CHEMICAL CANNON Materials: Baking Soda (2g), Vinegar (5ml), 1 test tube (200mm or greater) and stopper Drop the baking soda into the test tube containing the vinegar and quickly insert the stopper. Note: make sure that the test tube is not pointed towards the audience. While conducting this experiment you can explain the different forms of matter (solid, liquid and gases) and the fact that either one can be a product despite this form not being involved in the reaction. FLAMING CUSTARD Materials: Custard (~12g), Foot Pump, Tubing, Bunsen Burner, Retort Stand Custard does not burn while in lumps. However when dispersed in a fine powder it is very flammable – due to the exposed surface area. This may be demonstrated by placing a spatula full of Custard powder in a Bunsen burner. This will not burn (Although it may caramelise). Then demonstrate how the fine dispersed form burns. Connect a pump to the bottom end (small end) of a funnel via some rubber tubing. Place approximately 10 grams of custard into the funnel. Place a Bunsen burner about 6-12 inches above the mouth of the funnel at a 45o angle by using a retort stand. Ignite the Bunsen and adjust until a blue flame is present. Force air through the funnel via the pump. Note: the stronger the pump the more powder dispersed and hence the bigger the flames obtained HARD BOILED MAGIC Materials: 1 hard-boiled egg, 1 empty bottle, and matches Place the egg on top of the empty bottle. The egg should be a tight fit on the bottle. Explain to the students that if they forced the egg into the bottle that it would break. Then light a match and place it inside the bottle. Replace the egg so it acts like a lid. The lighting match will burn all the air in the bottle creating a vacuum sucking the egg into the bottle (without breaking it). The egg may be removed by blowing vigorously into the bottle.
MONEY TO BURN Materials: 3 beakers, Water( ~30mls) Ethanol (20mls) NaCl (~2g) Bunsen Burner, Two strips of Paper and a €5 note Beaker One contains 10mls of Ethanol Beaker Two contains 10mls of water Beaker Three contain 10mls of Water, 10mls of Ethanol and 2 grams of salt Soak the first strip of paper in beaker one. Place this in the Bunsen and observe a bright burning flame. Explain this is because alcohol is flammable. Soak the second strip in beaker two and place in the Bunsen. Note the paper will not burn because it is so wet. Then place the €5 note in beaker three and ask they audience what they expect. IF you place the note in the Bunsen the alcohol will burn off yielding a bright flame yet the note will not burn. The presence of the salt also increases the “brightness” of the flame. Note: Ensure the fiver is totally soaked and if necessary put in more water than Alcohol especially if the Fiver is not yours.
Blue Bottle Glucose (an aldehyde) in an alkaline solution is slowly oxidised by dioxygen, forming gluconic acid: CH2OH–CHOH–CHOH–CHOH–CHOH–CHO + 1/2 O2==> CH2OH–CHOH–CHOH–CHOH–CHOH–COOH (1.1) In the presence of sodium hydroxide gluconic acid is converted to sodium gluconate. Methylene blue speeds up the reaction, acting as an oxygen transfer agent. By oxidising glucose methylene blue itself is reduced and becomes colourless (formation of leucomethylene blue): If there is a sufficient air in the system, leucomethylene blue is quickly re-oxidised and the blue colour of solution is restored. On standing glucose reduces the dye and the colour of the solution disappears. In dilute solutions the reaction takes place at 40 - 60 oC. In more concentrated solutions (as used here by us) the process occurs at room temperature. The chemistry demonstrated by this experiment is used to determine accurately dissolved dioxygen in boiler feed water [1,2]. The indicator here is reduced by the glucose in the presence of base and is subsequently oxidised by the fresh dioxygen - a redox process. Preparation. Two one-litre Erlenmeyer flasks are half filled with tap water. 2.5 g of glucose are dissolved in one of the flasks (flask A), and 5 g of glucose in another (flask B). Demonstration. Just prior to the show 2.5 g of sodium hydroxide (NaOH) are dissolved in flask A and 5 g of NaOH in the flask B. During the presentation approximately 1 ml of 0.1% solution of methylene blue is added into each flask and flasks are stoppered and shaken to dissolve the dye. During this manipulation, the explanation of the chemistry involved can be given. The flasks are then set aside, and the blue colour of the solutions gradually disappears as glucose is oxidised by the dissolved dioxygen. The effect of concentration on the rate of reaction is also demonstrated here. It should be pointed out that the flask having twice the concentration uses the dissolved oxygen in about half the time as the other one. It should also be noticed that, having gone colourless, a blue zone remains close to the surface. This is due to oxygen diffusion from the air space within the flasks. After the solutions became colourless the blue colour can be restored by shaking the flasks. This experiment can be “revisited” several times during the show. Relevant references 1. L.S. Buchoff, N.M. Ingber and J.H. Brady, Anal. Chem., 1955, 27, 1401. 2. G.P. Alcock and K.B. Coates, Chem. and Ind., 1958, 554 3. H.W. Roesky and K. Möckel, Chemical Curiosities, trans. T.N. Mitchel and W.E. Russey, New York, VCH Publishers, Inc., 1996, p. 76. 4. C.P. Ellis, “Colour changes”, Chem. Britain, 1983, 326. 5. J.L. Lambert, M.J. Chejlava, G.T. Fina and N.L. Luce, “Sequential colour reactions: school colours and red-white-and-blue demonstration”, J. Chem. Educ., 1983, 60, 141. 6. A. Myers, “The blue bottle reaction and photosynthesis”, School Sci. Rev., 1981/1982, 63, 112. 7. Tested Demonstrations in Chemistry, ed. G.L. Gilbert, et al., Denison University, Granville, OH, 1994, vol. 1, pp. G-2, G-3. 8. B. Iddon, The Magic of Chemistry, BDH Chemicals, Poole, 1985, p. 3.
