Thursday 22/12/2011

1 Basic theory

Werner's Coordination Theory

When aqueous ammonia is added to a solution of cobalt dichloride, CoCl2, a

blue precipitate forms of the corresponding hydroxide, Co(OH)2, which dissolves on the addition of an excess of ammonia to give a solution that immediately begins to

absorb oxygen and turn brown. From the oxidized solution, the following

compounds can be isolated:

Composition

Colour

(I) CoCl3·6NH3 Orange-yellow

(II) CoCl3·5NH3·H2O Pink

(III) CoCl3·5NH3 Purple

A great many other compounds of this type can be prepared, (i) by starting with

other cobalt salts, and (ii) by carrying out further reactions with the compounds first obtained. Among these other compounds are two other chlorides:

(IV) CoCl3·4NH3 Violet (V) CoCl3·4NH3 Green

At first sight, the properties of compounds (I) - (V) are very puzzling:

(i) All of them fail to give a brown precipitate of Co(OH)3 when treated with sodium hydroxide solution - a property expected of compounds containing the

Co3+

ion. Only on boiling does a precipitate form.

(ii) All of them fail to give ammonium chloride when treated with concentrated

hydrochloric acid - a reaction that would be expected of a compound containing

ammonia. Only on boiling with sodium hydroxide is ammonia evolved.

(iii) While all of them give a precipitate of silver chloride when treated with

silver nitrate in aqueous solution, the amount that is precipitated in the cold is in some cases less than expected, as shown in the table below. In these cases the full

amount is only obtained by leaving the solution to stand for a long time or by boiling

it.

(iv) The five compounds give different numbers of ions in solution, as judged

from (a) the electrical conductivities of their solutions in comparison with those of

other electrolytes, (b) the extent to which they depress the freezing point of water. The number of ions judged to be present are shown in the table below.

Moles of AgCl

precipitated per

mole

Number of ions

present in solution

per CoCl3

(I) CoCl3·6NH3 3 4

(II) CoCl3·5NH3·H2O 3 4

(III) CoCl3·5NH3 2 3 (IV) CoCl3·4NH3 1 2

(V) CoCl3·4NH3 1 2

These properties, and those of many other compounds of a similar kind, were

brilliantly rationalized by Alfred Werner in 1893. In this year, at the age of only 26,

he proposed what is now referred to as his “coordination theory”, for which he was awarded a Nobel prize in 1913. Its principal postulates are as follows:

(1) An atom exhibits two types of valency, its ordinary valency (V), and a valency that determines the number of neighbouring atoms to which it is

bound (the “coordination number”).

(2) An atom’s ordinary valency is satisfied by other atoms or radicals; its

coordination number is satisfied by atoms, radicals, or molecules.

(3) Bonds to neighbouring atoms are directed towards fixed positions in space.

On the basis of postulates (1) and (2), Werner formulated compounds (I)(V) as shown diagrammatically below. Ordinary valency bonds are designated by black

lines () and bonds between neighbouring atoms by green lines (). The cobalt

atoms have their ordinary valency of three (as in CoF3) and are given a coordination number of six. The other atoms are given coordination numbers to match (e.g. four

for the nitrogen atom in NH3, leaving one for the cobalt).

Now if in these formulations the black lines are taken to be ionic bonds, and

the green and double lines are taken to be bonds of a non-ionic character, the

properties of compounds (I) - (V) given above are fully explained. Thus (I) and (II) would be expected to give three Cl

– ions in solution, (III) only two, and (IV) and (V)

only one. This leads to the customary formulations:

(I) [Co(NH3)6]Cl3 containing the [Co(NH3)6]3+

ion;

(II) [Co(NH3)5H2O]Cl3 containing the [Co(NH3)5H2O]

3+ ion;

(III) [CoCl(NH3)5]Cl2 containing the [CoCl(NH3)5]2+

ion;

(IV), (V)

[CoCl2(NH3)4]Cl containing the [CoCl2(NH3)4]+ ion.

The existence of two isomers of [CoCl2(NH3)4]Cl was explained by Werner on the basis of postulate (3). Indeed, the fact that only two isomers of this formula were

known led him to propose that the arrangement of atoms around the cobalt atom is

an octahedral one, since other arrangements would lead to more than two isomers. This is shown in the following table.

