) positions and trans for the opposite (180o apart) positions. It is occasionally necessary

to use a number system to designate the position of each ligand. For square-planar

complexes, groups 1-3 and 2-4 are in trans positions. Note that only two of the

transpositions need be numbered in the name of the second complex below. This is because

in a square-planar complex the other two ligands must then be in trans positions. Since

positions 2 and 4 are equivalent these numbers need not be mentioned.

M

1

2

3

4

Pt

Cl

NH3

NO2

NH3 Pt

NH3

Cl NO2

Br

trans-diamminechloronitroplatinum(II) 1-ammine-3-bromochloronitroplatinum(II) ion

-

Number system in square-planar complexes

The number system for octahedral complexes has the trans positions numbered 1-6, 2-4, and

3-5.

M

12

3

4

5

6

Rh

NH3

NH3

NH3

NH3 Br

Br

Pt

Cl

NH3

py NO2

Br

I

cis-tetraamminedibromorhodium(III) ion1-ammine-2-bromo-4-chloro-6-iodonitro(pyridine)platinum(IV)or trans-ammineiodo-trans-bromochloronitro(pyridine)platinum(IV)

+

2.2.5 Naming optical isomers

If a solution rotates plane-polarised yellow light (the NaD line) to the right, the solute is

designated a (+) isomer; if to the left, a (-) isomer.

(+) K3[Ir(C2O4)3] potassium (+)trioxalatoiridate(III)

(-) [Cr(en)3]Cl3 (-)tris(ethylenediamine)chromium(III) chloride

3 Isomerism in metal complexes

Isomerism can be divided into two broad divisions: structural isomerism and

stereoisomerism.

Structural isomerism: Ionization isomerism, hydration isomerism, coordination isomerism,

linkage isomerism, ligand isomerism, and polymerization isomerism. Stereoisomerism:

Geometric isom

3.1 Structural isomerismerism, conformational isomerism, and optical isomerism.

3.1.1 Structural isomerism: ionization isomerism

Ionization isomerism results from the interchange of negative ligand within the first

coordination sphere of a complex that has an anion outside the coordination sphere. Such

isomers yield different ions in solution.

Examples:

[Co(NH3)4ClNO2]I and [Co(NH3)4ICl]NO2: When both are dissolved in water the first gives a

complex as cation and I- as anion whereas the second gives a complex as cation and NO2

- as

anion. Other examples are [Co(NCS)2(en)2]Cl and [Co(NCS)Cl(en)2]NCS; Pt(NH3)3Br]NO2

and [Pt(NH3)3NO2]Br. The isomers can be readily distinguished by appropriate qualitative

analysis.

[Co(NH3)5Br]SO4 and [Co(NH3)5SO4]Br: These can also be distinguished by appropriate

qualitative analysis.

Develop qualitative analysis schemes to distinguish between each pair of the compounds

above.

3.1.2 Structural isomerism: hydration isomerism/solvate isomerism

Hydration isomerism results from the interchange of H2O and another ligand between the

first coordination sphere and the ligand outside. Here H2O can be a ligand or water of

crystallization.

Most common example is CrCl3.6H2O, which can give three possible isomers,

distinguishable by their colours:

[Cr(H2O)6]Cl3 Violet

[Cr(H2O)5Cl]Cl2.H2O Blue green

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

Other examples are [CoCl(en)2H2O]Cl2 and [CoCl2(en)2]Cl.H2O; [CrCl2(py)2(H2O)2]Cl and

[CrCl3(py)2H2O].H2O.

The isomers can be distinguished by quantitative precipitation of free chloride using silver

nitrate.

Describe how you would carry out the quantitative precipitation of free chlorides in the

isomers in 5.2. Write appropriate equations and show how you would use your results to

distinguish the isomers.

3.1.3 Structural isomerism: coordination isomerism

Coordination isomerism occurs in salts in which both cation and anion are complex ions.

Isomerism arises from interchange of ligands between the two complex ions.

Examples:

[Co(NH3)6][Cr(ox)3] and [Cr(NH3)6][Co(ox)3]

[PtII(NH3)4][Pt

IVCl6] and [Pt

IV(NH3)4Cl2][Pt

II(NH3)4]

[Co(NH3)6][Cr(CN)6] and [Cr(NH3)6][Co(CN)6]

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

3.1.4 Structural isomerism: linkage isomerism

Linkage isomerism occurs when one or more of the ligands can coordinate to the metal ion in

more than one way. Linkages with two or more possible sites of attachment to a metal are

known as ambidentate ligands. In principle they include NO2-, SCN

-, CN

-, S2O3

2-, CO,

CONH2, CS(NH2)2, and (CH3)2SO but only the first four have been shown to form linkage

isomers.

Examples

[Co(NH3)5NO2]2+

and [Co(NH3)5ONO)]2+

are isomers . In the first one NO2- bonds to Co by

N (a nitro complex) whereas in the second it bonds by O (a nitrito complex). The complexes

can be distinguished by using IR spectroscopy. For the O-bonded ligand, characteristic

absorption bands at 1065 and 1470 cm-1

are observed whereas corresponding bands for the N-

bonded ligand are at 1310 and 1430 cm-1

. The above complexes can be written as

[Co(NH3)5(NO2-N)]2+

and [Co(NH3)5(NO2-O)]2+

respectively.

[Co(NH3)5SCN]2+

has two isomers, [Co(NH3)5SCN-S]2+

and [Co(NH3)5SCN-N]2+

, where

SCN- bonds via S and N respectively. Other examples can be found in complexes of CN

-

where the ligand can bond by C (cyno complexes) or N (isocyano complexes).

Other examples: [dipyPd(SCN)2] and [dipyPd(NCS)2]; [Mn(CO)5SCN] and [Mn(CO)5NCS].

All linkage isomers can be readily identified by IR spectroscopy.

e.g. M-N=C=S has a characteristic band at 780 – 860 cm-1

whereas in M-S-CN it occurs at

690 – 720 cm-1

.

3.1.5 Structural isomerism: polymerization isomerism

Polymerization isomerism refers to complexes, which have the same empirical formulae but

different molecular masses.

Examples:

[PtCl2(NH3)2] and [Pt(NH3)4][PtCl4] and [Co(NH3)3(NO2)3] and [Co(NH3)6][Co(NO2)6]

3.1.6 Structural isomerism: Ligand isomerism

Ligand isomerism is due to isomerism in the organic compounds that constitute the ligands.

Since many ligands are organic compounds and the latter have a large number of possibilities

for isomerism, the resulting complexes can show isomerism from this source.

Examples of isomeric ligands are 1, 2-diaminopropane (pn) and 1,3-diaminopropane (tn) or

o-, m- and p-toluidine (CH3C6H4NH2). When these compounds are used as ligands the

resulting complexes are also isomeric.

e.g. [Co(pn)3 and [Co(tn)3 are isomers.

3.2 Stereoisomerism

3.2.1 Stereoisomerism: geometrical isomerism

Geometrical isomerism occurs when a metal complex of the same formula and same basic

geometry has a different arrangement of ligands. This occurs in 4-coordinate square-planar

and 6-coordinate octahedral complexes.

3.2.2 Geometric isomerism in square planar complexes

cis and trans isomerism occurs when ligands are arranged adjacent (90o) and opposite (180

o)

respectively. Common geometric isomers are found in complexes of the type MA2B2. They

are more common in Pt and Pd complexes. Only two isomers are known in square-planar

complexes of this type. If the compound is tetrahedral only one isomer is possible..

A second type is of the form MABCD. Three square-planar complexes of this type is

possible. If it is tetrahedral, then only one isomer (which is also potentially optically active,

but not resolvable) is possible. The first type is illustrated below:

Pt Pt Pt

Cl

Cl

Cl

Cl

Cl

Cl

NH3 NH3 NH3 NH3NH3 NH3

The trans-isomer isnon-polar

Asymmetric stretchIR active

Symmetric stretchIR inactive

Cl

Cl

Cl

Cl ClPt Pt Pt

Cl

NH3

NH3

NH3

NH3

NH3

NH3

The cis-isomer is polar

Asymmetric stretchIR active

Symmetric stretchIR active

The cis- and trans-isomers of [Pt(NH3)2Cl2] can be distinguished by IR spectroscopy. An IR

active vibration leads to a change in molecular dipole moment.

The three isomers of the second type are given below.

M M M

A

B

C

D

A

C

B

D

A

B

D

C

3.2.3 Geometric isomerism in octahedral complexes

Two types of geometric isomerism are recognized in octahedral complexes: the simple types,

which exist in complexes of the type MA2B4, in which A may be adjacent to each other or

opposite and the second for complexes of the type MA3B2

MA2B4 type

An example is [Co(NH3)4Cl2]+

Cl

NH3

NH3

NH3NH3

Cl

Co

Cl

NH3

NH3NH3

Co

Cl

NH3

cis-isomer trans-isomer

MA3B3 type

Two types are possible: (1) here the ligands of one type forms an equilateral triangle on one

of the faces (the facial isomer, abbreviated as fac) and (2) the ligands span three positions

such that two are opposite, ortrans, to each other (the meridionalisomer, abbreviated as mer).

The following complexes can display the fac-mer isomerism: [Ru(H2O)3Cl3]; [Pt(NH3)3Br3]+;

[Pt(NH3)3I3]; [Ir(H2O)3Cl3]; [Rh(CH3CN)3Cl3]; [Co(NH3)3(NO2)3]; and [M(CO)3PR3], (M =

Cr, Mo, W).

Cl

H2O

Cl

ClH2O

H2O

Ru

OH2

Cl

ClCl

Ru

H2O

H2O

fac-isomermer-isomer

3.2.4 Stereoisomerism: conformational isomerism

Conformational isomerism occurs if a metal complex can exist in two totally different

geometric forms. For example, [NiCl2(Ph2PCH2Ph)2] can exist as square-planar in the solid

state and tetrahedral in solution.

Ni

P

P

Cl

Cl

Square-planar

Ni

Cl

PCl

P

Tetrahedral

P = PH2PCH2Ph

Properties of geometric isomers:

One isomer is usually stable in solid. In solution it often isomerizes to a mixture or

the other isomer. For example, green trans-[Co(en)2Cl2]+ isomerizes to a mixture of

cis and trans. The purple cis-isomer also isomerises to cis and trans-isomers.

Isomers usually have different colours.

Their chemical properties are usually different; for example they react at different

rates.

3.2.5 Stereoisomerism: optical isomerism

Optical isomerism is concerned with chirality.

A molecule is chiral if it possesses a non-superimposable mirror image. Octahedral

complexes, like [Co(acac)3], which has three bidentate chelating ligands also possesses non-

superimposable mirror images. Chiral molecules (enantiomorphs) rotate plane polarized

light in opposite directions. This property is known as optical activity and the two mirror

images are known as optical isomers or enantiomers.

Enantiomers rotate plane polarized light to equal extents in opposite directions, the

dextrorotatory (dor+) enantiomorphs to the right and the laevorotatory (lor-) to the left at a

particular wavelength. The observation of optical activity depends upon the chemical

properties of the chiral molecules; if the two enantiomorphs interconvert rapidly to give an

equilibrium mixture containing equal amounts of the two forms, there will not be any overall

rotation. A mixture of equal amounts of enantiomorphs is called a racemic mixture.

A polarimeter is used to measure the rotation, , of an enantiomorph. The amount of rotation

depends on the wavelength of the light, temperature, and the concentration of the compound.

The specific rotation, [], for a chiral compound in solution is given by:

[] =

c x l where = observed rotation, l = path length of solution in the polarimeter (in dm) and c =

concentration (in g cm-1

). Light of a single frequency is used for specific rotation

measurements and a common choice is sodium D-line in the emission spectrum of atomic

sodium; the specific rotation at this wavelength is denoted as []D. The importance of

chirality has been found in the dramatic differences in the activities of different enantiomers

of chiral drugs.

Read: E Thall (1996) Journal of Chemical Education, Vol. 73, p.481 – “When drug

molecules look in the mirror”.

3.3 Nomenclature of chiral molecules

Nomenclature of chiral molecule is complicated.

In terms of the sign of the rotation of plane-polarized light:

- the rotation is denoted (+) or d for dextrorotatory and (-) or lfor laevorotatory

- the sign and magnitude of rotation could be incorporated ; (-)589 or (-)D (where D is

sodium D-line at a wavelength of 589 nm)

This method of naming defined in terms of an observable (the rotation) does not bear any

direct relationship with the absolute configuration of the molecule.

IUPAC recommended and Δ system

3.3.1 Definitions and notation of chiral complexes

Enantiomers are a pair of stereoisomers that are non-superimposable mirror images.

Diastereomersare stereoisomers that are not an enantiomeric pairs.

(+) and (-)prefixes: the specific rotation of enantiomers is equal and opposite. Enantiomers

are distinguished by the sign of []D. Two enantiomers of a compound A with []D values of

+12 and -12 are denoted as (+)-A and (-)-A.

d and lprefixes: sometimes (+) and (-) are denoted by dextro- andlaevo- for right and left

rotations respectively.

The +/- or d/l notation, as mentioned above, has nothing to do with the absolute configuration

of an enantiomer (the arrangement of the substituents or ligands). The following prefixes are

used for describing absolute configuration.

R and S prefixes: this convention is used for labeling chiral carbon atoms (tetrahedral with

four different groups attached) and is based on the Cahn-Ingold-Prelog notation.

The four groups attached to the chiral carbon atom are prioritized according to the atomic

number of the attached atoms, highest priority being assigned to highest atomic number, and

the molecule then viewed down the C-X vector, where X has the lowest priority. The R- and

S-labels for the enantiomers refer to a clockwise (rectus) and anticlockwise (sinister)

sequence of the prioritized atoms, working from high to low. Example: CHClBrI, view down

the C-H bond:

C

I

H

BrCl

C

I

H

ClBr

1

23

1

2 3

R S

This notation is used for chiral organic ligands, and also for tetrahedral complexes.

Δ and prefixes: The enantiomers of octahedral complexes containing three equivalent

bidentate ligands (tris-chelate complexes) are among those which are distinguished using Δ

(delta) and (lambda) prefixes. The octahedron is viewed down a 3-fold axis, and the

chelates then define either a right- or a left-handed helix. The enantiomer with right-

handedness is labeled Δ, and that with left-handedness is .

=

=

Depending on the metal ion and the number and nature of chelate rings enantiomers can be

separated.

3.3.2 Rules for optical activity

An optically active compound must not have:

A centre of inversion

A plane of symmetry

An improper axis S

The mirror image must not be superimposable.

Examples of stereoisomers

[Co(en)2Cl2] can exist as cis- and trans-isomers. The cis isomer is potentially optically active

because its mirror image is not superimposable.

N Cl

N Cl

Co

N

N

NCl

NCl

Co

N

N

cis-[Co(en)2Cl2]

But the trans-isomer has a plane of symmetry and a centre of inversion. It is therefore not

optically active.

N N

N N

Co

Cl

Cl

trans-[Co(en)2Cl2]

The first compound to be resolved was cis-[Co(en)2(NH3)Cl]2+

. It was resolved with

(+)bromo-π-camphor sulphate. In the solid state, the enantiomorphs are stable. However,

they often racemize in solution and therefore may not be possible to isolate isomers even if

you know they exist.

The best complexes to isolate and keep in solution are those where

M is inert (undergoes slow substitution); e.g. CoIII

(d6) and Cr

III (d

3) complexes. For

example [Co(en)3]Cl3 is potentially active and boiling for hours will not convert one isomer

to the other. On the other hand [Zn(en)3]2+

racemises very rapidly because Zn2+

is a d10

ion

and is very labile.

Ligands are bidentate or polydentate. Such complexes are stabilized by chelate effect.

Good examples are found in EDTA complexes.

The first purely inorganic complex to be resolved into its optical isomers was [CoL3]6+

,

where L+

ligand = cis-[Co(NH3)4(OH)2]+ complex. L

+ chelates through the two O-donor

atoms.

