CHEM1612 - Pharmacy Week 8: Complexes I

Dr. Siegbert Schmid

School of Chemistry, Rm 223

Phone: 9351 4196

E-mail: siegbert.schmid@sydney.edu.au

Unless otherwise stated, all images in this file have been reproduced from:

Blackman, Bottle, Schmid, Mocerino and Wille,     Chemistry, John Wiley & Sons Australia, Ltd. 2008

     ISBN: 9 78047081 0866

Lecture 22-3

Complexes

Blackman Chapter 13 and Sections 10.4, 11.8 Biologically important metal-complexes Complex ions Kstab

Coordination compounds Chelates Geometry of complexes Solubility and complexes Nomenclature Isomerism in complexes

Co(EDTA)-

Lecture 22-4

Metal Ions as Lewis Acids

M2+

H2O(l)

[M(H2O)4]2+

adduct

º M2+(aq)

(Hydrated M2+ ion)

Whenever a metal ion enters water, a complex ion forms with water as the ligand.

Metal ions act as Lewis acid (accepts electron pair). Water is the Lewis base (donates electron pair).

Lecture 22-5

Complex Ions Definition: A central metal ion covalently bound to two or more

anions or molecules, called ligands.

Neutral ligands e.g.: water, CO, NH3

Ionic ligands e.g.: OH-, Cl-, CN-

[Ni(H2O)6]2+, a typical complex ion.

Ni2+ is the central metal ion Six H2O molecules are the ligands overall 2+ charge.

Blackman Figure 13.12

Lecture 22-6

They consist of:• Complex ion (metal ion with attached ligands)• Counter ions (additional anions/cations needed for zero net charge)

Eg. [Co(NH3)6]Cl3 (s) [Co(NH3)6]3+(aq) + 3 Cl-(aq)

Coordination Compounds

Complex ion Counter ions

e.g. [Co(H2O)6][CoCl4]3 (s) [Co(H2O)6]3+(aq) + 3 [CoCl4]-

(aq)

In water coordination compounds behave like electrolytes: the complex ion exists as the cation and the 3 Cl- ions are separate.

Note: the counter ion may also be a complex ion.

Lecture 22-7

CoordinationCompound

ComplexIon

CounterIons

Ligands within the coordination sphere remain bound to the metal ion

Coordination compoundsFigure from Silberberg, “Chemistry”, McGraw Hill, 2006.

Lecture 22-8

e.g. Ag+(aq) + 2 NH3 Ag(NH3)2

+(aq)

Ligands must have a lone pair to donate to the metal. The ‘donation’ of the electron pair is sometimes referred to as a

“dative” bond.

Complex Ions

Lecture 22-9

A small and multiply-charged metal ion acts as an acid in water, i.e. the hydrated metal ion transfers an H+ ion to water.

6 bound H2O molecules5 bound H2O molecules

1 bound OH- (overall charge reduced by 1)

Acidity of Aqueous Transition Metal Ions

Acidicsolution

Figure from Silberberg, “Chemistry”, McGraw Hill, 2006.

Lecture 22-10

Free Ion Hydrated Ion Ka

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

Fe2+ Fe(H2O)62+(aq) 4 x 10-9

Pb2+ Pb(H2O)62+(aq) 3 x 10-9

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

AC

ID S

TR

EN

GT

H

Metal Ion HydrolysisEach hydrated metal ion that transfers a proton to water has a characteristic Ka value.

Lecture 22-11

M+ Coord no. M2+ Coord no. M3+ Coord no.Cu+ 2,4 Mn2+ 4,6 Sc3+ 6Ag+ 2 Fe2+ 6 Cr3+ 6Au+ 2,4 Co2+ 4,6 Co3+ 6

Ni2+ 4,6 Au3+ 4Cu2+ 4,6

Zn2+ 4,6

The number of ligand atoms attached to the metal ion is called the coordination number. varies from 2 to 8 and depends on the size, charge, and electron

configuration of the metal ion.

