Biophysics of Metalloenzymes

Topics and Themes:

1) (Metallo-) Proteins and Enzymes in the Cell

2) Some Principles of Coordination Chemistry

3) Methods for Investigation at Molecular Level

4) Overview on Metal Cofactors in Biology

5) Cofactor Assembly and Maturation

6) Biological Excitation-Energy and Electron Transfer

7) Proton Transfer at Metal Cofactors

8) Metal centers in Photosynthesis and Water Oxidation

9) Biological Hydrogen Catalysis

10) Metal Cofactors in Nitrogen Fixation

11) Carbon Oxide Conversion at Metal Sites

12) Molybdenum Enzymes

13) Oxygen Activation Reactions

14) Metal Centers in Human Diseases

15) Bioinspired Materials

Biophysics of Metalloenzymes M. Haumann

Biologically Relevant Elements

biomass constituting elements

further essential elements

essential in certain organisms

used in therapy and diagnostics

1st-row transition metals (V, Mn, Fe, Co, Ni, Cu, Zn) are most common in cofactors in proteins

2nd+3rd-row metals (Mo, W) used in certain enzymes

(Earth) alkali metals mostly structural functions

Main Elements of Life

All organisms (biomass) contain relatively few main elements:

.

Main elements of Life (carbon chemistry):

6 main elements constitute 96-98% of the human body, i.e. [%]:

oxygen (O) 61

carbon (C) 22.5

hydrogen (H) 9.5

nitrogen (N) 2.5

calcium (Ca) 1.7

phosphorus (P) 1.1

sulfur (S) 0.3

potassium (K) 0.2

sodium (Na) 0.1

All further elements occur in very small quantities (trace elements)

Fluor F 5 1931

Eisen Fe 4 17. Jh

Zink Zn 2.3 1896

Silicium Si 1 1972

Titan Ti 0.7

Brom Br 0.26

Kupfer Cu 0.07 1925

Zinn Sn 0.03 1970

Cadmium Cd 0.02 1977

Nickel Ni 0.015 1975

Iod I 0.015 1820

Selen Se 0.014 1975

Mangan Mn 0.012 1931

Arsen As 0.007 1975

Molybdän Mo 0.005 1953

Chrom Cr 0.002 1959

Cobalt Co 0.002 1935

Element Symbol Mass (g) discovered as essential

Trace Elements in the Human Body (70 kg)

The Transition Metals

Structural Properties

Atomic radii:

Ionization Energies

1st

2nd

d-Electron Configuration

Electron configurations of ions of the first-row transition metals

— the energy of the 3d orbitals is significantly less than that of the 4s orbital

e.g. Ti: 4s23d2 Ti3+: 3d1

but Zn: 4s23d10 Zn2+: 3d10

— these ions do not have 4s electrons (since the 3d orbitals are lower in energy)

Oxidation states and ionization energies

various ions formed by loosing electrons

e.g. Ti → Ti2+, Ti3+, Ti4+

3d4 most common

— to the right of the row the higher oxidation states are not observed because the

3d orbitals become lower in energy as the nuclear charge increases, making

electrons difficult to remove

e.g. Zn → Zn2+ {Zn3+, etc ← NOT OBSERVED}

4s23d10 observed

Observed Oxidation States

(4)

(4)

3

4 5 6(3)( )

Most Important 3d Transition Metals

Nickel, Ni

-mainly in the +2 (+3) oxidation state(s)

-aqueous solutions of Ni(II) salts contain Ni(H2O)62+ ion (hexaquo-complex)

Vanadium, V

The most common oxidation state is +5 as in V2O5 and VF5

Manganese, Mn–The only 3d metal that can exist in all oxidation states from +2 to +7–Manganese(VII) ion: MnO4

- permanganate ion (strong oxididant in solution)

Iron, Fe

-chemistry mainly involves +2 and +3 oxidation

states (but +4, +5, +6 in biology)

Copper, Cu

-many alloys contain copper (brass, bronze, sterling silver, 14-18 Karat Au)

-chemistry involves +2 oxidation state, but also some compounds with +1, +3

Zinc, Zn

+2 oxidation state only

Vanadium metal (center) in solution as

V2+(aq), V3+(aq), VO2+(aq), and VO2+(aq),

(left to right)Co2+ Mn2+ Cr3+ Fe3+ Ni2+

Coordination Chemistry

Color is a qualitative assay of valence levels (electronic structure)

Transition metal complexes often colored (absorb light in UV/vis)