Co

HO

OH

Co(NH3)4

3

6+

3.3.3 Methods for resolution of isomers

Most common methods used for least soluble diastereoisomers (cations and anions)

only, involve reacting the diastereoisomer with an optically active organic anion.

For example,

(±) C+ + optically active organic anion (+)A

- → (+)C

+.(+)A

- and (-)C

+.(+)A

-

Useful for diastereoisomeric salts (not mirror images).

Employing differences in solubility of the two diastereoisomeric salts

Example,

(±)[Co(en)3]3+

+ (+)tartrate2-

→ (+)[Co(en)3].(+)tartCl.5H2O precipitates but

(-)[Co(en)3].(+)tart stays in solution. This is a practical method.

The separated diastereoisomers must now be reconverted:

Example,

(-)[Co(en)3].(+)tart + NaI → (-)[Co(en)3]I3 (precipitated)

(+)[Co(en)3].(+)tart + xss KI → (+)[Co(en)3]I3

Other methods

Partial asymmetric synthesis

Example

(±)[Fe(phen)3]2+

+ (+)[SbO tartrate] → (-)[Fe(phen)3].(+)[SbOtart] ppt

(+) form still remains in solution.

This method only works if the metal complex has a certain amount of mobility. The

precipitate is fairly labile and on heating, the (+) form goes to (-) form, which is insoluble.

For anions the same technique is used except that organic cations are used as resolving

agents.

A metal complex which is optically active can itself be used as a resolving agent for

other metal complexes.

Examples,

Using (+)[Co(en)2(NO2)2]+ can be used to resolve [Co(en(ox)2]

- with 100% yield and

[Co(edta)]-. Once resolved these metal complexes can be used to resolve optically active

metal cations.

3.3.4 Resolving neutral metal complexes

The complex can be adsorbed on an optically active column like quartz, (+) lactose, or starch.

Example,

[Co(acac)3]0 on (+)lactose can be eluted with C6H6 or hexane. One isomer is preferentially

adsorbed and one is eluted.

4 Stereochemistry

Stereochemistry is that branch of chemistry that concerns with the structures of compounds.

Inorganic stereochemistry deals with central atoms having coordination numbers from two to

twelve.

4.1 Coordination number 2

Coordination number 2 is confined to complexes of CuI, Ag

I, Au

I, and Hg

I. The complexes

are all linear. Examples include ammine complex of AgI,

[NH3 → Ag NH3]+, and cyano complex of Ag

I, [NC Ag CN]

-. These ions have d

10

configuration. Other examples are [AgCl2]-, [HgCl2], [AuCl2]

- and Au(CN)2

-.

4.2 Coordination number 3 This is rather rare among metal complexes. Many complexes which appear to be 3-

coordinate as judged by their stoichiometry are found upon examination to have higher

coordination numbers. For examples, Cs[CuCl3] has an infinite chains –Cl-CuCl2-Cl-, a

coordination number of 4; K[CuCl3] has an infinite double chains of Cl4-Cu2Cl2-Cl4, a

coordination number of 6, having a distorted octahedral structure; and K[Cu(CN)2] has a

chain [-CN-Cu(CN)-CN-Cu(CN)-CN-) and is an example of true 3-coordination.

Four other examples of truly 3-coordinate complexes are the triiodomercurate(II) anion, HgI3-

, bis(thiourea)copper(I) chloride, [Cu(tu)2]Cl, tris(trimethylphosphinesulphidecopper(I)

perchlorate, [Cu(SPMe3)3]+ClO4

- and tris(triphenylphosphine)platinum(0), [{(Ph)3P}3Pt]

0.

All the examples feature ligands with bulky groups.

4.3 Coordination number 4

This coordination number is very common. The structures formed with this coordination

number can be divided into two: tetrahedral and square planar. There are, however,

intermediate structures and distortions.

4.4 Tetrahedral complexes

Tetrahedral structures are not stabilized by large CFSE (see later). The structure is favoured

by large ligands like Cl-, Br-, I

- and small metal ions of three types:

Those with a noble gas configuration ns2np

6 such as Be

2+, Mn

VII (e.g. MnO4

-)

Those with a pseudo-noble gas configuration ns2np

6(n-1)d

10, such as Cu

I, Zn

II Ga

III

and Ni0.

Those transition metal ions, which do not strongly favour other structures by virtue of

the CFSE such as CoII, d

7.

Specific examples for transition metal ions are MnO4-, Ni(CO)4 and [Cu(py)4]

+. Tetrahedral

complexes do not exhibit geometric isomerism, although they are potentially optically active

just like tetrahedral carbon.

Example: bis(benzoylacetonato)beryllium.

4.5 Square planar complexes

They are formed by very few metal ions. The best known are d8 species such as Ni

2+, Pd

2+,

Pt2+

, and Au3+

. There are a few complexes of Co2+

(d7) with bidentate ligands that are square

planar, but otherwise such complexes are rather scarce. Chlorophyll and other bio-complexes

are important exceptions to this rule where the geometry is dictated by the rigid porphyrin

structure. Square planar structure is favoured by non-bulky, strong field ligands with

sufficiently good -bonders to compensate for the energy „lost‟ through 4- rather than 6-

coordination. For example [Ni(CN)4]2-

is square planar whereas Ni2+

forms octahedral

complexes with H2O and NH3, and tetrahedral complexes with Cl-, Br

- and I

-. For the heavier

d8 metals steric requirements are relaxed and the effective field strength of all ligands is

increased. Under such conditions [PdCl4]2-

, [PtCl4]2-+

and [AuCl4]- are square planar.

Account for why Pd2+

and Pt2+

have stronger tendency to form square planar complexes

than Ni2+

.

There is only a small difference in the energy of square planar and tetrahedral complexes.

Both structures can therefore inter-convert easily. A number of Ni2+

complexes do inter-

convert readily.

In the M2[CuX4] series of complexes of CuII, variation of M

I and X gives complex anions

with stereochemistries ranging from square planar (e.g. (NH4)2 [CuCl4]) to almost tetrahedral

(e.g. Cs2[CuBr4].

Square planar complexes, of the formula [MA2B2], exhibit cis-trans isomerism (see next

chapter). If such complexes are neutral molecules, they may be readily distinguished by the

presence of a dipole moment in the cis isomer but none in the trans isomer.

4.6 Coordination number 5 The structures of 5-coordinate complexes lie between two limiting geometries: trigonal

bipyramidal and square pyramidal. These limiting structures are not markedly different. The

conversion of one structure into other requires a relatively slight distortion.

Trigonal bipyramidal Square pyramidal

Examples: [CdCl5]3-

is almost an ideal trigonal bipyramidal; [NiCl5]3-

is almost an ideal

square pyramidal.

4.7 Coordination number 6 This is the commonest and most important coordination number for transition metal

complexes. The geometry usually corresponds to six coordinated atoms at the corners of an

octahedron or a distorted octahedron.

A regular geometry

The octahedral geometry is often subjected to tetragonal distortion leading to elongation or

contraction of the axial bonds. Complexes of the type [MA6] can have regular geometry

whereas complexes of the types [MA5B], [MA4B2], etc cannot have regular octahedral

geometry because not all the bonds will have the same length.

5 BONDING

5.1 Valence Bond Theory (VBT); Revision

Valence bond theory assumes that bonding in coordination compounds is solely covalent.

The formation of coordinate compounds is described by the hypothetical sequence:

Removal of electrons from the metal to give the appropriate cation.

Hybridization of those atomic orbitals which will provide a set of equivalent hybrid

orbitals directed towards the ligands.

Where necessary, rearrangement of the metal‟s electrons in order to ensure that the

hybrid orbitals are empty.

Formation of covalent bonds by overlap of the hybrid orbitals with the ligand

orbitals containing the lone pair of electrons.

The common types of hybrid orbitals are given below:

Coordination No. Atomic orbitals Hybrid orbitals Geometry

2

3

4

4

5

6

spx

spxpy

spxpypz

spxpydx2

-y2

dx2

- y2spxpypz

spxpypzdx2

-y2dz

2

sp

sp2

sp3

sp2d or dsp

2

dsp3 or sp

3d

sp3d

2 or d

2sp

3

Linear

Trigonal planar

Tetrahedral

Square planar

Square pyramidal

or trigonal

bipyramidal

Octahedral

[Hybridization is mixing of orbitals to get hybrid orbitals. Hybridization is not a

phenomenon, but a mathematical manipulation. Hybrid orbitals have no physical existence in

reality]

Examples

Consider CoII complexes in 4- and 6- coordinate complexes:

Co

Co2+

sp3

XX XX XX XX

sp3 hybridization

3d7

4s2 4p

04d

0

- tetrahedaral

dsp2

sp3d

2

d2sp

3

XX XX

XX

XX XXXX XX XX XX

XX XX XX XX XX

XX XX

dsp2 hybridization

sp3d

2 hybridization - octahedral

d2sp

3 hybridization - octahedral

?

*

*

XX Lone pairs of electrons from the ligand.

* Rearranged electrons to create empty hybrid orbitals.

? The remaining electron from the CoII ion promoted to an unknown orbital.

- square planar

Note: Two types of octahedral complexes are apparent: one which uses only 3d orbitals and

the other which makes use of 4d orbitals for bonding. These have been described as

“covalent” or “inner orbital” and “ionic” or “outer orbital” respectively. In the inner orbital

case note the promotion of a single electron to an unspecified outer orbital. This is in

agreement with the case with which Co2+

complexes are oxidized to Co3+

.

Secondly the two forms of 6-coordinate and 4-coordinate complexes differ in the number of

unpaired, non-bonding electrons present. Since the magnetic moments of complexes are

dependent on the number of unpaired electrons, magnetic moments can provide useful

indications of structure and bond type.

Shortcomings of the valence bond theory

Other aspects of magnetic behaviours, such as variation of magnetic moment with

temperature are not explained by VBT.

Does not explain origin of colour in transition metal compounds. This is a major

shortcoming.

Does not explain why some ligands give rise to high-spin and some to low-spin

complexes.

Cannot distinguish two complexes with same number of unpaired electrons but with

different structures.

Examples

1. Discuss bonding in [Co(NH3)6]Cl3 and [CoF6]3-

which are diamagnetic and

paramagneticrespectively.

Co

Co3+

XX XX XX XX

3d 4s 4p 4d

[Co(NH3)6] XX XX

d2sp

3 hybridization; inner orbital octahedral complex; XX = NH3;

lone pair on N

Note that all electrons are paired, hence the complex is diamagnetic.

[CoF6]3-

XX XX XX

sp3d

2 hybridization; outer orbital octahedral complex; XX = F

-

XX XX XX

3d 4s 4p 4d

Note 4 unpaired electrons, paramagnetic. Magnetic moments can therefore be used to

distinguish between the two complexes.

2. Consider [Fe(H2O)6]3+

with 5 unpaired electrons and Fe(CN)6]3-

with one unpaired

electron. Use the valence bond theory to explain these observations.

3. There are three groups of complexes of Ni2+

; one group is diamagnetic and two

are paramagnetic. By considering 4-coordinate and 6-coordinate complexes, account for the

three groups of complexes.

6 Crystal Field Theory

6.1 Octahedral complexes

Crystal field theory (CFT) assumes that bonding in coordination compounds is electrostatic

(i.e. ionic). The d-orbitals fall into two groups: those whose orbitals point along the axes (dx2

-

y2 and dz

2) and those whose orbitals point between axes (dxy, dyz, dxz). The shapes of the d-

orbitals are given (see page 25).

In gaseous metal ion, the 5 orbitals are degenerate (i.e. of the same energy). If a spherically

symmetric field of negative charges is placed around the metal ion, the energy of all the

orbitals will be raised equally. This is as a result of the repulsion between the negative field

and the electrons of the orbitals. When the field results from the influence of real ligands the

symmetry must be less than symmetrical because of the finite number (usually 4 or 6) of

ligands involved. When, for example, six ligands approach to form an octahedral complex,

there are two types of d-orbitals, the dx2

-y2 and dz

2, which have lobes point along the x-, y-,

and z- axes and dxy, dyz, and dxz whose lobes point between axes. If 6 ligands are imagined to

approach the cation along x-, y-, and z- axes until they reach their final equilibrium positions

then as they approach the electrons in the metal d-orbitals will be repelled, i.e. the energy of

the d-orbitals will be increased, but not equally.

6Dq

-4Dq

=10DqBaricentre

eg

t2g

Free ion(gaseous state)

Average energy of d-orbitalsunder spherically symmetric field

Splitting of d-orbital energies in octahedral crystal field

Energyo

o

o3/5

-2/5

The dx2

- y2 and dz

2 orbitals, which point directly towards the ligands will suffer greater

repulsion (i.e. higher energy) than those orbitals which point between axes. The energies of

the d-orbitals split into two sets: two axial and three non-axial. The axial orbitals are referred

to as eg while the non-axial are t2g. Thus the eg set increases in energy while the t2g set

decreases in energy. Their separation is given the symbol Δo or 10Dq and referred to as

crystal field splitting energy (for octahedral field). The energy level splitting in an octahedral

environment is given at page 26.

The separation between eg and t2g is very important because it introduces the possibility of re-

arranging the metal d-electrons. The Hund‟s rule still holds in the sense that d-electrons still

tend to remain unpaired until each orbital is singly occupied before pairing can occur.

However, electrons would tend to occupy the orbitals of lower energy before the eg orbitals

are filled, particularly if the separation energy is large.

6Dq

-4Dq

=10DqBaricentre

eg

t2g

Free ion(gaseous state)

Average energy of d-orbitalsunder spherically symmetric field

Splitting of d-orbital energies in octahedral crystal field

Energyo

o

o3/5

-2/5

For d1 to d

3 the electrons go singly into the t2g orbitals; however, for d

4, d

5, d

6, and d

7 two

options are open, depending on the magnitude of Δo.

If Δo is large, i.e. the crystal field is strong the electrons are forced to pair in the lower

t2g set and the configuration is known as „spin-paired‟ or „low-spin‟.

If Δo is small, i.e. the crystal field is weak, the maximum number of electrons remain

unpaired and the configuration is known as „spin-free‟ or „high-spin‟.

This is summarized below:

d1

d2

d3 d

4d

5

lowspin

highspin

d6

d7

d8 d

9

lowspin

highspin

Note that both low- and high- spin for d9 are the same

These arrangements correspond to the „outer orbital‟ and „inner orbital‟ types of VBT:

Outer orbital High spin

Inner orbital Low spin

Thus the vacant orbital „created‟ to allow for vacant d-orbitals required for hybridization (i.e.

d2sp

3) in VBT is equivalent to the empty eg orbital in low-spin configuration in CFT.

For example, Co2+

octahedral complex in CFT would generate two configurations:

oo

eg

t2g

eg

t2g

Co

2+ d

7 high spin VBT outer orbital Co

2+ d

7 low spin VBT inner orbital

Also note that the electron promoted to unknown orbital in VBT is actually in the eg level in

CFT and that Δo for high-spin complexes is smaller than Δo for low-spin.

Other examples:

Use the crystal field theory to account for

(a) [Fe(H2O)6]3+

having 5 unpaired electrons whereas [Fe(CN)6]3-

has one unpaired

electron.

(b) [Co(NH3)6]3+

is diamagnetic whereas [CoF6]3-

is paramagnetic.

6.2 Formation of tetrahedral complexes

Bonding in tetrahedral complexes occurs essentially the same way as in octahedral complexes

except that electrons approaching between axis (now t2) will be more strongly repelled than

those along axis (now e). The d-orbital splitting is thus an inverted form of the octahedral.

Baricentre

Energy

t

t2

e

-3/5

t2/5

t

The separation energy between e and t2 levels is Δt. Both Δo and Δt are called crystal field

splitting energies for octahedral and tetrahedral geometries respectively. For the same metal,

at the same oxidation state, and the same ligands, Δt -4/9Δo.

