GATE Bioinorganic Chemistry — Metalloproteins, Oxygen Carriers and Electron Transfer
Bioinorganic chemistry looks like biology to a student who has not prepared it, and like coordination chemistry to one who has. That second view is the useful one. Every question in this unit is really asking about a metal centre: its oxidation state, its geometry, its spin state, its ligands, and what the spectroscopy tells you about them. If you can do crystal field theory and read an IR or UV–visible spectrum, you already have the tools. This guide organises the topic around that idea, with three worked calculations and a table of the systems worth knowing by name.
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Why particular metals, and the Irving–Williams series
Biology uses a small set of metals, chosen by availability and by chemistry. Na⁺, K⁺, Mg²⁺ and Ca²⁺ are structural and signalling ions with no redox role. Fe, Cu, Mn, Co, Ni, Mo are redox-active and appear in catalysis and electron transport. Zn²⁺ is special: it is d¹⁰, so it has no accessible redox chemistry and no crystal field stabilisation energy, and therefore no geometric preference — which makes it a superb Lewis-acid catalyst that can be four-, five- or six-coordinate as the enzyme requires.
The stability of high-spin octahedral M(II) complexes with a given ligand follows the Irving–Williams series:
The rise across the series comes from decreasing ionic radius and increasing CFSE; the peak at Cu(II) reflects the extra stabilisation from Jahn–Teller distortion of the d⁹ ion; and the drop at Zn(II) follows from its d¹⁰ configuration having zero CFSE. This series explains why proteins must bind weaker-binding metals selectively rather than simply relying on affinity.
Oxygen carriers — the three systems compared
| Protein | Metal site | Bound O₂ described as | Approx. ν(O−O) in the oxy form | Colour change on oxygenation |
|---|---|---|---|---|
| Haemoglobin / myoglobin | One Fe in a porphyrin, proximal histidine below | Superoxide-like, end-on bent Fe−O−O | ≈1105 cm⁻¹ | Purple-red → bright red |
| Haemocyanin | Two Cu, no porphyrin, histidine ligands | Peroxide, side-on bridging between two Cu | ≈750 cm⁻¹ | Colourless → blue |
| Haemerythrin | Two Fe, no porphyrin, carboxylate/oxo bridged | Hydroperoxide bound to one Fe | ≈844 cm⁻¹ | Colourless → burgundy |
The wavenumbers are the evidence, not decoration. Free O₂ has ν(O−O) near 1556 cm⁻¹, superoxide O₂⁻ near 1145 cm⁻¹ and peroxide O₂²⁻ near 740–800 cm⁻¹, because adding electrons into the π* orbital lowers the bond order from 2 to 1.5 to 1. Measuring ν(O−O) in the oxy protein therefore tells you how much charge the metal has transferred to the oxygen — which is precisely the reasoning a GATE question wants to see. Note also that despite the name, haemocyanin and haemerythrin contain no haem group.
Haemoglobin — the spin-state story
Deoxyhaemoglobin has high-spin Fe(II), d⁶, five-coordinate, bound to four porphyrin nitrogens and one proximal histidine. Being high spin, it has electrons in the eg (σ-antibonding) orbitals, which makes the ion larger, so it cannot fit inside the porphyrin hole and sits about 0.4 Å out of the plane, pulled towards the histidine.
When O₂ binds at the sixth site, the iron becomes low spin d⁶ (t2g⁶, eg⁰), which is smaller, and it moves into the porphyrin plane. That small movement drags the proximal histidine and, through it, the whole protein helix — which is the mechanical link that makes the other three subunits bind oxygen more easily. This is cooperativity, and it is why haemoglobin's binding curve is sigmoidal while monomeric myoglobin's is a simple hyperbola.
θ is the fractional saturation, p the partial pressure of O₂, P₅₀ the pressure at half-saturation and n the Hill coefficient. For myoglobin n = 1 (no cooperativity); for haemoglobin n is about 2.8, less than the 4 that perfect cooperativity would give.
Worked example 1 — how much difference does cooperativity make?
Compare the fractional saturation of myoglobin (n = 1) and haemoglobin (n = 2.8) at an oxygen pressure equal to twice their own P₅₀.
