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The work of Dr. Mahmoud Marsafy

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REMEMBER: you must do your own experiment risk assessment, ideas are illustrated but health and safety details are NOT included

19. Hydrogen chloride, ammonia and ammonium chloride experiments

Comment on the Micro-Scale Experimental Charts (1, 2 , 3 & 4) for the preparation of HCl (strong acid) & NH₃ ( weak base ) gases and their solubility's in water.

Most of the common gases can be generated by the appropriate reactions in a small glass test tube, heated when required. The tube is immediately stoppered, thus avoiding the need to perform the experiment in a fume hood (if the gas product is irritating or toxic). A sample of the gas (for further testing) can be obtained (without further purification or drying) by directly sucking the gas into a Beral pipette. Most of the basic reactions of the common gases can be performed on these samples.

Chart 1 demonstrates the preparation of HCl gas and its reaction with NH₃ gas giving the white clouds of NH₄Cl. It was sufficient to leave the tube stoppered lightly, thus causing the HCl gas to leak slightly and immediately reacting with the adjacent NH₃ gas ( from a soaked tissue paper strip with ammonia solution).

The relevant ionic equations for the reactions investigated are prominently displayed on the chart. The reaction generating HCl is a type of double replacement reaction and it is not a redox reaction.

It is interesting to note that NH₃ and HCl are reacting both in the gas phase to produce (synthesize) NH₄Cl (s) as a finely dispersed solid causing the appearance of the white clouds. Further this is a type of neutralization reaction without producing also water, and is not a redox reaction.

Chart 2 demonstrates the preparation of the HCl and NH₃ gases from the same compound (NH₄Cl salt). The great solubility's of the gases is clearly demonstrated by the rapid ascent of water in the pipettes containing the gas samples. The solubilities being 500 vols of HCl & 1000 vol of NH₃ per 1 vol of water at 0^oC.

The universal indicator added reveals prominently the acidic and basic character of these gases. This solubility effect was clearly demonstrated without any attempt to dry either of the two very soluble gases.

It is very interesting to note, that while HCl is a gas heavier than air (127%), while NH₃ is a gas lighter than air (59%), yet both gases were adequately sampled without the conventional special arrangement for appropriate air displacement to obtain these samples.

The reactions involved in the preparation of HCl and NH₃ gases are displayed as ionic equations on the chart and are both types of double replacement reactions generating these gases and are not redox reactions.

The chart displays also most of the appliances required for small scale experimentation. The stoppered tubes proved to be a very convenient container for the solid chemicals required for the reactions. It is recommended to use plastic tubes for these chemicals and they can be stored adequately for long periods without further precautions.

Concentrated sulphuric acid is safely and conveniently manipulated from a well stoppered small bottle with a Beral pipette. Heating is carried out with the flame of a micro-spirit Burner.

Chart 3 is a comparison between the conventional experimental setup for demonstrating the great solubility of NH₃ gas and the micro-scale alternative recommended.

The conventional display does not refer to the technique for filling the large glass flask with sufficiently dry ammonia gas. Probably NH₃ from a special gas cylinder was the source of ammonia gas. This involved technique is beyond the resources of most schools, and would involve a quite elaborate technique, with adequate safety precautions. The set-up also involves elaborate stands and clamps for adequate support.

The micro-scale technique very easily and convincingly generates ammonia, from the basic chemicals taught in all school chemistry syllabi. NH₃ gas is very safely and simply sampled in plastic pipettes. The solubility results obtained are similar to the conventional technique.

The dissolution of NH₃ in water is depicted on the chart as a reversible ionic equation emphasizing the weak base so characteristic of the aqueous ammonia solution.

Chart 4 compares the Micro-Scale investigation of the preparation of HCl gas and its reaction with ammonia gas resulting in white clouds of a fine suspension of NH₄Cl.

It is hoped that this presentation will motivate school chemistry teachers to adopt the recommended micro-scale preparation of gases in stoppered tubes. Most of the reactions of these gases can be further carried out on samples manipulated by a plastic pipette.

This micro-scale technique can be immediately applied in school laboratories without the need of elaborate conventional appliances and facilities, such as fume hoods and intricate apparatus assemblies.

20. The conversion of iron(II), Fe²⁺ into iron(III), Fe³⁺ ions

The investigations were performed with drops of the appropriate chemicals added in the preferred order into the cavities of a plastic strip as indicated in Chart A.

