Microscale chemistry experimentation with a coin and a can
Mahmoud K. El-Marsafy, Peter Schwarz

 Introduction
Microscale Chemistry Experimentation has a tradition of more than 75 years in Egypt [1]. Since the end of the eighties it is part of pre-service teacher training at Ain Shams University  Cairo. The materials used  are syringes, vials, dropper bottles and plastic surfaces. This art of chemical experimentation was demonstrated in the biggest auditorium of Cairo American University using an overhead projector [2]. The authors presented results of their joint research in an unpublished book [3], in Internet (http://www.micrecol.de) and during the IUPAC/UNESCO conferences of Cairo (1998), Budapest (2000) and Peking (2002).
Microscale Chemistry Experimentation (MCE)  works with volumes between 5 µl (drop from a syringe with smallest hypodermic needle) and 5 ml. The appliances are affordable at a very low cost and they are worldwide available. Soft drink cans, blisters, Tetra Paks, drinking straws (spatulae) can be found in households. Small pieces of wire or of sheet metals are obtained by  craftsmen. Chemicals like  table salt, sugar, baking powder are obtained from the kitchen.

These economic and ecological advantages of MCE  with packages are supplemented by  methodological and didactic ones:


1 Pfennig coins - once the smallest unit of German money- were made of steel coated by copper. These coins are useful not only as 2.0-gram weights on a scale [6] but also as anodes, cathodes and as source of  copper salt solution.  This solution will be obtained by one of   the most dramatic chemical changes present in any curriculum of the world:  
 "While reading a textbook of chemistry, I came upon the statement `nitric acid acts upon copper` and I determined to see what this meant.Having located some nitric acid  I had only to learn what `act upon` meant. In the interest of knowledge I was even willing to sacrifice one of the few copper cents then in my possession. I put one of them on the table, opened a bottle labeled `nitric acid`, poured some of the liqid on the copper coin, and prepared to make the observation. But what was this wonderful thing which  I beheld?  The cent was already changed, and it was no small change either. A greenish-blue liquid foamed and fumed over the cent and over the table. The air became colored dark red. How could I stop this? I tried by picking the cent up and throwing it out of the window. I learned another fact: nitric acid acts upon fingers. The pain led to another unpremeditated experiment. I drew my fingers across my trousers and discovered: nitric acid acts upon my trousers...." (and also upon fingers giving them a yellow color) [7].  
By doing this coin experiment on microscale all above abservations can be reproduced by using only  1 microdrop (50 microlitres) of nitric acid. All other recommended safety precautions (goggles, waste manipulation, ventilation etc) should be strictly aoolied. The corroding and dyeing effect of nitric acid may be demonstarted with an egg instead of the skin.
In many countries Cola cans are made from two different metals: Only the closure consists of aluminium while the rest of the can is steel („FE“) double coated with plastic.  The complete can after rinsing is used for the collapsing can experiment [9]. The can is also used to demonstrate the burning and explosion of hydrogen [10], [11]: The upper side of the can is cut off, and the can is restored into its previous shape again and its bottom is perforated by a tiny hole.  For further Microscale Chemistry Experimentatio the can is cut into small pieces. In the following experiments Cu, Fe and Al will be used to explore basic concepts of electrochemistry.
Beads of different colors and sizes will be used as models for atoms, ions and electrons [12]. 

Experiments:

1. Direct electron transfer from copper atoms in a coin to the nitrate ions

.... ...
Photo 1

Left: A Pfennig coin reacts with a drop of nitric acid releasing a brown gas + green-blue solution

Middle: The direct transfer of two electrons from a copper atom to two nitrate ions is visualized by beads of different sizes and colors. The orange bead with the two micro beads is a model of a copper atom. The two blue beads each of them welded with three red beads show models of two nitrate ions.
Right: The models illustrate the electrons (white micro beads) transfered from the copper atom to the nitrate ions.These two electron acceptors will immediately continue reactions resulting in the nitrogen dioxide observed.
                            Reduction:       2 [NO3] -(aq)  + 2e- +4 H + (aq)  -->  2 NO 2 (g) + 2 H2O(l)
                              Oxidation:                                           Cu(s)  --->  Cu2+(aq) +2 e-
                     Redox reaction:    Cu(s) + 2 [NO3]-(aq) + 4 H+(aq)  -->   Cu2+(aq) + 2 NO 2 (g) + 2 H2O(l)    
                                                 Cu(s) + 4 HNO 3   -->   Cu(NO 3) 2 + 2 NO2 + 2 H2O(l)  


2. Direct electron transfer from aluminium atoms of the can closure to the copper ions (microscope)

 
Photo 2:
  1 x 2 mm of the can closure (Al) in a drop of copper salt solution of the previous experiment

A drop of the copper nitrate / table salt solution obtained from the Pfennig coin is placed on a  microscopic slide. A tiny piece of aluminium - cut from the closure of a Cola can-  is pushed into that drop.
It can be observed that small trees in the colour of copper are branching out.

