Download Electronic Structure of Metals The “Sea of Electrons”

7/24/2012
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
Chapter 19: Transition Metals
and Coordination Chemistry
Filling d-orbital shells
3d
4d
5d
Figure 19.1
Filling f-orbital shells
1
7/24/2012
General Properties of Transition Metals
• Metallic luster
• High electrical and thermal conductivity (Ag, Cu)
• Wide range of melting points (e.g. W @ 3400°C,
Hg @ -39°C) 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)
More General Properties of Transition 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
2
7/24/2012
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!
Electron configurations of the neutral
transition metal elements
•
•
•
•
Figure 12.27
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)
3
7/24/2012
Orbital Occupancy of Period 4 Transition Metals
3d
4p
Unpaired
Electrons
Element
4s
Sc
↑↓
↑
Ti
↑↓
↑
↑
V
↑↓
↑
↑
↑
Cr
↑
↑
↑
↑
↑
↑
6
Mn
↑↓
↑
↑
↑
↑
↑
5
Fe
↑↓
↑↓
↑
↑
↑
↑
4
Co
↑↓
↑↓
↑↓
↑
↑
↑
3
Ni
↑↓
↑↓
↑↓
↑↓
↑
↑
2
Cu
↑
↑↓
↑↓
↑↓
↑↓
↑↓
1
Zn
↑↓
↑↓
↑↓
↑↓
↑↓
↑↓
0
1
2
3
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
4
7/24/2012
Standard Reduction Potentials
Consider the reduction half-reaction:
Mn+ + neM
Reduction potentials (E°) for 1st-row transition metals in
aqueous solutions:
Sc
Ti
V
Mn
Cr
Zn
Fe
Co
Ni
Cu
-2.08 V
-1.63 V
-1.2 V
-1.18 V
-0.91 V
-0.76 V
-0.44 V
-0.28 V
-0.23 V
0.34 V
reducing ability
Sc3+ + 3eTi2+ + 2eV2+ + 2eMn2+ + 2eCr2+ + 2eZn2+ + 2eFe2+ + 2eCo2+ + 2eNi2+ + 2eCu2+ + 2e-
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-
5
7/24/2012
Transition-metal complexes
are extremely colorful.
Color is influenced by:
metal ion (dn configuration),
oxidation state, and
coordinated ligands.
K3[Fe(CN)6]
[Co(NH3)5Cl]Cl2
6
7/24/2012
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
+2
V2+(aq)
+3
V3+(aq)
+4
VO2+(aq)
+5
VO2+(aq)
V0(s)
7
7/24/2012
Oxidation States of Vanadium
 Different colors are due to different numbers of
electrons in the highest-occupied MOs of each Vcontaining 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).
Vanadium Oxidation States
HVO42-
VO(H2O)52+
V(H2O)63+
V(H2O)62+
2 HVO42(aq) + 3 Zn(s) + 14 H3O+(aq) + 8 H2O(l)
 2 V(H2O)62+(aq) + 3 Zn(H2O)62+(aq)
8
7/24/2012
Vanadium Oxidation States
HVO42-
VO(H2O)52+
2 HVO42(aq) + HSO3(aq) + 7 H3O+(aq)
 2 VO(H2O)52+(aq) + SO42(aq) + 2 H2O(l)
Vanadium Oxidation States
V(H2O)62+
V(H2O)63+
2 V(H2O)62+(aq) + H2O2(aq) + 2 H3O+(aq)

