Introduction 1-4 Cover Page 1 Table of Contents 2 Acknowledgments 3 Safety Precautions 4 Flames 5-30 Calcium Hypochlorite 5-6 Chlorate Fire 7-8 Cold Fire 9-10 Flaming Gummy Bear 11-12 Green Fire 13 Ice Fireworks 14-15 Nitrocellulose 16-18 Permanganate Combustion 19-20 Purple Fire 21-22 Red Fire 23-25 Smoke Screen 26-27 Yellow Fire 28-29 Zinc Sulfide 30 Polymers: 31-38 Disappearing Styrofoam 31-32 Nylon 33-34 Silly Puddy 35-36 Slime 37-38 pH 39-42 Acid Breath 39-40 Cabbage 41-42 Gas Laws 43-49 Collapsing Can 43 Geyser 44-45 Sinking Bobbers 46-47 Siphon Fountain 48-49 Color Changes 50-58 Alchemy 50-51 Colorful Cloth 52-53 Crystal Lattice 54 Malonic Acid Oscillator 55-56 Patriotic Colors 57-58 Other 59-77 Carbon Snake 59-60 Contact Explosive 61-62 Endothermic 63-64 Kinetic Molecular Theory 65 Luminol 66-69 Magic Eggs 70-71 Peroxide Foam 72-73 Pickle 74-75 Sugarloaf 76-77

Safety Precautions: Any of these experiments can be dangerous and are designed to be performed by an experienced scientist. It is strongly recommended that the experimenter become familiar with the chemicals and reactions in these demonstrations before performing them in front of any audience. It is assumed that every demonstrator will wear safety goggles, gloves and a lab coat. We also recommend having a fire extinguisher and a first aid kit available for every demonstration. Every demonstration has a section regarding safety precautions. The following list defines some of the general warnings for each of the demonstrations. Additional safety precautions are listed as needed.

1. Fire Hazard: This refers to any experiment where an open flame is generated. A fire extinguisher must be available for these demonstrations. Be sure that the reaction is done well removed from other flammable materials and that the demonstrator is prepared if some flammable material escapes the reaction container.

2. Smoke Hazard: Noxious gases or large amounts of smoke are produced. We recommend any of these experiments to be done in a fume hood. If no fume hood is available, the demonstrator should practice the reaction beforehand to determine the appropriate reaction quantities.

3. Strong Acid/Base: These experiments contain a strong acid or base and should be performed near a safety shower. In the absence of a safety shower, several liters of a neutralization solution should be readily available (2~5M NaHCO3, sodium bicarbonate). Many concentrated acids and bases should only be used in a fume hood (HCl, NH4OH, H2SO4, Acetic acid, etc.) because of the vapors.

4. Volatility Hazard: These reactions have a volatile liquid that is extremely flammable. Avoid having an open flame anywhere nearby as these volatile liquids have been shown to ignite from a flame several meters away.

Calcium Hypochlorite Combustion Principles Illustrated:

-Spontaneous combustion

Required Equipment:

-One Pyrex dish -10mL syringe

Required Chemicals:

-Calcium hypochlorite (Ca(ClO)2) -Glycerin (C3H6O3)

Demonstration Description: In a fume hood, a small amount of liquid is added to a pile of white granules. After several seconds the granules begin to smoke, grow and eventually burst into a bright yellow flame. Demonstration Procedure: Weigh out 10 grams of calcium hypochlorite. Make a small impression in the powder to provide room for the addition of glycerin. Add 5mL of glycerin to the impression. Step back and allow the reaction to proceed. The calcium hypochlorite is typically in excess in this reaction. For the best results a large deep impression should be made to allow as much glycerin as possible to be involved in the reaction. The easiest form of application of the viscous glycerin is a syringe. The chemicals used and produced in this demonstration are especially unpleasant but the chemistry is very interesting and merits the hassle. Theory: This reaction illustrates the principle of spontaneous reactions. Notice how the required reactants for each step are generated as products of a different reaction.