Typical coordination numbers for some metal ions are:

Coordination number

Lecture 22-12

Coordination Number and GeometryRemember Valence Shell Electron Pair Repulsion Theory (VSEPR)?

F

SbF

F F

FN C O S

FF

F

:

: :

::

: ::

:

F: ::

F:

:

:

F: ::

Blackman Chapter 5

Lecture 22-13

Coordination number

Coordination

geometry

2 linear

4 square planar

4 tetrahedral

6 octahedral

Examples

[Ag(NH3)2]+

[AuCl2]-

[Pd(NH3)4]2

+

[PtCl4]2-

[Zn(NH3)4]2

+

[CuCl4]2-

[Co(NH3)6]3

+

[FeCl6]3-

Coordination Number and Geometry

Lecture 22-14

Ligands that can form 1 bond with the metal ion are called monodentate (denta – tooth) e.g. H2O, NH3, Cl- (a single donor atom).

Some ligands have more than one atom with lone pairs that can be bonded to the metal ion – these are called CHELATES (greek: claw)

Bidentate ligands can form 2 bonds

e.g. ethylenediamine

Polydentate ligands – can form more than 2 bonds

e.g. EDTA - (hexadentate, can form 6 bonds)

Ligands

Lecture 22-15

H2C CH2

NH2

Mx+

H2N

Ethylenediamine (en) has two N atoms that can form a bond with the metal ion, giving a five-membered ring.

Bidentate chelate ligandsMX+(en)

Blackman, Bottle, Schmid, Mocerino & Wille, Figure 13.10

Lecture 22-16

Ethylenediaminetetraacetate tetraanion (EDTA4-)

N N

O O

OOO

O O

O

EDTA forms very stable complexes with many metal ions. EDTA is used for treating heavy-metal poisoning, because it removes lead and other heavy metal ions from the blood and other bodily fluids.

Hexadentate ligand: EDTA

Co(III)

[Co(EDTA)]-

N=blueO=red

Lecture 22-17

Examples of ligands

Table from Silberberg, “Chemistry”, McGraw Hill, 2006.

Lecture 22-18

OH2

Fe

OH2

H2O OH2

H2O OH2

3+NH2

Fe

NH2H2N NH2

H2N NH2

3+O

Fe

NN OO O

OO

O

O

-

[Fe(H2O)6]3+ [Fe(en)3]3+ [Fe(EDTA)]-

monodentateligands

bidentateligands

hexadentateligands

The charge of a complex ion is the charge of the metal ion plus the charge of its ligands:

e.g. [Ni(H2O)6]2+ charge of complex ion is that of the Ni2+ ion.

eg [NiCl4]2- Ni2+ ion coordinated to four chloride (Cl-) ions giving overall (2-) charge.

Examples of ligands

Lecture 22-19

M(H2O)42+

M(H2O)3(NH3)2+

M(NH3)42+

NH3

The stepwise exchange of NH3 for H2O in M(H2O)42+.

3NH3

3moresteps

Lewis bases: water and ammonia

Ammonia is a stronger Lewis base than waterFigure from Silberberg, “Chemistry”, McGraw Hill, 2006.

Lecture 22-20

Equilibrium Constant Kstab

The complex formation equilibrium is characterised by a stability constant, Kstab (also called formation constant):

Ag+(aq) + 2 NH3 Ag(NH3)2

+(aq)

23

23stab ][NH [Ag]

])[Ag(NH K

Metal Ion + nLigand Complex

The larger Kstab, the more stable the complex, e.g.

nstab [Ligand] [Metal]

[Complex] K

Lecture 22-21

Metal ions gain ligands one at a time. Each step characterised by “stepwise stability constant” aka “stepwise

formation constant”. Overall formation constant = Kstab = K1 x K2…x Kn

Example:

Ag+(aq) + NH3(aq) Ag(NH3)+

(aq) K1 = 2.1 · 103

Ag(NH3)+(aq) + NH3(aq) Ag(NH3)2

+(aq) K2 = 8.2 ·

103

Ag+(aq) + 2 NH3(aq) Ag(NH3)2

+(aq) Kstab =

Kstab = K1 x K2 = [Ag(NH3)2+] = 1.7 · 107

[Ag+] [NH3]2

Stepwise stability constant

Lecture 22-22

Demo: Nickel complexes

Ni2+ forms three complexes with ethylenediamine:

1. Mix [Ni(H2O)6]2+ and en in ratio 3:1 → some [Ni(en)(H2O)4]2+and [Ni(H2O)6]2

Green blue-green

2. Mix [Ni(H2O)6]2+ and en in ratio 1:1 → mostly [Ni(en)(H2O)4]2+

light blue

3. Mix [Ni(H2O)6]2+ and en in ratio 1:3 → mostly [Ni(en)3]2+ purple

Ni2+

H O2

H O2

H O2

H O2

NH2

NH2

CH2CH2

Ni2+

H O2

H O2

H O2

H O2

en Ni( ) (aq)en2+

Lecture 22-23

Biologically Important Complexes Many biomolecules contain metal ions that act as Lewis acids.

Give some examples of naturally occurring complexes.

Heme

Chlorophyll

Vitamin B12

Enzyme Carbonic anhydrase

Lecture 22-24

Heme

Heme is a square planar complex of Fe2+ and the tetradentate ring ligand porphyrin (bonds to 4 donor N atoms). Present in hemoglobin, which carries oxygen in blood, and myoglobin, which stores oxygen in muscle.

Porphyrin ring

O2 bound to Fe2+

Myoglobin protein

Blackman Figure 13.37

Lecture 22-25

Chlorophyll

Chlorophyll is a photosynthetic pigment, that

gives leaves the characteristic green colour.

It is a complex of Mg2+ and a porphyrin ring system (four N atoms are the chelae).

Figure from

Silberberg, “C

hemistry”, M

cGraw

Hill, 2006.

Lecture 22-26

Dorothy Crowfoot HodgkinThe Nobel Prize in Chemistry 1964

Nobelprize.org

Vitamin B12

Image download from Wikipedia

Lecture 22-27

CO2(g) + 2H2O(l) H3O+(aq) + HCO3

- (aq)

Carbonic anhydraseTetrahedral complex of Zn2+.

Catalyses reaction between water and carbon dioxide during respiration. Coordinated to 3 N, fourth site left free to interact with molecule whose reaction is being catalysed (here with water).

By withdrawing electron density, makes water acidic to lose proton and OH- attacks partial positive C of CO2 much more vigorously. Cd2+ is toxic because it competes with zinc for this spot.

Figure downloaded from Wikipedia

Lecture 22-28

Exercise

M104(0.48)10

0.01

][CN10

][Ag(CN)][Ag

10]][CN[Ag

][Ag(CN)

22220.02-20.0

2

20.02-

2

stabK

0.01 moles of AgNO3 are added to a 500 mL of a 1.00 M solution of KCN. Then enough water is added to make 1.00 L of solution. Calculate the equilibrium [Ag+] given Kstab [Ag(CN)2]– =1020 M–2.

(careful with the direction of the equation represented by Kstab!)

Ag+ + 2CN– [Ag(CN)2]–

initial /M 0.01 0.500 0 change ~ -0.01 -0.02 0.01equilibrium /M x 0.480 0.01

Lecture 22-29

Metal complex formation can influence the solubility of a compound.

e.g. AgCl(s) + 2 NH3 [Ag(NH3)2]+ + Cl-

This occurs in 2 stages:

AgCl(s) Ag+ + Cl- (1)

Ag+ + 2 NH3 [Ag(NH3)2]+ (2)

Complex formation removes the free Ag+ from solution and so drives the dissolution of AgCl forward.