The occurrence of high-spin and low-spin configurations for d3, d

4, d

5 and d

6 ions is also

possible in principle but rather rare to have low-spin tetrahedral complexes because of the

small crystal field splitting energy involved.

Reasons why Δt is smaller than Δo

Four ligands are involved in tetrahedral complexes as against six in octahedral.

Larger crystal field is therefore expected for six ligands than for four.

The ligands in tetrahedral complexes are less efficiently directed (between axes) than

they are in octahedral (along axes). The ligands are therefore not aimed at any of the

orbitals but exert a somewhat larger influence on the t2g orbitals than the eg orbitals.

6.3 Formation of square-planar complexes

The square-planar geometry is usually conceived as an octahedral complex where the axial

ligands have been elongated to infinity (i.e. totally removed). The effect of this is to reduce

the energy of the dz2 orbital (which points along the vertical axis). This would have the effect

of lowering the energy of the dxz

dx2

-y2

dxy

dz2

dxz, dyz

Octahedral Square-planar

n+

n+

n+

L

L

L

L

L

L

L

L

L

L

L

L

L L

L L

z

xy

Octahedralcomplex

Removal of axialligands

Square-palnarcomplex

and dyz orbitals of the t2g level; that is the degeneracy of the eg and t2g levels

are removed.

In the square-planar complexes of the d7 ion, Co

2+, and the d

8 ions, Ni

2+, Pd

2+, and Pt

2+,

electrons are forced to pair in the lower four orbitals, leaving the top dx2

-y2 orbital vacant. In

the VBT this orbital is used for bonding. (Remember dsp2 hybridization?)

6.4 The ligand field theory (LFT)

LFT is a modified CFT which allows the possibility of covalent bonding in addition to

electrostatic bonding in the formation of metal complexes. This theory can account, at least

qualitatively, for the crystal field splitting caused by various ligands. Ligands like CO, CN-,

phen, and NO2- which provide the largest crystal fields, are all able to form π bonds with the

central metal ion/atom. This π bonding markedly increases the magnitude of Δo.

The magnitude of Δo is also strongly influenced by the oxidation state of the metal ion and

the type of d electrons present. The higher the oxidation state of the metal ion the larger the

crystal field splitting. For example, [Co(NH3)6]3+

is diamagnetic; [CoF6]3-

is paramagnetic.

Ions with high charge and small sizes can be approached more strongly by ligands than larger

cations of smaller charge.

It is also known that the greatest crystal field splitting is observed in complexes of 5d metals

and decreases from 4d to 3d. Thus Δo([Ir(NH3)6]3+

) > Δo([Rh(NH3)6]3+

) > Δo([Co(NH3)6]3+

)

The arrangement of the ligands in the order of the magnitude of their Δo is known as the

spectrochemical series and follows the order:

Weak field ligands Intermediate field ligands Strong field ligands

I-<Br

-<Cl

-<OH

-<RCO2

-,F

-< Ox

-<H2O<NH3<en< NO2

-<phen<CN

-CO

To account for this order the extreme assumption of pure covalency in the VBT and pure

electrostatic model of CFT must be abandoned for the existence of both modes of bonding in

metal complexes.

Values of Δo for selected metal complexes

Complex Δ/cm-1

Complex Δ/cm-1

[TiF6]3-

[Ti(H2O)6]3+

[V(H2O)6]3+

[V(H2O)6]2+

[CrF6]3+

[Cr(H2O)6]3+

[Cr(H2O)6]2+

[Cr(NH3)6]3+

[Cr(CN)6]3-

[MnF6]2-

[Fe(H2O)6]3+

[Fe(H2O)6]2+

17 000

20 300

17 850

12 400

15 000

17 400

14 100

21 600

26 600

21 800

13 700

9 400

[Fe(ox)3]3-

[Fe(CN)6]3-

[Fe(CN)6]4-

[CoF6]3-

[Co(NH3)6]3+

[Co(NH3)6]2+

[Co(en)3]3+

[Co(H2O)6]3+

[Co(H2O)6]2+

[Ni(H2O)6]2+

[Ni(NH3)6]2+

[Ni(en)3]2+

14 100

35 000

33 800

13 100

22 900

10 200

24 000

18 200

9 300

8 500

10 800

11 500

Some applications of CFT

CFT provides explanations for

Colour in metal complexes;

Magnetic properties of metal complexes;

Determining the structure of complexes.

7 Colour

7.1 Origin of colour

Colour in metal complexes arises from

Transitions between metal-centred orbitals possessing d-character (d-d transitions).

Transitions between metal- and ligand-centred MOs which transfer charge from

metal-to-ligand or ligand-to-metal (charge-transfer transitions).

7.2 Selection rules

Electronic transitions obey the following selection rules:

Spin selection rule: ΔS = 0

Transitions may occur from singlet to singlet, or triplet to triplet states and so on, but a

change in spin multiplicity is forbidden.

Laporte selection rule: There must be a change in parity:

allowed transitions: g u

forbidden transitions: g g and u u

This leads to the selection rule: Δl = ± 1

and thus allowed transitions are s → p, p → d, d→ f;

forbidden transitions are s → s, p → p, d → d, f → f, s → d, p → f, etc.

7.3 d-d transitions

d-d transitions are generally weak with 100 and broad. This is because they are forbidden

transitions. The selection rule Δl = ±1 is broken because for d orbital l = 2; Δl = 0. Colour

is due to transitions between energy levels, e.g. between t2g and eg. Absorption of a particular

quantum of energy (Δ = hυ) causes promotion of an electron from t2g level to eg level. The

value of Δ is such that the energy lies in the visible part of the spectrum.

At 1kJ 83.7 cm-1

; Δo for [Ti(H2O)6]3+ 244 kJ. For this ion Δo = 10Dq = 20 300 cm

-1. Δ

has been measured for most ligands on a wide range of metals. Wavelengths () can be

converted to wavenumbers ( ) by the equation:

1

c=

Absorption bands are described in terms of max corresponding to the absorption Amax. The

unit of is nm while that of wavenumbers is cm-1

. If the electronic spectrum is done in a

solution, the extinction coefficient max of the absorption must also be reported. max indicates

the intensity of the absorption and is related to Amax.

max =Amax

c x l(max in dm

3 mol

-1cm

-1)

Generally the electronic spectra of

d1

, d4, d

6 and d

9 complexes consist of one absorption;

d2, d

3, d

7, and d

8 complexes consist of three absorptions;

d5 complexes consist of a series of very weak, relatively sharp absorptions.

7.4 What does Δ depend on?

It depends on the nature of metal within a row. In the first row transition metals Δo

for M2+

is 10 000 cm-1

and for M3+

17 000 cm-1

. Δo therefore depends on the oxidation

state of the metal ion. Δo for second row transition metals is 50% larger than those of

first row, i.e. Δ increases down a transition metal triad (3rd row TM‟s > 2nd row TM‟s >

1st row TM‟s.

For example: [M(NH3)6]3+

, Ir > Rh > Co

41 000 34 000 23 000 cm-1

Greatly on the nature of the ligand. Δ increases along the spectrochemical series: I-<

Br-< Cl

-< F

-< OH

-< Ox

-< H2O < NH3< en < phen < NO2

-< CN

-

All these facts enable the colours of complexes to be predicted.

[Ti(H2O)6]3+

, 490 nm; [TiCl6]3-

, >490 nm; [Ti(CN)6]3-

, <490 nm.

Hole formalism

d10-n

configuration behaves in CF similarly to dn configuration.

7.5 Charge-transfer transition

There are two types: metal-to-ligand (M→L) oxidation

or ligand-to-metal (M → L) reduction, which occurs usually as s → p or p → d or d → p.

These are allowed transitions and are therefore very intense.

[Ti(H2O)6]3+

, a d1 ion, is purple whereas [Ti(H2O)6]

4+, a d

0 ion, is yellow. The purple colour

in the former is due to d-d transition, whereas the yellow colour in the latter is due to charge-

transfer transition.

8 Thermodynamic effect on crystal field splitting

The separation of d-orbitals into t2g and eg levels translates into making the t2g orbitals more

stable and the eg orbitals less stable than the degenerate orbitals. Consequently a t2g electron

increases the stability of an octahedral complex by 2/5Δo and an eg electron decreases it by

2/5Δo. The net effect of all the d electrons represents the additional stability which may be

thought to accrue because the CF splits the d-orbitals. This is known as Crystal Field

Stabilization Energy, CFSE. This extra stability reflects in experimental values of such

thermodynamic quantities as hydration energies, lattice energies, and standard reduction

potentials. The most complete set of data is the hydration energies of bivalent ions: the

experimental values of ΔHo lie on an irregular double-humped curve.

If the CFSE is, in each case, subtracted from ΔHo, the „corrected‟ values fall very neatly on a

smooth curve, passing through the points for the spherically symmetrical d0, d

5, and d

10 ions.

CFSE, therefore, offers a ready explanation of the Irving-Williams order of stability, MnII<

FeII< Co

II< Ni

II< Cu

II> Zn

II.

8.1 Calculation of CFSE

A total crystal field stabilization energy is calculated by adding up the Dq energies for each

electron after the low-spin or high-spin configuration has been established.

Example: Calculate the CFSE for Cr3+

and Fe3+

ions.

Draw the crystal field energy level diagrams for each and in the case of Fe3+

, draw for both

low-spin and high-spin configurations before calculating the CFSEs.

CFSE = [6Dq (No. of electrons in the eg level)]+[-4Dq (No. of electrons in the t2g

level)]

= Dq [6 (No. of electrons in the eg level)] – [4 (No. of electrons in the t2g level)]

Cr3+

, d3 ion:

-4Dq

+6Dq

CFSE = 6 Dq x 0 -4 Dq x 3 = -12 Dq

-4Dq

+6Dq

-4Dq

+6Dq

Weak field

Strong field For weak field

CFSE = 6 Dq x 2 – 4 Dq x 3 = 12 Dq – 12 Dq = 0 Dq

For strong field

CFSE = 6 Dq x 0 – 4 Dq x 5 = -20 Dq

The lower the CFSE, the more stable the complex. Thus low-spin Fe3+

is more stable than

high-spin Fe3+

. Also low-spin Fe3+

ion is more stable than Cr3+

ion.

9 Jahn - Teller distortion

A regular octahedral environment is the most stable one for a spherically symmetrical metal

ion surrounded by six donor atoms. For metal ions with certain d electron configurations

which are not spherically symmetric, the regular octahedral configuration is not the most

stable. This situation is expressed in Jahn – Teller theorem: Any non-linear molecule that is

in an electronically degenerate state will undergo distortion to lower the symmetry, remove

the degeneracy, and lower the energy. Jahn – Teller distortion occurs strongly in all cases

where the eg level is not uniformly occupied by electron. Much weaker distortion also occurs

if the t2g level is not uniformly occupied.

In the case of eg, any odd electron can occupy either the dz2 or the dx

2-y

2 orbital. If, however,

the complex undergoes distortion the eg level is split and the electron can occupy the lower of

the two orbitals (the dz2 orbital in the case of tetragonal elongation, or the dx

2-y

2 orbital in the

case of tetragonal compression).

In tetragonal elongation the ligands on the z axis move out and therefore interact less with

those orbitals which have a z component; i.e. the dz2, dxz, and dyz, and these orbitals attain

lower energy (i.e. stabilized). Those orbitals without a z component, i.e. dx2

-y2 and dxy, will

be raised a corresponding amount.

Elongation: dz2 attains lower energy

dx2

-y2

dz2

dxy

dxz, dyz1 band

3 possible bands

L

LL

L

L

L

a

e

Bond length a > e

Compression: dx2-y

2 attains lower energy

dx2

-y2

dz2

dxy

dxz, dyz

1 band

3 possible bands

L

L

L

L

L

L

a

e

Bond length a < e

Consequently an ion which is susceptible to Jahn – Teller distortion could produce additional

bands (maximum of three altogether). Some of the additional bands may often appear as

shoulders, depending upon the magnitude of the distortion.

Configurations d4, d

7, and d

9 are the most distorted. In high-spin d

4, low-spin d

7, and d

9 the

eg level is not evenly occupied and are therefore susceptible to strong Jahn – Teller distortion.

Uneven occupation of the t2g orbitals also results in distortion but much weaker than in

uneven occupation of eg orbitals. Thus Ti3+

, d1-

is less distorted than Cu2+

, d9.

Relatively weak distortions are expected for tetrahedral complexes with d3, d

4, d

8, and d

9

configurations. Why?

Other evidences of Jahn – Teller distortion

IR can usually detect weakening of M→L bonds as a result of distortion.

X-ray structure would reveal unusually long or short bond lengths.

10 Stability of coordination compounds

Most measurements of stability are done in aqueous solutions when the complex in question

is formed

by a ligand displacing water from the aquo complex of the metal ion. Metal complexes are

formed in solution by stepwise reaction, and equilibrium constants can be written for each

step. For example,

[Ag(H2O)x]+ [Ag(NH3)(H2O)x-1]

+ + H2O k1

[Ag(NH3)(H2O)x-1]+ + NH3

[Ag(NH3)2(H2O)x-2]+ + H2O k2

For simplicity we can ignore the water molecules that make up the hydration sphere of an

aqueous metal ion. Moreover, the solvent water molecules involved in the reaction are not

included in the equilibrium constants, k1 and k2 above are called stepwise-wise stability

constants.

k1 =[Ag(NH3)

+]

[Ag+][NH3]

; k2 =[Ag(NH3)2

+]

[Ag(NH3)+][NH3]

The larger the value of the constant, the greater the concentration of the complex species at

equilibrium. A second type of equilibrium constant, , called an overall stability constant can

be defined for the reactions above:

=[Ag(NH3)

+]

[Ag+][NH3]

; [Ag(NH3)2

+]

[Ag+][NH3]

2

Since the ks and s describe exactly the same chemical systems, they must be related to each

other:

[Ag(NH3)2

+]

[Ag+][NH3]

2=

[Ag(NH3)2+]

[Ag+][NH3][NH3]

.[Ag(NH3)

+]

[Ag(NH3)+]

[Ag(NH3)2+]

[Ag(NH3)+][NH3]

.[Ag(NH3)

+]

[Ag+][NH3]

=

= k2 . k1

Note that 1 = k1 and 2 = k2 . k1. By similar treatment it can be shown that

n = k1 k2 k3 …..kn.

The numerical value of stability constant describes the relative concentration of species at

equilibrium. Large stability constants indicate that the concentration of a complex is much

greater than the concentration of the components it is made.

A complex is said to be stable if the equilibrium constants describing its formation is large.

10.1 Factors that influence complex stability

The equilibrium constant of a reaction is a measure of heat released in the reaction and the

entropy change during the reaction.

-RT lnk = ΔG = ΔH – TΔS

The larger the heat evolved in a reaction the more stable are the reaction products. The

following factors affect stabilities of complexes:

The metal ion and its charge. For example, M3+

complexes are more stable than M2+

.

Recall: Mn2+

<Fe2+

<Co2+

<Ni2+

<Cu2+

>Zn2+

, which is the reverse of the order for the

cation sizes.

The relationship between metal and donor atoms. Class-a acceptors or hard acids

(metals in Group IA and IIA along with the inner transition metals and the early

members of the first row transition metals) form their most stable complexes with

ligands containing O, N or F donor atoms. Class-b acceptors or soft acids (other

metals form their most stable complexes with ligands containing heavier members of

N, O, and F groups). For example, Rh, Pd, Ag, Ir, Pt, Au and Hg are class-b

acceptors. The remaining transition elements are regarded as borderline.

The types of ligands. Bidentate ligands form more stable complexes than

monodentate ligands due to „chelate effect‟. Chelate effect is stabilized due to

chelation. The chelate rings of 5 and 6 are more stable. The stability of a metal

chelate is greater than that of an analogous non-chelated metal complex. For

example, the stability of [Ni(en)3]2+

is greater than that of [Ni(NH3)6]2+

. The more

extensive the chelation the more stable the system.

