Since 2003, I have been volunteering at the Royal University of Phnom Penh in Cambodia, both during my summer vacations and also during part of my two fellowship leaves. My students are chemistry majors at the university who attend the “hands-on” demonstration workshops. Many who attend my workshops will become chemistry teachers in government schools, and the demand for a place is so high that a lottery is held. The chemistry department provides me with two volunteer assistants each year, and with their help, I have presented hands-on demonstration workshops in three Regional Teacher Training Centers to more than 110 teacher trainees. The response has been very enthusiastic with requests for many more workshops coming from the trainees as well as the regional trainers.
These trainee-teachers do not want demos that employ expensive reagents nor do they want “flashy” demos which have no underlying chemistry. Therefore, my focus has been on using locally available no-cost/low-cost materials for teaching chemistry in schools that are often without running water, electricity, science labs or supplies.
In this series, to which I hope others will contribute, I will introduce some experiments that I have found to be feasible even in the most remote areas of a developing country. What interests me particularly is that one demo can serve several roles. I described one such case in an article some years ago.1 The demo that I will discuss here is another example.
DEMO.... Aluminum loses to copper2
Aluminum foil and copper(II) sulfate are two readily-available items in Cambodian markets.3 For this demo, copper(II) chloride works better, if available; however, if copper(II) sulfate is used, then a spoonful of table salt must be added.
A loosely crumpled aluminum foil ball (15 cm × 30 cm) is put in approximately 300 mL of 1 mol L−1 copper(II) chloride solution (or the same concentration of copper(II) sulfate with a spoonful of table salt). Within a few minutes, one sees copper particles beginning to form on the aluminum ball, bubbling is visible and the reaction container becomes very hot. Within an hour, large amounts of copper are noticed.
First of all, this reaction can be used to very vividly illustrate a single replacement (displacement) reaction and to provide a good example of balancing a chemical equation:
2Al(s) + 3CuCl2(aq) → 2AlCl3(aq) + 3Cu(s)
The reaction is an illustration of the activity series, and as an extension, the net ionic equation can be derived:
2Al(s) + 3Cu2+(aq) → 2Al3+(aq) + 3Cu(s)
And, from the net ionic equation, the redox concept can be developed, that the reaction can be represented as two half-reactions involving reduction and oxidation:
Al(s) → Al3+(aq) + 3e−
Cu2+(aq) + 2e−(aq) → Cu(s)
As shown in the photo below, another extension is to use the reaction to illustrate the principles of limiting reactant by performing parallel reactions, one with an excess of copper(II) ion and the other, an excess of aluminum metal.
More advanced explanations
For a more advanced background in chemistry, one can quantify the redox process by comparing the reduction potentials of the two metals:
Al3+(aq) + 3e− → Al(s) Eo = −1.66 V
Cu2+(aq) + 2e− → Cu(s) Eo = +0.34 V
Al(s) → Al3+(aq) + 3e− Eo = +1.66 V
Cu2+(aq) + 2e− → Cu(s) Eo = +0.34 V
The sum being:
2Al(s) + 3Cu2+(aq) → 2Al3+(aq) + 3Cu(s) Eo = +2.00 V
A spontaneous reaction.
Why does the beaker become hot? Thermochemistry has the answer. We can write the two half-reactions in terms of enthalpy of formation values:
Al(s) → Al3+(aq) + 3e− ΔfHo = −538 kJ mol−1
Cu(s) → Cu2+(aq) + 3e− ΔfHo = +65 kJ mol−1
The real driving force for the reaction, then, is the very exothermic formation of the aquated aluminum ion.
At first thought this seems unexpected, considering that in the gas phase to triply-ionize an aluminum atom is incredibly endothermic. The answer is that in aqueous solution, the tiny high-charge aluminum ion exists with six water molecules around it, [Al(OH2)6]3+(aq). The partially negative (δ−) oxygen atoms of each of the six water molecules are attracted extremely strongly to the 3+ aluminum ion (an opportunity to introduce polarity of covalent bonds and ion-dipole interactions). Thus the aqueous ion formation is highly exothermic, and, as a result, the overall reaction is exothermic.
