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Topic 10 – Transition Metals and Coordination Chemistry Hydrolysis of Metal Ions – aqueous metal ions are acidic

Metal cations act as Lewis acids while water is the Lewis base. Recap definitions:

• Arrhenius: H+(aq) + OH-(aq) ⇌ H2O(l)

– Acid: H+ producer in aqueous solution e.g. AH – Base: OH- producer e.g. MOH

• Brønsted - Lowry: H+ + A- ⇌ HA

– Acid: proton donor (H+) e.g. HCl

– Base: proton acceptor e.g. NH3

• Lewis: A + :B ⇌ A:B

– Acid: electron pair acceptor e.g. BF3 – Base: electron pair donor e.g. :NH3

The Coordination Bond A ligand donates an electron pair towards the metal ion to form a coordinate covalent bond.

• The lone pair is attracted towards the metal ion

• Ligands must have at least one lone pair • More than one ligand can bind to a metal ion

The Structure of Metal Complexes A coordination compound typically consists of a complex ion (a

transition metal ion with attached ligands) and counter ions -

anions or cations as needed to produce a compound with no net charge.

e.g. [Co(NH3)6]Cl3 The square brackets indicate the complex ion and the three Cl- ions

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the counter ions. There are 6 NH3 ligands. When dissolved in water the complex exists as the

cation and the 3 Cl- ions are separate.

[Co(NH3)6]Cl3 ⇌ [Co(NH3)6]3+(aq) + 3 Cl- (aq)

Acidity of Aqueous Metal Ions e.g. [Al (OH2)6]3+ ⇌ [Al (OH2)5(OH)]2+ + H+

Free Ion Hydrated Ion Ka

Acidity increases

Fe3+ Fe(H2O)63+(aq) 6 x 10-3

Cr3+ Cr(H2O)63+(aq) 1 x 10-4

Al3+ Al(H2O)63+(aq) 1 x 10-5

Be2+ Be(H2O)42+(aq) 4 x 10-6

Cu2+ Cu(H2O)62+(aq) 3 x 10-8

Pb2+ Pb(H2O)62+(aq) 3 x 10-8

Zn2+ Zn(H2O)62+(aq) 1 x 10-9

Co2+ Co(H2O)62+(aq) 2 x 10-10

Ni2+ Ni(H2O)62+(aq) 1 x 10-10

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Coordination Number and Geometry The number of ligands bonded to the metal ion:

This varies from 2 to 8 and depends on the size, charge and electron configuration of the metal ion.

Typical coordination number for some metal ions. M+ Coord no. M2+ Coord no. M3+ Coord no.

Cu+ 2,4 Mn2+ 4,6 Sc3+ 6 Ag+ 2 Fe2+ 6 Cr3+ 6

Au+ 2,4 Co2+ 4,6 Co3+ 6 Ni2+ 4,6 Au3+ 4

Cu2+ 4,6 Zn2+ 4,6

Coordination number

Coordination geometry

Example

2 linear [Ag(NH3)2]+ [AuCl2]-

4 square planar [Pd(NH3)4]2+

[PtCl4]2-

4 tetrahedral [Zn(NH3)4]2+

[CuCl4]2-

6 octahedral [Co(NH3)6]3+

[FeCl6]3-

A ligand is a neutral molecule or an ion having a lone pair of electrons that can be used to form a bond with a metal ion.

Unidentate (Monodentate) ligand – forms one bond (eg OH2, NH3, –OH, –Cl)

Bidentate ligand – can form two bonds (e.g. ethan-1,2-diamine (en))

Complexes with chelate ligands are usually more stable than those with monodentate ligands.

HN

HH

Mx+

H2C CH2

H2N NH2

Mx+

Unidentate Ligande.g. ammonia

Bidentate ligande.g. ethan-1,3-diamine (en)

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Polydentate ligand – can form more than two bonds

• Tridentate

• tetradentate

• pentadentate

• hexadentate

(e.g. EDTA, ethane-1,2-diaminetetraacetic acid used to treat severe lead poisoning)

[WS 33] Any molecule that possesses a lone pair of electrons can donate these to a transition metal cation.