                Reduction:            3 Cu2+(aq) +  6e- --> 3 Cu(s)
                Oxidation:                         2 Al(s) --> 2 Al3+(aq) + 6e -
       Redox reaction:        3 Cu2+(aq) + 2 Al(s) -->  3 Cu(s) +  2 Al 3+(aq)


3. Electron transfer from atoms of activated aluminium to water molecules

.

    Photo 3   Aluminium of a can closure activated by a Pfennig coin reacted with water

A can closure is connected to a Pfennig coin by a crocodile clamp and placed in a blister with sea water solution.
After some hours a white gelatinous precipitate (wp) appears below the can closure. Gas bubbles (b) can also be seen at the  rim of the coin.
The observations are the result of an indirect electron transfer from aluminium to water through copper.  Aluminium is the electron donor , water being reduced to hydrogen:
 
 Reduction:        6 H2O(l) + 6 e- --> 6 OH-(aq) + 3 H2(g

Oxidation:                      2 Al(s) --> 2 Al3+(aq) + 6e -
Redox reaction:   6 H 2 O(l) + 2 Al(s) --> 2 Al(OH)3(s)  + 3 H2(g)


4. Forced electron transfer from Cl - to Cu2+ ions during electrolysis

. .
Photos 4.1 - 4.3
 Electrolysis of a mixture of copper salt solution and table salt (1 + 2), testing for chlorine (3)

The solution from experiment 1 is rinsed into a blister by adding drops of concentrated table salt solution.
Left: Two pencil leads –connected with a 9-Volt battery- dip into this solution. Middle:  A layer with the color of  copper appears on the negative (left) electrode, gas bubbles are around the positive electrode.  
          Cathodic reduction:                Cu2+(aq) + 2e- --> Cu(s)
          Anodic oxidation:                           2 Cl-(aq)   --> Cl2 (g)   + 2e-
     Redox reaction:   Cu2+(aq) +   2 Cl -(aq)  --electrolysis--> Cu(s)  +  Cl2 (g)
Right: This photo shows the result of testing for chlorine:  A white tissue paper wetted with colorless potassium iodide solution was dipped into the right side of the blister. This reagent for chlorinet turns brown.  
A second redox reaction takes place during the test for chlorine:   Cl 2    + 2 l-  -----> 2 Cl -    + l2


5. Spontaneous electron transfer from Cu to Cl2 on carbon electrodes in  a Galvanic Cell

.

 Photo 5
Potential and current between Cu und Cl2 on carbon electrodes in a blister
After finishing the electrolysis of copper salt / table salt solutions in the previous experiment the battery is removed and the electrodes are connected to a digital multimeter. It reads a potential of 1.10 V. By switching over to the DC Amp range a current of about 2 mA could be detected . 
A spontaneous indirect electron transfer reaction between copper (electron donor) and chlorine (electron acceptor) takes place while a current is observed .
                         Reduction:      Cl2 (g)   + 2e-  --> 2 Cl-(aq)
                         Oxidation:            Cu(s)       --> Cu2+(aq) + 2e-
               Redox reaction:   Cu(s) + Cl2 (g)   --> Cu2+(aq) +   2 Cl-(aq)

6. Galvanic Cells with a Pfennig (Cu) and the two different metals (Al and Fe) from a Cola can

. .

Photo 6
Galvanic Cell potential between Cola iron and copper of a Pfennig
 

The skin can act as an electrolyte in Galvanic Cells. Potentials and currents could be measured while Cola can metals and a Pfennig coin were pressed between thumb and forefinger in three combinations.
Results: 0.25 V between Fe and Cu (photo), 0.25 V between Al and Fe and 0.40 V between Al and Cu.
In the Al/Cu cell the potential is nearly the sum of the potentials in the two other cells.
The electron transfer always takes place from the less noble metal (anode, negative electrode) to the more noble one (positive electrode, cathode).
 


7. Batteries  of Galvanic Cells  made from can metals in vials with salt bridges



Photo 7
Potential between a battery of three Galvanic Cells made of two Cola can metals (Al, Fe) 

Materials: Music cassette box as stand, six injection bottles (3 ml) full of salt water, three sanded pieces of Cola can iron (Fe) and closures (Al) fitting to these bottles, 2 insulated copper wires with crocodile clips to connect with the digital multimeter,  2 more cocodile clips to connect the adjacent Al and Fe-electrodes, 3 wool threads as salt bridges.
Experiment: Three identical Galvanic Cells (single and connected in series) are assembled with the red wool threads (S) forming salt bridges.
1. An aluminium electrode and an iron electrode are connected with the multimeter.
2. Two of these cells are connected in series to form a two cell battery.
3. The photo shows the assembly for the three cell battery and the potential obtained when the terminal Al and Fe electrodes are connected to the multimeter.
 Results:
1. The potential of a single is 0.24 Volt.   2. In a two cell battery the potential is doubled:  0.51 Volt.
3. In the three cell battery the potential is tripled as in the photo (0.86 Volt).
  The potentials in Galvanic Cells in series add up.

