First apologize for this kind silly question.One of my student ask me this question but I was unable to get a perfect answer from internet.Generally dilute HCl and impure zinc are using for hydrogen production in laboratory.My question is why do reaction stoped after some little time reaction if we use pure zinc and concentrated HCl?

Is it for coverup zinc by zinc sulphate? It will be thankful for any kind help.

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    $\begingroup$ Your title has little to do with your question. Make up your mind. We don't use concentrated HCl because the reaction would be too vigorous. As for why the reaction stops: because we make it stop. $\endgroup$ Dec 4, 2020 at 18:53
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    $\begingroup$ Sorry, what? zinc sulphide ?!? Where is that coming from? $\endgroup$
    – Karl
    Dec 4, 2020 at 18:58
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    $\begingroup$ As in, heats up and ignites the hydrogen - causing big exposions $\endgroup$
    – Gwyn
    Dec 4, 2020 at 20:07
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    $\begingroup$ And tries to get out of the reaction vessel $\endgroup$
    – Waylander
    Dec 4, 2020 at 20:12
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    $\begingroup$ Up vote: This is actually a very interesting question, NOT well addressed per my research, that involves, I suspect, some good understanding in various areas including electrochemistry and I even have a reference from a 2005 article in Physical Chemistry Chemical Physics Journal. Or, one can be content with Maurice's 'answer', which appears to be popular. $\endgroup$
    – AJKOER
    Dec 5, 2020 at 17:38

2 Answers 2


If metallic zinc $\ce{Zn}$ is in contact with concentrated $\ce{HCl}$ solution, $\ce{H2}$ is produced, but the reaction is exothermic : the solution gets hot. As $\ce{HCl}$ is not so highly soluble in hot water, a fraction of the dissolved $\ce{HCl}$ will be vaporized. So the gas produced will be a mixture of $\ce{H2}$ and $\ce{HCl}$, which is not wanted.

Anyhow, the reaction between pure $\ce{Zn}$ and pure concentrated $\ce{HCl}$ is not really rapid. And, after some time, when the solution contains a certain amount of $\ce{Zn^{2+}}$, the reaction rate decreases in such a manner that the reaction looks finished. It is rather surprising, and even deceiving.

However, if some metallic impurities are present in solution, like $\ce{Co^{2+}}$ or $\ce{Cu^{2+}}$, the zinc metal reduces these ions according to $$\ce{Zn + Co^{2+} -> Co + Zn^{2+}}$$or $$\ce{Zn + Cu^{2+} -> Cu + Zn^{2+}}$$ As a result, some small amounts of metallic cobalt or copper are deposited on the zinc metal. The contact of the two metals $\ce{Zn + Co}$ or $\ce{Cu}$ produces a galvanic cell, and the reaction rate increases a lot. $\ce{Zn}$ gets dissolved much quicker, producing electrons that go onto the copper or cobalt spots, where they reduce $\ce{H+}$ to produce a lot of $\ce{H2}$.

So the best way of producing H2 by the reaction $\ce{Zn + HCl}$ is to use not too concentrated solutions of $\ce{HCl}$ solution, with small amounts of copper or cobalt salts added into the acidic solution.

The same phenomena of galvanic cell happens if $\ce{HCl}$ is replaced by $\ce{H2SO4}$

  • $\begingroup$ HCl is still highly soluble in hot water, as about 20% HCl forms azeotrop with maximum boiling point approx 120 Deg C. So gas above heated diluted HCl has decreased the relative content of HCl, but it is relatively enriched for concentrated HCl. $\endgroup$
    – Poutnik
    Dec 5, 2020 at 6:33
  • $\begingroup$ Good anecdotal recount, however, adding a source reference is my recommendation (and then, you may be able to actually correctly account for some of the chemistry, or just read my answer). $\endgroup$
    – AJKOER
    Dec 5, 2020 at 16:10

As to why it is better not to employ concentrated Hydrochloric acid is likely due to the exothermic nature of its reaction with Zinc metal and the volatility of $\ce{HCl}$ itself fostering a possible loss of Hydrogen Chloride.

Further, the very presence of water may be beneficial (see Hydrogen formation in the reaction of Zn+(H2O)n with HCl) in allowing a Zinc ion to be moved into a $\ce{[Zn(H2O)6](2+)}$ hydration sphere which can even further interact in a ligand exchange reaction. Relatedly, to expound per Libre Text on the reaction:

If you add concentrated hydrochloric acid to a solution containing hexaaquacobalt(II) ions (for example, cobalt(II) chloride solution), the solution turns from its original pink color to a dark rich blue. The six aqua molecules are replaced by four chloro ions. The reaction taking place is reversible.

$\ce{[Co(H2O)6](2+) + 4 Cl- <=> [Co(H2O)_6Cl_4](2-) + 6 H2O }$

And further on the above reaction, to continue quoting:

Concentrated hydrochloric acid is used as the source of chloride ions because it provides a very high concentration compared to what is possible with, say, sodium chloride solution. Concentrated hydrochloric acid has a chloride ion concentration of approximately 10 mol dm-3. The high chloride ion concentration pushes the position of the equilibrium to the right according to Le Chatelier's Principle.

where I would expect a similar reaction sequence (as reported in the cited 2005 article from the Journal of Physical Chemistry Chemical Physics) with copper in place of cobalt.

Also, a cited improvement in reaction rate, relating to the use of impure zinc, is perhaps best illustrated in an related experiment where the acid is, albeit, $\ce{H2SO4}$ acting on Zinc (as a solid piece of $\ce{Zn}$ metal) in the presence of a source of a copper impurity. Note, the experiment employs three test tubes containing $\ce{Zn}$ + Acid, but one without any Copper presence, one with low surface area Copper turnings, and the 3rd with aqueous $\ce{CuSO4}$.

Surprisingly to some, it is claimed that the 3rd test tube may actually be the largest source of Hydrogen gas! Interestingly, in the 3rd test tube, there is no starting copper metal. However, there is cited a displacement formation of new Cu (black in color) metal by Zinc interacting with the cupric ions of $\ce{CuSO4}$. This newly formed black Copper, however, has a decidedly high surface area compared to both the piece of Zinc metal and the Copper turnings.

As such, my cited explanation of the accelerated reaction is based on the electrochemical cell formed with an anode of Zn metal, a high area cathode of Copper metal, all in an electrolyte of copper ions. The anodic corrosion of the zinc is now observably accelerated due to the now very favorable ratio of the low surface area zinc anode to the high surface area black Copper cathode.

Further, electrochemical reactions, in general as long as there is some reagent concentration presence, are NOT driven by the relative concentration considerations as occurs with standard chemical reactions (for the current reaction, see graph of reaction rate here depicting a flatting).

So apparently, there are perhaps several reasons including an electrochemical underpinning, as to why to preferentially employ dilute $\ce{HCl}$ along with impure $\ce{Zn}$ metal.


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