Calcium hypochlorite will spontaneously decompose according to the following reaction:

2Ca(ClO)2 + heat à 2CaO + 2Cl2 + O2 The presence of chlorine gas is confirmed by the potent chlorine odor that will quickly permeate the area surrounding the chemical. In the presence of glycerol, the CaO formed spontaneously by the decomposition of calcium hypochlorite will react with the glycerol:

3CaO + C3H6O3 à 3C(solid graphite) + 3Ca(OH)2 + heat This reaction is confirmed by the formation of a large black carbonaceous residue. The heat produced accelerates the decomposition of calcium hypochlorite. When a sufficient amount of heat is produced, the oxygen formed by the spontaneous decomposition of calcium hypochlorite will directly react with the glycerol or the solid graphite and ignite:

C3H6O3 + 3O2 à 3H2O + 3CO2

C + O2 à CO2 This first reaction produces a large amount of water. Calcium oxide is extremely reactive towards water and will react, generating intense heat:

CaO + H2O à Ca(OH)2 + heat This heat generated accelerates the further decomposition of calcium hypochlorite and the cycle continues. Safety Considerations: FIRE HAZARD

SMOKE HAZARD- The reactants and the products generate large amounts of potent gas. Calcium hypochlorite is a strong oxidizer and hygroscopic. Direct contact with the skin should be avoided.

Disposal:

Rinse all of the products thoroughly and discard the solid residue in a waste container.

Chlorate Fire Principles Illustrated:

- Instability of an oxidized halogen

- Reduction/oxidation (redox) reaction

Required Equipment:

-Pyrex dish -Mortar and pestle -Medicine dropper

Required Chemicals:

-Sugar (C12H22O11) -Potassium chlorate (KClO3) -Concentrated sulfuric acid

(H2SO4) Demonstration Description: A violent fire erupts when concentrated sulfuric acid contacts a small pile of powdered sugar and potassium chlorate. Demonstration Procedure: The preferred ratio for this experiment is to use 2 parts potassium chlorate for every 1 part of sugar. One gram of total powder is sufficient for a large audience. Using the mortar and pestle, grind the potassium chlorate until it is a fine, evenly dispersed powder. The sugar may also be ground for best results or powdered sugar may be used. Note: The sugar and potassium chlorate should be ground independently. The two powders should then be gently mixed together thoroughly and placed in a small mound in the center of the Pyrex dish. Addition of one or two drops of concentrated sulfuric acid will immediately ignite the mixture. It is also possible to dip a glass rod into the sulfuric acid and touch it to the powder. Theory: The exact reaction between potassium chlorate, sugar and sulfuric acid is not known. This experiment does illustrate a redox reaction. The flame accompanies the dehydration and oxidation of the powdered sugar and the accompanying reduction of chlorine. A possible reaction where sulfuric acid acts only as a catalyst is:

4KClO3 + 2C12H22O11 + 19O2 à 4KOH + 24CO2 + 20H2O + 2Cl2 Chlorates are particularly reactive due to the oxidation state of the chlorine. Oxygen is slightly more electronegative than chlorine and will essentially hoard the electrons away from chlorine, oxidizing the chlorine to a 5+ oxidation state. In the

presence of a good reducing agent, such as sugar, the chlorine will act as an oxidizing agent and react readily to gain electrons. Safety Considerations: FIRE HAZARD SMOKE HAZARD STRONG ACID Disposal: Most of the reactants are entirely consumed. The remaining product should be flushed down the drain.