Myoglobin, n = 1, p = 2 P₅₀:
θ = 2 ÷ (1 + 2) = 2/3 = 0.667, i.e. 66.7% saturated
Haemoglobin, n = 2.8, p = 2 P₅₀:
22.8 = e2.8 × 0.6931 = e1.9407 = 6.964
θ = 6.964 ÷ (1 + 6.964) = 6.964 ÷ 7.964 = 0.874, i.e. 87.4% saturated
What this shows. A cooperative carrier fills up much more sharply around P₅₀, so it can load nearly fully in the lungs and unload substantially in the tissues over a modest pressure change. A non-cooperative carrier cannot do both. That is the whole functional point of n > 1, and it also explains why myoglobin — which stores rather than transports — does not need cooperativity and has a lower P₅₀ so that it takes oxygen from haemoglobin.
Two further regulators are worth naming: the Bohr effect (lower pH and higher CO₂ in active tissue reduce oxygen affinity, releasing more O₂ where it is needed) and 2,3-bisphosphoglycerate, an allosteric effector that binds the deoxy form and lowers affinity.
Electron-transfer proteins
- Cytochromes — iron in a porphyrin, cycling between Fe(III) and Fe(II). They are classified a, b, c by the haem type and axial ligation. Unlike haemoglobin, both axial sites are occupied by protein ligands, so the iron cannot bind O₂ — the site is blocked deliberately, because a cytochrome's job is to move electrons, not to carry oxygen.
- Iron–sulfur proteins — [2Fe–2S], [3Fe–4S] and [4Fe–4S] clusters with cysteine thiolate ligands. Rubredoxin has a single Fe with four cysteines. These clusters tune their reduction potentials over a very wide range by changing the protein environment, which is how a single structural motif serves many different steps of a chain.
- Blue copper (type 1) proteins — plastocyanin and azurin. The copper is held in a distorted geometry, roughly trigonal with a long axial contact, by two histidines, a cysteine and a methionine. The intense blue colour is a ligand-to-metal charge transfer band from the cysteine sulfur to Cu(II) near 600 nm, with a molar absorptivity of the order of 5 × 10³ dm³ mol⁻¹ cm⁻¹ — about a hundred times stronger than a typical Laporte-forbidden d–d band, which is exactly how you can tell the two apart. The strained, "entatic" geometry lies between the preferences of Cu(II) and Cu(I), so very little reorganisation is needed when an electron arrives — and fast electron transfer is the result.
Worked example 2 — quantifying a metalloprotein by Beer–Lambert
A purified blue copper protein has ε = 4500 dm³ mol⁻¹ cm⁻¹ at its charge-transfer maximum. A solution in a 1.00 cm cell gives an absorbance of 0.360. Find the concentration.
Substituting:
c = 0.360 ÷ (4500 × 1.00) = 8.0 × 10⁻⁵ mol dm⁻³
Why this is the standard method. Because the charge-transfer band is so intense, a very dilute protein solution still gives a measurable absorbance, and the band is characteristic of that specific copper site. The same approach quantifies cytochromes from their Soret and α bands. Use the ε value supplied in the question — extinction coefficients differ between proteins and are determined experimentally for each one.
Zinc enzymes — Lewis acid catalysis
Carbonic anhydrase is the standard example and appears repeatedly. A Zn²⁺ ion is held by three histidine residues, and the fourth coordination site holds a water molecule. Coordination to the dipositive metal pulls electron density from the water and lowers its pKa dramatically — from about 15.7 for free water to roughly 7 — so at physiological pH a substantial fraction of the enzyme carries a bound hydroxide. That Zn−OH⁻ is the nucleophile that attacks CO₂ to give bicarbonate. Generating a strong nucleophile at neutral pH is the entire catalytic trick, and it works precisely because Zn(II) is a good Lewis acid without being redox-active.
Carboxypeptidase A uses zinc in the same spirit, polarising the substrate carbonyl and activating water for attack on the peptide bond.
Worked example 3 — energetics of a single electron transfer
An electron passes from a donor with reduction potential +0.03 V to an acceptor at +0.25 V, one electron at a time. Find ΔG°′.
Step 1 — the potential difference.
E°cell = E°(acceptor) − E°(donor) = 0.25 − 0.03 = 0.22 V
Step 2 — the free energy. n = 1, F = 96,485 C mol⁻¹
ΔG° = −1 × 96,485 × 0.22 = −21,227 J mol⁻¹ = −21.2 kJ mol⁻¹
The useful conversion to remember: for a one-electron transfer, 1 volt corresponds to about 96.5 kJ mol⁻¹. That single number lets you convert any reduction-potential difference into an energy in your head. Quote whichever potentials the question supplies — biological reduction potentials are conditional values that depend on the protein and the conditions, so they are given in the problem rather than memorised.