Metal may show paramagnetism depending on oxidation state and ligands

Some Transition Metal Properties

Typical metals, high electrical and thermal conductivities

• Similarities within a given period as well as within a given vertical group

due to the fact that the last electrons added to the transition metal elements are

inner electrons:

• d electrons in d – block transition metals

Differences in physical properties among the transition metals can be large

“hard” (non-polarizable) vs “soft” (polarizable)

Fe, Ti Cu, Au, Ag

ready to form oxides vs no reaction with O2

Cr, Ni, Co, Fe Au, Ag, Pt

Transition metals often show multiple oxidation states (and varying complex charge)

Ionic compounds with nonmetals

the cations are often complex ions, species in which the transition metal ion is

surrounded by a number of ligands

Coordination Compounds

A (typical) coordination compound consists of a metal ion and ligands

– it is an ionic compound, electrically neutral

– complex ion = transition metal ion + attached ligands

e.g. [Co(NH3)5Cl]Cl2

Co(NH3)5Cl2+ ← complex ion

2 Cl- ← counter ions (anions)

– coordination compounds ionize in solutions (like salts), solvent

[Co(NH3)5Cl]Cl2 (s) Co(NH3)5Cl2+(aq) + 2 Cl- (aq)

Coordination number of metal ions

– the number of bonds formed between a metal ion and the ligands in the complex

ion is termed the coordination number

– depending on the size, charge, and electron configuration of the transition metal

ion, the coordination number can be from 2 to 8

– In proteins mainly 3,4,5,6 ligands observed for 1st-row 3d metals

– many metal ions show more than one coordination number

Ligands and Bonding

Ligands

–a neutral molecule or ion having a lone pair that can be used to form a

bond with a metal ion

-Lewis bases by definition are ligands

all serve as σ-donors, some are π-donors as well, and some are π-

acceptors

-the metal ion is a Lewis acid

–a metal – ligand bond is called a coordinative covalent bond

-it results from a Lewis acid – base interaction in which a ligand

donates an electron pair to an empty valence orbital on a metal ion

- dative bond

• coordination number and geometry depend on metal and number of d-

electrons

In solution: ligands are usually water species (H2O, OH-)

In protein: ligands are sidechains of amino acids or non-protein molecules

(e.g. CN, CO, substrates)

Biophysics of Metalloenzymes M. Haumann SS2014

Hard and Soft Acids and Bases (HSAB)

Lewis-acids and -bases

Qualitative description of chemical reactions

Structure and reactivity of metal complexes

hard: particles with large charge density (large charge/radius ratio), low polarizability

soft: small charge density, high polarizability

The hardness of an acid increases with decreasing size, increasing charge, and smaller

polarizibility

The hardness of bases increases for decreasing size, decreasing polarizability, and decreasing

oxidizability

Reactions of hard acids with hard bases and soft acids with soft bases lead to more stable

compounds than hard – soft combinations

hard bases stabilize high oxidation states

soft bases stabilize low oxidation states

Coordination Complex

H2N

CoH2N

H2N

NH2

Cl

Cl

+

central atom (metal)

dativ ligands (2e- donor)

σ-ligands (1e- donor)

1

2

34

6

5

coordination number 6

oxidation state

III

+3

+

geometry: octahedral

H2N

CoH2N

H2N

NH2

Cl

Cl

+

complex charge number +1

[CoIII(NH2)4Cl2]+

Calculation of Formal Metal Oxidation State

[MnO4]- (permanganate) contribution total

Valence electrons on Mn(0) 7 7

Complex charge (attributed to the metal) +1 8

Remove 4 O22- ligands (closed-shell) -8 0

Oxidation state (difference between 7 and 0) Mn(VII) +7

[Cr(NH3)6]3+

Valence electrons on Cr(0) 6 6

Complex charge (attributed to the metal) -3 3

Remove 6 neutral NH3 ligands -0 3

Oxidation state (difference between 6 and 3) Cr(III) +3

Stabilization of Oxidation States

Geometries and Coordination Numbers

L M

L

L

LM

L

LML

L M

L

L

ML L

L

L

ML L

L

L M

L

L

L

tetraedrisch

quadratisch-planar

L ML

L

L

L

M

L

L L

L

L

ML

L L

L

L

L

L

LL

M

L

LL

trigonale Bipyramide

quadratische Pyramide

Oktaeder

trigonales Prisma

2 3 4 5 6

tetrahedral (Td) trigonal-bipyramidal octahedral (Oh)

square-planar (D4h) square-pyramidal (C4v) trigonal-prismatic

Geometry: mostly independent of ground state electronic configuration, difficult to predict, ligand

repulsion, metals with different d electron count can have same geometry

Steric: M-L bonds are arranged to have the maximum separation around the metal

(protein: matrix restaints)

