The reaction between reactive metals, such as zinc, and dilute acids, is well known.  The rate of such a reaction can be assessed by the rate at which bubbles of hydrogen gas are produced.
 Zn_(s)  + 2H¹⁺ _(aq) ----------> Zn²⁺ _(aq)  +  H_{2 (g)}

The choice of acids may also affect the rate of the reaction.  In this investigation, it is recommended that 1.0 M sulfuric acid (not hydrochloric acid) be used.  

1. Hydrochloric acid may react more readily with zinc than sulfuric acid, so the effect of catalysis is less obvious.  The faster reaction is associated with formation of an anionic complex, ZnCl₄^2- .  

2. The effectiveness of the experiment may depend upon purity of reagents.  The reaction between pure zinc and pure 1.0 M H₂SO₄ may be quite slow at room temperature, and the presence of copper makes a very distinct difference.  Certain impurities may obscure the difference  between the samples with and without copper.

3. This experiment is effective with quite small quantities of materials.  In the second (testing prediction) stage, the pieces of zinc with and without copper may be placed side by side in the same acid sample in a small flat dish.

4. More advanced students may wish to test the zinc and copper wire with probes of a voltmeter to show that an electrolytic cell has indeed been created.

5. Some students might consider whether the bubbles are originating from a reaction involving copper.  A piece of the same copper wire, placed near the zinc pieces without touching them, might serve as a useful control.  Textbooks say that no hydrogen is produced from copper and dilute sulfuric acid.  Is there any blueness forming near the copper?

The objective of this activity is to attempt to interpret observations made of the reaction between zinc metal and dilute sulfuric acid, with and without the addition of a copper catalyst.

If bubbles are produced faster from zinc and dilute sulfuric acid in the presence of copper metal, then it must be the copper metal that makes the difference. If another two small equal pieces of zinc are placed in fresh 1.0 M sulfuric acid, with one having a large enough strand of copper wire twisted around it, it may just be possible to see a difference in the way that hydrogen is produce. Zinc and copper together in an electrolyte (sulfuric acid) make an electrolytic cell. Zinc, the more reactive metal, is the anode, copper is the cathode, so hydrogen ions are likely to be reduced to hydrogen gas at the copper electrode.

Without the copper, there is no obvious electrolytic cell, so hydrogen gas would have to be produced by the direct reaction between zinc atoms and hydrogen ions at the surface of the metal.

Test the predictions made. It should be possible to see that with zinc and sulfuric acid alone, bubbles come from the surface of the zinc, but when copper wire is twisted around the zinc, (with a loose end of wire hanging free), hydrogen gas appears at the copper wire.

Mike Clark

The following comment is by Jeff Hughes.  Mike Clark's response follows.  

Mike,
I read with interest your demonstration 'Electrochemical catalyst'. I have used a similar demonstration with my classes for some years (after reading about the demonstration in Journal of Chemical Education).  However, I would query your explanation of why the rate of evolution increases. When I do the demonstration, I touch a copper wire to the surface of the zinc. The hydrogen evolves from the surface of the copper only. The demonstration works even better with a Pt wire. The explanation I give is that the over potential of hydrogen on copper is less than that on zinc. The over potential of H2 on Zn is about 0.72 volts (which nearly cancels out the standard potential of Zn/Zn2+ - hence this is why you do not see much evolution when the zinc is added to the acid) but H2 on Cu is .48V (at 1milliamp/cm2 current density) and H2 on Pt is close to zero.  The reason for the differences in over potentials is related to the ease of formation and evolution of H2 bubbles, which occurs in several steps:- (i) diffusion of H+ to the metal surface;  (ii) H+ + e => H;   (iii) formation of a layer of M-H i.e. H atoms adsorbed on the metal  surface;  (iv) H + H => H2;  (v) formation and evolution of a bubble from the surface

The metal is involved in step (iii) and the ease of formation of the hydride depends on the metal. Thus, I would maintain that there is no involvement of the Cu/Cu2+ couple in the reaction, as evidenced by the greater ease of evolution with Pt, although it has a much higher standard potential.

Regards,
Jeff Hughes
Applied Chemistry Department
RMIT

Dear Jeff,
Thanks for your comments on the zinc/copper/acid demonstration. I take your points. My interest in the demonstration is at secondary school level, where I use it as a means of demonstrating the involvement of a catalyst in a chemical reaction, in opposition to the idea that beginning students may get that a catalyst is "just there" and "somehow" alters the rate of a reaction. I have never tried to introduce the idea of over voltages at that level.
Regards,
Mike Clark