Cold Fire Principles Illustrated: -Heat of vaporization -Heat capacity of water Required Equipment:

-large beaker >1000mL -small beaker ~250mL -Flame source -Tongs

Required Chemicals:

- Isopropyl alcohol (Rubbing alcohol- C3H8O) -Small cotton cloth or paper towel -Sodium chloride (NaCl)

Demonstration Description: A moist towel is ignited and yellow flames surround the towel without burning it. Demonstration Procedure: Make 100mL of ~50% isopropyl alcohol. Standard rubbing alcohol is 70% isopropyl alcohol and can be diluted to 50% by adding 30mL of water to 70mL of alcohol. Add 0.01g of NaCl to the alcohol solution. Immerse the towel or cloth in the solution, remove and squeeze out the excess liquid. Using the tongs to hold the cloth, ignite the cloth with the flame source (match, lighter, Bunsen burner, etc.). The cloth should remain ignited for several minutes and self extinguish when the alcohol is gone. Keep the cloth suspended over a large beaker partially filled with water in the event of an unforeseen emergency. This demonstration has also been done using paper money with moderate success. The key to success is to use older bills that have increased absorptive capacity. Theory: This experiment illustrates several principles. In order for the alcohol to ignite it must first evaporate to the vapor phase. Much of the energy associated with the flaming alcohol vapors is absorbed by the liquid alcohol in the transition to the vapor phase. The

water also absorbs large amounts of energy as it rises in temperature and evaporates. This is an excellent illustration of the heat capacity of water. Any additional energy from the flames is given off to the surrounding air. The combustion reaction for isopropanol:

2C3H8O + 9O2 à 6CO2 + 8H2O The addition of NaCl serves to give the flame a yellow color characteristic of

burning sodium ions. Safety Considerations: FIRE HAZARD VOLATILITY HAZARD Disposal: The excess dilute alcohol can be flushed down the drain.

The Flaming Gummy Bear Principles Illustrated:

-Reduction/oxidation (redox) reaction -Activation energy barrier

Required Equipment:

-One large 50mL test tube -Test tube clamp -Bunsen burner -Tweezers

Required Chemicals:

-Small gummy bears -5 grams of potassium chlorate (KClO3)

Demonstration Description: A gummy bear is dropped into a large test tube containing a clear liquid. The gummy bear immediately bursts into an intense purple flame and continues to burn for about 30 seconds. Demonstration Procedure: Add 4-5 grams of potassium chlorate to a large 50mL test tube. Heat the test tube over a flame until the potassium chlorate completely melts. While the potassium chlorate is still in the liquid phase, use tweezers to drop a gummy bear into the test tube. Potassium perchlorate has been shown to be entirely ineffective for this experiment as it will decompose to potassium chlorate at melting temperatures prior to igniting with the gummy bear. Any high sugar compound may be substituted for the gummy bear and toothpicks have also been shown to work. Theory: The exact reaction between sugar and potassium chlorate is not known but a plausible reaction is shown:

2KClO3 + 2C12H22O11 + 21O2 à 2KOH + 24CO2 + 20H2O + 2HCl Chlorine and atmospheric oxygen are reduced by oxidizing carbon.

This experiment also shows the ability to overcome the activation energy barrier by application of heat. Every chemical reaction can be illustrated with a reaction

coordinate that diagrams the energy level of the system as the reaction proceeds:

Products

Reactants

ENERGY

Activation energy

Net reaction energy

Melted potassium chlorate

The reaction requires an initial input of energy followed by a large release of energy as the products are formed. The amount of energy required to begin the reaction is called the activation energy. In this demonstration, the heat required to melt the potassium chlorate supplies the activation energy. Once melted the system is “sitting” on top of the reaction coordinate and the addition of the gummy bear will provide the final reactant to allow product formation. Safety Considerations: FIRE HAZARD- Molten potassium chlorate is extremely flammable and should be dealt with carefully. SMOKE HAZARD

It is important not to use too much potassium chlorate as it has a tendency to splatter out of the tube. The gummy bear may also jump out of the tube.