Briggs-Rauscher Reaction:
Theory: The next step in the demonstrations is to move from single-step and multistep (clock) reactions to oscillating reactions. Here, after reaction is completed, the system returns to its “initial” state and the reaction starts over again. The particular example used (the Briggs-Rauscher (BR) reaction) is perhaps the most impressive of the chemical oscillations. Three colourless solutions are mixed in a beaker and the stirred batch goes through 15 or more cycles of colourless, to amber, to blue-black, before ending as a blueblack mixture with the odour of iodine. The reaction was developed by Thomas S. Briggs and Warren C. Rauscher of Galileo High School in San Francisco [1]. This reaction is a hybrid of two other oscillating chemical reactions, the Bray-Liebhafski reaction and the Belousov-Zhabotinsky reaction [2]. The simplified (but still rather lengthy) explanation of the chemistry involved is given in the end of this section.
Materials: 30% hydrogen peroxide Potassium Iodate Sulphuric Acid Malonic Acid Manganese Sulphate (hydrate) Soluble Starch
Safety: 30% (100 vols) hydrogen peroxide is corrosive and a strong oxidising agent, contact with skin and eyes must be avoided. Sulphuric acid can cause severe burns, when concentrated it is a powerful dehydrating agent and generates considerable heat when diluted with water. If spilled on to the skin, care must be exercised to avoid excessive heating when flushing with water. If you have an ice bath to hand use the water from that or you can quickly remove the bulk of the liquid with a dry cloth or tissue. Flush with plenty of water and then treat the area with sodium hydrogen carbonateNaHCO3. Spills on the floor or bench should be neutralised with NaHCO3 and rinsed thoroughly. Preparation: Prepare three solutions as follows. Solution A: Pour 400 ml of distilled water into a 2-litre beaker. Wearing gloves pour 410 ml of 30% hydrogen peroxide into the water. Dilute the solution to one litre with distilled water. Solution B: Place 43 g of potassium iodate (KIO3) and approximately 800 ml of distilled water in the second 2-litre beaker. Add 4.3 ml of concentrated H2SO4 to this mixture. Warm and stir the mixture until the potassium iodate dissolves. Dilute the solution to one litre with distilled water. Solution C: Dissolve 16 g of malonic acid and 3.4 g of manganese sulphate (hydrate) in approximately 500 ml of distilled water in the third 2-litre beaker. In the 100-ml beaker heat 50 ml of distilled water to boiling. In a separate beaker mix 3 g of soluble starch with about 10 ml of water and stir the mixture to form a slurry (or the starch can be suspended in a little alcohol but in this case extra care must be exercised when pouring into the boiling water). Pour the slurry into the boiling water and continue heating and stirring the mixture until the starch has dissolved (1-2 minutes). Pour this starch solution into the solution of malonic acid and manganese sulphate and dilute the mixture with distilled water to one litre.
Procedures: Set a 1.5-litre beaker on the magnetic stirrer and place the stirring bar into the beaker. Pour 500 ml of solution A and 500 ml of the solution B into the beaker and adjust the stirring rate to produce a large vortex in the mixture. Then pour 500 ml of solution oC into the beaker. The initially colourless solution will become amber almost immediately. Then it will suddenly turn blue-black. The blue-black will fade to colourless, and the cycle will repeat several times with a period which initially lasts about 15 seconds but gradually lengthens. After a few minutes the solution will remain blue-black.
