Metalloproteins: Structure and Function 1.Introduction 1.1. Metalloproteins: Functions in Biological Chemistry 1.2. Some fundamental metal sites in metalloproteins 2. Mononuclear zinc enzymes: Carbonic anhydrase 3. Metalloproteins reacting with oxygen 3.1. Why do aerobic organisms need metalloproteins? 3.2. Oxygen transport proteins & Oxygenases Hemoglobin, Myoglobin Cytochrome P Hemerythrin & Ribonucleotide Reductase R2 & Methane monooxygenase diiron subunits Hemocyanin & Tyrosinase 4. Electron transfer proteins 4.1. Iron-sulfur proteins 4.2. Blue copper proteins 5. Conclusion F8390

3. Metalloproteins reacting with oxygen 3.1. Why do aerobic organisms need metalloproteins? Cells of aerobic organisms need oxygen. First, oxygen is needed to gain energy from food (respiration) and for other processes. Second, toxic organic substances are eliminated from the body by oxidation, whereupon OH-groups are attached to the molecule (this specific process is called hydroxylation, in mammals it occurs mainly in the liver). This renders the toxic molecule water-soluble and it can be eliminated (through the urine in mammals). Cellular respiration C 6 H 12 O O 2  6 CO H 2 O  G 0 = -674 kcal/mol Elimination of xenobiotics. Example: hydroxalation of hexane by Cytochrome P450 Cytochrome P450 minor major minor

1. The solubility problem Water solubility of oxygen at 25 o C and pressure = 1 bar is at 40 mg/L water. This is not enough to guarantee the oxygen supply to mitochondria by mere diffusion. Cells of aerobic organisms use therefore oxygen transporters. Use of oxygen by aerobic organisms is hampered by two problems: 2. The kinetic problem Oxygen has two unpaired electrons in its ground state and forms therefore a triplet state. The overwhelming majority of organic molecules (such as glucose or n-hexane) have all electrons paired and occur therefore in the singlet state. The products of oxidation of organic molecules, CO 2 and H 2 O, are also in singlet states. According to the so-called Wigner-rule, processes in which the spin-state changes are « spin-forbidden », that is, they have a large kinetic barrier. The solution of the problem is binding of O 2 to a transition metal complex. In transition metal complexes, spin-state changes are less inhibited due to the spin-orbit coupling. The oxygen-bound metal complex can therefore transit from a triplet state to a singlet state, and then react with an organic substrate which has also a singlet ground-state.

Activation of O 2 with the help of a transition metal complex: Adduct formation from a pentacoordinated [FeL 5 ] 2+ complex and O 2

2s 2p O O 2 O Molecular orbital level diagram for O 2 : 3  g - state

 *2p  2p  2p v  *2p h z x y yz xz  *2p v  2p h Bonding and antibonding MOs formed by AOs 2p in the O 2 molecule antibonding bonding

O O z x y O O antibonding  2p v  2p h  *2p v  * 2p h Bonding and antibonding MOs forming a  bond in the O 2 molecule

Activation of O 2 with the help of a transition metal complex: Adduct formation from a pentacoordinated [FeL 5 ] 2+ complex and O 2

Splitting of d orbitals in an octahedral environment (6 equal ligands) Cetral transition metal atom Lone-pairs of ligands xyxzyz z2z2 x 2 -y 2 M

Splitting of d orbitals in an tetragonal environment (5 equal ligands) Cetral transition metal atom Lone-pairs of ligands xyxzyz z2z2 x 2 -y 2 xz 6 ligands octahedral field 5 ligands octahedral field M

Splitting of d orbitals in an tetragonal environment (5 equal ligands) Cetral transition metal atom Lone-pairs of ligands xyxzyz z2z2 x 2 -y 2 xz x 2 -y 2 z2z2 6 ligands octahedral field 5 ligands octahedral field M

Splitting of d orbitals in an tetragonal environment (5 equal ligands) Cetral transition metal atom Lone-pairs of ligands xyxzyz z2z2 x 2 -y 2 xz xy x 2 -y 2 z2z2 6 ligands octahedral field 5 ligands octahedral field M Fe II is a d 6 ion Remember for next slide

