Bio-inorganic and environmental chemistry

Metalloproteins are coordination complexes wearing a biological disguise: the "ligands" are amino-acid side chains (histidine imidazole nitrogen, cysteine thiolate sulfur) and organic macrocycles (porphyrin, corrin, chlorin rings) built into a protein scaffold instead of small inorganic molecules in a flask.

Haemoglobin and myoglobin (Fe). The active site is haem: an $\text{Fe}^{2+}$ ion at the centre of a tetradentate porphyrin ring (four nitrogen donors), a fifth site taken by a "proximal histidine" nitrogen, and the sixth site open for $\text{O}_2$. Deoxy-Fe(II) is five-coordinate and high-spin; on oxygen binding the field strengthens, iron becomes low-spin, and its smaller ionic radius pulls it into the porphyrin plane. That geometric shift is transmitted through the protein and drives cooperative $\text{O}_2$ binding across haemoglobin's four subunits — myoglobin, a single subunit, shows no such cooperativity.

Chlorophyll (Mg). A related macrocycle, the chlorin (a reduced porphyrin), but with central $\text{Mg}^{2+}$, a closed-shell $d^0$ ion. Chlorophyll's strong visible absorption therefore cannot come from d–d transitions; it comes from $\pi \to \pi^*$ transitions delocalised over the conjugated ring. Magnesium's job is purely structural — holding the macrocycle rigid and planar.

Vitamin $B_{12}$ (Co). Built on a corrin ring, smaller and less conjugated than porphyrin, coordinating cobalt. Its defining feature is a direct Co–C $\sigma$ bond, one of the very few stable metal–carbon bonds found in biology, which enables radical and methyl-transfer catalysis unavailable to the other three metals here.

Zinc in carbonic anhydrase. $\text{Zn}^{2+}$ is $d^{10}$ — a filled subshell, so no d–d transition and no colour. It sits in roughly tetrahedral coordination (CN 4): three histidine nitrogens plus a water molecule. Zinc's Lewis acidity polarises that water, lowering its $pK_a$ so it exists as nucleophilic $\text{Zn–OH}^-$ at physiological pH, which attacks $\text{CO}_2$ to give bicarbonate. Zinc contributes no redox chemistry — consistent with a closed $d^{10}$ shell.

Copper in electron-transfer proteins (plastocyanin, azurin). Copper cycles between $\text{Cu}^{2+}$ ($d^9$) and $\text{Cu}^{+}$ ($d^{10}$). The protein imposes a distorted "entatic state" geometry that minimises the reorganisation energy needed to flip between these two oxidation states rapidly, making copper an efficient single-electron shuttle in electron-transport chains. The unifying quantity behind all of this is crystal field stabilisation energy (CFSE), which for an octahedral field is

where: $n_{t_{2g}}$ and $n_{e_g}$ are the numbers of electrons in the lower ($t_{2g}$) and upper ($e_g$) split d-orbital sets, and $\Delta_o$ is the octahedral crystal field splitting energy. A filled $d^{10}$ shell ($n_{t_{2g}}=6$, $n_{e_g}=4$) always gives $\text{CFSE}=0$ regardless of $\Delta_o$ — the reason zinc and closed-shell $d^0$ magnesium are colourless and redox-silent, while partially filled $d^6$ (Fe) and $d^9$ (Cu) centres are colour- and spin-tunable.

Haemoglobin's function reduces to one CFT question: which spin state does $\text{Fe}^{2+}$ ($d^6$) adopt in five- versus six-coordinate geometry, and why does that state fix whether iron sits in or out of the porphyrin plane. This matters because the answer is the mechanical basis of cooperative oxygen binding, the single most-tested idea in this section.

Problem: Using CFSE, show quantitatively why $\text{Zn}^{2+}$ in carbonic anhydrase is colourless with zero CFSE, while $\text{Cu}^{2+}$ in blue copper proteins is strongly coloured, and connect this to their different biological roles.

Step 1. Write $\text{Zn}^{2+}$'s configuration: $[\text{Ar}]3d^{10}$, so $n_{t_{2g}} = 6$ and $n_{e_g} = 4$ — every orbital slot filled.

Step 2. Substitute into the CFSE formula: $\text{CFSE} = (-0.4\times 6 + 0.6\times 4)\Delta_o = (-2.4 + 2.4)\Delta_o = 0$. With all ten d-orbital slots occupied, no d–d transition is possible at any $\Delta_o$, so zinc complexes are colourless and zinc's biological role is purely Lewis-acid catalysis — polarising a bound water molecule, not any redox or spectroscopic property.

Step 3. Contrast with $\text{Cu}^{2+}$: $[\text{Ar}]3d^9$, a partially filled, asymmetrically occupied subshell. A low-energy visible-region transition is now possible (d–d or ligand-to-metal charge transfer from a cysteine thiolate), giving blue copper proteins their colour. More importantly, $\text{Cu}^{2+}$ sits just one electron from the closed-shell $d^{10}$ of $\text{Cu}^+$, so interconversion between the two costs little reorganisation energy.

Step 4. Read off the biological consequence. Zinc's filled $d^{10}$ shell (CFSE = 0) makes it colourless and redox-silent, restricting it to Lewis-acid catalysis, while copper's partially filled $d^9$ shell makes it coloured and gives it a low-barrier one-electron shuttle between $\text{Cu}^{2+}$ and $\text{Cu}^{+}$ — the same CFT contrast explains why these two metals serve entirely different roles.

• Assuming every metalloprotein's colour comes from d–d transitions. Chlorophyll is coloured because of macrocycle $\pi \to \pi^*$ transitions — Mg is $d^0$, so no d–d transition is even possible; always check the d-electron count first.
• Treating haemoglobin's coordination sphere as fixed. It genuinely switches between five-coordinate (deoxy, high-spin) and six-coordinate (oxy, low-spin) — a fixed picture misses the cooperativity mechanism entirely.
• Calling vitamin $B_{12}$'s ring a porphyrin. The corrin ring is related but smaller and less conjugated, and $B_{12}$'s chemistry is defined by its unusual Co–C bond, not by ring type.
• Assuming all first-row metal cofactors are redox-active. $\text{Zn}^{2+}$ ($d^{10}$, CFSE $=0$) is redox-silent; do not default to "metal in biology means electron transfer."

Ask the tutor: want a closer look at how the entatic state tunes copper's redox potential, or a full walk-through of haemoglobin's Bohr effect?