Electronic: d electron count combined with the complex electron count must be

considered when predicting geometries for TM complexes with non-bonding d electrons

e.g. CN = 4, d8 (16 e−) prefers square planar geometry, d10 (18e−) prefers tetrahedral geometry

linear

Preferred Coordination Numbers & Geometries

-metal atom size, steric interactions between ligands, electronic interaction metal & ligands

-high CN favored by high oxidation state (e− poor) metals and small ligands

Irving-Williams series

Irving-Williams Series: relative stabilities of complexes formed by a metal ion. For high-spin

complexes of the divalent ions of first-row transition metals, the complex stability constants are:

Mn(II) < Fe(II) < Co(II) < Ni(II) < Cu(II) > Zn(II)

This order was found to hold for a wide variety of ligands. Explanations:

(1) Ionic radius decreases for Mn2+ to Zn2+.

(2) Ligand Field Stabilization Energy (LFSE) increases from zero for Mn(II) to a maximum at

Ni(II). This makes the complexes increasingly stable. CFSE for zinc(II) is zero.

(3) Although the CFSE of Cu(II) is less than that of Ni(II), octahedral Cu(II) complexes are

subject to the Jahn-Teller effect, which affords a complex extra stability.

When the stability constants are quantitatively adjusted for this effect, they follow the trend.

None of these explanations can satisfactorily explain the broad scope of validity of Irving-

Williams series (both octahedral and tetrahedral complexes containing different ligands). The

covalent and electrostatic contributions to metal-ligand binding energies further contribute.

Valence Orbitals

-A transition metal ion has 9 valence atomic orbitals: 5 (n)d, 1 (n+1)s, 3 (n+1)p

-These orbitals are of appropriate energy to form bonding interaction with ligands

-Orbitals oriented orthogonal to each other create possibilities for ligand overlap

-For an σ –only bonding Oh complex, 6 σ bonds are formed and the d orbitals are non-bonding

-The non-bonding d orbitals often give TM complexes many of their properties

18-Electron Rule

The 18-valence-electron rule is used for predicting the formulae of stable metal complexes

The 9 valence orbitals of transition metals can accommodate 18 electrons as bonding or

nonbonding electron pairs

Combination of these 9 atomic orbitals with ligand orbitals creates nine molecular orbitals

that are either metal-ligand bonding or non-bonding

18 valence electrons = noble gas or closed-shell configuration -> most stable

The ligands in a complex determine the applicability of the 18-electron rule

Complexes that obey the rule often have strong-field ligands (CO) lowering the energies of

the resultant molecular orbitals so that they are occupied

Compounds that obey the 18 VE rule are typically "exchange inert“ but often reactive toward

electrophiles such as protons

Complexes with fewer than 18 valence electrons tend to show enhanced reactivity

Ligands excerting a weak ligand field increase the energies of orbitals so that they can be

non-occupied and complexes with less than 18 VE are stable (addition or removal of an

electron has little effect on complex stability)

Ligand Field Theory

spherical field → octahedral field

M ML

L L

L

L

L

6 L

"Oktaederfeld"

„ligand field“

LFT describes the bonding, orbital arrangement, and other properties of coordination complexes

-It represents an application of molecular orbital theory to transition metal complexes

d-Level Degeneracy

dxzdxy dyzdz2 dx2-y2

∆ = ligand field splitting energy

„degenerate“

(entartet)

electronic

levels

d-orbitals

Ligand Field Splitting Energy

• the nature of the metal ion

• the metal's oxidation state - a higher oxidation state leads to a larger splitting

• the arrangement of the ligands around the metal ion

• the nature of the ligands surrounding the metal ion. The stronger the effect of the

ligands, the greater the difference between the high and low energy 3d groups

LFSE (crystal field splitting, ∆) depends on:

d-Level Degeneracy and Geometry

t2g

eg

d-Levels and Distortion

octahedral square-planar

Ligand Removal

t2g

∆o

Oh

∆t

energ

y

eg*

ML

L L

LM

L

L L

L

L

L

z

- 2 L

1 eσ

2 eσ

L

ML

LL

b1g

a1g

eg+b2g

D4h non-bonding

t2*

e

Td

anti-bonding

occupation does not

cause a change in

bond order

occupation lowers the

stability of the complex

Bond lengths and Geometry

E

Ha.

