ntroduction to the 3-d block and transition Metals

The elements scandium to zinc (Z = 21 to 30) are known as the 3d block of elements or 3d-block of metals because here the first

of the possible d sub-shells is progressively filled (3d-block - first row of the d-blocks).

o See rules on how to work out electron configurations

The transition elements are group of industrially important metals mainly due to their strong inter-atomic metallic bonding giving them

generally high melting/boiling points and high tensile strength.

These-called 'transition metal characteristics' arise from an incomplete d sub-shell energy level but scandium and zinc are not true

transition metals i.e. Ti to Cu are the real transition elements (reasoning later).

Note that physically, zinc is low melting and a lower tensile strength compared to the others in the 3d block.

Although scandium is physically typical of a transition metal e.g. high melting point and high tensile strength, chemically, scandium only

forms a single and colourless triple charged ion (Sc3+). Therefore like zinc (only Zn2+), shows non of the typical characteristics of transition

metal chemistry e.g. variable oxidation state, coloured complex ions, catalytic properties of the metal or ion. This is all explained in detail

later.

Therefore probably the best definition of a transition metal is an element which forms at least one ion with an incomplete d sub-

shell containing at least one electron. How this relates to variable oxidation state and coloured complex ions is elaborated further

in section 10.2 and the subsequent sections on the individual metals (links below) and some of the. Zinc (Zn2+, [Ar]3d10) and scandium

(Sc3+, [Ar]3d0) cannot meet this criteria.

The presence of the partially-filled d sub-shells of electrons gives transition elements properties which are not in general possessed by the

main group elements, namely Groups 1-7 and 0, BUT, there are similarities with other metals, particularly in Groups 2, 3 and  4.

PLEASE NOTE the following about these Transition Elements notes:

o All the reactions are described with visual observations and full ionic equations whether redox reactions or not.

o I have made extended use of standard electrode potentials to indicate not only the relative oxidising/reducing power of a half-

cell reaction, but also to argue for the thermodynamic feasibility of a reaction.

However, if you have done little on using electrode potentials you should still be able to follow the reaction descriptions

and the accompanying equations.

Electrode potentials (half-cell potentials) and how to work out EØreaction (which must be a +ve voltage for the reaction to be

feasible) are all explained in Equilibria Part 7 Redox Equilibria and Electrode Potentials . How to use half-cell

reactions to work out full redox equations is explained onRedox Reactions Part II section 6 Constructing redox

equations.

o In the latest Periodic Table convention, the 3d-block elements are considered the 'head elements' of Groups 3-12.

Groups 1-2 remain unchanged but Groups 3-7 and 0 become Groups 13-18. I tend to retain the Groups 3-7 and 0

convention for the moment but future is 13-18!

'latest'

Group

number

3 4 5 6 7 8 9 10 11 12

Period 4 21, Sc 22, Ti 23, V 24, Cr 25, Mn 26, Fe 27, Co 28, Ni 29, Cu 30, Zn

Period 5 39, Y 40, Zr 41, Nb 42, Mo 43, Tc 44, Ru 45, Rh 46, Pd 47, Ag 48, Cd

Period 6 57, La 72, Hf 73, Ta 74, W 75, Re 76, Os 77, Ir 78, Pt 79, Au 80, Hg

Outer

electron

s(n = 3 to

5)

n

d1(n+1)s2

n

d2(n+1)s2

n

d3(n+1)s2

n

d5(n+1)s1

n

d5(n+1)s2

n

d6(n+1)s2

n

d7(n+1)s2

n

d8(n+1)s2

n

d10(n+1)s1

n

d10(n+1)s2

o There are actually many 'vertical' chemical similarities in a 'classic' periodic table way of thinking to justify this

latest 'numbering' of the Periodic Table. e.g.

In most cases the three elements quoted above, per vertical column, have the same outer electron configuration.

'Modern Group 3': Scandium and yttrium have very similar with a relatively simple M3+ ion chemistry.

'Modern Group 10': Nickel, palladium and platinum are good hydrogenation catalysts. They all tend to form more square

planar complexes than other transition elements.

'Group 11': Copper, silver and gold are relatively unreactive metals in terms of corrosion. They form linear complexes like

the cationic, [Ag(NH3)2]2+ or the anionic [CuCl2]- and [Au(CN)2]-. All three are extremely good conductors of heat and

electricity.

'Modern Group 12': Zinc and cadmium chemistry is mainly about the M2+ ion.