8. Electrolysis of table salt solutions with anodes made of two Cola can metals

.. ..
Photos 8.1 - 8.3
Electrolysis of salt water with aluminium anode (1) and with iron anode (2 and 3 )

Experiment 1 : Salt water and one drop of universal indicator solution (UI) are electrolysed in the cavity of a blister packing.  A strip of sanded cola can metal (Fe) is used as negative electrode, its closure (Al) as positive electrode. A strong flow of gas bubbles is observed at the Fe cathode. The green colour of UI turns to violet (alkaline).
                      Reduction:  4 H2O(l)  + 4 e - --> 2 H2(g) + 4 OH-(aq)
At the Al anode less gas bubbles come out, the green colour of UI turns to red (acidic).
                     Oxidation:            2 H2O(l)  --> O2 (g) + 4 H+(aq) + 4 e-
                    Redox reaction:    2 H2O(l)  --> 2 H2 (g) +  O2 (g)
Verification of this interpretation:
The liquids around anode and cathode are mixed after the end of electrolysis: The colours of UI return to green again. This proves that equivalent quantities of hydroxide and of oxonium ions have developed during cathodic reduction and anodic oxidation. These ions neutralised each other during mixing
Experiment 2 (2 and 3)
In this experiment the electrodes are exchanged: The iron from the can is used as the positive electrode, the closure (Al) as the negative electrode. Both are dipped into the cavity of the blister packing with salt water + UI again and are connected to a 9-Volt battery.
Gas bubbles are observed at the negative electrode (Al). The color of the UI changes to violet (alkaline).
                    Reduction: 2 H2O(l)  + 2 e - -->  H2(g) + 2 OH-(aq)
No gas is observed at the positive electrode (Fe), the UI remains green.
               Oxidation:                     Fe(s)  --> Fe2+(aq) +  e-
             Redox reaction:   2 H2O(l) + Fe(s)  --> Fe 2+(aq)  +   2 OH-(aq) + H2(g) .
Verification of this interpretation:
3: The electrolysis is repeated without UI. The liquids around anode and cathode are mixed:
     A green slimy precipitation appears:  Fe 2+ (aq) + 2 OH-(aq) --> Fe(OH)2(s)
     By adding drops of diluted sulfuric acid this precipitation disappears:
                                               Fe(OH)2(s) + 2 H+(aq) –neutralisation-->  Fe 2+(aq) + 2H2O
On addition of drops of KMnO 4 solution its color is lost proving the presence of Fe2+ ions:
        5Fe2+ (aq) + MnO4-(aq) +  8 H+(aq) ---direct redox reaction -->  5 Fe3+(aq) + Mn2+(aq) +  4 H2O
In order to test for iron(III) ions two drops of the solution are transferred into two cavities of a blister. One drop of potassium hexacyanoferrate(II) solution is added to one sample: A blue precipitate is observed. One drop of ammonium thiocyanate solution is added to the other sample. A red brown color appears.
     

Literature

[1]  E.C. Grey, Practical Chemistry by Micro-Methods, W. Heffer & Sons, Ltd., Cambridge 1925
[2]  M.K. El-Marsafy et al, The microscale chemistry laboratory technology its implications on the future education, Book of abstracts symposium on science…, S. 18 – 20, American University Cairo 1995
[3]  M.K. El-Marsafy,  P. Schwarz, Microscale chemistry experimentation, 119 Seiten englisch und arabisch, Kairo 1996
[4]  P. Schwarz, Chemische Unterrichtsexperimente im Internet, CHEMKON/8. 1, (2001) 47
[5] P. Schwarz, Nachhaltiges Erziehen durch mikrochemisches Experimentieren im Kindergarten, Chemie und Schule 2 (2001), S. 13f
[6] P. Schwarz, Beobachten, Messen, Experimentieren im Kindergarten der Sharm El Sheikh Language School, CHEMKON/8. 3, (2001) 169f
[7] Brown LeMay Bursten, Chemistry the central science, Prentice Hall International, Inc., New Jersey 2000
     referring to Ira Remsen, Chemistry Textbook 1901)
[8]   http://www.micrecol.de/water4.html
[9]   http://www.micrecol.de/water4a.html
[10] http://www.micrecol.de/afuel11.html
[11] V. Obendrauf, Döschen hüpf! Die pfeifende Dose im Schülerversuch, Chem. Sch. 15 (2000) Nr. 4, S. 8 - 12
[12] http://www.micrecol.de/elecE5c.html

About the authors

M.K. El-Marsafy  http://www.micrecol.de/ marsafyE.html   P. Schwarz  http://www.micrecol.de/ peterE.html