Ψa(1s)

LCAO

atomic orbital

Hb.

Ψb(1s)

Ψσ = N • c1•Ψa(1s) + c2•Ψb(1s)

MO

σ∗-orbital

σ-orbital

MO: Linear Combination of Atomic Orbitals

Bonding

σ-bond (rotational symmetry to binding axis)

-for example s+s, pz+pz, s+pz and dz2+dz2

-the molecular orbitals created by coordination result from the donation of 2 electrons by each

of 6 σ-donor ligands to the metal valence orbitals

-in Oh complexes, ligands approach along the x-, y- and z-axes, so their σ-symmetry orbitals

form bonding and anti-bonding combinations with the dz2 and dx2−y2 orbitals

-the dxy, dxz and dyz orbitals remain non-bonding orbitals

-weak bonding (and anti-bonding) interactions with the s and p orbitals of the metal also occur

-this gives a total of 6 bonding (and 6 anti-bonding) molecular orbitals

π-bond

-for example p-p, weaker than σ-bonds

(1) via ligand p-orbitals that are not used in σ bonding

(2) via π or π* molecular orbitals on the ligand

-the ligand π orbitals interact with the dxy, dxz and dyz metal d-orbitals (these orbitals are non-

bonding when only σ bonding takes place)

-π bonding of the metal-to-ligand type (also called π –backbonding) occurs when the LUMOs of

the ligand are anti-bonding π* orbitals.

-π bonding of the ligand-to-metal type occurs when the π orbitals on the ligands are filled. They

combine with the dxy, dxz and dyz orbitals on the metal and donate electrons to the resulting π-

symmetry bonding orbital between them and the metal

-as each of the six ligands have two orbitals of π-symmetry, there are 12 in total.

For octahedral geometry:

M

L

dz2

p

metal ligandcomplex

σ

eσ∗

e

σ

M

L

L

+ M + M L LM+ M M

L

L++

1 eσ + 1 eσ + 1/4 eσ + 1/4 eσ + 1/4 eσ + 1/4 eσ = 3 eσ

σ-Bonds

L

M

dxz

metal ligandcomplex

π∗

eπ

eπ∗

L

M

dxy

M M ML L

L

eπ + eπ + eπ + eπ = 4 eπ

+ + +

π-Bonds

σ

σ∗

dz2

π

π∗

dyz dxz

δ

δ∗

dx2-y2 dxy

Metal-metal bonds

Bond Strength - Overlap Integral

σ > π

distance dependence

angular dependence

π backbonding

π backbonding (π backdonation): electrons

move from an atomic orbital on the metal atom

to a π* antibonding orbital on a π-acceptor

ligand. It is common in transition metal

complexes with multi-atomic ligands such as

carbon monoxide. Electrons from the metal are

used to bond to the ligand, in the process

relieving the metal of excess negative charge.

Bonding of π-conjugated ligands to a transition

metal involves a synergic process with donation

of electrons from the filled π-orbital or lone

electron pair orbital of the ligand into an empty

orbital of the metal (donor–acceptor bond),

together with release (back donation) of

electrons from an nd orbital of the metal (which

is of π-symmetry with respect to the metal–

ligand axis) into the empty π*-antibonding

orbital of the ligand.

(Top) the HOMO and LUMO of CO. (Middle) an example of a σ-

bonding orbital in which CO donates electrons to a metal center

from its HOMO. (Bottom) an example where the metal center

donates electrons through a d-orbital to CO LUMO.

http://en.wikipedia.org/wiki/Pi_backbonding

(i) strengthens

the M-C

bond

(ii) weakens

the C-O

bond

Charge Transfer

high-spin vs low-spin

-Hund´s Rule vs LFSE + spin-pairing energy (2e in one orbital)

-strong-field ligands (CO, CN, NO2): large ∆, low-spin complex

-weak-field ligands (eg halogenides): small ∆, high-spin complex

[Fe(NO2)6]3− [FeBr6]

3−

-in an octahedral ligand field for

transition metals with 4,5,6 or 7

d-electrons high and low-spin

complexes

Spin State

d2-configuration

(t2g)2

Configuration

t2g

eg

(t2g)1(eg)

1

∆o

(eg)2

∆o

d-Level Energies and Spin State

Spin crossover compounds

LIESST

Thermodynamically

driven (+ vibrational

couplings) spin state

change, F(T, P, hν, etc.)