Known isomers vs. the number theoretically possible for three different

structures

Complex

type

Number of

known

isomers

Number of isomers theoretically

possible

MA5B One One One One

MA4B2 Two Three (1,2;

1,3; 1,4)

Three (1,2;

1,4; 1,6)

Two (1,2;

1,6)

MA3B3 Two Three

(1,2,3;

1,2,4; 1,2,5)

Three

(1,2,3;

1,2,4; 1,2,6)

Two

(1,2,3;

1,2,6)

The absence of a third isomer of [CoCl2(NH3)4]Cl does not of course prove that the octahedral structure is correct. The third isomer may simply be much less stable

or more difficult to isolate. Werner was able to prove, however, that the octahedral

structure is the correct one by preparing a compound that would be expected to exist in optically active forms if the structure were octahedral, but not if it was planar or a

trigonal prism, and showing that the compound can indeed be resolved into optically

active isomers. This conclusion has been verified by X-ray crystallography.

Most metals can form some coordination compounds, but few form as many as

trivalent cobalt. Others that do include trivalent chromium, and bi- and quadri-valent platinum. Trivalent chromium and quadrivalent platinum have the same coordination

number and geometry as trivalent cobalt; bivalent platinum has a coordination

number of four with the groups around it arranged in a square.

The complex part of a coordination compound may be a cation, an anion, or an

uncharged molecule. For example:

complex cation [Co(NH3)6]3+

complex anion [CoF6]3

, [Co(CN)6]3–

, [Co(NO2)63–

]

neutral complex [CoCl3(NH3)3]0, [Co(NO2)3(NH3)3]

0

Neutral complexes are frequently more soluble in nonpolar organic solvents than they are in water. When they do dissolve in water, they behave as non-electrolytes.

Werner's Coordination Theory - Revision Points

Postulates

1. In co-ordination compounds, central metal atoms exhibit primary valency and

secondary valency.

The primary valency is ionizable. Secondary valency is not ionizable.

The primary valency corresponds to oxidation state.

The secondary valency corresponds to coordination number. (the central metal ion

and ligands are not ionizable)

2. Every metal atom has a fixed number of secondary valencies (coordination

number(s)).

3. The metal atom tends to satisfy boths its primary valency as well as its secondary

valency. Primary valency is satisfied by negative ions (metal ion has a positive

charge) whereas secondary valency (coordination number) is satisfied either by

negative ions or by neutral molecules. (In certain case a negative ion may satisfy

both types of valencies).

4. The coordination number or secondary valencies are always directed towards the

fixed positions in space and this leads to definite geometry of the coordination

compound.

Valence bond theory

In 1927, valence bond theory was formulated and it argues that a chemical bond

forms when two valence electrons, in their respective atomic orbitals, work or function to hold two nuclei together, by virtue of effects of lowering system

energies. Building on this theory, the chemist Linus Pauling published in 1931 what some consider one of the most important papers in the history of chemistry: "On the

Nature of the Chemical Bond". In this paper, elaborating on the works of Lewis, and

the valence bond theory (VB) of Heitler and London, and his own earlier works, Pauling presented six rules for the shared electron bond, the first three of which were

already generally known:

1. The electron-pair bond forms through the interaction of an unpaired electron

on each of two atoms.

2. The spins of the electrons have to be opposed.

3. Once paired, the two electrons cannot take part in additional bonds.

His last three rules were new:

4. The electron-exchange terms for the bond involves only one wave function from each atom.

5. The available electrons in the lowest energy level form the strongest bonds.

6. Of two orbitals in an atom, the one that can overlap the most with an orbital from another atom will form the strongest bond, and this bond will tend to lie

in the direction of the concentrated orbital.

Building on this article, Pauling's 1939 textbook: On the Nature of the Chemical Bond would become what some have called the "Bible" of modern chemistry. This

book helped experimental chemists to understand the impact of quantum theory on

chemistry. However, the later edition in 1959 failed to adequately address the problems that appeared to be better understood by molecular orbital theory. The

impact of valence theory declined during the 1960s and 1970s as molecular orbital

theory grew in usefulness as it was implemented in large digital computer programs. Since the 1980s, the more difficult problems of implementing valence bond theory

into computer programs have been solved largely, and valence bond theory has seen

a resurgence.