2 V(H2O)63+(aq) + 4 H2O(l)
9
7/24/2012
Metal ions influence color
[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
The d-orbital electron count influences
compound color
Metal ions influence 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
No d electrons – no color.
Full d orbitals – no color.
10
7/24/2012
Ligands influence color
[Ni(H2O)6]2+
green
[Ni(en)(H2O)4]2+ [Ni(en)2(H2O)2]2+
green/blue
blue
[Ni(en)3]2+
purple
What’s responsible for these colors?
Color is a result of electron transitions
• MO Theory revisited:
– Recall our simple
molecular orbital
diagram…it only
involved s and p
orbitals
– Now, however, we
have d orbitals to
consider…
11
7/24/2012
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).
n
ground state
excited state
12
7/24/2012
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
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.
13
7/24/2012
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
Fig 19.6
Cu+
2, 4
Mn2+
4, 6
Sc3+
6
Ag+
2
Fe2+
6
Au+
2, 4
Cr3+
6
Co2+
4, 6
Co3+
6
Ni2+
4, 6
Au3+
4
Cu2+
4, 6
Zn2+
4, 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 Brønsted-Lowry
acids and bases in that they can describe aprotic species
(no acidic protons are donated/accepted).
14
7/24/2012
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
acid
[Co(OH2)6]2+
base
Some Lewis Acids and Bases
• A Lewis base can influence electron rearrangement in a
Lewis acid
+
acid
base
15
7/24/2012
Some Lewis Acids and Bases
• A Lewis acid can expand its valence shell to accommodate
a Lewis base
+
acid
2 Fbase
Coordination Chemistry
Since metal cations can acts as Lewis acids, and ligands have
electron pairs to donate…inorganic coordination compounds
are often formed by Lewis acid / base chemistry
d-block elements:
oxidation state
Mn+
M
ne–
Mn+ + 6 Lqacid base
Transition metals readily
ionize, and can lose
multiple electrons
net charge
[Mn+L6](n-6q)
metal complex
Once they are ionized, metal ions tend to surround
themselves with electron pair donors (Lewis bases)
16
7/24/2012
Coordination compounds
What are some aspects of coordination compounds we
should understand?
– Coordination number
– Ligands
– Isomers and chirality
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.
17
7/24/2012
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
18
7/24/2012
Chelates (chele/chela = claw)
EthyleneDiamineTetraAcetic acid (EDTA)
HOOC
H2
C
HOOC
N
H2
C
CH2
H2C
C
H2
N
COOH
C
H2
O
COOH
H2
C
H2 C
N
Mn+
n = 2 to 4
M
H2 C
C
H2
O
O
O
O
N
CH2
H2 C
O
Very strong 1:1 complexes with transition metals O
Metal cations are sequestered from solution
O
See Fig 19.8
Used for detoxification and as a preservative.
The ligands can have a dramatic influence on
a metal complexes properties
NH
H
N
N Fe
1–
CN
CN
NH
vs.
H
N
N Fe
OH2
OH2
OH2
CN
[Fe(TACN)(CN)3]–
2+
[Fe(TACN)(H2O)3]2+
• unreactive
• reactive
• all electrons paired
• four unpaired electrons
• yellow
• blue
• iron oxidation state= +2
• negative redox potential
• iron oxidation state= +2
• positive redox potential
19
7/24/2012
NH
H
N
N Fe
1–
NH
CN
CN
H
N Fe
OH2
OH2
OH2
CN
iron oxidation state= +2
N
2+
iron oxidation state= +2
-1 – (-3) = +2
+2 – (0) = +2
• iron oxidation state= total molecular charge – S(ligand charges)
• ligand charges: CN= –1
TACN= 0
H2O= 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.
Ligands and Isomers
• When ligands are involved, you can get isomers:
– cis- and trans- (square planar)
– optical isomers (tetrahedral)
– mer- and fac- (octahedral)
20
7/24/2012
Stereochemistry can dramatically
influence key properties
Anti-cancer agent
cis chlorides
NOT an anti-cancer agent
trans chlorides
21
7/24/2012
Origin of anti-cancer activity
Origin of anti-cancer activity
22
7/24/2012
Isomers (Preview)
2 or more chemical species with identical
composition but different properties
Naming
Coordination
Compounds
23
7/24/2012
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
pentaamminechlorocobalt(III) chloride
• K3Fe(CN)6
potassium hexacyanoferrate(III)
• [Fe(en)2(NO2)2]2SO4
bis(ethylenediammine)dinitroiron(III) sulfate
24
7/24/2012
Complex Ions and the Localized Electron Model
Bond Formation
Mn+
L
Metal Ion
(electron acceptor)
Unoccupied hybrid orbital
Ligand
(electron donor with a lone pair)
Mn+
L
Coordinate covalent bond
See pg. 958
Figs. 19.20
and 19.19
Hybridization (L.E.M.)
Linear: sp
Ag(CN)2-
Square planar: dsp2
Ni(CN)42-
No reliable way to predict
sq. planar vs. tetrahedral
Tetrahedral: sp3
CoCl42-
L.E.M. can’t predict important
properties of complex ions, like
color or magnetism…
25
7/24/2012
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
The Crystal Field Model
Ligands produce an electrostatic field around the metal ion
d-orbital energies split in the electrostatic field
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.
26
7/24/2012
Octahedral Complexes
Consider ligands as negative point charges…consider
the location of the electrons in the orbitals, which will
repel the negative charges of the ligands.
Co(NH3)63+
Fig 19.21
dxy
d z2
d x2-y2
dyz
dxz
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.
eg (dz2 and dx2-y2)
d
t2g (dxy, dyz, dxz)
The nature of
the ligands
affects this
difference
octahedral ligand field
27
7/24/2012
Orbital Energy Splitting
(in Octahedral Complexes)
Example:
Co3+ (3d6)
eg orbitals
eg orbitals