YHS(T) = [1 + exp(∆H/kBT – ∆S/kB)]-1

[ ] 1))/1/1(/exp(1)(

−−∆+=

CBHSTTkHTY

Mössbauer

Tc

HOMO - LUMO

HOMO

LUMO

20 30 40 50 60

-3

-2

-1

0

1

2

3

4

MO

en

erg

y / e

V

Fed contribution to MO / %

Fep Fe

d

occ α ß α ßuno α ß α ß

Fe(II) Fe(0)

frontier orbitals

MO Energies and Redox Potentials

PMe3

N

Fe

CO

CO

S

Fe

OC

CO

S

Me3P

Ph

12

Spectro-Chemical Series

Spectrochemical series of ligands

Ligands causing small ∆ to large ∆ values:

I− < Br− < S2− < SCN− < Cl− < NO3− < N3− < F− < OH− < C2O42− ≈ H2O < NCS− <

CH3CN < py (pyridine) < NH3 <en (ethylenediamine) < bipy (2,2'-bipyridine) < phen (1,10-

phenanthroline) < NO2− < PPh3 < CN− ≈ CO

Compounds on the left are weaker ligands, compounds at the right are stronger ligands

Explanation needs to consider covalency of bonds (not purely ionic like in crystal field theory)

Spectrochemical series of metals

Metal ions for increasing ∆, largely independent of the ligands:

Mn2+ < Ni2+ < Co2+ < Fe2+ < V2+ < Fe3+ < Cr3+ < V3+ < Co3+

Whether a ligand exerts a strong or weak field on a given metal ion can not be predicted

∆ increases with increasing oxidation state, ∆ increases down a group

Solutions of Co3+ ions in the presence of:

(a) cyanide ions

(b) nitrite ions

(c) phenanthroline

(d) diaminoethane

(e) ammonia

(f) glycine

(g) water

(h) oxalate ions

(i) carbonate ions

Jahn-Teller Distortion

energy

Molecules in a (highly) degenerate electronic ground state are unstable and try to lower their

energy by splitting of the energy levels due to lowering of the symmetry

Occupied MOs become lower in energy, unoccupied MOs become higher in energy

In Oh complexes, the JT effect is most pronounced when an odd number of electrons occupy

the eg orbitals, i.e. d9, low-spin d7 (FeI) or high-spin d4 (MnIII) complexes

Ligation Motifs

Bridging ligands!

TM Sites in Proteins

metal intrinsic properties vs distorted (amino acid) ligand environment

„Solvatization“, charge, protonation

Mixed ligands

Bond strength, covalency

Coordination number

Geometry

Degeneracy

Oxidation state

Spin state

Delocalization

Redox potential

Protein-bound transition metals are particularly prominent in small

molecule activation

Small molecules in biological catalysis:

N2, O2, H2, CO, NO, CN, CO2, H2O, H2O2, N2O, CH4, NH4...

„ubiquituous reservoirs of chemical energy“

-Powering biological systems,

-Building of more complex molecules

-Signaling agents

-Inert at ambient conditions

-Thermodynamically stable

-High kinetic barrier for activation

Fundamental questions:

-How do metal ions in enzymes coordinate to and modulate the reactivity of

small molecules?

-Are there general principles that govern small molecule catalysis in Biology?

-Can one use knowledge of metal/small-molecule chemistry in enzymes for

the design of new synthetic catalysts?

Small Molecules as Ligands and Substrates

Biophysics of Metalloenzymes M. Haumann WS2014/15

Metal Cofactors in ProteinsR

Numerous fascinating systemsb

Biophysics of Metalloenzymes M. Haumann WS2014/15

Summary

Elements of life

Transition metals

Coordination compounds

HSAB

Oxidation state

Geometry

18 VE rule

Ligand field theory

Bonding

Irving-Williams series

Spectrochemical series

Spin state

Spin crossover compounds

HOMO/LUMO

Jahn-Teller

Ligation motifs

TMs in proteins

Small molecule ligands

Biophysics of Metalloenzymes M. Haumann WS2014/15

Literature

Coordination Chemistry, Gispert, Wiley-VCH, 2008

Coordination Chemistry: Concepts and Applications, Comba & Kerscher, Wiley-VCH, 2015

Koordinationschemie, Gade, Wiley-VCH, 2010

Koordinationschemie: Grundlagen und aktuelle Trends, Weber, Springer Spectrum, 2014

Bioinorganic Chemistry -- Inorganic Elements in the Chemistry of Life, Kaim et al., Wiley, 2013