THE SHAPES OF COMPLEX METAL IONS

This page describes the shapes of some common complex metal ions. It goes on to

look at some simple examples of stereoisomerism (geometric and optical) in complex ions. If you aren't sure what stereoisomerism is, you will find a helpful link

further down the page.

Some simple shapes for complex ions

These shapes are for complex ions formed using monodentate ligands - ligands which only form one bond to the central metal ion. You will probably be familiar

with working out the shapes of simple compounds using the electron pair repulsion

theory. Unfortunately that doesn't work for most complex metal ions involving transition metals. The answer is just to learn the shapes you need to know about. As

you will see, it isn't difficult

6-coordinated complex ions

These are complex ions in which the central metal ion is forming six bonds. In the simple cases we are talking about, that means that it will be attached to six ligands.

These ions have an octahedral shape. Four of the ligands are in one plane, with the

fifth one above the plane, and the sixth one below the plane. The diagram shows four fairly random examples of octahedral ions.

It doesn't matter what the ligands are. If you have six of them, this is the shape they will take up. Easy!

4-coordinated complex ions

These are far less common, and they can take up one of two different shapes.

Tetrahedral ions

These are the ones you are most likely to need for A' level purposes in the UK.

There are two very similar ions which crop up commonly at this level: [CuCl4]2-

and

[CoCl4]2-

. The copper(II) and cobalt(II) ions have four chloride ions bonded to them rather than six, because the chloride ions are too big to fit any more around the

central metal ion

That's not very difficult to remember either!

A square planar complex

Occasionally a 4-co-ordinated complex turns out to be square planar. There's no easy

way of predicting that this is going to happen. The only one you might possibly

come across at this level is cisplatin which is used as an anti-cancer drug.

Cisplatin is a neutral complex, Pt(NH3)2Cl2. It is neutral because the 2+ charge of the original platinum(II) ion is exactly cancelled by the two negative charges supplied

by the chloride ions

The platinum, the two chlorines, and the two nitrogens are all in the same plane. We

will have more to say about cisplatin immediately below

Stereoisomerism in complex ions

Some complex ions can show either optical or geometric isomerism.

Geometric isomerism

This occurs in planar complexes like the Pt(NH3)2Cl2 we've just looked at. There are two completely different ways in which the ammonias and chloride ions could

arrange themselves around the central platinum ion:

The two structures drawn are isomers because there is no way that you can just twist

one to turn it into the other. The complexes are both locked into their current forms

The termscis and trans are used in the same way as they are in organic

chemistry. Trans implies "opposite" - notice that the ammonias are arranged opposite each other in that version, and so are the chlorines. Cis implies "on the

same side" - in this instance, that just means that the ammonias and the chlorines are

next door to each other.

Optical isomerism

You recognise optical isomers because they have no plane of symmetry. In the organic case, it is fairly easy to recognise the possibiliy of this by looking for a

carbon atom with four different things attached to it. It isn't qute so easy with the

complex ions - either to draw or to visualise!

The examples you are most likely to need occur in octahedral complexes which contain bidentate ligands - ions like [Ni(NH2CH2CH2NH2)3]

2+ or [Cr(C2O4)3]

3-.

The diagram below shows a simplified view of one of these ions. Essentially, they

all have the same shape - all that differs is the nature of the "headphones". I have

deliberately left the charges off the ion, because obviously they will vary from case to case. The shape shown applies to any ion of this kind.