Explanation to Oscillation: The following explanation of the chemistry involved in this demonstration is taken (with some minor alterations) from Shakhashiri‟s book [2]. In the BR reaction the evolution of oxygen and carbon dioxide gases and the concentrations of iodine and iodide ions oscillate. The somewhat simplified mechanism of this reaction can be represented by the following overall transformation: IO3- + 2 H2O2 + CH2(COOH)2 + H+==> ICH(COOH)2 + 2 O2 + 3 H2O (11.1) This transformation is accomplished through two component reactions: IO3- + 2 H2O2 + H+==> HIO + 2 O2 + 2 H2O (11.2) HIO + CH2(COOH)2==> ICH(COOH)2 + H2O (11.3) The first of these two reactions can occur via two different processes, a radical process and a nonradical process. Which of these two processes dominates is determined by the concentration of iodide ions in the solution. When [I-] is low, the radical process dominates; when [I-] is high, the nonradical process is the dominant one. The second reaction (eq. (11.3)) couples the two processes. The reaction consumes HIO more slowly than that species is produced by the radical process when that process is dominant, but it consumes HIO more rapidly than it is produced by the nonradical process. Any HIO which does not react by eq. (11.3) is reduced to I- by hydrogen peroxide as one of the
component steps of the nonradical process for reaction (11.2). When HIO is produced rapidly by the radical process, the excess forms the iodide ions, which shut off that radical process and start the slower nonradical process. Reaction (11.3) then consumes the HIO so rapidly that not enough is available to produce the iodide ion necessary to keep the nonradical process going, and the radical process starts again. Each of the processes of reaction (11.2) produces conditions favourable to the other process, and, therefore, the reaction oscillates between these two processes. The detailed explanation requires attention to the individual steps of the two processes. If iodide ions are present in sufficient concentration, the reaction follows the nonradical process, reaction (11.2). The iodide ions react rather slowly with iodate ions, IO3- + I- + 2 H+==> HIO2 + HIO (11.4) The iodous acid (HIO2) is further reduced to hypoiodous acid (HIO), HIO2 + I- + H+==> 2 HIO (11.5) The hypoiodous acid is then reduced by hydrogen peroxide, HIO + H2O2==> I- + O2 + H+ + H2O (11.6) The net transformation represented by eq. (11.2) is obtained by the stoichiometric addition of eq. (11.4) + eq. (11.5) + eq. (11.6). Because reaction (11.2) is slower than reaction (11.3) under these conditions, so much HIO is used up by reaction (11.3) that reaction (11.6) cannot replenish the I- consumed in reactions (11.4) and (11.5); the [I-] keeps diminishing. Once the iodide ions have been sufficiently depleted, the nonradical process becomes very slow, and the radical process for reaction (11.2) can take over. This process involves the five steps [3]. IO3- + HIO2 + H+==> 2 IO2· + H2O (11.7) IO2· + Mn2+ + H2O ==> HIO2 + Mn(OH)2+ (11.8) Mn2+ + H2O2==> Mn2+ + H2O + HOO· (11.9) 2 HOO·==> H2O2 + O2 (11.10) 2 HIO2==> IO3- + HIO + H+ (11.11) These steps, when combined in the stoichiometry of 2 (eq. (11.7)) + 4 (eq. (11.8)) + 4 (eq. (11.9)) + 2 (eq. (11.10)) + eq. (11.11), have the overall result given by eq. (11.2). A significant feature of this process is that, taken together, the first two steps (eqs. (11.7)
and (11.8)) are autocatalytic - they produce 2 HIO2 for each one consumed. Therefore, the rate of these steps increases as they occur. Because this radical process is autocatalytic, it causes a rapid increase in the concentration of HIO, which is produced by the disproportionation of HIO2 (eq. (11.11)). This process does not rapidly consume all the iodate in the solution, because the last step is second order in the catalytic species. Thus, as its concentration increases because of the autocatalytic nature of the early steps, HIO2 is ever more rapidly consumed in this last step, and the sequence of the reactions quickly reaches a steady state. Equations (11.8) and (11.9) indicate the function of the manganese catalyst. The manganese is oxidised in reaction (11.8) and reduced in reaction (11.9). Its catalytic effect in the reaction is accounted for through its providing the means for reducing IO2· radicals to HIO2, thereby completing the autocatalytic cycle of equations (11.7) and (11.8). The hypoiodous acid produced by the radical process reacts with malonic acid by reaction (11.3). However, the radical process is faster than reaction (11.3), and the excess HIO reacts with hydrogen peroxide by reaction (11.6) to create I-, which shuts off the radical process and returns the system to the slow nonradical process initiated by reaction (11.4). The dramatic colour effects arise because reaction (11.3) does not take place in a single step, but by the sequence of reactions (11.12) and (11.13). I- + HIO + H+==> I2 + H2O (11.12) I2 + CH2(COOH)2==> ICH(COOH)2 + H+ + I- (11.13) The solution turns amber from the I2 produced through reaction (11.12), when the radical process maintains [HIO] greater than [I-]. The excess HIO is converted to I- through the reaction with H2O2 (eq. (11.6)). The solution suddenly turns dark blue when [I-] becomes greater than [HIO], and the I- can combine with I2 to form a complex with the starch. With [I-] high, reaction (11.2) switches to the slow nonradical process. The colour then fades as reaction (11.3) consumes iodine faster than it is produced. When the system switches back to the rapid radical process, the cycle is repeated.