Complex Formation and solubility

Lecture 22-30

Example: AgBr(s) Ag+(aq) + Br-

(aq)

Calculate the solubility of AgBr in:

a) water

b) 1.0 M sodium thiosulfate (Na2S2O3)

c) 1.0 M NH3

(Ksp (AgBr)= 5.0·10-13, Kstab ([Ag(S2O3)2]3- )= 4.7·1013; Kstab(Ag(NH3)2+)= 1.7·107)

Complex ion formation affects solubility

AgBr(s) Ag+(aq) + Br-

(aq)

Ksp = x2 = 5.0·10-13 x = 7.1 ·10-7 M

Ksp = [Ag+][Br-]

x x

a) Solubility of AgBr in water

Lecture 22-31

AgBr(s) Ag+(aq) + Br-

(aq)

Koverall = Ksp x Kstab = = 5.0·10-13 x 4.7·1013 = 24

Ag+(aq) + 2S2O3

2-(aq) [Ag(S2O3)2]3-

(aq)

AgBr(s) + 2S2O32-

(aq) [Ag(S2O3)2]3-(aq) + Br-

(aq)

(1)

(2)

(1)+(2)

b) Solubility of AgBr in sodium thiosulfate

[Ag(S2O3)23-][Br-]

[S2O32-]2

Initial Conc.ChangeEquilibrium Conc.

1.0 M-2x

1.0 -2x

0+xx

0+xx

Substitute: Koverall = x2/(1.0 - 2x)2 = 24 x = 0.45

Solubility of AgBr in thiosulfate is 0.45 M (c.f. in water 7.1 x 10-7 M)

1.0 M Na2S2O3

Lecture 22-32

AgBr(s) Ag+(aq) + Br-

(aq)

Koverall = Ksp x Kstab = = 5.0·10-13 x 1.7·107 = 8.5·10-6

1.0 M NH3

Ag+(aq) + 2NH3(aq) [Ag(NH3)2]+

(aq)

AgBr(s) + 2NH3(aq) [AgNH3]+(aq) + Br-

(aq)

(1)

(2)

(1)+(2)

c) Solubility of AgBr in ammonia

[Ag(NH3)2+][Br-]

[NH3]

Initial Conc.ChangeEquilibrium Conc.

1.0 M-2x

1.0 - 2x

0+xx

0+xx

Substitute: Koverall = x2/(1.0-2x)2 = 8.5·10-6 x = 2.9·10-3 M

Solubility of AgBr in NH3 is 2.9·10-3 M (c.f. in thiosulfate 0.45 M)

Lecture 22-33

The One Pot ReactionStart with a AgNO3 aqueous solution. Add sequentially :

Ag+ + OH- AgOH(s) (brown)

2 AgOH(s) + HPO42- Ag3PO4(s) (yellow)

Ag3PO4(s) + HNO3 3Ag+ + NO3- + HPO4

2-

Ag+ + Cl- AgCl (s) (white)

AgCl(s) + 2NH3 [Ag(NH3)2]+ + Cl-

[Ag(NH3)2]+ + Br- AgBr (s)(green/white)

AgBr(s) + 2S2O32- [Ag(S2O3)2]3- + Br-

[Ag(S2O3)2]3- + I- AgI (s) (yellow)

AgI(s) + 2CN- [Ag(CN)2]- + I-

2 Ag(CN)2- + S2- Ag2S + CN-(black)

+ NaOH

+ Na2HPO4

+ HNO3

+ NaCl

+ NH3

+ KBr

+ Na2S2O3

+ KI

+ KCN

+ Na2S

Ksp = 10-7.70 M2

Ksp = 10-16 M3

Ksp = 1.8 x 10-10 M2

Kstab = 1.7 x 107 M-2

Ksp = 5 x 10-13 M2

Kstab = 2.5 x 1013 M-2

Ksp = 8.3 x 10-17 M2

Kstab = 6.3 x 1019 M-2

Ksp = 8 x 10-51 M3

Lecture 22-34

Rules for nomenclature of coordination compounds:

Name cation, then anion, as separate words.Examples:

[Pt(NH3)4Cl2](NO2)2 tetraamminedichloridoplatinum(IV) nitrite

[Pt(NH3)4(NO2)2]Cl2 tetraamminedinitritoplatinum(IV) chloride

Name the ligands then the metal, all in same word.