Other systems worth knowing by name
| System | Metal centre | Function | Point examiners like |
|---|---|---|---|
| Chlorophyll | Mg in a chlorin ring | Light harvesting in photosynthesis | Mg is not redox-active; it organises the ring |
| Vitamin B₁₂ / coenzyme B₁₂ | Co in a corrin ring | Rearrangement and methyl transfer | A genuine metal–carbon bond in biology |
| Nitrogenase | FeMo cofactor plus an Fe protein | Reduces N₂ to NH₃ | Requires ATP; produces H₂ alongside NH₃ |
| Cytochrome P450 | Fe haem with a cysteine thiolate axial ligand | Inserts an O atom into C−H bonds | The thiolate ligand is what enables the oxidation |
| Catalase / peroxidase | Fe haem | Destroys H₂O₂ | High-valent iron-oxo intermediate |
| Superoxide dismutase | Cu/Zn, Mn or Fe | Converts O₂⁻ to O₂ and H₂O₂ | Zn is structural; Cu is the redox centre |
| Photosystem II oxygen-evolving complex | Mn₄Ca cluster | Oxidises water to O₂ | Cycles through several oxidation states |
| Ferritin | Iron oxide-hydroxide core | Iron storage | Stores iron safely; free Fe would generate radicals |
| Na⁺/K⁺-ATPase | No redox metal centre | Ion gradient across membranes | Moves 3 Na⁺ out for 2 K⁺ in, per ATP |
Metal toxicity and chelation
Heavy metals are toxic largely because they bind strongly to thiol groups and displace the native metal from an enzyme's active site. The chemical remedy is a ligand that binds the toxic metal more strongly than the protein does and forms a complex that can be excreted. Applying hard–soft acid–base reasoning predicts the pairings: soft metal ions such as Hg(II) and Pb(II) are best chelated by soft sulfur donors, which is why dithiol ligands are used, while a hard ion is better matched by oxygen and nitrogen donors such as those of EDTA. A chelating ligand always binds more strongly than the equivalent number of separate ligands — the chelate effect, driven mainly by the favourable entropy of releasing several small ligands when one multidentate ligand binds. Selectivity matters as much as strength: a chelator that also strips essential zinc, copper or calcium creates a new problem while solving the old one. (This is chemistry, not clinical advice; any actual treatment is a matter for a qualified doctor.)
Common mistakes that cost marks
- Saying oxyhaemoglobin contains Fe(III) and O₂. The accepted description is Fe(III) bound to superoxide, formulated as an Fe(II)−O₂ adduct with substantial charge transfer; the diamagnetism of the oxy form is the evidence. What is certainly wrong is calling it a simple Fe(II) with unchanged O₂. Methaemoglobin, with genuine Fe(III) and no bound O₂, cannot carry oxygen at all.
- Assuming haemocyanin and haemerythrin contain haem. Neither does; the names are historical.
- Attributing the blue colour of type 1 copper to a d–d transition. It is a cysteine-S → Cu(II) charge-transfer band; the very high molar absorptivity is the proof.
- Giving zinc a redox role. Zn(II) is d¹⁰ and redox-inert in biology; its function is Lewis acidity and structure.
- Mixing up myoglobin and haemoglobin curves. Hyperbolic for myoglobin, sigmoidal for haemoglobin; myoglobin also has the lower P₅₀.
- Forgetting why the iron moves into the porphyrin plane. It is the high-spin to low-spin change and the resulting decrease in ionic radius, not the mass or charge of the O₂.
- Treating the Irving–Williams order as ending at Cu. It peaks at Cu(II) and then falls at Zn(II).
Preparation map
| Sub-topic | Question style | Priority |
|---|---|---|
| Oxygen carriers compared | Match protein, metal, ν(O−O) and colour | High |
| Spin state and geometry in Hb | Explain the trigger for cooperativity | High |
| Electron-transfer proteins | Identify the site from its spectrum or ligands | High |
| Zinc enzymes | Why the bound water's pKa falls | High |
| Beer–Lambert on metalloproteins | Concentration from absorbance | Medium — quick numerical |
| Irving–Williams and the chelate effect | Order stabilities; explain the trend | Medium |
| Nitrogenase, B₁₂, P450, OEC | Name the metal and the reaction | Medium |
The one calculation this unit asks you to do is Beer–Lambert. Quantifying a metalloprotein from the absorbance of its charge-transfer or Soret band — and going the other way to find ε from a known concentration — is standard practice and a standard exam question. The free Beer–Lambert calculator handles A, ε, c and path length in any direction, so you can check your rearrangement and your powers of ten before you commit an answer.
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