This micro-technique yields identical results as compared with the conventional technique utilizing test tubes. Moreover, since only drops of reagents are required for generation of Chlorine and Bromine, it is quite safe to carry out the reactions without use of the fume hood.

The multi-cavity plastic strip is the ordinary common packaging strip for pharmaceutical tablets. It incorporates 10 cavities. Each cavity allows the addition of a total of 10 to 15 drops of the required reactant solutions. After the reactions the strip is easily cleaned by: first emptying into a waste container, followed by flushing with water, and wiping dry. Such plastic strips can be used repeatedly for many tests.

Chart A is a summary of the planning for this investigation. The reactants added (in drops) to each cavity is clearly indicated by the lines connecting to the appropriately labeled plastic dropper bottles containing the various reagents.

Chart B is an image of the results of the oxidation of Fe²⁺ ion by some different oxidizing reagents in each cavity, as indicated

An aqueous, slightly acidified, clear, almost colorless, dilute Iron(II) sulphate solution , is important for this investigation to avoid any initial precipitation, which invariably occurs if the acid (dilute H₂SO₄) is omitted in the preparation. Such a solution has been found to keep unchanged for a long time.

· Two drops of this (Fe²⁺) solution were added in each cavity followed by drops of the indicated oxidizing reactant, with stirring. Immediately after observing the oxidation reaction, drops of NaOH were added to test for the resulting Fe3+ ion. In every case, a distinct characteristic red-brown precipitate of Fe(OH)₃ developed immediately except for the iodine oxidant. However the red-brown precipitate was obtained after adding extra drops of iodine to the initial precipitate which was coloured somewhat green.

· The original Fe²⁺ ion reacts with NaOH solution yielding a muddy green precipitate of Fe(OH)₂ as indicated in the topmost left reaction cavity.

Cl₂ was generated by drops of H₂SO₄ added to drops of bleach solution, which is considered as a solution containing the ClO^- ion [Chlorate (I)]. However Bleach solution itself alone is a strong oxidizing agent. If drops of Bleach solution are added to drops of the Fe²⁺ solution, oxidation to Fe³⁺ will immediately occur, and simultaneously a red-brown Fe(OH)₃ precipitate develops, since the bleach solution itself is sufficiently alkaline. The bleach reagent utilized was a commercially household product obtained from local supermarkets

Br₂ was generated in the same reaction cavity by adding drops of KBr solution followed by drops of bleach and acid. The generated chlorine displaced the bromide ions generating bromine. Oxidation with bromine could also be carried out by directly adding drops of bromine water, instead of the displacement reaction utilized.

The iodine solution utilized for oxidation of Fe²⁺ solution was prepared by dissolving the sparingly soluble iodine crystals in water, after addition of potassium iodide. This brown solution contains the I₃^- anion which is the oxidizing species. However, when the reaction is carried out , the precipitate first formed is a mixture of a greenish precipitate mixed with some brown- red precipitate , as well. Addition of more drops of iodine solution is required to convert it predominantly into a brown-red precipitate. This is evidence that the iodine solution is a weak oxidizing agent.

Oxidation by H₂O₂ was performed by the sequential addition of 2 drops of the Fe²⁺ solution, then two drops of H₂O₂ solution, resulting in oxidation to Fe³⁺, as evidenced by a yellowish coloration developing. On further addition of one or two drops of NaOH the red-brown Fe(OH)₃ precipitate occurs with, at the same time causes catalytic decomposition of H₂O₂ releasing gas bubbles (O₂)

The oxidation by MnO₄^- and Cr₂O₇^2- was performed with the addition of drops of both reagent solutions without further acid addition.

The oxidation with concentrated nitric acid requires its addition in excess, without heating. However it is better to add drops of a more concentrated solution of NaOH for testing the Fe³⁺ ion instead of the usual (1M) NaOH solution.

Note: All the above oxidation reactions were carried out on the soluble iron(II) ions in aqueous solution. However the insoluble iron(II) hydroxide precipitate can also be oxidized directly by adding the same oxidizing reagents. The order of reagent additions in this case would be :

2 drops of Fe²⁺ solution + 3 drops of NaOH (1M) + drops of the oxidizing reagents.

Thus the dirty green precipitate of Fe(OH)_2(s) , will be transformed directly into the red-brown precipitate of Fe(OH)_3(s).

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