The above reaction steps constitute a skeleton mechanism for the BR oscillating reaction. Upon initial mixing of the solutions, IO3- reacts with H2O2 to produce a little HIO2. The HIO2 reacts with IO3- in the first step of the radical process (eq. 11.7). The autocatalytic radical process follows, rapidly increasing the concentration of HIO. The HIO is reduced to I- in a reaction with H2O2 (eq. 11.6). The large amount of HIO reacts with I-, producing I2 (eq. 11.12). The I2 reacts slowly with malonic acid, but the concentration of HIO, I2 and I- all increase, because reaction (11.2) is faster than reaction (11.3). As [I-] increases, the rate of its reaction with HIO2 (eq. 11.5) surpasses that of the autocatalytic sequence of reactions (11.7) and (11.8). The radical process is then shut off, and the accumulation of reduced iodine is consumed by reaction (11.3) operating through the sequence of reactions (11.12) and (11.13). Eventually [I-] is reduced to such a low value that reactions (11.7) and (11.8) become faster than reaction (11.5), and the radical process takes over again. This oscillating sequence repeats until the malonic acid or IO3- is depleted. Chloride ion concentrations in excess of 0.07 M suppress the oscillations [4]. Therefore the vessels used for the preparation of the solutions must be clean and distilled water must be used for all preparations.
Discussion of Experiment: This experiment is very entertaining. The continues colour change during the reaction is generally unexpected by the audience and this adds to the experiment. This experiment is a must in any Science Road Show and would have to be included in this Road Show as it is one of the best experiments that was carried out.
JUG of Mystery Action: Water is poured from a jug into a series of six empty glasses. The glasses become filled with liquids coloured: 1) red, 2) White, 3) Blue, 4) Black, 5) Green, 6) Amber. Materials: One jug containing 5g of ferric ammonium sulphate in 500ml of water: Approximately 1/2g of the following solids dissolved in a few mls of water: (1) Potassium Thiocyanate, (2) Barium chloride, (3) Potassium ferrocyanide, (4) Tannic Acid, (5) Tartaric Acid, (6) Sodium Hydrogen Sulphite
Why: (1) Thiocyanate ion forms a deep red colour with iron(III), (2) Barium ion fomrs a white cloudy precipitate with sulphate ion, (3) ferrocyanide ion fomrs a deep blue compound with the iron (4) Tannic Acid forms a balck complex with iron(III), (5) Tartaric Acid forms a greenish complex with iron(III), (6) Sodium Hydrogen Sulphite forms an amber product with iron(III)
Note: Use Ferric not ferrous in the jug
The clock reaction
Theory: This experiment starts the theme of “clock” reactions. The demonstration is known as the “Old Nassau Reaction”, a clock reaction which turns orange and then black (and has therefore also been named the “Halloween Reaction”).
The reaction in this experiment takes place in several steps. First, sodium metabisulphite reacts with water to form sodium hydrogen sulphite: Na2S2O5 + H2O ==> 2 NaHSO3 Hydrogen sulphite ions reduce iodate(V) ions to iodide ions: IO3- + 3 HSO3-==> I- + 3 SO42- + 3 H+ Once the concentration of iodide ions is large enough that the solubility product of HgI2 (4.5 x 10-29 mol3 dm-9) is exceeded, orange mercury(II) iodide solid is precipitated until all of the Hg2+ ions are used up (provided that there is an excess of I- ions). Hg2+ + 2 I-==> HgI2 (orange or yellow) If there are still I- and IO3- ions in the mixture, the iodide-iodate reaction IO3- + 5 I- + 6 H+==> 3 I2 + 3 H2O Takes place and the blue starch-iodine complex is formed, I2 + starch ==> complex (blue or black) Materials: Sodium Metabisulphite Mercury(II) Chloride Potassium iodate
Safety: All soluble mercury salts are poisonous and should be treated accordingly.
