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Chemistry review questions on Chapter 19
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7/24/2012
1
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
Chapter 19: Transition Metals
and Coordination Chemistry
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+
Vanadium Oxidation States
2 V(H2O)62+(aq) + H2O2(aq) + 2 H3O
+(aq)
2 V(H2O)63+(aq) + 4 H2O(l)
Vanadium Oxidation States
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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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.
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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
Some Lewis Acids and Bases
acid base
+
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A Lewis acid can expand its valence shell to accommodate a Lewis base
Some Lewis Acids and Bases
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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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
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?
Weak vs. strong field ligands
d
If we need to fill the d orbitals with four electrons, where does the fourth electron go?
Pairing the electron requires energy pairing energy (P)
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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?
Occupying an eg orbital requires energy
Weak vs. strong field ligands
d
If we need to fill the d orbitals with four electrons, where does the fourth electron go?
< P = Weak field > P = Strong field
Examples: [Cr(OH2)6]2+ [Cr(CN)6]
4-
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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?
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
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
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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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49
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)