t2g orbitals
t2g orbitals
Weak Field
Strong Field
Figs 19.22 and 19.23
Transition Metal Ion Properties
Weak Field
eg orbitals

t2g orbitals
Strong Field
eg orbitals

t2g orbitals
Example:
Co3+ (3d6)
High spin compounds
yield maximum number
of unpaired electrons:
(Paramagnetic )
Low spin compounds
yield minimum
number of unpaired
electrons:
(Diamagnetic)
28
7/24/2012
Spectrochemical series
CN- > NO2- > en > NH3 > H2O > OH- > F- > Cl- > Br- > IStrong-field
ligands
Large 
small 
Weak-field
ligands
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
CN- is a strong-field
ligand
[Fe(CN)6] 4- is
diamagnetic
eg orbitals
Strong Field

t2g orbitals
29
7/24/2012
Examples
d5 complex – high spin
Examples
d5 complex – low spin
30
7/24/2012
Examples
d1 - d3 complexes – only one spin configuration
Examples
d8 – d10 complexes – only one spin configuration
31
7/24/2012
Why do we see the colors we do…energy is absorbed.
[Ti(OH2)6]3+ or [Ti(OH)6]3- ion
eg orbitals
eg orbitals


t2g orbitals
t2g orbitals
Ground electronic state
photon absorption
Excited electronic state
 = 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
See Table 19.16
Absorbed Wavelength
Observed Color (complementary)
Greenish yellow
Yellow
Red
Violet
Blue
Green
Colored compounds used in tattoos:
http://pubs.acs.org/cen/whatstuff/85/8546sci4.html
32
7/24/2012
Appears
dz2 dx2-y2
absorbs
green
increasing energy
dxy
dyz
dxz
G B I V
dz2 dx2-y2
absorbs
blue
hn
dxy
dyz
dxz
dz2 dx2-y2
absorbs
violet
hn
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
Energy
R O Y
hn
dxy
dxz
dz2
Fig 19.27
dyz
tet
dx2-y2
Always weak field, high spin.
33
7/24/2012
Example:
2-
Cl
Co
Cl
Cl
Cl
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+
Energy
(2) Fill electrons in d orbitals from bottom up
dxy
dxz
dz2
dyz
tet
dx2-y2
Square Planar and Linear Complexes
Fig 19.29
34
7/24/2012
 is influenced by:
• The Mn+ oxidation state
 (M3+) >  (M2+) >  (M+)
Example,
=
Fe(II)(NH3)62+
12,800 cm-1
vs.
Fe(III)(NH3)63+
26,000 cm-1
• The row in which Mn+ lies in periodic table
 (3rd row) >  (2nd row) >  (1st row)
 is influenced by:
• The identity of the ligands
Example,
[Fe(II)L6]2+
L=
Δ=
H2 O
8,900
CN–
30,000
Cl–
5,900 cm-1
Spectrochemical series
35
7/24/2012
The spectrochemical series
Ligands
I- < Br- <S2- < SCN- < Cl- < ONO- < N3- < F- < OH- <
C2O42- < O2- < H2O < NCS- < CH3C=N < py < NH3 < en
< bpy < phen < NO2- < PPh3 < CN- < CO
Metal ions
Mn2+ < Ni2+ < Co2+ < Fe2+ < V2+ < Fe3+ < Co3+ < Mo3+
< Rh3+ < Ru3+ < Pd4+ < Ir3+ < Pt4+
Ligands influence color
[Ni(H2O)6]2+
[Ni(en)(H2O)4]2+
[Ni(en)2(H2O)2]2+
[Ni(en)3]2+
Appears:
green
green/blue
blue
purple
Absorbs:
red
red / orange
orange
yellow
O
increasing d–orbital splitting
36
7/24/2012
Weak vs. strong field ligands
If we need to fill the d orbitals with four
electrons, where does the fourth electron go?
d
Weak vs. strong field ligands