3 O [L 5 Fe]  3 [L 5 FeO 2 ] spin-allowed: n° of unpaired electrons unchanged (only the two unpaired valence electrons shown) One of the  * orbitals of O 2 overlaps with the d z 2 orbital of Fe and forms a bond; the other  * orbital is non-bonding 1 [L 5 Fe] dz 2 empty From last slide 3O23O2

3 O [L 5 Fe]  3 [L 5 FeO 2 ] spin-allowed: n° of unpaired electrons unchanged (only the two unpaired valence electrons shown) One of the  * orbitals of O 2 overlaps with the d z 2 orbital of Fe and forms a bond; the other  * orbital is non-bonding process spin-forbidden but rendered possible by spin-orbit coupling 3O23O2 1 [L 5 Fe] 3 [L 5 FeO 2 ] Stabilization! 1 [L 5 FeO 2 ]

In transition metal complexes, spin-orbit coupling renders spin-forbidden transitions possible. Metal complexes can therefore activate (triplet) oxygen for reactions with (singlet) organic molecules. [ML n ] m+ + 3 O 2 1 [ML n O 2 ] m+ + 1 [Substrate] 1 [Oxidation products] Metal-oxygen adducts can also be used as oxygen carriers! 2. Oxygen transport proteins & oxygenases

Oxygen transport proteins: O 2 binding in active sites Hemoglobin (vertebrates, some invertebrates) Hemocyanin (molluscs, some arthropods) Hemerythrin (some marine invertebrates) Lippard: Bioinorganic Chemistry, 1994

2e - 2 in vertebrates Respiration: Reduction of O 2 to H 2 O by cytochrome c catalyzed by the enzyme cytochrome-oxidase Hemoglobin & Myoglobin

153 amino acids

Orbital energy difference is smaller than spin-pairing energy. Unpaired electrons are antiferromagnetically coupled  diamagnetic state Now let us make the energy difference between d xz,yz and  * n larger Weiss Nature 1964, 202, 83-84; 203, 183 Mechanism of O 2 binding to Fe in myoglobin and hemoglobin

Orbital energy difference is larger than spin-pairing energy  electron will go to d yz Pauling Nature ,

peroxide superoxide dioxygen Fe(II)-O 2, Fe(III)-O 2 -, or Fe(IV)-O 2 2- ? What experimental data can be used to determine whether oxygen in oxyhemoglobin resembles more to Fe(III)-O 2 - or to Fe(II)-O 2 ?

Stretching frequencies and bond lengths in dioxygen species Species O-O [cm -1 ]d O-O [A] O O O O Mb-O M-O M- O Oxymyoglobin resembles Fe III -O 2 -

O 2 versus CO discrimination Quantum chemical calculations indicate that the terminal O-atom is more negatively charged in the O 2 complex than in the CO-complex. Hydrogen bonding with distal histidine favors O 2 -binding against CO-binding. Isolated heme has a 10 4 times larger affinity for CO than for O 2. In myoglobin & hemoglobin, this discrimination factor drops to Why?

Hemoglobin/Myoglobin: History 1936 L.Pauling: Oxyhemoglobin is diamagnetic 1938-M. F. Perutz & J. C. Kendrew: x-ray studies on Hb & Mb 1939 L. Pauling: measurements of pK a  2 histidines, 1 coordinated + 1 farther from Fe proximal histidine pK a =6.8 distal histidine pK a = J. C. Kendrew: Crystal structure of myoglobin at 2 A 1964J. J. Weiss: Oxyhemoglobin is a Fe(III)-(O 2 - ) complex with antiferromagnetic coupling 1964 L.Pauling: Fe(II)-O 2 complex, postulate of H-bond 1981 S. E. V. Phillips, B. P. Schoenhorn: H-bond confirmed by neutron diffraction

Practical training - Download from the pdb database the structure of human hemoglobin 1HHO The coordinate file contains one half of the tetrameric structure, with one alpha and one beta subunit. - Display the structure using VMD - Highlight (using Graphics/Representation, Selected atoms „name FE“, drawing method CPK) the iron atoms of each subunit - Identify the residue numbers of both heme residues and highlight them - Identify the proximal and distal histidines - Measure the distances O-NE2 and Fe-NE2 distances for both histidines, for both hemes - Carry out the same for the CO-Hemoglobin complex 1HCO - Interpret your observations