Preparation: The following three solutions need to be prepared. A. Make a paste of 4 g of soluble starch with a few mils of water. Pour onto this 500ml of boiling water and stir. Cool to room temperature, add 13.7 g of sodium metabisulphite (Na2S2O5) and make up to 1 l with water. B. Dissolve 3 g of mercury(II) chloride in water and make the solution up to 1 l with water. C. Dissolve 15 g of potassium iodate (KIO3) in water and make the solution up to 1 l with water.
Procedure: Mix 50 ml of solution A with 50 ml of solution B. Then pour into this mixture 50 ml of solution C. After about 5 seconds the mixture will turn an opaque orange colour as insoluble mercury iodide precipitates. After further 5 seconds the mixture suddenly turns blue-black as a starch-iodine complex is formed. The second colour change (orange to black) is not normally expected by the audience and comes as a real surprise. This experiment can be extended in several ways. Diluting all the solutions by a factor of two increases the time taken for the colour changes to occur. Using a smaller volume of solution B speeds up the reaction. The effect of changing the amounts and concentrations of the various reactants cannot always be predicted simply because of the complexity of the system. For example, if the volume of solution B is doubled, the appearance of the orange colour is delayed and the blue colour fails to appear at all. If using mercury salts is not desirable, a somewhat simpler clock reaction can be performed. This is known as iodine clock reaction or Landolt reaction. The experiment is performed by mixing equal volumes of two solutions, one containing 2 g dm-3 KIO3 and H2SO4 0.03 M; the second - 0.4 g dm-3 of NaHSO3 in starch (2 g dm-3) previously dissolved in boiling water. The initially colourless mixture suddenly turns dark blue. There are several extensions to this reaction as well.
VOLCANO Materials: 10g grams of Ammonium dichromate Theory The orange crystals of ammonium dichromate, if heated to a sufficient temperature, start to decompose producing voluminous green chromium(III) oxide. Once started the reaction is self-supporting because it does not require any external oxidising or reducing agents. In fact, both the oxidant (Cr+6) and the reductant (N-3) are present in the same molecule, (NH4)2Cr2O7==> Cr2O3 + 4 H2O + N2 (13.1) Normally this demonstration is performed as the classical Volcano experiment, so called because ash produced occupies a far greater volume than that of the original dichromate and so builds a „mountain‟ of dark green ash with sparks and read heat emanating from a „Caldera‟ at its summit. Because the chromium(VI) and (III) compounds, like so many other things, are now suspected carcinogens this demonstration should no longer be carried out on an open (sand) tray. Do not conduct this experiment when extraction is not available.
Safety: This experiment should be carried out on an open tray only in a well-ventilated area. Chromium(VI) compounds and chromium(III) oxide are irritants for the skin and eyes and especially the respiratory tract if inhaled. In addition, some chemical suppliers note that chromium(VI) compounds, including ammonium dichromate, are suspected carcinogens, the risk being related to the length of exposure. Gloves and safety goggles must be worn when handling (NH4)2Cr2O7. Procedure: Place a pile of ammonium dichromate on a sand tray (or a ceramic tile) near a source of extraction ventilation. Wet the tip of the cone of ammonium dichromate with a combustible liquid (i.e., ethanol) and to set light to it by a gas burner (or a match).
Water-Wine-Milk-Beer:
Materials: Saturated Sodium Bicarbonate 20% Sodium Carbonate Phenolphthalein Indicator Saturated Barium Chloride Sodium Dichromate Concentrated Hydrochloric Acid.
Procedure: Begin with a tumbler containing "water". The solution is not really water. Prepare ahead of time a glass 3/4 full of water. To this add 20 to 25mL of saturated sodium bicarbonate and 20% sodium carbonate solution. (pH = 9) Pour the "water" into a wineglass. The wineglass should contain a few drops of phenolphthalein indicator. Pour the "wine" into a second tumbler. The second tumbler should contain 10mL of saturated barium chloride solution. Pour the "milk" into a beer mug. The beer mug should contain a very few crystals of sodium dichromate. Just prior to doing the demonstration; add 5mL of concentrated hydrochloric acid.
COLOURED CARNATIONS Prepare "carnations" from tissues and pipe cleaners. Moisten with phenolphthalein solution and allow to dry Spray with dilute ammonia or other weak alkali when the white carnations turn pink. For blue flowers, use thymolphthalein solution For yellow flowers, use p-nitrophenol.