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Having conducted an investigation involving multiple metals, I have found that nickel refuses to react with concentrated hydrochloric acid. I have even left the reaction overnight and still have seen no mass change in the solid piece of nickel. Is the reaction simply extremely slow or does it not react despite being above hydrogen in the reactivity series?

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    $\begingroup$ Define "concentrated". $\endgroup$ – Shaka Boom Jul 13 '16 at 18:29
  • $\begingroup$ I think that the nickel passivation is the problem. Not sure if the nickel is reacting with water to form oxides or HCl (to form chlorides), but I would bet that nickel which was ground to powder in an inert atomphere, would react with dry HCl (gas) quickly, and at lower temperature. $\endgroup$ – Ben Welborn Jul 14 '16 at 18:00
  • $\begingroup$ What is the color of HCl now? Does it remain colorless? Solutions of NiCl2 are yellow or green, so you don't have to weight Ni to see if reaction is going. $\endgroup$ – sixtytrees Jul 17 '16 at 21:09
  • $\begingroup$ Aye the HCL remains the same colour. $\endgroup$ – Yazan Jul 18 '16 at 17:07
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Try 30% and heat it to 80C. All of the chloride should be exhausted after a day or less. You can then evaporate the solution to generate crystals of nickel chloride. Any work with nickel compounds should always be done in a well-ventilated hood.

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  • $\begingroup$ Thank you for this. I was using 30% hydrochloric acid I believe and I have done the same with zinc, magnesium, tin, and aluminium. All had immediate noticeable reactions at even 10 degrees celsius. My question is why, despite having a greater reactivity, does it react so slowly? Does it have something to do with a difference in its oxide form? $\endgroup$ – Yazan Jul 14 '16 at 12:39
  • $\begingroup$ One last question: Would a similar result be expected of all other acids or would it react faster with them? $\endgroup$ – Yazan Jul 14 '16 at 15:49
  • $\begingroup$ The acid should be in a well-ventilated hood too. $\endgroup$ – Oscar Lanzi Feb 28 '18 at 11:02
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The reaction is given here.

$$\ce{Ni + 2HCl → NiCl2 + H2}$$

Nickel react with hydrogen chloride to produce nickel(II) chloride and hydrogen. Hydrogen chloride - diluted solution. This reaction takes place slowly.

In the wikipedia page of nickel, the reaction is explained elaborately:-

Nickel(II) chloride is produced by dissolving nickel or its oxide in hydrochloric acid. It is usually encountered as the green hexahydrate, the formula of which is usually written $\ce{NiCl2•6H2O}$. When dissolved in water, this salt forms the metal aquo complex $\ce{[Ni(H2O)6]^2+}$. Dehydration of $\ce{NiCl2•6H2O}$ gives the yellow anhydrous $\ce{NiCl2}$.

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In my PM refining lab I frequently use straight HCl 32% to remove surface nickel from my metallic products (post-first melt) containing predominately Au/Ir. The nickel easily bubbles away into a green solution that eventually turns yellow. I then pour it off into my nickel-saver beaker. If the nickel is extremely susceptible to HCl, the nickel might be in the form of an oxide.

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This is called passivation of metal surface. For a similar reason aluminum doesn't dissolve in water, but Al/Hg amalgam reacts with water. In case of Al Al2O3 film prevents the reaction. In case of Ni you probably form a thin layer of NiCl2 that isn't very soluble in concentrated HCl. Try diluting HCl by adding 2-3 volumes of water. Stirring should help. Fine powder will react faster than big pieces of metal.

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