Green Fire Principles Illustrated:

-Reduction/oxidation (redox) chemistry -Photon emission

Required Equipment:

-Mortar and pestle -Pyrex dish

Required Chemicals:

-Barium nitrate (BaNO3) -Potassium chlorate (KClO3) -Sucrose (table sugar) or shellac powder

(C12H22O11) Demonstration Description: A small pile of powder is ignited and immediately bursts into a bright green flame. Demonstration Procedure: For the best results mix together 2 grams of barium nitrate, 2 grams of potassium chlorate and 0.6 grams of shellac. All three components should be finely ground SEPARATELY and mixed after grinding. This product is easily ignited and produces a very intense green flame as well as a fair amount of smoke. If you do not have shellac gum, you can produce a similar flame with 1 gram of sugar added to 2 grams of potassium chlorate and 1 gram of barium nitrate. As before, finely grinding each component separately and mixing afterwards allows for the best results. The flame is neither as intense nor as green as that produced using shellac. Ignition of the powders can be problematic, but we found a candle attached to a meter stick to be quite effective. Theory:

These compounds undergo an oxidation/reduction reaction where the sugar or shellac is oxidized and the nitrogen and chlorine atoms are reduced. Multiple products are formed and the exact chemical equation is not known. The following unbalanced equation is a reasonable possibility using sugar as the reducing agent:

Ba(NO3)2 + KClO3 + C12H22O11 + O2 à KOH + CO2 + H2O + Cl2 + Ba(OH)2 + N2

The high energies of this reaction result in the excitation of electrons, which upon returning to their ground state will emit light. The net effect of the various light emissions from these elements will result in the observable green color. Safety Considerations: FIRE HAZARD SMOKE HAZARD Do not store these mixtures for extended periods of time in any quantity. Disposal: Any remaining product can be rinsed and washed down the drain.

Ice Fireworks Principles Illustrated:

-Reduction/oxidation (redox) chemistry -Spontaneous combustion -Excitation and emission -Catalysis

Required Equipment: -Mortar and Pestle -Pyrex Dish Required Chemicals:

-Zinc dust (Zn) -Ammonium nitrate

(NH4NO3) -Ammonium chloride

(NH4Cl) -Ice (H2O) -Barium nitrate

(Ba(NO3)2) Demonstration Description: A brilliant green flame and smoke are produced when an ice chip is added to a small pile of powder. Demonstration Procedure:

Grind, SEPARATELY, 4 g of NH4NO3, 1 g of NH4 and 0.5 g of Ba(NO3)2. Carefully mix these compounds and place them in the Pyrex dish. Carefully mix 4 grams of zinc dust into the mixture and form a cone with the mixed powder. Place an ice chip on the pile. Within a few seconds the reaction should start, producing a brilliant green flame and a fair amount of smoke. Water initiates the reaction so care should be maintained to avoid inadvertent combustion. Theory: Decomposition of ammonium nitrate is catalyzed by the chloride ions in ammonium chloride. Water initiates the redox reaction between zinc and nitrate. The reaction becomes autocatalytic as water is produced from the decomposition of nitrate:

NH4NO3 à N2O + 2H2O This reaction generates a large amount of heat and the resulting high temperature

causes the ammonium nitrate to melt. The liquid ammonium nitrate leads to the dissolution and oxidation of zinc, the flame producing reaction. A reasonable equation for the reaction is:

Zn + NH4NO3 à N2 + ZnO + 2H2O The flame that is observed in this reaction is due to the exothermic nature of the

process. The green flame produced is due to the excitation of the valence electrons in the barium metal, and the photon that is then emitted when the electrons relax. Safety Considerations: FIRE HAZARD SMOKE HAZARD Do not store these mixtures for extended periods of time in any quantity. Disposal: Any remaining product can be rinsed and washed down the drain.

