Chapter 19 Review Questions

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Chapter 19: Transition Metals

and Coordination Chemistry

19.1 Survey of transition metals

19.2 1st-row transition metals

19.3 Coordination compounds

19.4 Isomerism

19.5 Bonding in complex ions: The localized electron model

19.6 The crystal field model

19.7 The molecular orbital model

19.8 The biological importance of coordination complexes

Figure 19.1



Filling d-orbital shells

3d

4d

5d

Filling f-orbital shells

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General Properties of Transition Metals

Metallic luster

High electrical and thermal conductivity (Ag, Cu)

Wide range of melting points (e.g. W @ 3400C, Hg @ -39C) and hardness

Wide range of reactivity toward O2 Fe3O4 - magnetite (magnetic recording material)

Fe2O3 rust (scales off complete corrosion) Oxides of Cr, Co, and Ni- very hard, protective

Coinage metals (Au, Ag, Pt, Pd) do not react

readily with O2 (noble metals)

Easily oxidized Readily form ionic complexes

e.g. Fe(H2O)62+, [Co(NH3)4Cl2]

+

Many coordination compounds are colored Many coordination compounds are paramagnetic

More General Properties of Transition Metals

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Some important aspects of transition metal ions: 1. The valence electrons are in d orbitals 2. The d orbitals do not have a large radial extension 3. The d orbitals are, therefore, mostly nonbonding in

complexes of transition metal ions For these reasons, the effects of redox changes are substantially smaller for transition metals than for main group elements

Review Section 12.13!

Figure 12.27

Electron configurations of the neutral

transition metal elements

3d start to fill after 4s is full

Cr and Cu are exceptions to trend: both are 4s1 3dn

Neutral TM: 3d and 4s orbitals similar in energy

3d orbitals for TM ions much less E than 4s, so 4s electrons leave first (1st row TM ions do not have 4s electrons)

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Orbital Occupancy of Period 4 Transition Metals

Element 4s 3d 4p

Unpaired

Electrons

Sc 1

Ti 2

V 3

Cr 6

Mn 5

Fe 4

Co 3

Ni 2

Cu 1

Zn 0

When you oxidize a transition metal, remove s electrons first!

Oxidation States

See Table 19.2 for common oxidation states of the 1st-row transition metals

+1 up to +7 are observed, with +2 and +3 most

common Highest O.S. is loss of all 4s and 3d electrons

As the oxidation state is increased, the d orbitals are

stabilized, and the metals get harder to oxidize further

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Standard Reduction Potentials

Consider the reduction half-reaction: Mn+ + ne- M

Reduction potentials (E) for 1st-row transition metals in aqueous solutions:

Sc3+ + 3e- Sc -2.08 V Ti2+ + 2e- Ti -1.63 V V2+ + 2e- V -1.2 V Mn2+ + 2e- Mn -1.18 V Cr2+ + 2e- Cr -0.91 V Zn2+ + 2e- Zn -0.76 V Fe2+ + 2e- Fe -0.44 V Co2+ + 2e- Co -0.28 V Ni2+ + 2e- Ni -0.23 V Cu2+ + 2e- Cu 0.34 V

red

ucin

g a

bility

See Table 19.3 (opposite signs b/c reduction vs. oxidation potentials)

Oxidation Potentials

(opp. sign from standard reduction potentials)

Consider the oxidation half-reaction: M Mn+ + ne-

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[Co(NH3)5Cl]Cl2

K3[Fe(CN)6]

Transition-metal complexes

are extremely colorful.

Color is influenced by:

metal ion (dn configuration),

oxidation state, and

coordinated ligands.

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Oxidation States of Mn

2 MnO4-(aq) + 5 H2C2O4(aq) + 6 H

+(aq)

2 Mn2+(aq) + 10 CO2(g) + 8 H2O(l) * Observe several intermediates (mixtures of MnO4

-, lower O.S. of

Mn, and Mn(III)-oxalate complexes)

Table 19.6

Oxidation State influences color

VO2+(aq) +4

VO2+

(aq)

+5

V3+(aq) +3

V2+(aq) +2

V0(s)

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Different colors are due to different numbers of

electrons in the highest-occupied MOs of each V-

containing polyatomic ion.

V +4

is the most common oxidation state. V +5

is

easily converted to V+4

by the mild reducing agent

NaHSO3(aq). An excess of the stronger reducing

agent Zn(s) is required to convert V+5

to V +2

, which

is then easily oxidized to V +3

by dilute (0.5%)

H2O2(aq).