From modern 'Group 3 to 7' the maximum known oxidation state known (albeit in some pretty unstable compounds at

times) is equal to the 'new' group number i.e. Sc/Y/La (+3) to Mn/Tc/Re (+7).

The discontinuity of atomic/proton number from lanthanum to hafnium on period 6 is due to the insertion of the 4f-block

elements 58Ce to 71Lu.

Comparison of certain properties of the 3d block of metals and other elements for Z = 1 to 38 particularly the preceding Group 1

metal potassium and the Group 2 metal calcium.

o Periodicity plots for elements Z = 1 to 38  Look for Z = 21 (Sc) to 30 (Zn)

 Note: There is no direct link back to here, so use <== 'back' on browser bar.

o Melting/boiling points: Generally higher than other elements in period 4.

o 1st ionisation energy: The 3d block 1st ionisation energies tend to increase from left to right and fit in with the general pattern for

period 4.

o Pauling electronegativity: The 3d-block values range from a relatively low 1.3 to 1.9 and fit in with the general pattern of

increasing value across period 4.

o Atomic radius: 3d-block elements have similar values and significantly less than for potassium and calcium.

o Electrical/thermal conductivity: The 3d-block are quite good conductors of electricity/heat and very good in the case of copper

(ditto silver Ag below Cu).

o Density: 3d-block range from 3.0 to 8.9g/cm3 and significantly more than for potassium (0.86) and calcium (1.5).

o Periodicity plots for elements Z = 1 to 96  if you want to look for the 4d and 5d blocks!

Other comparison points of the elements titanium to copper (true transition metals) with nearby metals.

o Potassium (+1), calcium (+2) and scandium (+3) only have one oxidation state in compounds, whereas Ti to Cu have compounds

in at least at least three oxidation states, even if some are not very stable!

 

10.2. Introduction - information & general characteristics of  3d block Metals Sc-Zn

Data Table 1 - summary of selected properties - concentrating only on the 3d-block

Z and symbol 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn

property\name scandium titanium vanadiumchromiu

m

manganes

eiron cobalt nickel copper zinc

melting point/oC 1541 1668 1910 1857 1246 1538 1495 1455 1083 420

density/gcm-3 2.99 4.54 6.11 7.19 7.33 7.87 8.90 8.90 8.92 7.13

atomic radius/pm 161 145 132 125 124 124 125 125 128 133

M2+ ionic radius/pm na 90 88 84 80 76 74 72 69 74

M3+ ionic radius/pm 81 76 74 69 66 64 63 62 na na

common oxidation states +3 only +2,3,4 +2,3,4,5 +2,3,6 +2,3,4,6,7 +2,3,6 +2,3 +2,+3 +1,2 +2 only

outer electron config. 3d14s2 3d24s2 3d34s2 3d54s1 3d54s2 3d64s2 3d74s2 3d84s2 3d104s1 3d104s2

Electrode potential M(s)/M2+(aq) na -1.63V -1.18V -0.90V -1.18V -0.44V -0.28V -0.26V +0.34V -0.76V

Electrode potential M(s)/M3+(aq) -2.03V -1.21V -0.85V -0.74V -0.28V -0.04V +0.40 na na na

Electrode potential M2+(aq)/M3+

(aq)

na -0.37V -0.26V -0.42V +1.52V +0.77V +1.87V na na na

CLICK for a more detailed data table 2 summary

General Physical Characteristics

The transition metals are the most important structural metals for industry due to their strength arising from the strong inter-atomic forces

(see metal bonding and alloy structure).

The strong bonding is due to small ionic radii and at least 3 delocalised 3d or 4s electrons contributing to the bonding which accounts for

their high tensile strength, malleability (can be readily beaten into shape) and ductility (can be drawn into wire).

They are silvery-grey solids apart from the dark orange of copper.

They generally have high melting/boiling points and densities and readily mix with themselves or other elements to give a huge variety of

alloys with a wide range of uses based on varied hardness, strength, malleability and anti-corrosion properties. 

There is a general, but small, contraction of the atomic/ionic radii across the series as the atomic/proton number rises, i.e. an increasing

positive attractive force on the outer electrons of the same sub-shells (3d and 4s).