Following from the exothermicity, as any entropy change is small (with no gases being formed or consumed), the reaction has a negative (spontaneous) free energy change and a corresponding positive (spontaneous) net Eo (as shown above).
The chloride ion mystery
As an extension of this experiment, one can show how there is no reaction of aluminum with copper(II) sulfate until chloride ion is added to the solution. Thus the chloride ion must be playing an active role in the process as a catalyst. So, superficially, one can leave it at that.
However, according to collision theory, chemical reactions usually occur when two species interact at a time (for a few reactions, three-body collisions are necessary). In the overall ionic equation written above, two aluminum atoms and three copper(II) ions react together. Thus there must be a series of reaction steps occurring on the aluminum surface.
Though surface reactions are often complex, the necessity of chloride ions seems to provide a clue. Several hypotheses have been proposed, most of which do not seem very likely.4 One explanation, not previously considered, relates to the reaction process in which a copper(II) ion would adhere to the aluminum metal surface. The first probable step would be the reduction of copper(II) initially to copper(I) on the aluminum. But copper(I) compounds are very insoluble, and a layer of copper(I) oxide or sulfate could possibly form over the surface of the aluminum, precluding further reaction. However, in the presence of chloride ion, the soluble dichloridocuprate(I) ion is formed, [CuCl2]−(aq), enabling reduction to continue stepwise through to the copper(0) metal.
The bubble problem
As noted above, small amounts of gas are produced. The explanation is that the copper(II) ion solution is slightly acidic and thus there is a small proportion of oxidation of aluminum by the hydrogen ions:
2Al(s) + 6H+(aq) → 2Al3+(aq) + 3H2(g)
This is a good explanation — as far as it goes. Teachers would probably leave the explanation at this point. But there can always be the very curious student who demands to know why/how the copper(II) ion is acidic. The answer comes back to the fact that the aquated copper(II) ion is also surrounded by water molecules — but much more weakly than those around the aluminum ion. A hydrogen ion (or more correctly, a hydronium ion) can be lost in an equilibrium that lies much to the left, but that produces enough hydronium ion to cause a small proportion of the reaction:
[Cu(OH2)6]2+(aq) + H2O(l) ⇌ [Cu(OH2)5(OH)]+(aq) + H3O+(aq)
This reaction, involving readily available materials, can be used as a focus for a whole range of chemistry activities. However, to delve deep into the chemistry, it is not as simple as one would think! Nevertheless, it is wise for the chemistry teacher to have the knowledge base beyond that of the essentials, prepared for those extra-inquisitive students who want to understand the how and the why.
PART 2 has two demos using aluminum cans, sodium hydroxide, copper(II) sulfate and table salt, all of which are commonly available at the market.
Geoff Rayner-Canham, Grenfell Campus, Memorial University, Corner Brook NL, Canada, is thanked for working with me to develop the conceptual background to, and relevance of, the demonstrations.
- M. Hauben and G. Rayner-Canham, “Reaction in a Bag: Explanations for an Endothermic/Exothermic Gas-Producing Demonstration,” Chem 13 News, pages 14-15 (February 1996).
- Copper(II) sulfate has a variety of uses, the most common being as a fungicide and algaecide in farming. See: https://en.wikipedia.org/wiki/Copper(II)_sulfate#As_a_herbicide.2C_fungicide_and_pesticide
- One commonly cited explanation is that the chloride ion reacts with the surface layer of aluminum oxide over the aluminum metal to give the tetrachloridoaluminate ion, [AlCl4]−(aq) ion, revealing the reactive metal surface for reduction of copper(II) ions. However, this seems unlikely, partially because the chloride ion concentration does not have to be particularly high for the reaction to occur, yet formation of the complex requires high concentration of the chloride ion. More importantly, it goes against common sense, for aluminum-hulled boats seem to be quite unreactive in seawater.