Aetiologically iron-associated metabolic acidosis can be grouped together with cyanide, metformin, propofol and all the other mitochondrial toxins which cause lactic acidosis.
Iron attacks the mitochondria by lipid peroxidation of their delicate membranes; it also appears to interfere with the enzymes of the electron transport chain and with the components of Krebs cycle, all of which decreases the amount of pyruvate processed aerobically.
Now, iron is usually mentioned as an important cause of metabolic acidosis, and there is a warm spot reserved for it in the “MUDPILES” mnemonic. An impressionable person might be inclined to believe that iron contributes to the high anion gap metabolic acidosis by dissociating into unmeasured anions, much like the toxic alcohols. However, that would be wildly inaccurate, because iron is a cation.
The acidosis here is multifactorial. Some textbooks (Fowler’s Handbook on the Toxicology of Metals) suggest that the acidosis is mainly due to the physicochemical effects of the iron ion itself. Other sources (Goldfranks Manual of Toxicologic Emergencies) attribute the acidosis to a raised lactate, of which not all is generated by direct effects of the iron, but rather due to the fluid loss (from an ulcerated gut), cardiogenic shock (due to the myocardial mitochondrial toxicity) and fulminant hepatic failure. On top of that, a fair portion of the lactic acidosis is due to the direct mitochondrial toxicity of iron in all tissues.
So, how much iron does it take?
The molar mass of iron is 55.8g per mole; thus 1 mmol of iron weighs 0.0558g, or 55.8mg.
Most iron replacement tablets contain between 1 and 2 mmol of elemental iron.
The discerning suicide victim typically takes at least 60mg/kg of iron, which ends up being about 75mmol of elemental iron for a 70kg person. Generally, that is enough to cause all sorts of life-threatening problems; the serum iron concentration at which one might consider taking action is about 90 micromol/L. Adventurous toddlers have been known to survive serum concentrations as high as 2992 micromoles/L, or around 2 mmol/L.
90 micromoles per litre is a very low concentration in comparison to the other ionic species in the bloodstream. However, the capacity of blood to bind iron (by saturating transferrin) is even lower. The glorious oracle of Wikipedia reports that the iron-binding capacity of total blood is 45-66 micromoles/L. Thus, if one has 90 micromoles per litre of iron, about 30 micromoles in every litre are free ions.
Now, this free ferrous (divalent) iron does not persist in its divalent state. It converts rapidly to a ferric (trivalent) state, and then reacts with body water to produce ferric hydroxide:
Fe3+ +3(H2O) = Fe(OH)3 + 3(H+)
Thus, for every 30 micromoles of iron ions there are 90 micromoles of hydrogen ions generated, which the body fluid buffers need to titrate. Or rather, this is the classical view of the acid-base disturbance.
One must remember that the ferric hydroxide is actually rather insoluble at body fluid pH.
We can observe the acid-base balance of the iron overdose in terms of changes in strong ion concentrations. Divalent ferrous (Fe II) iron is a strong fully-dissociated cation; when it enters the trivalent oxidation state and combines with water, it becomes an insoluble ferric hydroxide, and is essentially removed from the equation.
Thus, the strong cation becomes a weak cation; the conversion of iron from Fe(II) into Fe(III) decreases the strong ion difference.