Ligand logks Co2+

Ni2+

Cu2+

Zn2+

NH3 logk1

logk2

logk3

logk4

logk5

logk6

2.1

1.6

1.1

0.8

0.2

-0.6

2.8

2.2

1.7

1.2

0.8

0.03

4.2

3.5

2.9

2.1

-0.5

-

2.4

2.4

2.5

2.2

-

-

en logk1

logk2

logk3

6.0

4.8

3.1

7.5

6.3

4.3

10.6

9.1

-1.0

5.7

4.7

1.7

trien logk1

logk2?

8.1 10.7 16.0 8.9

6-en logk1 15.8 19.3 22.4 16.2

trien = diethylenetetramine 6-en = pentaethylenehexamine

NH HN

H2NNH2

= H2NCH2CH2(NHCH2CH2)4CH2CH2NH2

Why is Co2+

<Ni2+

<Cu2+

>Zn2+

?

Why 6 NH3, 3 en, 2 trien and one 6-en?

Why is logk1> logk2> logk3> ……?

Why is logk6 for [Co(NH3)6]2+

negative?

Why is logk5 and logk6 negative or non-existent for [Cu(NH3)6]2+

?

Why is log6 (for NH3 complexes) < log3 (for the en complexes) < log1 (for 6-

encomplexes) in each metal?

For steric reasons large bulky ligands form less stable metal complexes than do analogous

smaller ligands. For example H2NCH2CH2NH2 forms more stable complexes than

(CH3)2NCH2CH2N(CH3)2.

Examples:

Complex s

[Ni(NH3)6]2+

1 2 3 4 5 6

5 x 102

6 x 104

3 x 106

3 x 107 1.3 x 10

8 1.0 x 10

8

[Ni(en)3]2+

1 2 3

5 x 107 1.1 x 10

14 1.4 x 10

18

[Ni(dien)2]2+

1 2

6 x 1010

8 x 1018

dien = H2NCH2CH2NHCH2CH2NH2Ref.: Basolo & Johnson, Coordination Chemistry, p. 89.

For NH3 6s; for en 3s; for dien 2s. Why?

s [Ni(dien)2]2+

> [Ni(en)3]2+

> [Ni(NH3)6]2+

. Why?

The successive stability constants for [Cu(H2O)5NH3]2+

, [Cu(H2O)4(NH3)]2+

,

and[Cu(H2O)3(NH3)3]2+

are x, y, and z respectively. Write equations for each step

andcalculate the overall formation constant for [Cu(H2O)3(NH3)3]2+

.

11. Introduction to Magnetochemistry

11.1 Introduction

Magnetochemistry is the study of the magnetic properties of materials. By "magnetic

properties" refers to whether they will be attracted or repelled by a magnet. In this course we

are interested in the magnetic properties of metal complexes. A study of magnetic properties

of metal complexes could be used to derive the oxidation state of and the arrangement of

electrons in the central metal ion. It can also be used to derive the stereochemistry of the

complex. Magnetochemistry is therefore very informative in the study of coordination

chemistry,

Magnetism arises from moving charges, such as an electric current in a coil of wire. In a

material which does not have a current present, there are still magnetic interactions. Atoms

are made of charged particles (protons and electrons) which are moving constantly.

The processes which create magnetic fields in an atom are:

1. Nuclear spin. Some nuclei, such as a hydrogen atom, have a net spin, which creates a

magnetic field.

2. Electron spin. An electron has two intrinsic spin states (similar to a top spinning)

which we call up and down or alpha and beta.

3. Electron orbital motion. There is a magnetic field due to the electron moving around

the nucleus.

Each of these magnetic fields interacts with one another and with external magnetic fields.

However, some of these interactions are strong and others are negligible.

11.2 Magnetism

Lenz's Law (~1834), states that: when a substance is placed within a magnetic field, H, the

field within the substance, B, differs from H by the induced field, 4I, which is proportional

to the intensity of magnetization, I.

That is:

B = H + 4I ------------------------------------------------ (Eq. 1)

where B is the magnetic field within the substance, H is the applied magnetic field ,

and I is the intensity of magnetisation.

This can also be written as

B/H = 1 + 4I/H,

or B/H = 1 + 4

where B/H is called the magnetic permeability of the material and is the magnetic

susceptibility per unit volume, (I/H)

By definition, in a vacuum is zero, so that B=H.

It is usually more convenient to measure mass (gram) susceptibility, g, which is related to

the volume susceptibility through the density.

g =/ ----------------------------------------------------------- (Eq. 2)

where is the density.

Finally to get our measured quantity on a basis that can be related to atomic properties, we

convert to molar susceptibility

m =g * MW (MW = molecular weight of the sample)

Normal paramagnetic substances obey the Curie Law

= C/T ----------------------------------------------------------------- (Eq. 3)

where C is the Curie constant.

Thus a plot of 1/ versus T should give a straight line of slope 1/C passing through the origin

(0 K).

Many substances give a straight line that intercepts just a little above 0 K and these are said to

obey the Curie-Weiss Law:

= C/(T+) --------------------------------------------------------------- (Eq. 4)

where is known as the Weiss constant.

The constant C is given by the Langevin expression, which relates the susceptibility to the

magnetic moment:

m =N2/3kT ---------------------------------------------------------------- (Eq. 5)

where N is Avogadro number, k is the Boltzmann constant, and T the absolute temperature.

Re-writing Eq, 5 gives the magnetic moment as

= 2.828 √mT = 2.828(m.T)½

-------------------------------------------- (Eq. 6)

Many transition metal salts and complexes are paramagnetic due to partially filled d-

orbitals.

The experimentally measured magnetic moment () (Eq. 6) can provide some important

information about the compounds themselves:

No of unpaired electrons present

Distinction between high-spin and low-spin octahedral complexes

Spectral behaviour, and

Structure of the complexes.

11.3 Sources of Paramagnetism

Orbital motion of the electron generating ORBITAL MAGNETIC MOMENT (l)

Spin motion of the electron (on its own axis) generating SPIN MAGNETIC

MOMENT (s)

If a transition metal ion contains n unpaired electrons (multi-electron system), the total orbital

and total spin motions are, respectively, given by:

L = l1 + l2 + l3 + …………….

S = s1 + s2 + s3 + ……………

where l = orbital angular momentum, s = spin angular momentum

The magnetic moment due to the contribution of orbital and spin motions is given by

L+S = [4S(S+1)+ L(L+1)]½

B.M. ------------------------------------- (Eq. 7)

For TM-complexes, the magnetic properties arise mainly from the exposed d-orbitals. The d-

orbitals are perturbed by ligands.

The rotation of electrons about the nucleus is restricted which leads to L = 0; i.e. the

orbital contribution to magnetic moment is quenched; thus

s = [4S(S+1)]½

B.M. -------------------------------------------- (Eq. 8)

The spin-only formula can be derived as follows:

S = n (1/2) = n/2; n = no of unpaired electrons

Hence

s = [4S(S+1)]½

B.M.

= [4(n/2)(n/2+1)]½ B.M.

= [n(n+2)]½

B.M.

This is called Spin-Only Formula.

11.4 When does orbital angular momentum contribute?

There must be an unfilled/half-filled orbital similar in energy to that of the orbital occupied

by the unpaired electrons. If this is so, the electrons can make use of the available orbitals to

circulate or move around the centre of the complexes and hence generate L and L

Essential Conditions:

1. The orbitals should be degenerate (t2g or eg)

2. The orbitals should be similar in shape and size so that they are transferable into one

another by rotation about the same axis (e.g. dxy is related to dx2-y2 by a rotation of 45o

about the z-axis.

3. Orbitals must not contain electrons of identical spin.

For an octahedral complex:

Condition t2g set eg set

1 Obeyed Obeyed

2 Obeyed Not obeyed

3 Since 1 and 2 are satisfied

condition 3 dictates whether

t2g will generate l or not

Does not matter since

condition 2 is already not

obeyed.

These conditions are fulfilled whenever one or two of the three t2g orbitals contain an odd

no. of electrons.

Exercise: Work-out all possible dn LS and HS cases with orbital contribution.

HS:

LS:

11.5 Other Reasons for Orbital Contribution:

Although normally develops from GS, sometimes ES also may contribute, especially the

GS-ES energy difference is very small

Example:

Take Ni2+

octahedral; d8; GS: t2g

6eg

2 no l

ES: t2g5eg

3l contributes

Similarly,

Take Co2+

tetrahedral; d7 GS: e

4 t2

3nol

ES: e3t2

4l contributes

Therefore, obs>s for both Oh Ni2+

and Td Co2+

.

11.6 Magnetic Properties of lanthanides

4f electrons are too far inside 4fn5s

25p

6

(compared to the d electrons in transition metals)

Thus 4f normally unaffected by surrounding ligands

Hence, the magnetic moments of Ln3+

ions are generally well described from the

coupling of spin and orbital angular momenta ~ Russell-Saunders Coupling to give J

vector

spin orbit coupling constants are large (ca. 1000 cm-1)

ligand field effects are very small (ca. 100 cm-1)

o only ground J-state is populated

o spin-orbit coupling >> ligand field splittings

o magnetism is essentially independent of environment

Magnetic moment of a J-state is expressed by:

J = L+S, L+S-1,……L-S

For the calculation of g value, we use

minimum value of J for the configurations up to half-filled;

i.e. J = L-S for f0-f

7 configurations

maximum value of J for configurations more than half-filled;

i. e. J = L+S for f8-f

14 configurations

For f0, f

7, and f

14, L = 0, hence J becomes S

Lanthanide Ions and their Magnetic Moments

Ion Config. of

Ln3+

Ground

state

No. of

unpaired

electrons

g µ

(Calculated)

µ

(experimental)

La3+

Ce3+

Pr3+

Nd3+

Pm3+

Sm3+

Eu3+

Gd3+

Tb3+

Dy3+

Ho3+

Er3+

Tm3+

Yb3+

Lu3+

4f0

4f1

4f2

4f3

4f4

4f5

4f6

4f7

4f8

4f9

4f10

4f11

4f12

4f13

4f14

1S0

2F5/2

3H4

4I9/2

5I4

6H5/2

7F0

8S7/2

7F6

6H15/2

5I8

4I15/2

3H6

2S7/2

1F0

0

1

2

3

4

5

6

7

6

5

4

3

2

1

0

-

6/7

4/5

8/11

3/5

2/7

-

2

3/2

4/3

5/4

6/5

7/6

8/7

-

0

2.54

3.58

3.62

2.68

0.34

0

7.94

9.72

10.63

10.60

9.57

7.63

4.50

0

0

2.3 -2.5

3.4 – 3.6

3.5 – 3.6

-

1.5 – 1.6

3.4 – 3.6

7.8 - 8.0

9.4 - 9.6

10.4 – 10.5

10.3 – 10.5

9.4 – 9.6

7.1 – 7.4

4.4 – 4.9

0

Sample Landè Calculation for a Ln3+

ion

e.g. Pr3+

[Xe]4f2; find Ground State from Hund's Rules

S = 1/2 + 1/2 = 1 L = 3 + 2 = 5

J = 6, 5, 4; J =4 is chosen for f2

g = (3/2) + [1(1+1)-5(5+1)/2(4)(4+1)] = 0.8

J = 3.577 B.M. Experiment = 3.4 - 3.6 B.M

Landé formula fits well with observed magnetic moments for all but SmIII

and EuIII

ions.

Moments of these ions are altered from the Landé expression by temperature-dependent

population of low-lying excited J-state(s).

Use of Ln3+

Magnetic Moments

NMR Shift Reagents - paramagnetism of lanthanide ions is utilized to spread resonances in 1H NMR of organic molecules that coordinate to lanthanides

11.7 Magnetic States of Matter

Diamagnet - A diamagnetic compound has all of its electron spins paired giving a net spin of

zero. Diamagnetic compounds are weakly repelled by a magnet.

Paramagnet - A paramagnetic compound will have some electrons with unpaired spins.

Paramagnetic compounds are attracted by a magnet. Paramagnetism derives from the spin

and orbital angular momenta of electrons. This type of magnetism occurs only in compounds

containing unpaired electrons, as the spin and orbital angular momenta is cancelled out when

the electrons exist in pairs.

Compounds in which the paramagnetic centres are separated by diamagnetic atoms within

the sample are said to be magnetically dilute.

If the diamagnetic atoms are removed from the system then the paramagnetic centres interact

with each other. This interaction leads to ferromagnetism (in the case where the

neighbouring magnetic dipoles are aligned in the same direction) and antiferromagnetism

(where the neighbouring magnetic dipoles are aligned in alternate directions).

These two forms of paramagnetism show characteristic variations of the magnetic

susceptibility with temperature.

In the case of ferromagnetism, above the Curie point the material displays "normal"

paramagnetic behaviour. Below the Curie point the material displays strong magnetic

properties. Ferromagnetism is commonly found in compounds containing iron and in alloys.

For antiferromagnetism, above the Neel point the material displays "normal" paramagnetic

behaviour. Below the Neel point the material displays weak magnetic properties which at

lower and lower temperatures can become essentially diamagnetic. Antiferromagnetism is

more common and is found to occur in transition metal halides and oxides, such as TiCl3 and

VCl2.

A worked example

Account for the magnetic moments of (Et4N)2[NiCl4] recorded at 80 and 300 K.

80K 300K

3.25 3.89 B.M.

Ni2+

is a d8 metal ion.

The formula suggests a 4 coordinate complex and we can assume that the complex is

tetrahedral with a d electron configuration of e4 t2

4 therefore the spin-only magnetic moment

can be calculated as 2.83 B.M.

Why did we ignore the possibility of it being square-planar?

The free ion Russell-Saunders ground term is 3F (L=3 and S=1) which will give rise to a

lowest energy T term in a tetrahedral field and hence the resultant magnetic moment is

expected to be temperature dependent and have a direct orbital contribution.

The observed values may be quite different then to the calculated spin only magnetic

moment.

The value of S+L can be calculated as: S+L = [4S(S+1)+L(L+1)]

= (8+12) = 20 = 4.47 B.M.

If you use the spin-only formula s = n(n+2) = 8 = 2.8 BM

Now go back and check above the observed magnetic moments at the given temperatures.

What do you conclude?

From the observed values it can be seen that the magnetic moment of the d8 Ni

2+ complex is

intermediate between the s and S+L values (probably due to partial quenching of the orbital

angular momentum contribution) and is dependent on temperature.

12. LANTHANIDES AND ACTINIDES

Introduction

You would recall that elements in the Periodic Table could be classified into blocks: s-, p-, d-

and f-blocks. The d- and f-blocks are also called transition elements, while the term

transition metals are generally retained for the d-block metals, the f-block is generally

referred to as inner transition elements, which appears firstly after lanthanum and secondly

after actinium, in the Periodic Table. The elements from cerium to lutetium are known as

lanthanides, after lanthanum. Because of its chemical similarity to these elements, lanthanum

is usually included with the lanthanides. Scandium and yttrium also share similar chemical

properties with the lanthanides and are both included when discussing the chemistry of the

lanthanides.

The second series of f-block starts from thorium to lawrencium and are called actinide series.

For the same reason as in lanthanum, actinium is also considered with the actinides series

when discussing their chemistry.

The transition elements have similar characteristic properties including the following:

Solid (metals) except mercury, which is a liquid.

Conductors of electricity and heat

Form alloys with one another and with metallic main group elements.

High melting and boiling, except mercury.

Mostly reactive with mineral acids to form salts (some are inert to conc. oxidizing

acids

Display variable oxidation states, accessible by simple laboratory reactions.

However, there are exceptions in the lanthanides, where the chemistry is dominated

by the Ln3+

ions.

Like the d-block, the early members of the actinide series (protactinium to americium)

display variable oxidation states but the latter members behave like the lanthanides,

dominated by An3+

ions.