If it does so, it becomes a ligand. In general, a metal cation bonds or coordinates to 4 or 6 ligands to form a complex ion. The complex ion is written in square brackets. If the charges on the metal

cation and the ligand do not balance, counter ions are needed. These are not bonded to the metal. The complex ion and the counter ions together make a coordination compound.

ON

O

O

NO

O OH

O

O

Polydentate ligande.g. EDTA

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Questions: 1. Complete the following table. In each case, draw the complex ion with tapered and dotted

bonds.

Geometry Coordination

compound Complex ion

Ligand(s) (underline the donor

atom)

Counter ion

octahedral [Co(NH3)6]Cl3

Co

NH3

NH3

NH3H3NNH3H3N

3+

[Co(NH3)6]3+

NH3 Cl-

octahedral [Fe(OH2)6]Br3

tetrahedral K2[NiCl4]

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The chelate effect The higher stability of complexes containing polydentate ligands over that for analogous complexes

involving monodentate ligands is called the chelate effect. It is of huge importance in biology and in medicine. The reaction below strongly favours the chelate complex:

[Ni(NH3)6]2+ + 3 en [Ni(en)3]2+ + 6NH3 K = 2.0 × 109

Questions: 1. What bonds are being made and broken in the reaction, and how many of each.

2. Given your answer to Q1, predict whether ΔrH will be (a) positive, (b) negative or (c) close to

zero.

3. By considering the number of reactant and product molecules, predict the sign of ΔrS.

4. Using your answer to Q2 and Q3, explain why the reaction occurs (i.e. the thermodynamic

basis of the chelate effect.)

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Naming Metal Complexes 1) Naming Ligands: The normal chemical name is used unless the ligand is negatively charged in which cases “o” is used as the suffix.

Ligand Name Formula

Neutral Ligands aqua H2O

ammine NH3

Anionic Ligands

fluorido F–

chlorido Cl–

bromido Br–

iodido I–

cyanido CN–

hydroxido OH–

2) Naming Metal Ions: (a) If the complex is neutral or positively charged the normal metal name is used.

(b) If the complex is negatively charged, ‘ate’ is added to the metal name.

e.g. Co: cobaltate, Zn: zincate, etc. There are a few special cases of metal names for anionic complexes: argentite (Ag), ferrate (Fe), cuprate (Cu), plumbate (Pb).

3) Specifying the number of ligands: The number of ligands of any one type is indicated with the appropriate Greek prefix.

monodentate: di, tri, tetra, penta, hexa, etc. polydentate: bis, tris, tetrakis

4) Ordering the names: Ligands are named first and are listed in alphabetical order (Note: prefixes do not affect the order). 5) Specifying the oxidation state: The oxidation state of the metal is indicated by roman numerals.

e.g. Fe3+ is given as iron(III)

6) Naming complex salts: If cations or anions are present they are named as separate words and are not numbered. (There

are no differences to the rules for naming simple salts.) e.g. [CoCl2(NH3)4]Cl tetraamminedichloridocobalt(III) chloride

7) Indicating the presence of solvent molecules: Water of crystallisation (hydration) is indicated separately at the end of the name.

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e.g. [CoCl2(NH3)4]Cl·2H2O tetraamminedichloridocobalt(III) chloride-2-water

Question: Try naming/drawing:

tetraammineaquacyanidocob

alt(III) chloride

Na[Al(OH)4

[Ni(OH2)6]Br2

Potassium

amminepentafluoridoplatinate(IV)

[Co(en)2(OH2)2]Cl3

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Isomerism in Metal Complexes

• Structural Isomerism: different atom connectivity

Coordination sphere isomerism: [CrBr2(OH2)4]Cl [CrBrCl(OH2)4]Br

• Stereoisomerism: same atom connectivity but different arrangement of atoms in space (a) Geometric isomerism:

(b) Optical isomerism:

trans isomer : superposable cis isomer: non-superposable and optically

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

Equilibria Involving Complexes

(From

Silberberg, McGraw Hill, 4th Ed.)