Nitrocellulose (Gun cotton) Principles Illustrated:

-Reduction/oxidation (redox) chemistry -High energy organic products

Procedure A Required Equipment:

-Three 250mL beakers -Ice bath -Paper towels -Large Pyrex dish

Required Chemicals:

-70mL concentrated sulfuric acid (H2SO4)

-30mL concentrated nitric acid (HNO3)

-100mL of 1M sodium bicarbonate solution (NaHCO3)

-5g dry absorbent cotton Procedure B Required Equipment:

-One 500mL beaker -250mL ice bath -Three 250mL boiling water baths -Hot plate -Large ice bath -Tongs/tweezers -Paper towels

Required Chemicals:

-75mL concentrated sulfuric acid (H2SO4) -75mL concentrated nitric acid (HNO3) -5g dry absorbent cotton (NaHCO3)

Demonstration Description: A piece of cotton is placed in a large Pyrex dish and ignited with a lighted candle on a long stick. The cotton rapidly burns and leaves no residue. Demonstration Procedure: Procedure A:

This procedure is faster but produces a more impure nitrocellulose. Place a 250mL beaker filled with cold water, in the ice bath. In the other 250mL beaker, combine nitric and sulfuric acids by slowly adding nitric acid to sulfuric acid with gentle swirling.

When the mixing is finished, place about one-third of the cotton into the acid solution. Make sure that the cotton is completely submerged and saturated with the acid. Gently agitate the cotton with a glass stirring-rod for one minute.

Next, remove the 250mL beaker from the ice bath, and place the cotton strips into that beaker. Let the strips soak for one minute with slight mixing. Remove the cotton and rinse it under running water for 2-4 minutes. The rinsing step is done to ensure that no residual acid is left in the cotton that will interfere with the demonstration. Place the cotton into another 250mL beaker containing 100mL of the sodium bicarbonate solution. If bubbles are produced, remove the cotton from the solution and continue rinsing until the cotton can be submerged into the sodium bicarbonate without producing excessive bubbling. Rinse a final time, squeeze dry and spread the cotton over paper towels to dry. Repeat the same procedure for the other two-thirds of the cotton. Allow the cotton to dry completely (2-3 days) before attempting to light it. Procedure B:

This procedure is slower but produces a more pure form of nitrocellulose. Begin by slowly mixing two equal volumes of acid together in a 500mL beaker and allow the acid solution to cool to room temperature in an ice bath. Place 5g of cotton into the beaker and submerge completely. Allow the cotton to soak for approximately 30 minutes with occasional stirring. Remove the cotton and squeeze out as much excess acid as possible. Submerge the cotton directly into an ice bath and stir for two minutes to dilute the residual acid in the cotton. Wash the cotton for two more minutes in running water, and then place the cotton in a boiling water bath. After 30 minutes, move the cotton to a fresh boiling-water bath. Three separate, boiling-water baths should be used for 30 minutes each. The purpose is to remove all of the acid from the cellulose. Test the last water bath for acidity, and if necessary, boil it again. If the acid is no longer apparent, rinse the cotton with cold water, squeeze dry, and lay flat on paper towels for 2-3 days to dry.

To light, place the nitrocellulose in the Pyrex dish. Using a long candle, touch the bottom of the nitrocellulose with a flame. A small amount of smoke and ash will be produced after burning.

It is recommended to oven dry the cotton for several hours prior to this procedure. DO NOT OVEN DRY THE COTTON AFTER THIS PROCEDURE.

An interesting variation is to ignite the cotton in a large (3-4L) beaker. The smoke with be contained within the beaker and will continue to burn for a short time after the nitrocellulose is gone.

Theory: Nitrocellulose or guncotton is cellulose trinitrate.

CH

CH

CH CH

CH

O

O2NO ONO 2

H2C

ONO 2

n The nitrate groups have a highly oxidized nitrogen atom, which leads to extreme

instability and generates an extremely exothermic decomposition reaction: 4C6H7O2(NO3)3 + 9O2 à 24CO2 + 14H2O + 6N2

Other products of incomplete combustion are also produced, including NO and CO. In this demonstration, the guncotton is not explosive because it is not contained. Safety Considerations:

STRONG ACID SMOKE HAZARD- The acid fumes are noxious. FIRE HAZARD- The cotton may create a small amount of floating ash.

Disposal: Do not store nitrocellulose for long times or in large amounts and always burn it

before disposing. Neutralize the acids with a weak base and flush down the drain.