Oxidation States of Vanadium

Vanadium Oxidation States

2 HVO42(aq) + 3 Zn(s) + 14 H3O

+(aq) + 8 H2O(l)

2 V(H2O)62+(aq) + 3 Zn(H2O)6

2+(aq)

HVO42- V(H2O)6

2+ VO(H2O)52+ V(H2O)6

3+

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2 HVO42(aq) + HSO3

(aq) + 7 H3O+(aq)

2 VO(H2O)52+(aq) + SO4

2(aq) + 2 H2O(l)

HVO42- VO(H2O)5

2+


2 V(H2O)62+(aq) + H2O2(aq) + 2 H3O

+(aq)

2 V(H2O)63+(aq) + 4 H2O(l)


V(H2O)62+ V(H2O)6

3+

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

2+ [Co(H2O)6]2+ [Ni(H2O)6]

2+ [Cu(H2O)6]2+

d3 d6 d7 d8 d9

Metal ions influence color

The d-orbital electron count influences

compound color

Metal ions influence color

No d electrons no color.

Full d orbitals no color.

[Mg(H2O)6]2+ [Al(H2O)6]

3+ [Ca(H2O)6]2+ [Sc(H2O)6]

3+ [Zn(H2O)6]2+

d0 d0 d0 d0 d10

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Ligands influence color

[Ni(H2O)6]2+ [Ni(en)(H2O)4]

2+ [Ni(en)2(H2O)2]2+ [Ni(en)3]

2+

green green/blue blue purple

Whats responsible for these colors?

Color is a result of electron transitions

MO Theory revisited:

Recall our simple molecular orbital

diagramit only involved s and p

orbitals

Now, however, we have d orbitals to

consider

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MO Theory - Part I The d orbitals reach only a very

short distance from the

nucleus they are essentially non-bonding orbitals

An octahedral dn complex has (12+n) electrons to fill in. The

first 12 go in the bonding

orbitals.

MO Theory Part II

The movement of electrons between these levels is the source of the

chemical properties of transition

metal complexes (color, magnetic

properties, reactivity).

ground state excited state

n

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19.4 Isomerism





Coordination Compounds

The real bulk of inorganic chemistry occurs in the reactions of coordination compounds (or complexes).

A coordination compound contains a complex ion and counter ion

Complex ion: a central metal ion surround by one or more ligands

Counter ion: ion that balances the charge of a complex ion to form a neutral compound

Ligands are ions or molecules that have an independent existence: NH3, H2O, CO, 2,2-bipyridine (bpy), etc.

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Ligand: A neutral molecule or ion

having a lone pair that can be used to

form a bond to a metal ion

Typical Coordination Numbers

Cu+ 2, 4

Ag+ 2

Au+ 2, 4

Mn2+ 4, 6

Fe2+ 6

Co2+ 4, 6

Ni2+ 4, 6

Cu2+ 4, 6

Zn2+ 4, 6

Sc3+ 6

Cr3+ 6

Co3+ 6

Au3+ 4

Fig 19.6 See Table 19.12

Lewis Acids and Bases

To understand how coordination compounds form, we

need to understand Lewis acids and bases

A Lewis acid is an electron pair acceptor

A Lewis base is an electron pair donor

Lewis acids and bases are different from Brnsted-Lowry

acids and bases in that they can describe aprotic species

(no acidic protons are donated/accepted).

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Some Lewis Acids and Bases

Molecules with an incomplete octet can act as Lewis acids

acid base

Metal cations act as Lewis acids

Co2+ + 6 H2O [Co(OH2)6]2+

acid base

A Lewis base can influence electron rearrangement in a Lewis acid


acid base

+

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A Lewis acid can expand its valence shell to accommodate a Lewis base


2 F- +

acid base

d-block elements:

M Mn+

ne

oxidation state Transition metals readily ionize, and can lose multiple electrons

Mn+ + 6 Lq-

Once they are ionized, metal ions tend to surround themselves with electron pair donors (Lewis bases)

[Mn+L6](n-6q)

acid base metal complex

net charge

Coordination Chemistry

Since metal cations can acts as Lewis acids, and ligands have

electron pairs to donateinorganic coordination compounds are often formed by Lewis acid / base chemistry

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What are some aspects of coordination compounds we

should understand?