The presence of partially filled d and f orbitals in the f-block elements often results in

some transition element in having odd number of electrons that make them t be

paramagnetic.

Compounds are often coloured due to the presence of partially –filled d orbitals in the

d transition metals, and d or f in f transition elements.

12.1. The Lanthanides

General introduction

Position in the Periodic Table: Z = 58 – 71

Electron configuration

The electron configuration can be written as 4f2-14

5s25p

65d

0-16s

2 or [Xe]4f

2-145d

66s

2.

The third outer shell, (4f orbital) is being filled with electrons, while the number of electrons

in the outer shell (6s2) and in the penultimate (5s

25p

6) shell of most of the lanthanides is

identical. Even though the energies of the 4f and 5d are similar, the 4f states predominate in

the chemistry of the lanthanides.

Table1. Electronic configuration, oxidation states, radii and colour of the lanthanides

Element Name Atomic

No.

Abundan

ce/ppm

Electron.

Config.

Oxidn.

States

Radius

(pm)*

M3+

Colour

M3+

La

Ce

Pr

Nd

Pm

Sm

Eu

Gd

Tb

Dy

Ho

Er

Tm

Yb

Lu

Lanthanum

Cerium

Praseodymium

Neodymium

Promethium

Samarium

Europium

Gadolinium

Terbium

Dysprosium

Holmium

Erbium

Thulium

Ytterbium

Lutetium

57

58

59

60

61

62

63

64

65

66

67

68

69

70

71

66

9.1

40

4.5 x 10-

20

7.0

2.1

6.1

1.2

4.5

1.4

3.5

0.5

3.1

0.8

[Xe]5d16s

2

[Xe]4f15d

16s

2

[Xe]4f36s

2

[Xe]4f46s

2

[Xe]4f56s

2

[Xe]4f66s

2

[Xe]4f76s

2

[Xe]4f75d

16s

2

[Xe]4f96s

2

[Xe]4f10

6s2

[Xe]4f11

6s2

[Xe]4f12

6s2

[Xe]4f13

6s2

[Xe]4f14

6s2

[Xe]4f14

5d16s

2

3

3, 4

3, 4

3

3

2, 3

2, 3

3

3, 4

3

3

3

3

2, 3

3

116

114.3

112.6

110.9

109.3

107.9

106.6

105.3

104

102.7

101.5

100.4

99.4

98.5

97.7

Colourless

Colourless

Green

Lilac

Pink

Yellow

Pale pink

Colourless

Pale pink

Yellow

Yellow

Lilac

Green

Colourless

Colourless

Sc

Y

Scandium

Yttrium

21

39

[Ar]3d14s

2

[Kr]4d15s

2

3

3

74.5**

90**

Colourless

Colourless

* For coordination number 8.

** For coordination number 6.

Consequently the 5d1 electron is transferred to the 4f state in their atoms in contrast to

lanthanum, where the single electron remains as a d electron. The d electron is also not

transferred in the Gd atom where it is preceded by a stable f7 configuration.

The exceptions are:

Cerium, for which the sudden contraction and reduction in energy of the 4f orbitals

immediately after La is not yet sufficient to avoid occupancy of the 5d orbital.

Gd, which reflects the stability of the half-filled 4f orbital.

Lu, at which point the shell is fully filled.

This, however, only has marked effect on the aqueous chemistry of cerium, which is

otherwise dominated by +3 oxidation state. At +3 state the configuration varies regularly

from 4f1 (Ce

III) to 4f

14 (Lu

III).

The elements can be subdivided into two sub-families according to the filling of the 4f

orbitals:

The first seven (Ce-Gd) in which each of 4f orbital is filled by one electron are called

the cerium sub-family

The other seven (Tb-Lu) in which the 4f orbitals are filled by a second electron are

called terbium sub-family.

In Gd and Lu, as in La, the extra electron over the stable f7 and f

14 configurations respectively

is located on the 5d state.

The f electrons have no significant effect on the chemical properties of most lanthanides,

except in few instances where, when slightly excited, one (seldomly 2) electron is transferred

to the 5d state. Their properties are therefore mostly determined by the 5d16s

2 electrons.

Consequently they have great similarity to Group III d-elements (Sc, Y, and La). The

greatest similarity to the lanthanides is in Y and La whose atomic and ionic radii are close to

those of the lanthanides.

As evident above, the difference in electron structure on the atoms is in the third outermost

shell, which does not greatly affect their chemical properties; all the lanthanides are

extremely alike. Because of this special similarity of properties the lanthanides are classed

with lanthanum, yttrium and scandium.

Despite their similarity, the elements have their differences.

Their uniform properties are due to lanthanide contraction, which accounts for the gradual

decrease in atomic and ionic sizes in the Ce-Lu series. For any particular property in which

4fn configuration is maintained across the series a regular variation is observed. However,

variations in properties where the 4fn configuration is not maintained are highly irregular.

12.2. Lanthanide contraction

Lanthanide contraction consists of a significant and steady decrease in size of atoms and ions

with increasing atomic number. This is due to the poor shielding of the 4f and 5f electrons,

as a result of the shapes of the orbitals, which results in a steady increase in effective nuclear

charge and consequent reduction in size of atoms and ions as one proceeds from La to Lu.

The successive shrinkages is called the lanthanide contraction. This has remarkable effect on

the radii of subsequent elements, which are usually smaller than expected. For example ZrIV

and HfIV

have almost the same radii despite the atomic number of 40 and 72 respectively.

The importance/consequences of lanthanide contraction

The normal +3 oxidation state resembles one another much more closely than

members of any series of elements and the separation of some of the lanthanides were

used to be very difficult.

The reduction in size from one LnIII

to the next makes their separation possible, but

the smallness and regularity of the reduction makes the separation difficult.

By the time Ho is reached the LnIII

radius has been sufficiently reduced to be almost

identical with that of YIII

which is why this much larger element is always found

associated with the heavier lanthanides.

As mentioned when we discussed the second and third row transition metals, the

comparable size of metals in the second row and the corresponding ones in the third

row of transition metals is due to lanthanide contraction.

Lanthanide contraction and actinide contraction

The actinide contraction initially parallels that of lanthanides; however, the

contraction is smaller than expected from curium on. This is probably due to poorer

shielding by 5f electrons in the elements.

The lanthanide curve consists of two very shallow arcs with a discontinuity at the

spherically symmetrical Gd3+

(4f7) ion. A similar discontinuity is not observed at

Cm3+

.

The 4f electrons appear to be more deeply buried within the atom that they are not affected

by the environment to any great degree. On the other hand the 5f electrons, at least in the

earlier elements of the series, Th to Bk, are available for bondng, allowing oxidation state up

to +7. These elements therefore resemble d electrons of the transition metals.

12.3. Names:

Rare earths, used originally to describe almost any naturally occurring but unfamiliar oxide,

reflecting the difficulty encountered in separating the elements.

Lanthanons or lanthanoids, which arises from their relationship to lanthanum, element

57,(Group IIIA).

The elements are far from being rare, however, and hence the preference for lanthanides,

lanthanons or lanthanoids. We will adopt lanthanides in this course, with general symbol Ln

to refer to the 14 elements Ce – Lu, inclusive.

Both lanthanides and actinides are also collectively known as inner transition metals. La

and Ac, strictly Group III elements, are classified with lanthanides and actinides respectively

because they share similar properties with lanthanides and actinides respectively.

The lanthanides comprise the largest naturally-occurring group in the Periodic Table.

“Yttria” was isolated in 1794 by J Gadolin. The discovery of the lanthanides is summarized

in Table 2.

12.4 Terrestrial abundance and distribution

The only rare of the lanthanides is the unstable 147

Pm (promethium), t½, 2.62 y), which occurs

as traces in uranium ores. Cerium is the 26th

most abundant of all elements. It is half as

abundant as chlorine and 5 times as abundant as lead. Thulium, the rarest after Pm, is more

abundant than iodine and as common as bismuth. Although there are over 100 minerals

known to contain lanthanides, the only two of commercial importence are monazite (a

mixture La, Th, Ln orthophosphate), and bastnaesite (a La, Ln fluorocarbonate, MIII

CO3F).

Monazite is widely but sparsely distributed in many rocks but, because of its high density and

inertness, it is concentrated by weathering into sands on beaches and river beds. Ilmenite

(FeTiO3) and cassiterite (SnO2) are often found with monazite. La, Ce, Pr and Nd make up

about 90% of monazite with Y and the heavier elements taking about 10%.

Monazite and other minerals carrying lanthanides in the +3 state are usually poor in Eu,

which has a relatively strong tendency to assume the +2 state and therefore found with the

calcium group of minerals.

12.5. Preparation of the elements

The ores are dressed to yield minerals of better than 90% purity. They are then broken down

(“opened”) by either acidic or alkaline attack. Variations can be introduced depending on the

extent to which the metals are to be separated from each other.

Separations are enhanced by

Different solubility of Ln2(SO4)3.Na2SO4.xH2O for the light and heavy lanthanides;

Low solubility of the hydrous oxide of thorium.

Because only small amounts of thorium and heavy lanthanides are present in bastnaesite, its

chemical treatment is less complex.

Acid opening

Monazite

Digest with conc. H2SO4

at 200 o

C. Extract with waterInsoluble residues

Solution of sulphates of Ln, La, Th

Precipitate of crude basic salts of Th

Partial neutralisation with NH3 (aq)

Solution of (Ln/La)(SO4)3

Light lanthanides ppt as (ln/La)2(SO4)3 Na2SO4.xH2O

Solution of sulphates of heavy lanthanides

Na2SO4

Alkali opening

Monazite

Digest with 73% NaOH at140

oC. Extract with water

Residue of crude ThO2

Add to boiling HCl until pH 3.5

Solution of impure (Ln/La)Cl3

Solution of (Ln/La)Cl3

3BaCl2 + Ln2(SO4)3, i.e.stoichiometric amounts

Slurry of impure hydrous oxides

Precipitate of BaSO4 carrying down RaSO4

12.5.1 Separation of individual elements

By change of oxidation state

A few lanthanides have oxidation states other than the characteristic +3, and the difference in

properties of the +2 and +4 states makes separation fairly easy.

Ce is removed by oxidizing the lanthanon solution with MnO4- or bromate in alkaline

medium, whereby Ce3+

is oxidized to CeIV

. Being less basic than LnIII

, CeIV

is more

easily hydrolyzed and precipitated as a basic salt or as the hydroxide (or hydrated

CeO2), leaving the other LnIII

ions in solution.

Eu can be obtained in the +2 state either by electrolytic reduction with a mercury

cathode or by using zinc amalgam followed by precipitation as EnSO4.

These two metals are first removed from the Ln3+

solution before embarking on separating

the others. The other metals are difficult to separate because of the negligible difference in

their ionic sizes as the series is traversed.

Ion-exchange technique

Separation by this technique is carried out commercially on a large scale.

The ion-exchange behaviour depends primarily on the hydrated ionic radius. The smaller the

hydrated size of the ion the more tightly it bounds to ligands. La has the largest ionic size

and Lu the smallest. La therefore has the smallest hydrated ion and Lu the largest hence the

order of elution is Lu → La. The separation is effected by use of appropriate complexing

agents at appropriate pHs. The ion of smallest radius also forms the strongest complexes and

hence the preference for the aqueous phase is enhanced. One of the best complexing agents

is α-hydroxyisobutyric acid, (CH3)2CH(OH)COOH, at 25 o and pH 3.55, but EDTA and other

hydroxo or amino carboxylic acids can also be used.

A typical ion-exchange resin is sulphonated polystyrene, which may be denoted by H-Resin,

since it corresponds to an insoluble strong acid. A solution containing the lanthanides is

applied, and the acid formed is washed down,

3 H-Resin + Lnaq3+

Ln(Resin)3 + 3 H3O+

The column is eluted with a weakly acidic solution of the complexing agent, e.g. citric acid

(H-Cit), buffered with ammonia to a constant pH = 5, when the new equilibrium,

Ln(Resin)3 + 3 H-Cit 3 H-Resin + LnCit3

is set up. As the buffered citric acid flows down the column, the concentration of the

lanthanide ions changes and the equilibrium reverses several times. The heavier ions (smaller

hydrated ions) are more strongly complexed by the citrate ion and so will tend to spend more

times in solution and less on the resin. Because the complexes possess a lower positive

charge than the initial Ln3+

, they are less tightly held by the resin than Ln3+

and are washed

down the column first and will eventually be eluted. The elements are usually dilute. They

are concentrated by precipitation of the oxalate and the exchange repeated to give pure

samples.

Other methods of separation

Precipitation

The substance with the lowest solubility product is precipitated most rapidly and most

completely. Addition of OH- ions to the aqueous solution of the lanthanide nitrates results in

the precipitation of the hydroxides, the weakest base, Lu(OH)3 being precipitated first and the

strongest base La(OH)3 last. Only partial separation is achieved using this method.

Thermal reaction

On heating the nitrates there is a temperature when the least basic would be converted to the

oxide. The mixture is leached with water when the insoluble oxide remains. The oxide is

then converted to the nitrates again when the process is repeated.

Fractional crystallization of their simple salts, e.g. nitrates, sulfates, bromates, perchlorates

and oxalates, including some double salts of the type 2L(NO3)3.3Mg(NO3)2.24H2O.

Solubility decreases from La to Lu. The re-crystallization process is repeated several times.

Complex formation

The oxalates of the lanthanons are insoluble, but they can be held in solution by a chelating

agent, e.g. EDTA. The stability of the M-EDTA complexes varies. The metal-complexes are

destroyed in order of increasing stability by addition of acid which then precipitate as

insoluble oxalates. Separation is not complete hence the oxalates are re-dissolved and the

process is repeated several times.

Solvent extraction

The ratios of the partition coefficients of La(NO3)3 and Gd(NO3)3 between a solution of the

metal ions in strong HNO3 and tributylphosphate is 1:1.06. Though this difference is small a

very large number of partition can be performed using a continuous counter – current

apparatus. Kg quantities of 95% pure Gd have been obtain by this method.

Change in oxidation state

As mentioned earlier a few lanthanons have oxidation states other than the characteristic +3

and the difference in properties of the +2 and +4 states makes separation fairly easy. For

example Ce can be removed form lanthanide mixtures by oxidizing the solution with MnO4-

or bromate in alkaline medium. CeIV

is smaller and has higher charge and less basic than

Ce3+

. It is therefore precipitated as Ce(OH)4, CeO2, or a basic salt, leaving the M3+

ions in

solution.

Alternatively Ce3+

can be readily extracted from the other M3+

lanthanides in HNO3 solution

using tributylphosphate. 99% pure Ce can be obtained in one stage from a mixture containing

40% Ce.

Europium can be obtained in the +2 state either by electrolytic reduction with a mercury

cathode or zinc amalgam, followed by the precipitation of EuSO4.

Table 2. The discovery of the oxides of Group IIIA and the Lanthanide elements (Earnshaw

& Greenwood, p 1229)

Element Discoverer Date Origin of name

Ce Cerium

La Lanthanum

Pr Praeseodymium

Nd Neodymium

Sm Samarium

Eu Europium

Y Yttrium

Tb Terbium

Er Erbium

Yb Ytterbium

C G Mosander

C G Mosander

C A von Welsbach

C A von Welsbach

L de Boisbaudran

E A Demarcay

C G Mosander

C G Mosander

C G Mosander

J C G de Marignac

1839

1839

1885

1885

1879

1901

1843

1843

1843

1878

The steroid, Ceres

Greek lanthanein, to escape

notice.