10.2b. General Chemical Characteristics and electron configurations

21 Scandium, Sc 1s22s22p63s23p63d14s2[Ar]3d 4s

22 Titanium, Ti 1s22s22p63s23p63d24s2[Ar]3d 4s

23 Vanadium, V 1s22s22p63s23p63d34s2[Ar]3d 4s

24 Chromium, Cr 1s22s22p63s23p63d54s1[Ar]3d 4s

25 Manganese, Mn 1s22s22p63s23p63d54s2[Ar]3d 4s

26 Iron, Fe 1s22s22p63s23p63d64s2[Ar]3d 4s

27 Cobalt, Co 1s22s22p63s23p63d74s2[Ar]3d 4s

28 Nickel, Ni 1s22s22p63s23p63d84s2[Ar]3d 4s

29 Copper, Cu 1s22s22p63s23p63d104s1[Ar]3d 4s

30 Zinc, Zn 1s22s22p63s23p63d104s2[Ar]3d 4s

The chemistry is dominated by the behaviour of the 3d electrons. The 3d block corresponds to the filling of the

3d sub-shell of electrons, best appreciated by the 'box diagrams' of their electron structure.

Each half-arrow is an electron, which tend to singly occupy the sub-orbitals as much as possible to minimise

repulsion (Hund's Rule of maximum multiplicity). 

The outer electrons of the element are either in the 3d or 4s sub-shell. The 4s sub-shell is initially filled by

potassium [Ar]4s1 and calcium [Ar]4s2.

The electron arrangement for each element from Sc to Zn is also given at the start of each individual metal section in terms of s, p and d notation.

All 10 elements, Sc to Zn are 3d block elements (the filling of the 3d sub-shell) BUT a true transition element is one in which there is an

incomplete d sub-shell holding at least one electron in one or more chemically stable ions (Ti to Cu). For 3d block metals this means at

least one stable ion with the configuration within the range [Ar]3d1 e.g. Ti3+ to [Ar]3d9 e.g. Cu2+ and so excludes scandium and zinc. Zinc

only forms Zn2+, [Ar]3d10 and scandium only forms Sc3+, [Ar]3d0, so neither can meet this criteria for a true transition metal. See theory of colour in

transition metal complexes.

There are two apparent anomalies in the electron configuration sequence from left to right as the 3d sub-shell energy level is filled:

Cr is not 3d44s2 and Cu is not 3d94s2

because inner half-filled or fully-filled filled 3d sub-shells seem to be a little lower in energy, marginally more stable.

The total number of outer 3d/4s electrons is equal to the maximum oxidation state from Sc(+3) to Mn(+7) and there are many stable compounds

exhibiting these maximum oxidation states.  After Mn there is significantly less stability of species with the metal in oxidation states above +3 for

Fe and Co, and above +2 for Ni, Cu and Zn.

The four 'classic' chemical characteristics (but NOT unique to transition metals) are ...

(1) Complex formation: Appendix 2  offers an introduction as well as numerous examples 'en route' particularly from Ti to Cu.

(2) Formation of coloured ions: Appendix 4  offers an introduction to the origin of the colour in transition metal complex ions as well as

examples 'en route' from colourless 'non-transition' Sc3+ complexes, coloured TiII, III, IV to CuII 'true transition' complexes and finally colourless 'non-

transition' Zn2+ complexes at the end of the 3d-block.

(3) Variable oxidation state - variable valency:

From Sc to Mn the maximum oxidation state is determined by the total maximum number of 3d and 4s electrons. After that, things get very

complicated but the maximum tends to fall down to +2 for zinc after +3 for Fe and Co (there are some higher oxidation state species, but

not that common and not that stable in aqueous media).

o See Table of examples of compounds/complexes for different oxidation states of the 3d-block elements.

The relative ease of oxidation state change for Ti to Cu AND the maximum oxidation state formed by Sc to Mn, is partly explained by

considering the ionisation energies involved and a comparison with Group 1, 2 and 3 metals helps too.

o In the sequences below the atoms and ionised species are all in the gaseous state as is the convention for ionization energy data.

o The energies (kJmol-1) required to remove the next most loosely bond electron to give the next more highly charged ion (the next

higher oxidation state) are shown as a sequence.

o Only for the first example, potassium, are the full formal equations shown.

o The successive ionisation energy sequences for Group 1 (potassium), Group 2 (calcium), the 3d-block (e.g. titanium) and Group 3

(gallium) are now considered for period 4.