Coordination number

Ligands

Isomers and chirality

Coordination compounds

Coordination Number (the number of ligands around the central atom)

Coordination number is influenced by

the size of the central atom

the bulk (or lack thereof) of the ligands

electronic interactions between metal and ligand

Coordination numbers can vary widely

2 and 3 (rare); 4, 5, and 6 (most common); others

Polymetallic complexes are possible, too.

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Ligands

Ligands can bond in one or more sites on the metal ion:

1 (monodentate): NH3, CO, H2O, I, Cl, etc.

2 (bidentate): acac, bpy, en, dppe

3 (tridentate): dien

4, 5, 6 (polydentate): cyclam, Cp, 18-crown-6

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H2C

NH2C

CH2

N

H2C

M

O

O

O

O

CH2

O

H2C

O

O

O

Chelates (chele/chela = claw)

EthyleneDiamineTetraAcetic acid (EDTA)

N

H2C

CH2

N

H2C

CH2

CH2

H2C COOH

COOH

HOOC

HOOC

Mn+

n = 2 to 4

Very strong 1:1 complexes with transition metals

Metal cations are sequestered from solution

Used for detoxification and as a preservative.

See Fig 19.8

The ligands can have a dramatic influence on a metal complexes properties

[Fe(TACN)(CN)3]

vs.

[Fe(TACN)(H2O)3]2+

unreactive reactive

FeNN

N H

H

CN

CNCN

1

FeNN

N H

H

OH2

OH2OH2

2+

all electrons paired four unpaired electrons

yellow blue

iron oxidation state= +2 iron oxidation state= +2

negative redox potential positive redox potential

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FeNN

N H

H

OH2

OH2OH2

2+

iron oxidation state= +2

iron oxidation state= total molecular charge

ligand charges: CN= 1 TACN= 0

H2O= 0

FeNN

N H

H

CN

CNCN

1

iron oxidation state= +2

-1 (-3) +2 (0)

Since the metals are identical, the oxidation states are identical, and only the ligands differ, the ligands must be responsible for

the differing properties.

= +2 = +2

S(ligand charges)

Ligands and Isomers When ligands are involved, you can get isomers:

cis- and trans- (square planar)

optical isomers (tetrahedral)

mer- and fac- (octahedral)

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Anti-cancer agent

cis chlorides

Stereochemistry can dramatically

influence key properties

NOT an anti-cancer agent trans chlorides

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Origin of anti-cancer activity

Origin of anti-cancer activity

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Isomers (Preview)

2 or more chemical species with identical

composition but different properties

Naming

Coordination

Compounds

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Naming Coordination Compounds 1. Cation named before anion

2. Ligands named before metal ion

3. o is added to the end of anionic ligand names (chloro-, bromo-, iodo-, etc.). Neutral ligands retain their name (except

H2O, NH3, CO, NO)

4. Use prefixes (mono-, di-, tri-, tetra-, penta- and hexa-) for the

number of simple ligands; (bis-, tris-, tetrakis-, etc. for multiple

complex ligands)

5. Metal oxidation state is denoted with roman numerals in

parentheses.

6. Ligands are named in alphabetical order

7. If the complex ion has a negative charge, add ate to the metal name (vanadate, ferrate, etc.). Sometimes the Latin

name is used.

Naming Examples

[Co(NH3)5Cl]Cl2

K3Fe(CN)6

[Fe(en)2(NO2)2]2SO4

pentaamminechlorocobalt(III) chloride

potassium hexacyanoferrate(III)

bis(ethylenediammine)dinitroiron(III) sulfate

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Mn+ L

Metal Ion(electron acceptor)

Ligand(electron donor with a lone pair)

Unoccupied hybrid orbital

Coordinate covalent bond

Mn+ L

Complex Ions and the Localized Electron Model

Bond Formation

See pg. 958

Hybridization (L.E.M.)

Linear: sp

Square planar: dsp2

Tetrahedral: sp3 CoCl42-

Ni(CN)42-

Ag(CN)2-

No reliable way to predict

sq. planar vs. tetrahedral

L.E.M. cant predict important properties of complex ions, like

color or magnetism

Figs. 19.20

and 19.19

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19.4 Isomerism





The Crystal Field Model

d-orbital energies split in the electrostatic field

Ligands produce an electrostatic field around the metal ion

Electron occupancy of d orbitals depends on the magnitude

of splitting

Crystal field model does NOT explain complex geometry

or bonding

Why care?

CFM explains how color and magnetism can arise in

complex ions by considering the d orbitals of the transition

metal.

