Molecules and Reactions Foiled Again Hypothesis

Preamble

Since 2003, I have been volunteering at the Majestic University of Phnom Penh in Kingdom of cambodia, both during my summer vacations and also during part of my two fellowship leaves. My students are chemical science majors at the university who attend the "hands-on" demonstration workshops. Many who attend my workshops volition become chemistry teachers in government schools, and the need for a place is so high that a lottery is held. The chemical science section provides me with ii volunteer administration each year, and with their help, I take presented easily-on sit-in workshops in three Regional Teacher Preparation Centers to more than 110 teacher trainees. The response has been very enthusiastic with requests for many more workshops coming from the trainees besides equally the regional trainers.

These trainee-teachers do not desire 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-price/depression-cost materials for didactics chemical science in schools that are often without running water, electricity, science labs or supplies.

In this series, to which I hope others will contribute, I volition introduce some experiments that I take institute to be feasible even in the most remote areas of a developing country. What interests me particularly is that one demo tin can serve several roles. I described one such case in an article some years ago.¹ The demo that I volition discuss hither is some other example.

DEMO.... Aluminum loses to copper^two

Aluminum foil and copper(II) sulfate are two readily-bachelor items in Cambodian markets.³ For this demo, copper(II) chloride works better, if available; notwithstanding, if copper(2) sulfate is used, then a spoonful of common salt must be added.

A loosely crumpled aluminum foil ball (15 cm × 30 cm) is put in approximately 300 mL of ane mol L^−ane copper(2) chloride solution (or the same concentration of copper(Ii) sulfate with a spoonful of tabular array salt). Within a few minutes, ane sees copper particles get-go to form on the aluminum ball, bubbling is visible and the reaction container becomes very hot. Within an hr, large amounts of copper are noticed.

First of all, this reaction tin can be used to very vividly illustrate a single replacement (deportation) reaction and to provide a good example of balancing a chemical equation:

2Al(south) + 3CuCl_two(aq) → 2AlCl₃(aq) + 3Cu(s)

The reaction is an illustration of the activity series, and as an extension, the internet ionic equation tin exist derived:

2Al(s) + 3Cu^two+(aq) → 2Alⁱⁱⁱ⁺(aq) + 3Cu(s)

And, from the net ionic equation, the redox concept can be adult, 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, some other 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 metallic.

Photo: the reaction of aluminum metal with copper(Two) ion. Left beaker, with excess of copper(2) ion. Right chalice, with backlog of aluminum metal.

More than avant-garde explanations

For a more advanced background in chemistry, one can quantify the redox process by comparing the reduction potentials of the ii metals:

Al³⁺(aq) + 3e⁻ → Al(s)                                             E^o = −1.66 Five

Cu²⁺(aq) + 2e⁻ → Cu(s)                                           Due east^o = +0.34 V

To give:

Al(s) → Al³⁺(aq) + 3e⁻                                              E^o = +one.66 5

Cu²⁺(aq) + 2e⁻ → Cu(s)                                           E^o = +0.34 Five

The sum beingness:

2Al(s) + 3Cu²⁺(aq) → 2Alⁱⁱⁱ⁺(aq) + 3Cu(s)                         Eastward^o = +2.00 V

A spontaneous reaction.

Why does the beaker get hot? Thermochemistry has the answer. We can write the 2 half-reactions in terms of enthalpy of formation values:

Al(s) → Alⁱⁱⁱ⁺(aq) + 3e⁻             ΔfH^o = −538 kJ mol^−ane

Cu(s) → Cu²⁺(aq) + 3e⁻           ΔfH^o = +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, because that in the gas stage to triply-ionize an aluminum atom is incredibly endothermic. The reply 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 half dozen 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 germination is highly exothermic, and, equally a result, the overall reaction is exothermic.

Following from the exothermicity, equally whatever entropy change is pocket-size (with no gases being formed or consumed), the reaction has a negative (spontaneous) free energy change and a corresponding positive (spontaneous) cyberspace East^o (equally shown above).

The chloride ion mystery

Every bit an extension of this experiment, i tin can show how at that place 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 agile role in the process as a goad. So, superficially, one tin leave information technology at that.

All the same, according to collision theory, chemical reactions usually occur when 2 species interact at a time (for a few reactions, iii-torso collisions are necessary). In the overall ionic equation written in a higher place, two aluminum atoms and three copper(II) ions react together. Thus there must exist 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 footstep would exist 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 maybe form over the surface of the aluminum, precluding farther 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 chimera problem

Every bit 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 pocket-sized proportion of oxidation of aluminum past the hydrogen ions:

2Al(s) + 6H⁺(aq) → 2Al³⁺(aq) + 3H₂(g)

This is a expert caption — as far every bit information technology goes. Teachers would probably get out the explanation at this point. Only there can always exist the very curious student who demands to know why/how the copper(2) ion is acidic. The answer comes dorsum to the fact that the aquated copper(Two) ion is too surrounded by h2o molecules — but much more than weakly than those around the aluminum ion. A hydrogen ion (or more correctly, a hydronium ion) tin 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)_vi]²⁺(aq) + H_twoO(50) ⇌ [Cu(OH_ii)_five(OH)]⁺(aq) + H_threeO⁺(aq)

Commentary

This reaction, involving readily available materials, can exist used equally a focus for a whole range of chemistry activities. Withal, to delve deep into the chemical science, it is non as unproblematic as one would remember! Withal, information technology 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 ii demos using aluminum cans, sodium hydroxide, copper(II) sulfate and tabular array table salt, all of which are commonly available at the market.

Acknowledgements

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.

References

1. M. Hauben and G. Rayner-Canham, "Reaction in a Handbag: Explanations for an Endothermic/Exothermic Gas-Producing Demonstration,"Chem 13 News, pages 14-15 (February 1996).
2. https://world wide web.flinnsci.com/media/622135/95000.pdf
3. Copper(II) sulfate has a variety of uses, the most common existence every bit a fungicide and algicide in farming. See: https://en.wikipedia.org/wiki/Copper(II)_sulfate#As_a_herbicide.2C_fungicide_and_pesticide
4. 1 normally cited explanation is that the chloride ion reacts with the surface layer of aluminum oxide over the aluminum metal to requite the tetrachloridoaluminate ion, [AlCl_four]⁻(aq) ion, revealing the reactive metal surface for reduction of copper(Two) 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 loftier concentration of the chloride ion. More importantly, it goes against mutual sense, for aluminum-hulled boats seem to be quite unreactive in seawater.

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Source: https://uwaterloo.ca/chem13-news-magazine/december-2013-january-2014/feature/chemistry-using-minimum-cost-resources-part-1

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