The objective of this activity is to attempt to interpret observations madeof the reaction between zinc metal and dilute sulfuric acid, with and withoutthe addition of a copper catalyst.
The reaction between reactive metals, such as zinc, and dilute acids, iswell-known. The rate of such a reaction can be assessed by the rateat which bubbles of hydrogen gas are produced.
Zn(s) + 2H1+(aq) ----------> Zn2+ (aq) + H2 (g)
The choice of acids may also affect the rate of the reaction. In thisinvestigation, it is recommended that 1.0 M sulfuric acid (not hydrochloricacid) be used.
If two small, equal pieces of granulated zinc are placed into small tubesand covered with 1 to 2 mL of dilute (1.0 M) sulfuric acid, bubbles shouldbe seen, although they may not come quickly, especially at first when theacid is added to the zinc. As soon as a fine stream of bubbles is visiblein both tubes, add to one of them either some small pieces of copper metal,or about 1.0mL of 0.5 M copper (II) sulfate solution. Does the additionof the copper make any difference to the rate at which bubbles are produced?
If copper (II) sulfate solution is used, it should be noted that the bluecolor disappears fairly quickly, and the zinc darkens in color. Thisis due to a displacement reaction, in which copper metal is precipitatedonto the zinc: Zn(s) + Cu2+(aq)-----------> Zn2+ (aq) + Cu(s)
If there is indeed a difference in the rate at which bubbles are produced,then it may be possible to explain how the copper catalyst works.
Predicting a possibility.
If bubbles are produced faster from zinc and dilute sulfuric acid in thepresence of copper metal, then it must be the copper metal that makes thedifference. If another two small equal pieces of zinc are placed infresh 1.0 M sulfuric acid, with one having a large enough strand of copperwire twisted around it, it may just be possible to see a difference in theway that hydrogen is produce. Zinc and copper together in an electrolyte(sulfuric acid) make an electrolytic cell. Zinc, the more reactivemetal, is the anode, copper is the cathode, so hydrogen ions are likely tobe reduced to hydrogen gas at the copper electrode.
Without the copper, there is no obvious electrolytic cell, so hydrogen gaswould have to be produced by the direct reaction between zinc atoms and hydrogenions at the surface of the metal.
Testing the possibility
Test the predictions made. It should be possible to see that with zincand sulfuric acid alone, bubbles come from the surface of the zinc, but whencopper wire is twisted around the zinc, (with a loose end of wire hangingfree), hydrogen gas appears at the copper wire.
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, ZnCl42- .
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.
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.
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.
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. Text books say that no hydrogen is produced from copper and dilute sulfuric acid. Is there any blueness forming near the copper?
The following comment is by Jeff Hughes. Mike Clark's repsonse follows.
I read with interest your demonstration 'Electrochemical catalyst'. I haveused a similar demonstration with my classes for some years (after readingabout the demonstration in Journal of Chemical Education). HoweverI would query your explanation of why the rate of evolution increases. WhenI 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 demonstrationworks even better with a Pt wire. The explanation I give is that theoverpotential of hydrogen on copper is less than that on zinc. The overpotentialof H2 on Zn is about 0.72 volts (which nearly cancels out the standard potentialof Zn/Zn2+ - hence this is why you do not see much evolution when the zincis 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 inoverpotentials 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 hydridedepends on the metal. Thus I would maintain that there is no involvementof the Cu/Cu2+ couple in the reaction, as evidenced by the greater ease ofevolution with Pt, although it has a much higher standard potential.
Applied Chemistry Department
Thanks for your comments on the zinc/copper/acid demonstration. I take yourpoints. My interest in the demonstration is at secondary school level, whereI use it as a means of demonstrating the involvement of a catalyst in a chemicalreaction, in opposition to the idea that beginning students may get thata catalyst is "just there" and "somehow" alters the rate of a reaction. Ihave never tried to introduce the idea of overvoltages at that level.