I was thinking about obtaining mercury from mercury(II) nitrate. Is it possible? I've tried putting chunks of Fe into it, I've got the logic from the voltaic table (where Fe is on the left side of Hg, so I think it would get it out of the compound). Can anyone please correct me and show me the right way? I've tried searching on Google, but found nothing.

I'm just a rookie in chemistry, and was interested in it when my teacher announced about a talent show which I'll perform some cool chemistry there (I'm going to perform using fires, I've tried the colored fires and the alcohol carboy). I'm also planning to make a mirror, if anybody has some ideas please comment below.

  • 1
    $\begingroup$ I personally don't think this question deserves downvotes. There was some sloppy terminology, but it's more or less fixed now. Besides, OP demonstrated knowledge of reactivity series and decided to consult with a broader audience before conducting a potentially dangerous experiment. This IMO deserves some respect. $\endgroup$
    – andselisk
    Nov 11 '19 at 16:58
  • 1
    $\begingroup$ I think it would be way safer for you to to make a mirror using the silver mirror test for aldehydes. $\endgroup$
    – Waylander
    Nov 11 '19 at 18:06

The process you described would be more appropriately called "reduction of mercury(II) to elemental mercury". Unfortunately, the trick with iron likely won't work (something more inert like copper would be a better choice though).

Mercury(II) oxide is weakly basic, so mercury salts in general would easily undergo hydrolysis and form basic oxosalts in aqueous solution unless acidified. Mercury(II) nitrate quickly hydrolyses to poorly soluble yellow mercury(II) oxide upon dilution or addition of alkali:

$$\ce{Hg(NO3)2(s) + H2O(l) -> HgO(s) + 2 HNO3(aq)}\label{rxn:R1}\tag{R1}$$

which can be converted back to nitrate by adding excessive amount of nitric acid, which in turn won't leave any chance for iron not to be oxidized, so you end up with both metals in solution.

I didn't want to dive deep into the discussion of how mercury(II) salts hydrolyze and leave it simple, but after receiving critique from Maurice claiming $\ce{Hg(OH)NO3}$ to be the “real” product, I think I allow myself to add a paragraph or two. Raman spectroscopy and x-ray scattering studies in late 1960 demonstrated that hydrolysis of mercury(II) salts produces series of polynuclear oxo-bridged species of types $\ce{Hg2OH(H2O)2^3+},$ $\ce{Hg3O(H2O)3^4+}$ or $\ce{Hg4O(OH)(H2O)3^5+}$ [1, 2]. Formation of mercury(II) hydroxide nitrate $\ce{Hg(OH)NO3}$ as a hydrolysis product was taught in 1940–1950s era, and it stuck in several textbooks published later probably because it's listed in all editions of Pauling's General chemistry up to 1988. However, not only it's oversimplified (granted, reaction \eqref{rxn:R1} is also an oversimplification in a sense it's a boundary case), but is also an incorrect concept.

The most up-to-date summary of what is really happening when $\ce{Hg(NO3)2}$ is dissolved in water can be found in Mercury handbook [3, p. 115]:

$\ce{Hg(NO3)2}$ solutions are only stable in the presence of a certain amount of nitric acid, which prevents hydrolysis. $\ce{Hg(NO3)2}$ quickly hydrolyzes in excess water and produces a precipitate of $\ce{Hg3O2(NO3)2 · H2O}$ or, when boiled in dilute solutions, forms mercury(II) oxide $(\ce{HgO}).$

Regarding the reduction of mercury(II), there are two pathways: $\ce{Hg^0}$ may be obtained from mercury(II) nitrate using a dry or a wet method.

Speaking of a dry method, the most straightforward way to obtain metallic mercury from nitrate is to heat $\ce{Hg(NO3)}$ in a distillation apparatus (b.p. of mercury is 357 °C). Above $\pu{400 °C}$ it readily decomposes:

$$\ce{Hg(NO3)2(l) ->[\pu{400 °C}] Hg(g) + 2 NO2(g) + O2(g)}\label{rxn:R2}\tag{R2}$$

\eqref{rxn:R2} is a brutto-reaction; the nitrate first decomposes to the red mercury(II) oxide at lower temperatures (which, in turn, decomposes to the elements):

$$\ce{2 Hg(NO3)2(l) ->[\pu{350 °C}] 2 HgO(s) + 4 NO2(g) + O2(g)}\tag{R3}$$

Wet method suggest mild conditions and a reaction in solution. For example, formic acid (also used in silver refining) being a strong reducing agent would cause a precipitate from ammoniacal mercury(II) nitrate solution.

Since nitrogen dioxide, mercury vapors as well as mercury salts and oxides are highly toxic, the reactions must be carried in a fume hood which makes it a poor fit for a talent show. Considering the possible risks and your level or preparation (no offence), I'd strongly advise to take extreme caution doing mercury chemistry and avoid public demonstrations until you become more experienced.

