As the title implies, what is the molecular basis of cyanide toxicity? I did some searching around at the CDC and it only states that it prevents cells from using oxygen. I also read how it could take as little as $131~\mathrm{ppm}$ to kill. How can a simple $\ce{CN-}$ ion do so much damage?

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    $\begingroup$ I think this question is much, much, much more interesting from a chemical point of view. I think it is well on topic here and decide not to close it. I hope someone will find the time to write up a better explanation than what is currently found on [bio.se]. $\endgroup$ Nov 19, 2015 at 4:52

2 Answers 2


Cyanide is a pretty good ligand for coordination compounds. The electron pair on carbon (which, incidentally, also carries the Lewis structure’s formal charge) is located in the HOMO — much like as in $\ce{CO}$, whose molecular orbitals can be found in this answer by Martin (replace oxygen with nitrogen to arrive at $\ce{CN-}$) — making it a good Lewis base and a good σ donor.

It is remarkable for being a strong-field ligand, i.e. one that often creates low-spin complexes even for 3d-metals (which generally prefer high-spin if there are no competing reasons for low-spin). While one could draw analogies to $\ce{CO}$ or even $\ce{NO+}$ whose π-antibonding orbitals are strongly Lewis acidic and participate in metal to ligand backbonding, it turns out that for cyanide the main reason is a high covalency of the metal–ligand-σ-bond due to the relevant orbitals being very similar in energy. Thus, cyanide is unique in that it both stabilises higher oxidation states by being an anion and low-spin complexes due to its binding modalities.

One might think, in analogy to $\ce{K4[Fe(CN)6]}$, that cyanide preferentially binds to haemoglobin’s iron(II) centre and thereby inhibits oxygen transport, but that is in fact not the case as this centre requires a less-oxidised iron(II). Oxygen (due to the resulting redox-mechanism) and carbon monoxide are much better ligands for haemoglobin.

What cyanide does strongly inhibit is cytochrome c oxidase, the last enzyme in the respiratory cycle which collects electrons from four cytochrome c’s and transfers them to an oxygen molecule to create two water molecules:

$$\ce{4 Fe^2+-cytochrome~$c$ + 4 H+ + O2 -> 4 Fe^3+-cytochrome~$c$ + 2 H2O}$$

This is done by a plethora of metal centres consisting of two haems, to cytochromes and two copper centres. I can’t tell you exactly which one, but there will be at least one which needs accessable iron(III) at some point in the catalytic cycle to which $\ce{CN-}$ can bind well and strongly inhibiting the catalysis of the enzyme.

Since the inhibited process is directly responsible for cellular respiration by transferring electrons to molecular oxygen (and being the only enzyme dealing with molecular oxygen in the respiratory chain), inhibition of this protein will cause the breakdown of the entire respiratory chain and lead to cellular suffocation.

Considering that an amount of $200~\mathrm{mg}$ can be toxic according to Wikipedia, and also considering that a human being contains some 10 trillion cells also according to Wikipedia, that’s still some 400 million cyanide ions per cell demonstrating the efficiency of binding and inhibition.

Note that popular antidotes also directly address the fact that cyanide is a strong ligand:

  • One antidote, hydroxycobalamin, presents the cyanide ion with a $\ce{Co^{III}}$ centre and a weakly bound, easily displaced hydroxido ligand. Hydroxide is displaced by cyanide creating cyanidocobalamin, a variant of vitamin B12 that the body has no problems dealing with. It is ejected via the kidneys. Injected into the circulatory system, it scavenges cyanides before they can enter cells.

  • Another antidote, also injected into the circulatory system, is 4-dimethylaminophenol (not to be abbreviated as DMAP to prevent confusion with 4-dimethylaminopyridine). It is an oxidising agent designed to oxidise haemoglobin’s iron(II) to methaemoglobin (containing iron(III)) thereby creating a much better acceptor for cyanide. Since there is a lot of iron(II) (a quick estimate with $8~\mathrm{mmol/l}$ gives $40~\mathrm{mmol}$ haemoglobin in human blood) in the blood for oxygen transport, small amounts of this can be sacrificed to capture cyanide before that reaches the respiratory cycle.


Cyanide is made up of a nitrogen triple bonded to a carbon with a lone pair, [C≡N]−. (Carbon has 4 valence electrons and Nitrogen has 5; pairing will result in a triple bond, a pair of electrons around the nitrogen, and an extra electron pairs with the 4th electron around the carbon). Carbon does not like to have a 5th electron, it is unstable, even the electron charge is shared between the Carbon and Nitrogen. The resulting negative charge (and triple bond) means that cyanide will be very reactive and want to form another bond.

As a result, cyanide is particularly deadly because it will bond with iron, (usually Fe3+ or Fe2+). In our body, iron is found in protein in the electron transport chain. When cyanide inhibits the electron transport chain by binding to iron, the cell can no longer produce energy.

Cyanide is also deadly when it bonds with hydrogen to create hydrogen cyanide, or HCN. HCN can be inhaled, where it will bond with iron, and result in cell death.

  • $\begingroup$ "The resulting negative charge (and triple bond) means that cyanide will be very reactive and want to form another bond." Not really - octet rule is satisfied in CN-. $\endgroup$ Nov 19, 2015 at 8:30
  • $\begingroup$ As ortho said. Even more: cyanide is actually pretty stable. You can store it on the shelf for years. $\endgroup$
    – Jan
    Nov 19, 2015 at 9:13

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