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)
37
7/24/2012
Weak vs. strong field ligands
If we need to fill the d orbitals with four
electrons, where does the fourth electron go?
Occupying an eg
orbital requires
energy – 
d
Weak vs. strong field ligands
If we need to fill the d orbitals with four
electrons, where does the fourth electron go?
d
 < P = Weak field
Examples:
[Cr(OH2)6]2+
 > P = Strong field
[Cr(CN)6]4-
38
7/24/2012
Weak vs. strong field ligands
If we need to fill the d orbitals with four
electrons, where does the fourth electron go?
d
“High-spin”
Examples:
[Cr(OH2)6]2+
“Low-spin”
[Cr(CN)6]4-
Demo: Nickel Complexes
Ni(H2O)62+(aq) + 6 NH3(aq) → Ni(NH3)62+(aq) + 6 H2O(l)
(octahedral)
(octahedral)
Ni(NH3)62+(aq) + 3 en(EtOH) → Ni(en)32+ + 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).
39
7/24/2012
Demo: Ammines
Cu(H2O)42+(aq) + 4 NH3(aq) → Cu(NH3)42+(aq) +
4 H2O(l)
Spectator Ion: SO42−
Ni(H2O)62+(aq) + 6 NH3(aq) → Ni(NH3)62+(aq) +
6 H2O(l)
Spectator Ion: NO3−
Co(H2O)62+(aq) + 6 NH3(aq) → Co(NH3)62+(aq) +
6 H2O(l)
Spectator Ion: Cl−
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
40
7/24/2012
Classes of isomers
Fig 19.9
1
1
3
2
4
Coordination Isomers:
[Cr(NH3)5SO4]Br and [Cr(NH3)5Br]SO4
SO4
Br
41
7/24/2012
Fig 19.10
2
Linkage Isomers:
NO2- can bond to the
metal through one of the
oxygens or through the
nitrogen
yellow
red
[Co(NH3)5(NO2)]Cl2
Pentaamminenitrocobalt(III)
chloride
[Co(NH3)5(ONO)]Cl2
Pentaamminenitritocobalt(III)
chloride
42
7/24/2012
Stereoisomers:
3
Cis
Geometrical isomers
Cis = together
Trans = across, opposite
Trans
Fig 19.11
3
Chloride ligands
Cis
Trans
green
violet
Fig 19.12
43
7/24/2012
a facial isomer (fac) where
the three identical ligands
are mutually cis
a meridional isomer (mer)
where the three ligands
are coplanar
44
7/24/2012
4
Optical Isomers
Figure 19.15
Mirror image of hand
Objects that are not
superimposable
until you make a
mirror image are
called chiral.
Zumdahl: hands are “nonsuperimposable mirror images”
45
7/24/2012
4
Figure 19.16
Isomers I and II
for [Co(en)3]3+
Nonsuperimposable
mirror images!
Geometric Isomers not always Optical Isomers
3
[Co(en)2Cl2
Trans isomer
Achiral Complex
]+
4
Cis isomer
Chiral Complex
Fig 19.17
46
7/24/2012
Achiral Complex
Chiral Complex
(I and III are enantiomers)
Chiral Amino Acids
C
N
C
C
*
C
O
O
D-Alanine (unnatural)
N
C
*
O
C
O
L-Alanine (natural in proteins)
* denotes “chirality center”, where the C noted has
4 different substituents (-CH3, -H, -COOH, -NH2)
47
7/24/2012
BIOINORGANIC CHEMISTRY
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.
48
7/24/2012
BIOINORGANIC CHEMISTRY
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.
Hemoglobin Molecule
Figures 19.33,19.36
Heme
• Sickle cell anemia (importance of structure)
• High-altitude sickness (how hemoglobin works)
• Toxicity of CO and CN- (ligand strength)
49