If your visual imagination will cope, you may be able to see that this ion has no

plane of symmetry. If you find this difficult to visualise, the only solution is to make the ion out of a lump of plasticene (or a bit of clay or dough) and three bits of

cardboard cut to shape

A substance with no plane of symmetry is going to have optical isomers - one of which is the mirror image of the other. One of the isomers will rotate the plane of

polarisation of plane polarised light clockwise; the other rotates it anti-clockwise

Geometry and magnetic nature of some complexes

Atom/ion/complex

(1)

Configuration

(2)

Oxidation

state of

metal (3)

Type

of

hybrid

ization

(4)

Geometry

shape

(5)

No. of

unpai

red

electr

ons

(6)

Magnetic

nature

(7)

Ni2+

(d8)

+2 2

Paramagne

tic

[NiCl4]2–

+2 sp3

Tetrahedr

al 2

Paramagne

tic

[Ni(CN)4]2+

+2 dsp2

Square

planar 0

Diamagneti

c

Ni 0 2 Paramagne

tic

Ni(CO)4

0 sp3

Tetrahedr

al 0

Diamagneti

c

[Ni(NH3)6]2+

+2 sp

3d

2(

outer)

Octahedra

l 2

Paramagne

tic

Mn2+

(d5)

+2 5

Paramagne

tic

[Mn(CN)6]4–

+2 d

2sp

3(I

nner)

Octahedra

l 1

Paramagne

tic

[MnCl4]2

+2 sp

3

Tetrahedr

al 5

Paramagne

tic

Cu2+

(d9) +2 1

Paramagne

tic

[CuCl4]2–

+2 sp3

Tetrahedr

al 1

Paramagne

tic

[Cu(NH3)4]2+

+2 dsp2 Square

planar

1 Paramagne

tic

(1) (2) (3) (4) (5) (6) (7)

Cr3+

(d3)

+3 3

Paramagneti

c

[Cr(NH3)6]3+

+3 d

2sp

3 (

Inner) Octahedral 3

Paramagneti

c

[Cr(H2O)6]3+

+3 sp

3d

2 (

Outer) Octahedral 3

Paramagneti

c

CO3+

(d6) +3 4

Paramagneti

c

[CoF6]3–

+3 sp

3d

2 (

Outer) Octahedral 4

Paramagneti

c

[Co(NH3)6]3+

+3 d

2sp

3 (

Inner) Octahedral 0 Diamagnetic

CO2+

(d7) +2 3

Paramagneti

c

[Co(H2O)6]2+

+2 sp

3d

2 (

Outer) Octahedral 3

Paramagneti

c

Fe2+

(d6) +2 4

Paramagneti

c

[Fe(CN)6]4–

+2 d

2sp

3 (

Inner) Octahedral 0 Diamagnetic

[Fe(H2O)6]2+

+2 sp

3d

2 (

Outer) Octahedral 4

Paramagneti

c

[Fe(NH3)6]2+

Same +2 sp

3d

2 (

Outer) Octahedral 4

Paramagneti

c

Fe3+

(d5) +3 Octahedral 5

Paramagneti

c

[Fe(CN)63–

+3 d

2sp

3 (

Inner) 1

Paramagneti

c

Fe

o 4 Paramagneti

c

Fe(CO)5

o dsp3

(Inner)

Trigonal

bipyramida

l

o Diamagnetic

The Spectrochemical Series

The value of the ligand field splitting parameter, i.e. the amount by which the

degeneracy of the d-orbitals is disturbed by the effect of the electrostatic field generated by the ligands, depends upon the identity of the ligands.

The ligand field splitting parameter can be measured by recording the optical

absorption spectrum of the complex. The first absorption maximum in the spectrum

is assigned to the t2g-to-eg transition, and the frequency of the absorption corresponds to the value of the splitting parameter. If one investigates the series of

compounds [CoX(NH3)5]n+

, where X is I–, Br

–, Cl

–, H2O

and NH3, one finds that the colors range from purple, for I–, through pink, for Cl

–, to

yellow, for NH3. This means that the frequency of the transition increases, and so the ligand field splitting parameter increases across the series.

The series obtained is independent of the identity of the metals at the center of the

complex, and so the ligands can be arranged into the Spectrochemical Series:

I–< Br

–< S

2–< SCN

–< Cl

–< NO

3–< F

–< OH

–< C2O4

–< H2O < NCS

–< CH3CN < NH3<

en < bipy < phen < NO2< PPh3 < CN–< CO

When a ligand causes there to be a large ligand field splitting parameter, it is said to generate a strong field, and when there is a small ligand field splitting parameter, it

is said to be a weak field ligand. For example, CN– is a strong field ligand and I

– is a

weak field ligand.