Note: chemical reactions are adopted from [4, p. 310–312]


  1. Cooney, R.; Hall, J. Raman Spectra of Mercury(II) Nitrate in Aqueous Solution and as the Crystalline Hydrate. Aust. J. Chem. 1969, 22 (2), 337. https://doi.org/10/b6t3h2.
  2. Johansson, G.; Haugsten, K.; Rasmussen, S. E.; Svensson, S.; Koskikallio, J.; Kachi, S. An X-Ray Investigation of the Hydrolysis Products of Mercury(II) in Solution. Acta Chem. Scand. 1971, 25, 2787–2798. https://doi.org/10/bn5j2g. (PDF)
  3. Kozin, L. F.; Hansen, S.; Kit, M. Mercury Handbook: Chemistry, Applications and Environmental Impact; RSC Publ: Cambridge, 2013. ISBN 978-1-84973-409-7.
  4. R. A. Lidin, V. A. Molochko, and L. L. Andreeva, Reactivity of Inorganic Substances, 3rd ed.; Khimia: Moscow, 2000. (in Russian)
  • 1
    $\begingroup$ I'd not do it directly. If I remember correctly (and reading wikipedia + basic inorg. chem) the mercuric nitrate can be turned into mercurous nitrate treating it with dilute nitric acid. The mercurous nitrate can then be exposed to UV light and yield elemental mercury and mercuric nitrate. In either case: Making a "mirror" with mercury would require very low temperatures, since it just melts and runs off or it will amalgamate. Then there is as the facts on the ground - the compounds involved are very toxic, or hazardous or both... OP: Find a silver solution instead... $\endgroup$ Nov 11 '19 at 20:53
  • $\begingroup$ @StianYttervik Yes, sort of, but both conversion to mercury(I) nitrate and photodecomposition are reversible reactions, so someone should take care of eliminating mercury metal from the system. I agree that pyrolysis of mercury(II) nitrate is an unpleasant and dangerous experiment, I just listed it as the most straightforward one (I added a note on that, too). And I completely agree about your suggestion for OP to look at silver, thank you! $\endgroup$
    – andselisk
    Nov 12 '19 at 8:20
  • $\begingroup$ Thanks everyone, and also thanks to andselisk. Thanks for warning me about that, I am making the mirror with silver. But when I came across the mercury nitrate, I was so curious on obtaining the mercury. Since we've just learned about electrochemistry, I'm just wondering if the voltaic table stuff really works. But it turns out no. Thank you so much andselisk and everyone, not only that I know I was wrong, I learnt a whole lot of new stuff! I'll reconsider heating the mercury nitrate if it's dangerous, once again thanks! $\endgroup$
    – Godlixe
    Nov 12 '19 at 22:54

Here is a possibly interesting safe path to try for a Zinc (or Aluminum) Hg amalgam, which can be used in organic synthesis (e.g., for the Clemmensen reduction) courtesy of the field of Hydrometallurgy. It starts by first placing a sheet of, say, pure Zinc in HCl to create some surface absorbed H atoms (or Al in NaOH, where purity is important to avoid toxic hydrogenated gases) in a fume hood. Remove and rinse with distilled water. Add the treated sheet of metal to a dilute solution of $\ce{Hg(NO3)2}$. This is expected to result in a product of Mercury(I) nitrate in dilute nitric acid. Reaction mechanics:

$\ce{H•(surface)⇌e-(aq) + H+(aq)}$

$\ce{Hg(2+)(aq) +e-(aq)->Hg(+)}$

$\ce{H+ + NO3- = HNO3}$

Apparently, the H• radical functional behaves (per its seemingly reversible formation reaction: e- + H+ = H• ) as apparently an (e-, H+) operating pair. See this source, 'Hydrometallurgy 2008: Proceedings of the Sixth International Symposium', p. 818, which cites a commercial leaching process, to quote Equation (5):

" PbS + 2 H• = Pb + H2S (5) "

which one could view functionally as:

$$ Pb(+2)S(2-) + 2 (e-, H+) = Pb + H2S (g) $$ Further, per Wikipedia on Mercury(I) nitrate the action of light (or less safe, boiling) results in the liberation of mercury, to quote:

"If the solution is boiled or exposed to light, mercury(I) nitrate undergoes a disproportionation reaction yielding elemental mercury and mercury(II) nitrate:[2]

$$ Hg2(NO3)2 ⇌ Hg + Hg(NO3)2 " $$ where the Hg(NO3)2 can be further attacked by hydrogen atoms.

Note, if an amalgam is not formed, one just has successfully removed Mercury from the nitrate.


Contrary to Andselisk's opinion, mercury nitrate $\ce{Hg(NO3)2}$ is not decomposed into $\ce{HgO}$ in water. Mercury nitrate is hydrolyzed into a basic salt, whose formula is $\ce{Hg(OH)NO3}$, and the equation is:

$$\ce{Hg(NO3)2 + H2O -> HNO3 + Hg(OH)NO3}$$

To get pure mercury from mercuric nitrate, Andselisk recommends heating their substance to 400 °C. This technique will produce huge clouds of nitrogen dioxide which is brown, toxic and corrosive. An easy way of avoiding this drawback is to mix thoroughly mercury nitrate $\ce{Hg(NO3)2}$ with an equal mass of sodium carbonate $\ce{Na2CO3}$, and heat this mixture. It will produce carbon dioxide $\ce{CO2}$ instead of nitrogen dioxide $\ce{NO2}$. The equation of the reaction is:

$$\ce{2 Hg(NO3)2 + 2 Na2CO3 -> 4 NaNO3 + 2 Hg + 2 CO2 + O2}$$

It should be mentioned that, above 380 °C, $\ce{NaNO3}$ will be slowly decomposed. So it is not recommended to heat to such a high temperature. The reaction starts at 150 °C. This operation should be done in a retort. The mercury will condense in the spout of the retort, and may be recovered. But mind! Mercury is toxic. And all operations with mercuric compounds are forbidden by the law.

  • $\begingroup$ Regarding the products of hydrolysis, see my updated answer. As for the reaction of nitrate with sodium carbonate, it would be nice to see a reference. I'm also wondering why a retort is required as I have never been in an operating chem lab with a retort would be used. "all operations with mercuric compounds are forbidden by the law" – are you sure about that? $\endgroup$
    – andselisk
    Nov 12 '19 at 8:12

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Not the answer you're looking for? Browse other questions tagged or ask your own question.