NiN N

NNOH2

OH2

2+

NiN N

NNOH2

OH2

2+ 2+

NiH2O N

NH2ON

N

NiN OH2

OH2NN

N

2+

N

Co

N

N

N

N

NN

Co

N

N

N

N

N

mirror plane

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Unlike the covalently bonded polyatomic ions such as NO3-, SO4

2-, etc., which do not dissociate

into their components, metal complexes in solution are in an equilibrium:

e.g. [M(OH2)4]2+ + 4 NH3 ⇌ [M(NH3)4]2+

+ 4 H2O

Like any other equilibrium, we can treat this quantitatively by assigning an equilibrium constant, known as the stability constant, Kstab

e.g. [M(OH2)4]2+ + 4 NH3 ⇌ [M(NH3)4]2+

+ 4 H2O

Kstab = [ [M(NH3)4]2+ ]

[ [M(OH2)4]2+ ] [NH3]4

e.g. [Cu(NH3)4]2+, Kstab = 5.6 x 1011 [Zn(NH3)4]2+, Kstab = 7.8 x 108 Kstab = stability constant or formation constant. The larger Kstab, the more stable is the complex.

[WS34] Geometrical isomers - same bonds but different arrangement of atoms in space The common geometries of coordination complexes are octahedral, tetrahedral and square planar, as represented in the table below.

Octahedral (6 coordinate) Tetrahedral (4 coordinate) Square Planar (4 coordinate)

Co

NH3

NH3

NH3H3NNH3H3N

3+

[Co(NH3)6]3+

2-

Hg

I

I II

[HgI4]2-

AuClClClCl

-

[AuCl4]-

In each one of these, each bond is the same. The tapered and dotted bonds are used to try to indicate the three-dimensional structure of the complex.

Because each bond is the same, there are no isomers possible for octahedral complexes

containing 5 ligands of one sort and 1 ligand of a different sort. Similarly, no isomers are possible

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for tetrahedral and square planar complexes containing 3 ligands of one sort and 1 ligand of a

different sort.

Questions: 1. Two isomers are possible for an octahedral complex like [NiCl2(NH3)4] which has 4 ligands of one sort and 2 of another. Draw these isomers and make sure you are happy that there are no other

isomers.

2. For 4-coordinate complexes, two geometries are found: tetrahedral and square planar. [PtCl2(NH3)2] is found in two isomeric forms, one of which has considerable anti-tumour activity. Is it

tetrahedral or square planar?

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Optical isomers – non-superimposable mirror images Optical isomerism or chirality is not just possible in organic chemistry! If a molecule is not the same

(“non-superimposable”) on its mirror image, it is chiral and can exist in two enantiomeric forms.

The picture below shows the structure of the [Ni(en)3]2+ complex ion. The ligand “en” is NH2CH2CH2NH2 and, as you can see, this bonds to the metal through both N atoms.

3. Carefully draw the mirror image of the complex. Is it the same (“superposable”) or different (“non-

superimposable”) to the form already drawn?

Ni

NH2

NH2

NH2H2NNH2H2N

4. The complex ion [NiCl2(en)2] is octahedral and exists in three isomeric forms. Draw these isomers.

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Transition or d-block Metals

Filling of Atomic Orbitals and Orbital Energy Ordering

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Oxidation States of the d-block metals

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Properties of the Period 4 Elements

4d and 5d elements have very similar sizes due to the ‘lanthanide contraction’

One of the most characteristic chemical properties of these elements is the occurrence of multiple oxidation states, often associated with different colours.

Ion Ox. state Colour

VO3- V yellow

VO2+ IV green

V3+ III blue V2+ II violet

CdS: yellow Cr2O3: green TiO2/ZnO2 : white Mn3(PO4)2: purple Co2O3/Al2O3: blue Fe2O3: ochre

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Electron Transfer Reactions

e.g. 5 Fe2+ + Mn7+ → 5 Fe3+ + Mn2+

Colourful Complexes

Aqueous solutions of the Co(III) complexes (from left to right see lecture slides): [Co(NH3)5OH2]3+, [Co(NH3)6]3+, trans-[CoCl2 (en)2]+, [CoCO (en)2O2]+ and [CoCl(NH3)5]2+.