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Octahedral Complexes Consider ligands as negative point chargesconsider the location of the electrons in the orbitals, which will

repel the negative charges of the ligands.

d d

dxy dyz dxz

z2 x2-y2 Co(NH3)63+

Fig 19.21

Close

overlap,

higher

energy

Ligands influence properties

The ligands on a metal complex influence the energy of the d orbitals.

Orbitals that point directly at ligands (dz2 and dx2-y2) are higher in energy.

Orbitals that point between ligands (dxy, dyz and dxz) are lower in energy.

d

octahedral ligand field

eg (dz2 and dx2-y2)

t2g (dxy, dyz, dxz)

The nature of

the ligands

affects this

difference

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Orbital Energy Splitting

eg orbitals

t2g orbitals

Strong Field

eg orbitals

t2g orbitals

Weak Field

Example:

Co3+ (3d6)

(in Octahedral Complexes)

Figs 19.22 and 19.23

Transition Metal Ion Properties

eg orbitals

t2g orbitals

Weak Field

eg orbitals

t2g orbitals

Strong Field Low spin compounds

yield minimum

number of unpaired

electrons: (Diamagnetic)

High spin compounds

yield maximum number

of unpaired electrons:

(Paramagnetic )

Example:

Co3+ (3d6)

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Spectrochemical series

CN- > NO2- > en > NH3 > H2O > OH

- > F- > Cl- > Br- > I-

Strong-field

ligands

Weak-field

ligands Large small

Example: Is [Fe(CN)6] 4- paramagnetic or diamagnetic?

Fe oxidation state: from ion and ligand charges,

(-4) (-6) = +2: Fe2+

Number of 3d electrons on Fe2+ : 8 2 = 6

eg orbitals

t2g orbitals

Strong Field

[Fe(CN)6] 4- is

diamagnetic

CN- is a strong-field

ligand

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Examples

d5 complex high spin

Examples

d5 complex low spin

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Examples

d1 - d3 complexes only one spin configuration

Examples

d8 d10 complexes only one spin configuration

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Why do we see the colors we doenergy is absorbed.

eg orbitals

t2g orbitals

[Ti(OH2)6]3+ or [Ti(OH)6]

3- ion

Ground electronic state

eg orbitals

t2g orbitals

Excited electronic state photon absorption

= photon energy = hn = hc/

= wavelength of absorbed light (nm) = 119,626/(kJ mol-1)

Large small complex absorbs blue end of spectrum Small large complex absorbs red end of spectrum

Visible spectrum width = 400 700 nm = 300 170 kJ mol-1

See Fig. 19.26

Absorbed Wavelength

Observed Color (complementary)

Greenish yellow Yellow

Red Violet

Blue Green

See Table 19.16

Colored compounds used in tattoos:

http://pubs.acs.org/cen/whatstuff/85/8546sci4.html

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R G O B Y V I

increasing energy

absorbs

blue

hn absorbs

green

Appears

absorbs

violet

dz2 dx2-y2

dxy dyz dxz

hn

dz2 dx2-y2

dxy dyz dxz

hn

dz2 dx2-y2

dxy dyz dxz

Tetrahedral Complexes

None of d-orbitals point

directly AT the ligands

Small orbital splitting

and splitting order is

reversed

tet = (4/9) oct

tet

En

erg

y

dz2 dx2-y2

dxy dxz dyz

Fig 19.27 Always weak field, high spin.

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Example: Cl

Co

ClCl

Cl

2-

How many unpaired electrons are there in this complex?

(1) Determine the number of electrons on the metal ion:

CoCl42-: (-4) (-2) = +2 7 electrons on Co2+

tet

En

erg

y

dz2 dx2-y2

dxy dxz dyz

(2) Fill electrons in d orbitals from bottom up

Square Planar and Linear Complexes

Fig 19.29

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is influenced by:

The Mn+ oxidation state

(M3+) > (M2+) > (M+)

Example,

Fe(II)(NH3)62+ vs. Fe(III)(NH3)63+

= 12,800 cm-1 26,000 cm-1

The row in which Mn+ lies in periodic table

(3rd row) > (2nd row) > (1st row)

The identity of the ligands

is influenced by:

Example, [Fe(II)L6]2+ L = H2O CN

Cl = 8,900 30,000 5,900 cm-1

Spectrochemical series

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The spectrochemical series

Ligands

I- < Br-

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Weak vs. strong field ligands

d

If we need to fill the d orbitals with four electrons, where does the fourth electron go?


d


Pairing the electron requires energy pairing energy (P)