o Gp1: K(g) == +418 ==> K+(g) == +3070 ==> K2+

(g) 

they would be formally written as:

for the 1st ionisation energy: K(g) - e- ==> K+(g)  

and for the 2nd ionisation energy: K+(g) - e- ==> K2+

(g)  

o Gp2: Ca(g) == +590 ==> Ca+(g) == +1150 ==> Ca2+

(g) == +4940 ==> Ca3+(g) 

o 3d-block: e.g. Ti(g) == +661 ==> Ti+ == +1310 ==> Ti2+ == +2720 ==> Ti3+ == +4170 ==> Ti4+ == +9620 ==> Ti5+

o Gp3: Ga(g) == +577 ==> Ga+ == +1980 ==> Ga2+ == +2960 ==> Ga3+ == +6190 ==> Ga4+ 

So, for Groups 1, 2 and 3, the ionisation energy dramatically rises after the outer shell of s or p electrons are removed, i.e. a very stable

electronic noble gas structure ([Ar], 1s22s22p63s23p6) is left. This gives a maximum positive stable oxidation state equal to the group

number. The energy required (very endothermic) to make Na2+, Ca3+ and Ga4+ is too high to be compensated by exothermic bond

formation with other elements like oxygen or chlorine etc.

o Also note that intermediate lower oxidation states Ca+ and Ga2+ (and  Ga+?)are not very stable either.

o I'm afraid ionisation energies and electron arrangements are not the only factors to be considered, you also need to study the

Born Haber Cycle in some detail to prove this, but not here and not usually on a pre-university course!

For the transition metals, at first, the successive ionisation energies rise relatively gradually, due to the 3d/4s electron levels being of

similar energy. When all the outer s and d electrons are removed to leave an [Ar] core, there is, as with Groups 1-3 etc., a dramatic rise as

an electron must be removed from the inner very stable noble gas (argon) core.

o Therefore Ti has a maximum oxidation state of +4, but +2 and +3 species are also formed, but NOT +5 compounds.

o This does mean however, across the 3d-block, there is the potential for very high oxidation states if there are enough 3s and 3d

electrons that can be energetically favourably removed or become involved in stable bonding e.g. Mn has a maximum oxidation

state of +7 by 'removing  *  or 'sharing' the outer 3d54s2 electrons. (see data table).

Similarly you can argue that the maximum oxidation states for vanadium would be +5 and chromium +6, as is indeed is

the case!

After manganese, things get complicated and there is a general decrease from Mn (+7) to Zn (+2) in the maximum

possible higher oxidation states, and many higher oxidation state compounds of Fe, Co, Ni and Cu are unstable and

uncommon.

o  *  Of course e.g. in manganese (VII) compounds, 7 electrons are not removed to give an Mn7+ ion, but all 7 outer electrons are

involved in the bonding and, unlike calcium and gallium, true transition metals form many stable compounds of the 'intermediate'

oxidation states e.g. manganese forms +2, +3, +4, +6, +7 oxidation sate compounds.

o  This is due to closeness of the energies of the 3d sub-shell electrons and the stabilising influence of ligand molecules like

water or ammonia and ligand ions like chloride or cyanide. Vacant 3d orbitals (and 4s/4p orbitals too) can accept pairs of electrons

to for stable dative covalent bonds.

I'm afraid arguments for the characteristic variable oxidation states of transition metals based on ionisation energies and

the similar energies of the 3d orbitals is a bit limited, but better than nothing!

I have done detailed notes on oxidation state/oxidation number and redox reactions - a lot of which come up in

transition metal chemistry.

(4) Catalytic activity by the elements and their compounds:

Doc Brown's Chemistry  Advanced Level Inorganic Chemistry Periodic Table Revision Notes - Transition

Metals

 Appendix 2 Complexes - introduction: ligands, bonding, co-ordination number and

charge on complex ions

What is a complex ion? What is a ligand? What do the terms monodentate ligand, bidentate ligand and

polydentate ligand mean? What is the co-ordination number of a complex ion?  The structure of transition

metal (3d-block) complexes is described with displayed formula diagrams and explainations include the

formation of central metal ion - ligand dative covalent bonds. What shapes can complexes be? e.g.

octahedral, tetrahedral, square planar and linear examples are presented.