Greek praseos + didymos, leek

green + twin or green twin (from

its green salts)

Greek neos + didymos, new twin

The mineral, samarskite

Europe

Ytterby, Sweden

Ytterby, Sweden

Ytterby, Sweden

Ytterby,Sweden

Sc Scandium

Ho Holmium

Tm Thulium

Gd Gadolinium

Dy Dysprosium

Lu Lutetium

Pm Promethium

L F Nilson

P T Cleve

P T Cleve

J C G de Marignac

L de Boisbaudran

G Urban

C A von Welsbach

C James

1879

1879

1879

1880

1886

1907

1947

Scandinavia

Latin Holmia: Stockholm

Latin Thule, “most northerly

land

Finnish chemist, J Gadolin

Greek dysprositos, hard to get.

Latin Lutetia: Paris

Greek, dysprositos, hard to find

Latin Lutetia: Paris

After Prometheus, the only Ln

that has never been found in

nature.

12.6. Chemical reactivity

The lanthanides are remarkably similar in chemical properties because of their near identical

configurations of 6s2+

+ an additional lightly held electron in either 5d or 4f orbital (the

energies of the 5d and the 4f orbitals are very nearly equal for te lanthanides of low atomic

numbers).

The elements are highly electropositive with M3+

/M potential varying from -2.25 V (Lu) to -

2.52 V (La). The chemistry is therefore predominantly ionic and of M3+

ions.

They are more reactive than the d-transition metals and are therefore closer to alkali or

alkaline earth metals than to most of the transition metals.

They all react with water with evolution of hydrogen

Ionic compounds are common and the coordination chemistry is quite different from, and less

extensive than, that of the d-transition metals.

Coordination numbers are generally high and stereochemistries are frequently ill-defined and

the complexes are labile.

Only strongly chelating ligands yield products that can be isolated from aqueous solution in

association with coordinated water.

Coordination numbers below 6, considered to be unusual, are found only with very bulky

ligands. The complexes therefore are prepared in non-aqueous systems.

Coordination numbers of 10 and over require chelating ligands with small “bites”, such as

NO3- or SO4

2-. Such complexes are found in large, lighter lanthanides.

The chemical properties of the lanthanides, however, differ

Because the successive addition of electrons to the 4f (or 5d) orbitals of the atoms on the

aufbau principle while maintaining the electrical neutrality of the atom, causes variations in

the distribution of the negative charge and

Because the effective nuclear charge (Zeff)ion of the M3+

increases steadily with increasing

atomic number and hence the electronegativities of the ions increase, and the radii of the ions

decrease.

The effect of ionic size on properties of compounds is noticeable as one traverses the series.

Salts become somewhat less ionic as the Ln3+

radius decreases across the series. The

hydroxides become less ionic and therefore become less basic. At the end of the series

Yb(OH)3 and Lu(OH)3, though mainly basic, can be made to dissolve with difficulty in hot

conc. NaOH.

The hydrated ions, [Ln(H2O)x]3+

, become increasingly susceptible to hydrolysis as one goes

across the series and need to be stabilized by addition of acids to their solutions.

There are no consistent trends noticeable in their aqueous or non-aqueous solutions. Some

distinctions can, however, be made between the cerium sub-family and the terbium sub-

family. The oxalates, double sulphates, and double nitrates of the cerium sub-family are

rather less soluble and the basic nitrates are more soluble than those of the terbium sub-

family.

Other oxidation states of +2 and +4 occur in Eu2+

and CeIV

. The two oxidation states are

stable in water and, even though they are strongly reducing and strongly oxidizing agents

respectively, have well established aqueous chemistry. LnIV

(Ln = Pr, Tb) and LnII (Ln = Nd,

Sm, Dy, Tm, and Yb) are also known in the solid state but are not stable in water.

Europium and ytterbium are particularly similar to the alkaline earth elements. They have the

lowest enthalpies of vaporization and the largest atomic radii of the lanthanides, more similar

to barium than to typical lanthanides.

One difference between the lanthanides and the transition metals lies in the sum of the first

three ionization energies which is 3 500 - 4 200 kJ mol-1

for the lanthanides compared to 5

230 kJ mol-1

for Cr3+

and 5 630 kJ mol-1

for Co3+

. The heat of atomization necessary to break

up the metal lattice is higher in the d transition series (because the d electrons are available

for bonding) than in the alkali, alkaline earth, and lanthanide metals.

12.6.1. Oxidation states

Since the 5d16s

2 electrons are mainly valence electrons in the lanthanides their most stable

oxidation state is +3. The elements adjacent to lanthanum (4f0), gadolinium (4f

7) and

lutenium (4f14

), however, have variable oxidation states. Thus in addition to +3 state, Ce has

+4 state arising from the transition of the 4f2 electrons to the 5d level; praseodymium with

4f36s

2 may also display +4 (by promoting 2 electrons to the d orbital) oxidation state although

this is less characteristic than in Ce. Eu with configuration 4f76s

2 can also have oxidation

state of +2.

Table 3. Oxidation states

La

+3

Ce

+3, +4

Pr

+3, +4

Nd

+3

Pm

+3

Sm

+3, +2

Eu

+3, +2

Ga

+3

Tb

+3, +4

Dy

+3, +4

Ho

+3

Er

+3

Tm

+3, (+2)

Yb

+3, +2

Lu

+3

Summary of the some basic reactions of the lanthanides

Ln

H2300 - 400

0C

X2

HeatLnX3

LnH2; LnH3

N2 (C, Si, P, S, and other non metals)

LnN

Heat

H2O

Ln(OH)3 + H2

O2

Heat

Ln2O3

(but CeO2

Ln2(CO3)3 + H2

H2O + CO2

Ln = Lanthanides

X = Halides

Reactions of Lanthanides

Complexes

The lanthanide ions have high charge but they are rather large (85 -100 pm) compared with

the normal transition elements (Cr3+

= 60, Fe3+

= 64 pm) and consequently they do not form

complexes very readily. The most common and stable complexes are those with chelating

oxygen ligands such as citric acid, oxalic acid, EDTA, and acetyacetone. These complexes

have water/solvent molecules attached to the central metal ion, and coordination numbers 7,

8, and 9 are very common. A variety of stereochemistries and coordination numbers are

found, e.g.,

Coordination number 7: [Er(NCS)6]3-

octahedral

Coordination number 10 [Ce(NO3)5]2-

trigonal bipyramidal; each NO3- is bidentate.

Complexes with monodentate ligands are much less stable than the chelates and undergo

dissociation easily in aqueous solution. Only en and NCS- complexes of N-donor ligands are

known and they are readily decomposed by water.

Ce4+

is smaller and has greater charge density. [Ce(NO3)6]2-

is formed in non-aqueous solvent

N2O4. It is 12 coordinate.

They have no complexes with π-bonding ligands, and this is attributable to the non-

availability of the f-orbital for bonding.

However, the existence of high coordination numbers would suggest that it is either that the f-

orbitals are involved in some bonding or that bond orders are less than one since involvement

of s, p, d orbitals would give a maximum of coordination number of 9.

12.7. Spectral and magnetic properties

Colour

The lanthanide ions that have unpaired electrons are coloured and are paramagnetic. Note

from Table 1 that the colour of 4fn ≈ 4f

14-n.

There is a fundamental difference between the spectra of the f-elements and those of the d-

elements. The difference arises from the fact that the 4f electrons are effectively shielded

from the influence of the external forces by the overlap of 5s2 and 5p

6 orbitals whereas the d-

electrons of the d transition metals are exposed directly to the influence of neighbouring

groups.

Consequently the 4fn configurations are only slightly affected by the chemical environments

of the ions and remain practically invariant for a given ion in all of its compounds.

On the other hand the free ion ground term of the d metals ions are subjected to the influence

of the chemical environment. It is therefore subjected to the effect of crystal field first before

any spin coupling comes into play.

As a result, electronic transitions between f orbitals give rise to extremely narrow absorption

bands, quite unlike the broad bands resulting from d-d transitions, and the magnetic

properties of the ions are little affected by their chemical surroundings.

Absorptions are observed in the visible or near UV regions of the spectrum except La3+

with

no f electrons and Lu3+

with no empty f orbitals.

Neither of the two ions shows any absorption bands in the UV/visible region.

The colours observed are due to transitions between f levels, f-f transitions.

Since the f-levels lie deep enough in the atom to be shielded from much perturbation by the

environment these transitions appear in the visible and near UV as sharp bands.

The sharp bands are useful for characterizing the lanthanides and for quantitative estimations.

The bands are weak because they are forbidden and the transitions only result in electron re-

distributions within the f orbitals.

The positions of the bands shift as the configurations changes giving rise to the visible

colours of the different ions given in Table 1.

Although Ce4+

is intensely coloured, the colour is not due to f-f transitions but to the charge

transfer transition between the ion and coordinated ligands.

Paramagnetism

In all cases where the f orbital is partially occupied the compounds are paramagnetic. These

elements differ from the d-block elements in that their magnetic moments do not obey the

simple spin-only formula:

µ = √n(n+2) B.M. 1

n = number of unpaired electrons or

µ = g√s(s+1) 2

s is the absolute value of the spin quantum number and g is a constant called the

gyromagnetic ratio ≈ 2.00.

The magnetic effect arising from the motion of the electron in its orbital (orbital contribution)

as well as that arising from the electrons spinning on its axis contribute to the paramagnetism

of the lanthanides (in the normal transition metal this orbital contribution, to a first

approximation, is usually quenched out by interaction with the electric fields of the

environment).

The magnetic moments of the lanthanides calculated by incorporating the orbital contribution

is given by:

µ = g√J(J+1) 4

g = 1 +S(S + 1) - L(L+1) + J(J+1)

2J(J+1)where

5

3

2 2J(J+1)

+ S(S+1) - L(L+1)=

The orbital moment is given by

µ = √L(L+1) 6

and the spin moment is given in equation 2.

The moment, taking care of orbital contribution is given by

µ = √4S(S+1) + L(L+1) 7

If the ground states of transition metal ions or lanthanides ions are S states (L = 0), there is no

orbital contribution, hence

µ = √4S(S+1) 8

The calculated magnetic moments of the lanthanide ions using equation 5 are given in Table

4.

Example: Calculate the magnetic moment of Ce3+

given that the ground term is 2F5/2

Ce3+

has configuration 4f1, S = ½, L = 3, J = L - S = 5/2

From equation 5 calculate g = 6/7

µ = 6/7√5/2(7/2) = 6/7√35/4 = 18/7

Table 4. Magnetic moments of lanthanide ions (Ln3+

)

Ion Ground state g µ (Calculated) µ

(experimental)

La3+

Ce3+

Pr3+

Nd3+

Pm3+

Sm3+

Eu3+

Gd3+

Tb3+

Dy3+

Ho3+

Er3+

Tm3+

Yb3+

Lu3+

1S0

2F5/2

3H4

4I9/2

5I4

6H5/2

7F0

8S7/2

7F6

6H15/2

5I8

4I15/2

3H6

2S7/2

1F0

-

6/7

4/5

8/11

3/5

2/7

-

2

3/2

4/3

5/4

6/5

7/6

8/7

-

0

2.54

3.58

3.62

2.68

0.34

0

7.94

9.72

10.63

10.60

9.57

7.63

4.50

0

0

2.3 -2.5

3.4 – 3.6

3.5 – 3.6

-

1.5 – 1.6

3.4 – 3.6

7.8 - 8.0

9.4 - 9.6

10.4 – 10.5

10.3 – 10.5

9.4 – 9.6

7.1 – 7.4

4.4 – 4.9

0

Note that except Sm3+

and Eu2+

the calculated values agree with experimental values.

Russell Saunders coupling

Spin-orbit coupling, which gives rise to a resultant angular momentum associated with an

overall quantum number J is much larger in the lanthanides than the crystal field and the

effect must be considered first.

J can take values J = L+S, L+S-1, ---L-S (or S-L if S<L), each corresponding to a different

energy, so that a term (defined by a pair of L and S values) is said to be split into a number of

component states (each defined by the same S and L values plus a value of J). The ground

state of the ion is that with J = L – S (or S – L) if the f shell is less than half-filled, and that

with J = L + S if the f shell is more than half filled. It is indicated simply by adding this

value of J as a subscript to the symbol for the ground term.

13. ACTINIDES

These are elements 90 – 103, corresponding to the filling of 5f orbitals.

Atomic

No.

Element Symbol Configuration Oxidation state

89

90

91

92

93

94

95

96

97

98

99

100

101

102

103

Actinium

Thorium

Protactinium

Uranium

Neptunium

Plutonium

Americium

Curium

Berkelium

Californium

Einsteinium

Fermium

Mendelevium

Nobelium

Lawrencium

Ac

Th

Pa

U

Np

Pu

Am

Cm

Bk

Cf

Es

Fm

Md

No

Lr

[Rn]6d17s

2

[Rn]6d27s

2

[Rn]5f26d

17s

2 or

5f16d

2

[Rn]5f36d

17s

2

[Rn]5f56d

07s

2

[Rn]5f66d

07s

2

[Rn]5f76d

07s

2

[Rn]5f76d

17s

2

[Rn]5f86d

17s

2 or

5f97s

2

[Rn]5f10

6d07s

2

[Rn]5f11

6d07s

2

[Rn]5f12

6d07s

2

[Rn]5f13

6d07s

2

[Rn]5f14

6d07s

2

[Rn]5f14

6d17s

2

III*

(III) IV*

(III) IV V*

III IV V VI*

III IV V* VI VII

III IV* V VI VII

II III* (IV) V VI

III (IV)

III IV

(II) III*

(II) III*

(II) III*

(II) III*

II* III

III*

* Most important oxidation states. Well characterized but slightly less important.

( ) Unstable oxidation states or in doubt.

13.1. Occurrence

They are all radioactive. Many of them do not occur naturally. Uranium and thorium are

available as ores and actinium, protactinium, neptunium and plutonium are available in small

amounts in these minerals. Thorium occurs up to 30% in monazite mixed with the

lanthanides. Their isolation is difficult and are now made artificially. Plutonium is available

in large quantity from fuel materials of uranium reactors.

All elements with higher atomic number than uranium are called transuranium elements.

13.2. Preparation of the elements

Neptunium and plutonium, named after planets, were made in 1940 by McMillan and

Abelson, and by Seaborg, McMillan, Kennedy, and Wahl respectively. Both are now

obtained from spent uranium fuel from nuclear reactors.

Two general methods that are to some extent complementary have been used for preparing

the transuranium elements:

Capture of neutrons, followed by β-decay which increases the atomic numbers by one

and

Capture of the nuclei of light elements, ranging from He to Ne, which increases

atomic number by several units in one step.

Successive neutron capture-β-decay sequences

The discovery of neptunium was based on the following sequence:

234

U (n,

U

23 min.

239Np

2.3d

239Pu

With the discovery of high power nuclear reactors an extension of the above sequence has

been used to produce other elements:

238U

(n, 239U

239Np

(n,

239Pu

(n,

240Pu

241Pu

24 360 y

241Am

The most stable isotope of Np,

237Np, is obtained by:

235

U

238U

(2n,

(n,2n)

237U

6.75d

93Np

237(2.2 x 10

6 y)

Only Pu is normally recovered since 239

Pu has fission properties similar to 235

U and can be

used as a fuel or in nuclear weapons. Some 237

Np is used to prepare 238

Pu (86.4 y), which is

used as a power source for satellites.

This method suffers from diminishing returns as each element has to be made from the one

before. For example in the irradiation of plutonium-239 less than 1% of the original sample

appears as californium-252 after the capture of 13 neutrons,

239Pu →

241Pu→

245Cm →

247Cm →

251Cf →

252Cf

% of original sample 100 30 10 1.5 0.7 0.3

The yield of heavier elements is controlled by:

The half lives of the various isotopes;

The ability to absorb neutrons.

The yield of heavier nucleus falls off sharply because successive neutron capture depends on

the build-up of intermediate elements and also because of decrease in nuclear stability with

increasing atomic weight. The heaviest elements are therefore best obtained by means of

cutting down on the step-by-step addition of neutrons. One means was provided by H-bomb.