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d


Occupying an eg orbital requires energy


d


< P = Weak field > P = Strong field

Examples: [Cr(OH2)6]2+ [Cr(CN)6]

4-

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d


Examples: [Cr(OH2)6]2+ [Cr(CN)6]

4-

High-spin Low-spin

Demo: Nickel Complexes

Ni(H2O)62+(aq) + 6 NH3(aq) Ni(NH3)6

2+(aq) + 6 H2O(l)

(octahedral) (octahedral)

Ni(NH3)62+(aq) + 3 en(EtOH) Ni(en)3

2+ + 6 NH3(aq)

(octahedral) (octahedral)

Ni(en)32+(aq) + 2 Hdmg(EtOH) + 2 H2O(l)

Ni(dmg)2(s) + 3 en(EtOH) + 2 H3O+(aq)

(octahedral) (square planar)

Note: If any green precipitate forms, it is Ni(OH)2(s).

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Demo: Ammines

Cu(H2O)42+(aq) + 4 NH3(aq) Cu(NH3)4

2+(aq) +

4 H2O(l)

Spectator Ion: SO42

Ni(H2O)62+(aq) + 6 NH3(aq) Ni(NH3)6

2+(aq) +

6 H2O(l)

Spectator Ion: NO3

Co(H2O)62+(aq) + 6 NH3(aq) Co(NH3)6

2+(aq) +

6 H2O(l)

Spectator Ion: Cl






19.4 Isomerism





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Classes of isomers

Fig 19.9

1 2 3 4

Coordination Isomers:

[Cr(NH3)5SO4]Br and [Cr(NH3)5Br]SO4

1

Br SO4

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Linkage Isomers:

NO2- can bond to the

metal through one of the

oxygens or through the

nitrogen

Fig 19.10

yellow

red

2

[Co(NH3)5(NO2)]Cl2 Pentaamminenitrocobalt(III)

chloride

[Co(NH3)5(ONO)]Cl2 Pentaamminenitritocobalt(III)

chloride

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Cis

Trans

Cis = together

Trans = across, opposite

Fig 19.11

Stereoisomers:

Geometrical isomers 3

Chloride ligands

Fig 19.12

green violet

3

Cis Trans

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a facial isomer (fac) where

the three identical ligands

are mutually cis

a meridional isomer (mer)

where the three ligands

are coplanar

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Figure 19.15

Mirror image of hand

Optical Isomers 4

Objects that are not

superimposable

until you make a

mirror image are

called chiral.

Zumdahl: hands are nonsuperimposable mirror images

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Figure 19.16

Isomers I and II

for [Co(en)3]3+

Nonsuperimposable

mirror images!

4

Chiral Complex Achiral Complex

Fig 19.17

Geometric Isomers not always Optical Isomers

4 3

Trans isomer Cis isomer

[Co(en)2Cl2]+

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Chiral Complex

(I and III are enantiomers)

Achiral Complex

C

C

C

N O

O

D-Alanine (unnatural)

* C

C

C

N

O

O

L-Alanine (natural in proteins)

*

* denotes chirality center, where the C noted has 4 different substituents (-CH3, -H, -COOH, -NH2)

Chiral Amino Acids

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TMs serve as the active site within many large biological

molecules.

Key is ability of TM metals to

Coordinate with and release ligands Easily undergo oxidation and reduction

Human body contains only 0.01% TM by mass, divided

among 3d Cr, Mn, Fe, Co, Ni, Cu, Zn and 4d Mo. Nature has

used the most abundant TMs:

3d abundance >> 4d/5d.

Fe is most abundant 3d element and the most used

biologically.

Mo is the most abundant 4d/5d element.

BIOINORGANIC CHEMISTRY

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Functions of these trace metals:

Electron Carriers. TM have >1 stable oxidation state.

Oxidized form can pick up electrons; reduced form can

release electrons elsewhere as pH or other conditions

change.

Oxygen Carriers. TM have >1 stable CN. At different O2

partial pressures, can bind or release this metabolically

crucial small molecule.

Catalysts (Enzymes). Flexibility of both oxidation state

and CN allows TM to bond reactants close together,

allowing reaction under milder conditions than normal.

Critical for organisms, which must carry out all metabolic

reactions near STP.

BIOINORGANIC CHEMISTRY

Hemoglobin Molecule

Heme

Figures 19.33,19.36

Sickle cell anemia (importance of structure) High-altitude sickness (how hemoglobin works) Toxicity of CO and CN- (ligand strength)

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