  GCSE/IGCSE Periodic Table Revision Notes  *    GCSE/IGCSE Transition Metals Revision Notes

INORGANIC Part 10 3d block TRANSITION METALS sub-index: 10.1-10.2 Introduction 3d-block Transition

Metals * 10.3 Scandium * 10.4 Titanium * 10.5Vanadium * 10.6 Chromium * 10.7 Manganese * 10.8 Iron *

10.9  Cobalt * 10.10 Nickel * 10.11 Copper * 10.12 Zinc * 10.13 Other Transition Metals e.g. Ag and Pt *

Appendix 1. Hydrated salts, acidity of hexa-aqua ions * Appendix 2. Complexes & ligands * Appendix

3. Complexes and isomerism * Appendix 4.Electron configuration & colour theory * Appendix 5. Redox

equations, feasibility, Eø * Appendix 6. Catalysis * Appendix 7. Redox equations  * Appendix 8.Stability

Constants and entropy changes * Appendix 9. Colorimetric analysis and complex ion formula * Appendix

10 3d block - extended data * Appendix 11 Some 3d-block compounds, complexes, oxidation states &

electrode potentials * Appendix 12 Hydroxide complex precipitate 'pictures' , formulae and equations

Advanced Level Inorganic Chemistry Periodic Table Index * Part 1 Periodic Table history * Part 2 Electron

configurations, spectroscopy, hydrogen spectrum, ionisation energies * Part 3 Period 1 survey H to

He * Part 4 Period 2 survey Li to Ne * Part 5 Period 3 survey Na to Ar * Part 6 Period 4 survey K to Kr and

important trends down a group * Part 7 s-block Groups 1/2 Alkali Metals/Alkaline Earth Metals * Part 8  p-

block Groups 3/13 to 0/18 * Part 9 Group 7/17 The Halogens * Part 10 3d block elements & Transition Metal

Series * Part 11 Group & Series data & periodicity plots * All 11 Parts have their own sub-indexes near the

top of the pages

Appendix 2. Complexes - introduction: ligands, bonding, co-ordination number and

charge on complex ions

A complex is formed by the combination of a central metal ion surrounded by, and bonded to,

neutral molecules or ions acting as 'ligands' (bits stuck on or appendages).

o If you have already read Appendix 1. you should note that it is riddled with complex ions and the

central metal ion does NOT have to be a transition element. The two ligands involved were H2O and

OH-.

A ligand is an atom, ion or molecule which can act as an electron pair donor (Lewis base) and usually

forms a dative covalent or 'co-ordinate' bond with the central metal ion.

o The lone pair donation is usually from an O, N or halogen atom of the ligand in this covalent co-

ordinate bonding.

o The central metal ion acts as a Lewis Acid, that is, an electron pair acceptor from the ligand by

way of vacant 3d, 4s, 4p  and even 4d orbitals for the 3d-block transition elements.

o The ligand acts as a Lewis Base, that is, an electron pair donor e.g. neutral ligands

like H2O: (water, aqua in complex name) or :NH3 (ammonia, ammine in complex name) and

negatively charged ligands like :OH- (hydroxide, hydroxo in complex name), Cl- (chloride ion, chloro

in complex name) and :CN-(cyanide ion, cyano in complex name).

o ...

A an example of the bonding in a complex ion is shown in the above diagram. The negative

cyanide ion is a monodentate ligand (forms one bond per ligand) and donates an electron

pair into a vacant 3d, 4s or 4p orbital in the iron(III) ion to form six dative covalent bonds.

The resulting ion has the formula [Fe(CN)6]3-, the overall charge of 3- is the aggregate of 6-

(cyanide ions) plus 3+ (iron ion)

The co-ordination number of 6, which means there are 6 central metal ion - ligand bonds.

It doesn't necessarily mean six ligands, you can get a co-ordination number of 6 from three

co-ordinated bidentate ligands (2 bonds per ligand), two tridentate ligands and from EDTA

just one ligand can form 6 dative covalent bonds with a central metal ion. More on this

below.

The most common complex ion you will come across is the hexaaqua cation of many

metals.

It has the general formula [M(H2O)6]n+

n, the charge on the central metal ion and hence the overall charge on the complex

ion n is usually 2 or 3 e.g.

n = 2 for titanium(II), vanadium(II), iron(II), cobalt(II), nickel(II), copper(II) and also

the Group 2 alkaline Earth metals magnesium, calcium etc.

and n is 3 for scandium, titanium(III), vanadium(III), chromium(III), iron(III), cobalt(III)

and also aluminium from Group 3.

The six neutral water ligands form 6 dative covalent bonds with the central

metal ion because the bonding pair of electrons comes from donation of a lone pair

from the oxygen atom of the water molecule.

Therefore the co-ordination number is 6 and it has a symmetrical octahedral shape.

The O-M-O bond angles are all 90o or 180o.