Number of ligands as Greek prefixes (di-, tri-, tetra-, penta-, hexa-), except ligands that already have numerical prefixes which use Latin prefixes (bis, tris, tetrakis…) e.g. bis(ethylenediamine) for (en)2

Nomenclature

Lecture 22-35

Nomenclature II Oxidation state in Roman numeral in parentheses after name of metal

e.g. [Ag(NH3)2]NO3 diamminesilver(I) nitrate

Anionic ligands end in '-ido';

Neutral ligands named as molecule, except those listed here:

FluoridoChloridoBromidoIodidoHydroxidoCyanido

New IUPAC Nomenclature: all anions ending in – ‘ide’ become -‘ido’.

(Please modify accordingly

pp.518-519 of your book)

Lecture 22-36

Ligands named in alphabetical order (but prefixes do not affect the order) e.g. [Co(NH3)5Cl]SO4 pentaamminechloridocobalt(III) sulfate

Anionic complexes end in ‘-ate’ e.g. K3[CrCl6] potassium hexachloridochromate(III)

Some metals in anionic complexes use Latin -ate names:

Nomenclature of Ligands

Not IronateNot CopperateNot LeadateNot SilverateNot GoldateNot Tinnate

Lecture 22-37

Nomenclature - Exercises

[Co(H2O)6]CO3

hexaaquacobalt(II) carbonate

[Cu(NH3)4]SO4

tetraamminecopper(II) sulfate

(NH4)3[FeF6]ammonium hexafluoridoferrate(III)

K4[Mn(CN)6]potassium hexacyanidomanganate(II)

Lecture 22-38

Example 1:

Find O.N. of Co in : [Co(NH3)5Cl]SO4 pentaamminechloridocobalt(?) sulfate

[Co(NH3)5Cl]2+ ammine is neutral, chloride is -1

O.N. -1 = +2 (sum of O.N.s = overall charge)

O.N. = +3

Assigning oxidation numbers

Example 2:

Find O.N. of Mn in :K4[Mn(CN)6] potassium hexacyanidomanganate(?)

[Mn(CN)6]4- (CN) is -1 overall

O.N. + 6x(-1) = -4 (sum of O.N.s = overall charge)

ON = +2

Lecture 22-39

You won’t be asked to draw formulae of complicated biological complexes.

You should be able to use the naming rules to write formulae from names and names from formulae.

About naming complexes

Lecture 22-40

Isomerism in ComplexesComplexes can have several types of isomers:

Structural Isomers: different atom connectivities

1. Coordination sphere isomerism

2. Linkage isomerism

Stereoisomers: same atom connectivities but different arrangement of atoms in space

3. Geometric isomerism

4. Optical isomerism

Lecture 22-41

CoordinationCompound

ComplexIon

CounterIons

Ligands within the coordination sphere remain bound to the metal ion

Coordination compoundsFigure from Silberberg, “Chemistry”, McGraw Hill, 2006.

Lecture 22-42

Coordination Isomers Ligands and counter-ions exchange place:

Example:

[Pt(NH3)4Cl2](NO2)2 tetraamminedichloridoplatinum(IV) nitrite

[Pt(NH3)4(NO2)2]Cl2 tetraamminedinitritoplatinum(IV) chloride

Two sets of ligands are reversed:

[Cr(NH3)6][Co(CN)6] NH3 is a ligand for Cr3+

[Co(NH3)6][Cr(CN)6] NH3 is a ligand for Co3+

ligands counterions

Lecture 22-43

Linkage isomers Occur when a ligand has two alternative donor atoms. Example 1:

NCSThiocyanate ion

H3N Co

NH3

NH3

N

NH3

H3N

C S

2+

H3N Co

NH3

NH3

S

NH3

H3N

C N

2+

and

cyanate ion NCOcyanato NCO:→

isocyanato OCN:→

Thiocyanato NCS:→

Isothiocyanato SCN:→

Pentaammineisothiocyanatocobalt(III) pentaamminethiocyanatocobalt (III)

Lecture 22-44

NO2- nitro O2N:→

nitrito ONO:→

Linkage Isomers Example 2:

N N

O

OOO

[Co(NH3)5(NO2)]Cl2

Pentaamminenitrocobalt(III) chloride

[Co(NH3)5(ONO)]Cl2

Pentaamminenitritocobalt(III) chloride

Figure from

Silberberg, “C

hemistry”, M

cGraw

Hill, 2006.

Lecture 22-45

Isomerism in ComplexesComplexes can have several types of isomers:

Structural Isomers: different atom connectivities

1. Coordination sphere isomerism

2. Linkage isomerism

Stereoisomers: same atom connectivities but different arrangement of atoms in space

3. Geometric isomerism

4. Optical isomerism

Lecture 22-46

Square planar complex. Four coordinate: cis- and trans-[Pt(NH3)2Cl2]

Stereoisomers: Geometric Isomers

No

anti-tumour

effect

cisplatin –

highly effective

anti-tumour agent

Figure from

Silberberg, “C

hemistry”, M

cGraw

Hill, 2006.

Lecture 22-47

Stereoisomers: Geometric Isomers

2 Cl next to each other

Octahedral complex. Six coordinate: cis- and trans- [Co(NH3)4Cl2]+

violet

green2 Cl axial to each other

Lecture 22-48

[NiClBrFI]2-

Stereoisomers: Optical Isomers

When a molecule is non-superimposable with its mirror image. Example: four different substituents about tetrahedral centre. Same physical properties, except direction in which they rotate the

plane of polarized light.

Lecture 22-49

cis-[Co(NH3)4Cl2]+ cis-[Co(en)2Cl2]+

ClCo

NH3

NH3H3N

Cl

NH3

ClCo

NH2

NH2H2N

Cl

NH2

++

Has no optical

isomers

Has optical

isomers

Stereoisomers: Optical isomers Metal atoms with tetrahedral or octahedral geometries (but not

square planar) may be chiral due to having different ligands. For the octahedral case, several chiralities are possible, e.g.

1. Complex with four ligands of two types.

Lecture 22-50

[M(en)3]n+ complexes have optical isomers:

Notsuperimposable

H2NCo

NH2

NH2H2N

NH2

NH2

NH2Co

NH2

NH2NH2

H2N

H2N

3+ 3+

Mirrorplane

Stereoisomers: Optical isomers

2. Having three bidentate ligands of only

one type - gives a propeller-type structure.

www.pt-boat.com

Lecture 22-51

Octahedral complex - stereoisomerism

rotation of I by 180° gives III ≠ II

Mirror

image

Cis-

Dichlorido

Bis(ethylendiamine)cobalt(III) ion

Figure from

Silberberg, “C

hemistry”, M

cGraw

Hill, 2006.

Lecture 22-52

Octahedral complex - stereoisomerism

rotation of I by 90° gives III = II

Mirror

image

Trans-

Dichlorido

Bis(ethylendiamine)cobalt(III) ion

Figure from

Silberberg, “C

hemistry”, M

cGraw

Hill, 2006.

Lecture 22-53

Question Does the square planar complex ion [Pt(NH3)(N3)BrCl]- have optical

isomers?

Br

Pt

N=N=N

NH3

Cl

Br

Pt

NH3

ClN=N=N

This complex has no optical isomers because it can be superimposed

on its mirror image.

Lecture 22-54

Summary

Concepts: Complex formation Stability constant and stepwise stability constant Acidity of some metal ions in solution Coordination compounds and geometry Nomenclature of coordination compounds Isomerism in Complexes

Calculations Complex Formation Equilibria in solution: complex formation + solubility