A staring material can be bombarded with species that contains several nuclear particles. For

example the α-particle bombardment is the easiest way, and many actinides, such as 248

Cm, 249

Bk, 249

Cf, and 256

Md were first made in this way: e.g. 253

Es99 + 4He2 →

256Md101 +

1n0

α-particle bombardment requires the target to be the element with an atomic number of two

less than the desired element.

Both methods discussed above suffer from limitations, in so far as in both cases the yields

decrease rather rapidly as the atomic number increases. In the α-particle bombardment the

decreasing yields of successive elements arise first because the amounts of available target

materials are smaller, and also the probability of fission in the compound nuclei increases

quite rapidly with its charge.

Two modifications of the above methods have been used to produce isotopes of elements up

to 105:

Nuclear explosions

There is a vast flux of fast moving neutrons in an atomic explosion that can lead to

simultaneous addition of a number of neutrons before the intermediate nuclei can decay.

Thus einsteinium and fermium were first discovered in the fall out products of the first

atomic bomb.

Similar process to nuclear explosion occurs in certain stars called supernovae. Supernovae

arise as a result of a gigantic nuclear explosion in the star, creating neutron fluxes many

orders of magnitude greater than even the most powerful man-made devices.

Bombardment involving heavy ions (pioneered by the Russians)

Ions involved include B5+

, C6+

, N7+

, or O8+

. For example,

238U +

12C (

250Cf)

*

246Cf + 4 n

fission products In this way it is possible to „leap up‟ several elements in one step:

238

U92 + 12

C6 → 246

Cf98 + 4 1n0

238U92 +

16O8 →

250Fm100 + 4

1n0

246Cm96 +

12C6 →

254No102 + 4

1n0

252Cf98 +

11B5 →

257Lw103 + 6

1n0

13.3 General Chemical Properties of the Actinides

As discussed above there appears to be a competition between 5fn7s

2 and 5f

n-16d

17s

2

configurations. For the elements in the first half of the f shell it appears that less energy is

required for the promotion of 5f → 6d than for 4f → 5d promotion in the lanthanides; there is

thus a greater tendency to supply more bonding electrons with the corollary of higher

valences in the actinides. The second half resembles the lanthanides more closely.

Furthermore the 5f orbital is more accessible to bonding than the 4f. In the actinide series a

situation arises in which the energies of 5f, 6d, 7s, and 7p orbitals are about comparable over

a range of atomic numbers (especially U to Am). Bonding can therefore involve any or all of

them. In the chemistry this situation is indicated by the fact that the actinides are much more

prone to complex formation than are the lanthanides, where the bonding is almost extremely

ionic. Indeed the actinides can even form complexes with certain π-bonding ligands as well

as forming complexes with halides, sulphate, and other ions.

Since the energies of 5f, 6d, 7s, and 7p levels are comparable the energies involved in an

electron shifting from one to another may lie within the range of chemical binding energies.

It is therefore difficult to place the electronic structure of the elements in compounds and in

solutions as the ligands vary. It is also impossible to say which orbitals are being utilized in

bonding or to decide meaningfully whether the bonding is covalent or ionic.

Roughly the series falls into two families:

The first seven members have properties that are similar to the d-transition series while the

latter seven are similar to the lanthanides.

13.3.1 Oxidation states

The most stable oxidation state of the elements up to uranium is the one involving all the

valence electrons. Thus Np forms Np(VII), using all its 5f and 7s electrons. This oxidation

state is oxidizing and the most oxidation state is Np(V). Pu also forms Pu(VII) and Am up to

Am(VI) but the most stable oxidation state drops to Pu(IV) and Am(III) respectively. This

variable oxidation states make these early members of the series similar to the 3d-transition

series. (Recall that for the 4d and 5d series the highest oxidation states are more stable and

for the 3d series the lower oxidation states are more stable and dominate the aqueous

chemistry).

The latter elements tend to be most stable in the III state like the lanthanides.

This pattern of higher oxidation state stabilities has more in common with the d-series than

with the lanthanides. Compare Mn(VII) or Ru and Os(VIII) which states become more

oxidizing across the series.

However the III state becomes predominant from Cm → Lr, similar to the lanthanides.

Note that although Bk(IV) is strongly oxidizing it is more stable that Cm(IV) and Am(IV),

thus showing a parallel to Tb where the IV state corresponding to the f7 configuration has

some stability. Am(II) is not formed in aqueous solution but known in chloride melts. This

shows some slight resemblance to Eu which attains f7 configuration in its fairly stable II state.

No(II) is stable due to the attainment of the f14

configuration, analogous to Yb2+

in the

lanthanides.

The trend in ionic radii is similar to that shown by the lanthanides hence one can refer to the

actinide contraction arising from similar increase in effective nuclear charge due to poor

shielding by the f electrons.

13.3.2. Difference between 4f and 5f orbitals

The chief difference between the two depends upon the relative energies and spatial

distribution of the orbitals. The 4f orbitals populated in the lanthanides are sufficiently low in

energy that the electrons are seldom ionized or shared (hence the rarity of the LnIV

species).

Furthermore the 4f electrons are buried so deeply within the atom that they are unaffected by

the environment to any great degree. In contrast the 5f electrons, at least in the earlier

elements of the actinides, Th to Bk, are available for bonding allowing oxidation states up to

+7. In this respect these electrons resemble d electrons of the normal transition elements.

However, the heavier members resemble the lanthanides in displaying mainly oxidation state

of +3.

13.3.4. Electronic spectra of the lanthanides and actinides

As discussed earlier the electronic spectra of the lanthanides are typically sharp. The

absorption of actinides may be conveniently divided into two groups:

Am3+

and heavier actinides which have spectra that resemble the lanthanides;

Pu3+

and lighter actinides that have spectra similar in some ways but which have a

tendency towards broadening of absorption peaks, somewhat like the broadening seen

in the d-block metal ions.

Apparently the greater exposure of the 5f orbitals in the lighter actinide elements results in

greater ligand-metal orbital interaction and some broadening from vibrational effects. As the

nuclear charge increases the 5f orbitals behave more like the 4f orbitals in the lanthanides and

the spectra of the heavier actinides become more lanthanide-like.

13.3.5. Magnetic properties

The magnetic properties of the actinide ions are considerably harder to interpret than those of

the lanthanide ions. The experimental magnetic moments are usually lower than the values

calculated by using Russell-Saunders coupling and this appears to be due both to ligand field

effects similar to those operating in the d transition series and to inadequacy of this coupling

scheme. Since 5f can participate to some extent in covalent bonding ligand effects are to be

expected.

13.3.6. Electronic spectra

The electronic spectra of the actinide compounds originate from three types of electronic

transitions:

(i) 5f → 5f transitions:

These are orbitally forbidden, but the selection rule is partially relaxed by the action of the

crystal field in distorting the symmetry of the metal ion. Because the field is stronger than for

the lanthanides the bands are more intense by about a factor of 10 and, though still narrow,

are about twice as broad and are more complex than those of the lanthanides. They are

observed in the visible and uv regions and produce the colours of aqueous solutions of simple

actinide salts.

(ii) 5f → 6d transitions:

These are orbitally allowed and give rise to bands which are therefore much more intense

than those of type (i) and are usually rather broader. They occur at lower energies than do the

4f → 5d transitions of the lanthanides but are still normally confined to the ultraviolet region

and do not affect the colour of their ion.

(iii) Metal → ligand charge transfer:

These again are fully allowed transitions and produce broad, intense absorptions usually

found in the ultraviolet but sometimes trailing into the visible region. They produce the

intense colours which characterize many actinide complexes, especially those involving the

actinides in the high oxidation states with readily oxidisable ligands.

The date of discovery and origin of the names of the actinides

Atomic

No.

Element Symbol Date Origin of name

93

94

95

96

97

98

99

100

101

102

103

Neptunium

Plutonium

Americium

Curium

Berkelium

Californium

Einsteinium

Fermium

Mendelevium

Nobelium

Lawrencium

Np

Pu

Am

Cm

Bk

Cf

Es

Fm

Md

No

Lr

1940

1940

1944

1944

1949

1950

1952

1952

1955

1958

1961

The planet Neptune

The planet Pluto (next planet)

America (by analogy with Eu, named after Europe)

P & M Curie, by analogy with Gd, named after

Gadolin

Berkeley, by analogy with Tb, named after Ytterby

California, location of the laboratory.

Einstein, relativistic relation between mass and

energy

Fermi, construction of self-sustaining nuclear

reactor

Dimitri Mendeleev (Periodic table of the elements)

Alfred Nobel (benefactor of science)

Earnest Lawrence (developer of the cyclotron)

14 INORGANIC REACTION MECHANISM

14.0 Preliminary comments

14.1 Inert and labile compounds:

Compounds that undergo substitution reactions at room temperature spontaneously are

said to be kinetically labile.

Examples: Addition of NH3 to aqueous solution of copper(II) in excess results in

instantaneous formation of a deep blue solution.

[Cu(H2O)6]2+

+ 4 NH3 → [Cu(NH3)4(H2O)2]2+

+ 4 H2O

blue deep blue

Addition of a thiocyanate solution to an iron(III) solution results in a change in colour from

pale blue to red:

[Fe(H2O)6]3+

+ SCN- → [Fe(H2O)5(SCN)]

2+ + H2O

very pale blue red

Compounds like these are said to be labile.

Taube suggested a reaction half-life of one minute or less as the criterion for lability.

Those where substitution takes hours/days are said to be kinetically inert.

Inert here does not suggest that reaction will not take place; it simply implies that the reaction

is very slow.

An example is the acid hydrolysis of [Co(NH3)6]3+

:

[Co(NH3)6]3+

+ 6 H3O+ → [Co(H2O)6]

3+ + 6 NH3

This reaction is very slow and hence [Co(NH3)6]3+

is an inert complex.

14.2. Inertness and stability:

[Cu(H2O)6]2+

, Fe(H2O)6]3+

and [Co(NH3)6]3+

are all thermodynamically stable because their

equilibrium constants for formation are very large. However, whereas the Cu2+

and Fe3+

are

labile (because they undergo substitution reactions readily, [Co(NH3)6]3+

is inert because it

undergoes substitution very slowly.

14.3. Differences between stability and inertness:

thermodynamically stable complexes have large, positive free energies of

reaction, G.

Inert complexes have large positive free energies of activation, G*

14.4. General rules guiding lability/inertness:

Labile complexes:

Complexes with central metal atom having d-electrons in the eg orbitals, e.g.

[Ga(C2O4)3]3-

, d10

; [Co(NH3)6]2+

, d7; [Cu(H2O)6]

2+, d

9; [Ni(H2O)6]

2+, d

8 (weak field) and

[Fe(H2O)6]3+

, d5 (high spin).

Complexes containing less than 3 electrons in the d orbitals, e.g. [Ti(H2O)]3+

, d1;

[V(phen)3]3+

, d2; and [Ca(EDTA)]

2+, d

0.

Inert complexes:

Octahedral low-spin d4, d

5, and d

6 complexes, e.g. [Fe(CN)6]

3-, d

5; [Co(NO2)6]

3-, d

6;

[PtCl6]2-

, d6.

Octahedral d3 complexes. e.g. [Cr(H2O)6]

3+, d

3.

Complexes with d8 configurations generally react somewhat faster but slower than the d

7,

d9, or d

10 complexes. Many square planar d

8 complexes are inert.

If you recall the Ligand Field Theory discussed last year and this year, you will realise that

the LFSE helps to see the picture clearly.

Summary:

Slow reactions (inert) Intermediate Fast reaction (labile)

d3, low spin d

4, d

5, and d

6,

strong field d8 (square

weak field d8

d1, d

2, high spin d

4, d

5, and

d6

planar) d

7, d

9, d

10

14.5 Preliminary Comments

Most reactions involve the coming together of reagents and/or the separation of the products

in individual molecular acts, but more stable and change into the products.

The transition state constitutes an energy barrier that the individual reacting species must

cross in order to complete the reaction, so that at a particular temperature the rate of reaction

will be determined by the height of the barrier.

The rate of a reaction is conveniently expressed quantitatively in terms of the half-life. The

half-life of a reaction is the time taken for half of the reactant to be consumed or the time for

half of the products to be formed. For example, the half life of the reaction:

[CoCl(NH3)5]2+

+ H2O → [Co(H2O)(NH3)5]3+

+ Cl-, t½ = 113 h

G0 for the reaction

G=/ for the reaction applicable only on the molecular scale

G

Reagent

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

[CoCl(NH3)5]2+

113 226 339 450 563 h

Concn.

Time

Reaction coordinate

Product

14.6 Definitions

Reaction rate is the rate of change of concentration of a substance involved in the reaction

(being either substrate or product or both).

Rate constants are expressed as a function of temperature (or time) factor.

Rate expression is the functional relation between rate and concentration; it provides the

important clue about the mechanism.

14.7 Elucidation of mechanism

The most important evidence in the elucidation of a reaction mechanism is the experimentally

determined rate equation.

For reaction

A + B → C + D

the rate law is given as

= ka[A] + kab[A][B] + kab'[A][B][H

+]

-1

Other additional information are, however, very essential too. In particular, attention should

be paid to

the exact nature, including the stereochemistry, of the reactants and products.

the presence of equilibria.

the stoichiometry of the reaction.

14.8 Confirmation of a mechanism from other than kinetic evidence

Confirmation of a mechanism suggested on the basis of a rate equation may be obtained

from:

Detailed knowledge of the nature of the reactants.

Detection (direct or indirect) of suspected intermediates.

Detailed knowledge of the nature of the products.

One example only:

The copper(I) reduction of iron(III),

Cu(I) + Fe(III) → Cu(II) + Fe(II)

The rate law experimentally determined is

][

)]()][([)([

H

ICuIIFed

dt

IIIFed

The inverse dependence on [H+] suggests that a hydrolysis product is the reagent. It is known

that iron(III) solution at a given pH contains the hydroxopentaaquairon(III) ion and that Cu(I)

does not undergo any measurable hydrolysis under the given conditions indicates that it is the

hydrolysed iron(III) species which is the reactant. The mechanism is then represented as:

[Fe(H2O)6]3+

+ H2O → [Fe(H2O)5OH]2+

+ H3O+

[Fe(H2O)5OH]2+

+ Cu(I) Products

15. Substitution in Octahedral Complexes

Simple substitution is defined as the replacement of a ligand in the coordination sphere by

another coming from the environment by a path that involves nothing more complicated than

a temporary change in the coordination number of the reaction centre. Thus making and

breaking of bonds are involved.

Substitution reactions are generally classified as having either dissociative or associative

mechanisms.

Consider: [A5MX] + Y- [A5MY] + X

-

Two reaction pathways are possible:

Extreme dissociative mechanism, the D or SN1 mechanism, can be represented as:

or

Co

X

A A

A

AA

-XCoA A

A

AA

Co

Y

A A

A

AA

+Y

Slow Fast

Intermediate

[A5MX]n+

[A5M](n+1)+ + X

-

Y-

[A5MY]n+

-X

slow

fast

This involves a rate-determining (slow) loss of X- to give an intermediate with a coordination

number reduced by one. Subsequent addition of Y to form the product is fast.

SN1 = means substitution nucleophilic unimolecular.

The extreme associative mechanism, the A or SN2 mechanism, can be represented as:

This involves the rate-determining slow addition of Y- to form an intermediate with a

coordination number increased by one. Subsequent loss of X- is fast.

SN2 = Substitution nucleophilic bimolecular.

Both D and A mechanisms consist of two steps: bond breaking and bond making.

Most reactions are, however, intermediate in character: both bond breaking and bond

making occur simultaneously. Such concerted processes are said to occur by I (for

Interchange) mechanism. Here, a cooperative interchange occurs in which X- leaves as

Y- arrives. Bond breaking takes place as bond formation is being executed.

The I mechanism is further divided into two: ID and IA

Dissociative Interchange (ID) mechanism: breaking of metal-ligand bonds has a greater

effect on the rate and activation energy than formation of new bonds.