As the metal ion is the same in all cases, the variety of colours arises from the different ligands surrounding the Co(III) ion.

Absorbed and Observerved Colours In many transition metal complexes d-electrons absorb energy in the visible range when they are promoted to higher energy excited states.

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Colourful Compounds

Paramagnetism of Transition Metals Unpaired electrons confer the property of being attracted to a magnet, a property retained when

compounds are formed. “Alnico” magnets, e.g., are an alloy containing Al, Ni and Co.

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[WS 35] The table below lists common ligands and their charges. You should notice that these charges are just those that you are used to seeing on these ligands when you have encountered

them elsewhere.

Neutral Ligands Anionic Ligands

OH2 NH3, NH2CH2CH2NH2 (en), pyridine (py) F-, Cl-, Br-, I-, OH-, O2-, CN-, SO4

2-, CH3CO2-,

EDTA4-

If you know the formula and charge of the complex ion or the formula of the coordination

compound, you can work out what oxidation number the transition metal cation must have as:

(i) the sum of the charges of the metal cation and the ligands adds up to give the charge of the

complex ion, and (ii) the sum of the charges of the complex ion and the counter ions adds up zero.

Examples (a) K2[NiCl4] contains 2K+ counter ions and a [NiCl4]2- complex ion. The latter contains 4Cl- so it

must contain Ni2+.

(b) The complex ion [Cr(en)3]3+ contain three neutral “en” ligands. The complex ion has a +3 charge so must contain Cr3+.

(c) [Cr(en)2Cl2] has no counter ions so the complex ion is neutral. “en” is a neutral ligand and there are 2Cl-, it must contain Cr2+.

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Questions: 1. Complete the ‘oxidation number’ column of the table overleaf by working out the oxidation number

of each of the transition metal cations.

The electronic configuration of transition metal cations The number of valence electrons on an atom is equal to its group number. In a cation, the oxidation

number is equal to the number of these electrons which have been removed.

Transition metal cations have a configuration 3dz where z is the number of valence electrons left:

z = number of valence electrons on atom – charge of cation = group number – oxidation number

For example: (a) Ni is in group 10 so Ni2+ has (10 – 2) = 8 valence electrons left: it has a d8 configuration.

(b) Cr is in group 6 so Cr3+ has (6 – 3) = 3 valence electrons left: it has a d3 configuration. (c) Cr is in group 6 so Cr2+ has (6 – 2) = 4 valence electrons left: it has a d4 configuration.

Questions: 1. Complete the ‘d configuration’ column in the table by working out z for each of the transition metal cations.

To minimize repulsion, electrons occupy orbitals singly before pairing up.

There five d-orbitals and, as each orbital can accommodate two electrons, there is space for a

maximum of ten electrons The electrons are added to these orbitals using Model 3. This has been completed for the complexes in the first three rows of the table. If process leads to unpaired electrons,

the complex is paramagnetic and is attracted towards magnetic field. If there are no unpaired electrons, the complex is diamagnetic and is repelled by magnets.

Questions: 1. Show the electron configuration for the transition metal cation using the box notation in the table.

2. Indicate if the complex is paramagnetic or not in the final column of the table.

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Coordination Compound or Complex

Oxidation Number

d Configuration

Electron Arrangement Paramagnetic?

K2[NiCl4] +2 d8

Yes

[Cr(en)3]3+ +3 d3

Yes

[Cr(en)2Cl2] +2 d4

Yes

[Fe(OH2)6]Br3

K3[Fe(CN)6]

[Ni(en)3]I2*

Na2[Zn(OH)4]

NH4VO3

VO2+

V3+

V2+

*en is NH2CH2CH2NH2 and can bond through lone pairs on both N atoms.