The ligand may attach itself by one or more bonds. The suffix '...dentate', prefixed by

mono/uni/bi/ploy/multi e.g. monodentate (unidentate), bidentate, orpolydentate (multidentate) is used to

denote the number of bonds each ligand makes with the central metal ion.

The total number of ligand bonds to the central metal ion is called the co-ordination

number.

o It is not the number of ligands, unless it is a monodentate ligand.

o There is no firm rules relating shape to a particular ligand.

o The six ligands don't have to be the same e.g.  ...

... which is the dichlorotetraaquachromium(III) ion. This octahedral complex with a co-

ordination number of 6, and note this has an overall ion charge of (2 x - from 2Cl-) + (3+

from Cr3+) = +, water is an electrically neutral ligand ...

... and in equations the complex ion would be written as [Cr(H2O)4Cl2]+

Examples of unidentate/monodentate ligands:

o

o e.g. above are shown two complexes with electrically neutral ligands: water H2O:,

ammonia :NH3 and primary aliphatic amines like butylamineCH3CH2CH2CH2NH2, 

o These ligands often form octahedral shaped complexes with a co-ordination

number of 6.

o e.g. negative ligands: chloride Cl-, cyanide CN-,

o The chloride ion Cl- forms the tetrahedral e.g. the

tetrachlorocuprate(II) complex ion ...

o [CuCl4]2-, note the overall charge is (2+) + (4 x -) = 2- and the co-ordination number is 4.

o The chloride ion can be too bulky to form an octahedral complex or a square planar complex,

though there is no firm rules relating complex shape to ligand.

o and CN- square planar e.g. the tetracyanonickelate(II) complex ion ...

o [Ni(CN)4]2-, note the overall charge is (2+) + (4 x -) = 2- and the co-ordination number is 4.

Note that [Cu(H2O)4]2+, in the hydrated salt CuSO4.5H2O, the tetraaquacopper(II) ion, with

the less bulky water molecule ligand, forms a blue square planar complex, whereas with the

larger chloride ion, a tetrahedral complex is formed.

o A linear shaped complex is formed between a silver ion the ligands ammonia or

cyanide.

cationic [H3N-Ag-NH3]+  and anionic [NC-Ag-CN]-

o [Ag(NH3)2]+ is formed in 'ammoniacal' silver nitrate solution used in the test for aldehydes. The

diamminesilver(I) ion has co-ordination number of 2 and an overall charge of a single + because the

ammonia molecule is an electrically neutral ligand.

Examples of bidentate ('two toothed') ligands:

o neutral ligands: diamines like 1,2-diaminoethane (ethane-1,2-diamine) H2NCH2CH2NH2 (bonds via

lone pair :N).

o negative ligands: ethanedioate ion C2O42-, (bonds via lone pair on the :O-). The L

represents where the dative covalent bond forms.

o shows three bidentate ligands co-ordinated to a central metal ion (co-ordination

number 6, 'octahedral' in bond arrangement).

o Examples: [Cr(H2NCH2CH2NH2)3]3+, H2NCH2CH2NH2 is often represented in shorthand by en,

and the complex simply written as [Cr(en)3]3+.

o Bidentate ligands are the first of what are called polydentate ligands and such complexes are

sometimes called chelates from the Greek for 'crab's claw' and the complex formation described as

a chelation process.

More examples of multi/polydentate ligands:

o EDTA4- (old name 'EthyleneDiamineTetraAcetic acid') forms six bonds with a central metal ion and

tends to displace all other ligands.

[Ni(NH3)6]2+(aq) + EDTA4-

(aq)   [Ni(EDTA)]2-(aq) + 6NH3(aq)

o The haemoglobin molecule acts as a multi/polydentate ligand with iron(II) ions in blood chemistry.

in an extremely simplified form the structure is: [protein-FeII-O2]

One ligand can replace another depending on the relative bond strengths in a reaction called ligand

exchange reaction.

When a bidentate or polydentate ligand is added to a pre-existing complex of monodentate ligands, it is

highly likely a more stable complex will be formed.

o The principal reason for this, (ignoring bond strengths), is the positive entropy

change accompanying the 'release' of 4 or 6 small molecules which offer a greater variation of ways

of arranging the particles or energy distribution.

If the ligands are easily exchanged, the complex is described as 'unstable' and if the ligands are more

strongly bound, the complex would be described as stable.

Complex ion stability is also related to the oxidation state of the transition metal in the presence of a

particular ligand.