Co

X

A A

A

AA

CoA

AAA

Co

Y

A A

A

AA

Slow Fast

Intermediate

X

A

Y

+Y -X

Associative Interchange (IA) mechanism: bond making is more important than bond

breaking.

Summary of different mechanisms of substitution reactions in Octahedral

Complexes

15.2 Divisions of Rate of Reaction

Class Rate of exchange

I 108 s

-1

II 105 - 10

8 s

-1

III 1 - 104 s

-1

IV ~10-8

s-1

D MECHANISM A MECHANISM

SN1 SN

2

I MECHANISM

ID IA

Class I: Rate constant ~10-8

s-1

. Here, purely electrostatic forces bind complexes. Ions

of alkali and alkali earth metals fall in this class. Ratio r

Z 2

for these ions

range up to

~10 x 1028

C2 m

-1. This class is difficult to study.

Class II Exchange of H2O is fast and first order rate constant. This class can be studied

using relaxation techniques. Here the equilibrium is perturbed by a fast

variation of pressure and temperature, and the response of the system is used

to estimate the rates of reaction. Dipositive transition metal ions, Mg2+

, and

Ln3+

are known. Here bonding is stronger than in Class I; but LFSE is

relatively small. r

Z 2

for these ions ranges

10 - 30 x 10-28

C2 m

-1

Class III H2O exchange is relatively slow than in I and II. First order rate constant is 1 -

104 s

-1. Reactions can be followed by more or less conventional kinetic

technique. Tripositive transition metal ions are examples, and stabilised, to

some extent, by LFSE. r

Z 2

ratio is greater than 30 x 10-28

C2 m

-1.

Class IV H2O exchange is slow, an example of inert complexes' behaviour. First order

rate constants range from 10-1

- 10-9

s-1

. Metal ions are comparable in size to

Class III ions and exhibit considerable LFSE, e.g. Cr3+

(d3) with CFSE of

12Dq; Co3+

(d6 low spin) with LFSE 24Dq and low spin d

8 ions like Pt

2+.

NOTE:

It is not the LFSE that makes reaction to be very slow but the loss upon formation of the

intermediate (activated complex). For example, if LFSE of the initial complex is >> than

that of the activated complex then the complex will be inert; but if the difference is small

then the complex will be labile.

The rate is also affected by the charge of the Mn+

ion, Mn3+

< M2+

.

Lability increases: Co3+

< Cr3+

< Mn3+

< Fe3+

etc.

Lability decreases: [AlF6]3-

> [SiF6]2-

> [PF6]-> SF6 (inert).

15.3 Complications

Solvent interaction

Ion pair formation

Nature of the ligands (Leaving and entering ligands)

15.4 The Rate Laws

D Mechanism:

][][

55

MXAkdt

MXAd

A Mechanism

]][[][

55

YMXAkdt

MXAd

Unfortunately, however, a particular rate law does not prove that the reaction is SN1 or SN

2.

For example, where there is solvent interaction, which is very common in reactions

conducted in aqueous solution, the keq value will not be a true value, and will have to be

corrected. Consequently for a reaction like

the rate law is

]][[][

255

OHMXAkdt

MXAd

= k/[A5MX]

since [H2O] is constant because water is the solvent. This reaction can then be SN1 or SN

2.

Generally for Co3+

, which is the ion of interest, substitution follows the D mechanism.

Ion pair formation:

Common when reacting cations and anions; and also in non-aqueous solvents.

So, the equilibrium constant, K1p, should be taken into consideration (in certain cases).

[A5MX] + H2OSlow

[A5M(H2O)] + X-

[A5M(H2O)] + Y-

Fast

[A5MY] + H2O

[A5MX]n+

K1p

k-1

{[A5MX]Y}n-m ....

kK1p[A5MX][Y]

1+ k-1 [Y]-d[A5MX] =

Rate is dependent on whether k-1[Y] << 1 or not.

15.5 Hydrolysis

Hydrolysis is generally classified according to the conditions. In acid solution the

process is termed acid hydrolysis.or aquation and is illustrated by the equation:

[A5MX] + H2O [A5M(H2O)] + X-

In a basic solution the process is termed base hydrolysis and illustrated by the equation:

[A5MX]n+

+ OH-

[A5M(OH)]n+

+ X-

Depending on pH of the reaction mixture it follows that the product of a given hydrolysis

can be a mixture of both the aqua and the hydroxo complex. For a typical complex for

which the base hydrolysis is observable the rate law is:

The first term (kA) refers to the acid hydrolysis and the second term (kB) refers to the base

hydrolysis.

15.5 Acid hydrolysis

At pH 0-3 base hydrolysis is negligible and the second term vanishes. The behaviour is then

first order.with respect to the complex and independent of acid concentration.

The rate law provides no information as to the role of water and does not enable a distinction

between a dissociative, associative or concerted process to be obtained. Mechanistic

information must, therefore, be obtained from other than kinetic sources.

For example, the rate of hydrolysis decreases with increase in the thermodynamic bond

strength of the Co-X bond indicating that this bond is broken initially before Y comes in.

Acid hydrolysis is easy to follow and pH dependent.

Substitution in octahedral complexes is predominantly dissociative, except probably when

the ion is large (like in the 2nd

and 3rd

row transition metal ions, where a dominant

associative character can be acquired).

In such larger ions there is more room for attack or lower nucleophilic attack and hence

permit association.

-d[A5MX]

dt= kA[A5MX] + kB[A5MX][OH

-]

The identity of the living group X has a large effect in dissociatively activated reactions

because their rate depends on the cleavage of the M---X bond. X is the only variable in

the reaction

[CoX(NH3)5]2+

+ H2O [Co(NH3)5(OH2)]3+

+ X-

There is a linear relationship between the logs of the rate constants and equilibrium

constants of the reaction.

Specifically, lnk = lnK + c

A plot of logk vs. logK gives a straight line, with intercept c. The straight line obtained

from the equation indicates the existence of a linear free energy relation (LFER). The

LFER of unit slope shows that changing X has the same effect on G for the conversion

of Co---X to the activated complex as it has on rG for the complete elimination of X

-.

A slope <1.0, indicates an associative character in Rh(III).

There is no important trans effect in octahedral complexes. Both cis and trans

ligands affect rates of substitution in proportion to the strength of the bonds they form

with the metal.

Another example is that the rate of substitution for the process

[Co(L-L)2Cl2]+ + H2O [Co(L-L)2(H2O)Cl]

2+ + Cl

-

increases with an increase in size of the ligand, L-L. This is explained in terms of

dissociative process when an increase in size of the non-involved ligand discourages

association to give a seven-coordinate intermediate.

Note:

By increasing the bulkiness of the ligand, one decreases the solvation effect, and

then the transition state slows down the reaction rate.

It must go through a D-mechanism.

The inductive effect rises due to increase of alkyl group, by increasing charge

effect.

H2NCH2CH2NH2

L L k x 105 s

-1

3.2

H2NCH2CHNH2

CH3

6.2

H2N CH CHNH2

CH3

CH3

15.0

H2N CH CH

CH3 CH3

NH242.0

H2N C C

CH3 CH3

NH2

CH3 CH3

3300

Since the base strengths of the substituents do not change by more than a 1.5

factor, the increase in rates is not due to inductive effect and solvation effect, but

possibly due to steric strains.

15.7 The conjugate base mechanism

The behaviour of base hydrolysis appears to be associative mechanism, but

subsequently has been found to have a conjugate base mechanism (called SN1CB for

substitution, nucleophilic, unimolecular, conjugate base). The reactions depend on

amine, ammine, or aqua ligands that can lose protons to form amido or hydroxo

species that are then more likely to lose one of the orher ligands.

K

1. [Co(NH3)5Cl]2+

+ OH-

[Co(NH3)4(NH2)Cl]+ + H2O

(pre-equilibrium)

k2

2. [Co(NH3)4(NH2)Cl]+

[Co(NH3)4(NH2)]2+

+ Cl- (slow)

3. [Co(NH3)4(NH2)]2+

+ H2O [Co(NH3)5(OH)]2+

(fast)

Overall:

[Co(NH3)5Cl]2+

+ OH- [Co(NH3)5(OH)]

2+ + Cl

-

In the first step, an NH3 ligand acts as a BrØnsted acid, resulting in the formation of its

conjugate base, the NH2- ion, as a ligand. Because NH2

- is a strong -donor it greatly

accelerates the loss of Cl- ion.

Octahedral substitution is greatly accelerated by OH- ions when ligands with acidic protons

are present.

Basolo revealed that the role of OH- is catalytic, by removing a proton from the base ligand.

The rate law is

Rate = k[CoCl(NH3)52+

][OH-]

The presence of [OH-] in the rate equation shows that it plays a rate-determining role.

However, it is not because [OH-] attacks the metal centre but rather because it deprotonates a

coordinated NH3 ligand to form a conjugate base; hence the name of the mechanism. A pre-

equilibrium is first established, followed by loss of Cl- to give a reactive amido species, and

finally formation of the product in a fast step.

If the equilbrium constant for equation (1) is K then the rate law consistent with mechanism

is given by

Rate = ][1

]][)([ 2

532

OHK

OHClNHCoKk

Two supporting evidences for this mechanism are

If NH3 is replaced by pyridine or another tertiary amine, base hydrolysis is very much

slower.

The exchange of H (in NH3) for D in alkaline D2O is much faster than the rate of base

hydrolysis.

The isotope ratio (18

O/16

O) in the product in 18

O-enriched water is the same as that in

water regardless of the leaving group (X = Cl-, Br

-, NO3

-). If an incoming water molecule

had a large influence (associative mechanism), the concentration of 18

O should be larger

in the product;

RNH2 compounds react faster than NH3 compounds, showing that steric crowding

favours the 5-coordinate intermediate (eq. 2 above).

The rate constants and dissociation constants for the compounds form a linear free energy

relationship (LFER), where a plot of lnkOH versus lnKOH is linear.

15.8. Anation reactions

Involve the replacement of coordinated water:

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

+ X- [Co(NH3)5X]

2+ + H2O

They may undergo ion-pairing formation, leading to a slow rate of water

displacement, or

Purely undergoing dissociative pathway.

e.g.:

the intermediate has long enough half-life to distinguish it from ion-paring process.

Rate = -dt

OHCNCod ])([ 2

25

= ][

]][)([

21

2

2521

Xkk

XOHCNCokk

once X-is in large concentration, the rate law approximates a first order, i.e. pseudo-

first order reaction condition and the kobsd resolves as following:

][

][][

][1

21

21

Xkk

Xkksubstratek

dt

substratedobsd

By taking series of runs with varying [X-], k1 will tend to be

= ~1.6 x 10-3

s-1

and therefore the ratio of k2/ k-1 is a measure of the nucleophilic

power of the different substituents, and k-1 is independent of the substituents.

k2 k1/k-1 can be found from graph.

The order of k2 is as follows:

OH-> N3

-> NCS

-> I

-> Br

->NH3> H2O at 40

oC and = 1.0.

16. Substitution Reactions in Square Planar Complexes

A large number of relatively inert square planar complexes have been known for a

long time, e.g. PtII, Pd

II, Ni

II, Au

III, Rh

I and Ir

I. They all have d

8 configuration, which

gives stability to square planar complexes. We will concentrate our efforts on PtII,

which shows the general features which are common to all the others.

Let us consider [PtA2LX], where X = leaving group, L = cis or trans to X (i.e. X

defines geometry).

The general process:

[Co(CN)5H2O]2- k1

k-1

[Co(CN)5]2-

+ H2O

+ X-

[Co(CN)5X]3-

k2

[PtA2LX] + Y-

[PtA2LY]SN

2

+ X-

PtII mechanism is invariably SN

2

This is predictable:

Add Y-; it is easy! The 5

th position is open, there is no steric hindrance to the

approach of Y-.

There is strong evidence for the existence of a 5-coordinate intermediate, e.g.

There is usually a retention of configuration:

cis-[PtA2LX] + Y-cis-[PtA2LY] + X

- (100% retention)

trans-[PtA2LX] + Y- trans-[PtA2LY] + X

- (100% retention)

16.1 Mechanism

The SN2mechanism is nota one step process. The rate is independent of X, i.e. it is the

rate of bond formation, which is important, not the rate of bond breaking.

16.2. Kinetics of the reaction:

In general a 2-term rate law is observed:

Rate = k1[complex] + k2[complex][Y-]

1st order term 2

nd order term

16.3 Explanation of the two terms

XPt

Y-

Ni(CN)4]2-

+ excess CN-

H2O[Ni(CN)5]

3-

Evidence from Uv-Vis

Direct attack on PtII.

Initial attack by solvent on PtII, and then replacement of solvent by Y

-.

The rate law = k2[complex][Y-]

Rate = k1[complex][S] = k'[complex]

because the concentration of the solvent is virtually constant.

Combining both gives the 2-term rate law, where the first term is for the SN2 reaction

involving the solvent and the 2nd

term is the SN2 involving the nucleophile, Y

-.

The two rate laws can be combined to give:

Pt

A

A

L X + Y- slow

RDSPtL

A

AY

X

Pt

A

A

L Yfast

+ X-

intermediate

Pt

A

A

L X + Sslow

RDSPtL

A

AX

S

Pt

A

A

L S + X-

intermediate

+Y-

slow

+ Y-

fast

Pt

A

A

L Y + S

PtL

A

AY

S

Rate = dt

complexd ][ = {k1 + k2[Y

-]}[complex]

Summary

Both reactions proceed via an associative process (A), involving a trigonal bipyramid

transition state. Chemical justification for this transition state includes:

Many five coordinate transition metal complexes are known, e.g., Fe(CO)5,

[CoL2(CO3)] +, [Ni(CN)5]

3-

ML3X complexes are sterically and electronically unsaturated and have space for Y

to coordinate.

Evidence

Rate law is consistent with associative mechanism

Charge on the metal centre - has no effect on the rate of reaction

Steric effect-significant increase in rate was observed for less hindered ligands; trans-

is faster than cis

C6H5->2-Me-C6H4

->2,4,6-Me3-C6H2

-

Note trans>cis for the substitution reaction

[Pt(PEt3)RCl] + Y- → [Pt(PEt3)RY] + Cl

-

16.4 The trans effect

It is observed that during the substitution reactions of square planar metal complexes, some

ligands preferentially direct the substitution trans to themselves. i.e., the choice of leaving

group is determined by the nature of ligand trans to it.

“The trans effect is defined as the effect of a coordinated ligand upon the rate of substitution

of ligands opposite to it” or

“The Trans effect can be defined as the effect of a ligand over rate of substitution of another

ligand positioned trans to it in the square planar complexes”.

Where 'T' is the trans directing group and „Nu‟ is the nucleophilic ligand which preferentially

substitutes the ligand 'X' which is trans to ligand 'T'.

For Pt(II) compounds the order of trans effect H2O ~ OH- ~ NH3 ~ NR3< Cl

- ~ Br

-< SCN

- ~ I

-

~ NO2- ~ C6H5

-< CH3

-< PR3 ~ AsR3 ~H

-< olefins ~ CO ~ CN

-.

trans-effect in synthesis

The Trans effect can dictate the product formed in the substitution reactions. For example the

Trans effect manifests in the synthesis of cisplatin, cis-diamminedichloridoplatinum(II). It is

prepared by substituting the two chloro groups of PtCl42-

by ammonia molecules.

Pt

Cl

Cl

Cl

Cl

NH3Pt

NH3

Cl

Cl

Cl

2- -

NH3

PtH3N

Cl

Cl

NH3

cis-platin

The trans product is obtained by starting from Pt(NH3)42+

. In this case the second Cl group is

substituted preferentially at trans position to the first one.

Pt

NH3

H3N

NH3

NH3

Cl-

Pt

NH3

H3N

Cl

NH3

2+ -

Cl-

PtH3N

Cl

NH3

Cl

trans isomer