Permanganate Combustion Principles Illustrated:

-Reduction/oxidation (redox) chemistry -Spontaneous combustion -Excitation and emission

Required Equipment:

-Large Pyrex dish -Disposable pipette

Required Chemicals:

-Potassium permanganate (KMnO4)

-Glycerin (C3H8O3) or Brake fluid (various petrol ethers and esters) Demonstration Description: A small amount of liquid is added to a pile of powder. After several seconds the powder begins to smoke and eventually bursts into a bright purple flame. Demonstration Procedure: Weigh out 5-10 grams of potassium permanganate. Make a small impression in the powder to provide room for the liquid. Add enough liquid to fill the impression. Step back and allow the reaction to proceed. The potassium permanganate is typically in excess in this reaction. For the best results a large deep impression should be made to allow as much liquid as possible to be involved in the reaction. Brake fluid works much better than the glycerin. The time to initiate the reaction is much shorter and the flame is much more intense. The advantage of the glycerin is that the reaction is more easily defined. Theory: This reaction illustrates the principle of spontaneous or chain reactions. The heat builds up during the reaction until enough heat is present to generate a flame. The reactions involved are not well characterized, but a probable reaction for the glycerin would be:

14KMnO4 + 4C3H8O3 à 7K2CO3 + 7MnO2 +5CO2 + 16H2O. Potassium permanganate is inherently unstable because of the extreme cationic state of the manganese. The oxygen atoms effectively remove a total of seven electrons

from the manganese, leaving it in a +7 oxidation state. The large positive nature of this atom allows for high-energy electrostatic interactions with the electrons of the oxygen atoms. In the presence of a reduced substance like glycerol or brake fluid, lower energy bonds replace these high-energy manganese-oxygen interactions and the differential energy is emitted as light and heat. Safety Considerations: FIRE HAZARD

Potassium permanganate and the products of this reaction will stain your skin and clothing and should be treated with care. Brake fluid should be treated as a strong solvent and direct exposure should be avoided.

Disposal: Discard the insoluble residue in a waste container. The unreacted permanganate and soluble products can be flushed down the drain.

Purple Fire Principles Illustrated:

-Reduction/oxidation (redox) chemistry -Photon emission

Required Equipment:

-Mortar and pestle -Pyrex dish

Required Chemicals:

-Cupric sulfate (CuSO4) -Potassium chlorate (KClO3) -Sulfur (S or S8)

Demonstration Description: A small pile of powder is ignited and immediately bursts into a bright bluish-purple flame. Demonstration Procedure: For the best results mix together equal parts of cupric sulfate, potassium chlorate, and sulfur. The cupric sulfate and potassium chlorate should be finely ground SEPARATELY and mixed after grinding. Add the sulfur as the last step and mix carefully to break up any sulfur clumps. This product is easily ignited and produces a very intense bluish-purple flame as well as a fair amount of smoke. Anhydrous pentahydrate cupric sulfate will work and is much easier to grind. Sulfur is messy and should not be ground. Theory: These compounds undergo an oxidation/reduction reaction where the sulfur is oxidized and the sulfate and chlorate ions are reduced. Multiple products are formed and the exact chemical equation is not known. Assuming that the majority of the redox chemistry occurs between the potassium chlorate and sulfur, and the other chemicals are mainly responsible for the color, a reasonable reaction is as follows:

4KClO3 + 6S + O2 à 2K2O + 6SO2 + 2Cl2

The sulfate ion can be reduced to form sulfur dioxide and at high temperatures it is possible that this reaction occurs to some degree.

The high energies of this reaction result in the excitation of electrons, which emit light upon returning to the ground state. The net effect of the various light emissions from these elements will result in the observable bluish-purple color. Purple is typically observed when potassium is burned and blue is associated with burning copper. Safety Considerations: FIRE HAZARD SMOKE HAZARD- Sulfur dioxide is particularly unpleasant. Do not store these mixtures for extended periods of time in any quantity.

Disposal: Any remaining product can be rinsed and washed down the drain. Multiple salts are deposited following the flame and are difficult to remove.

Red Fire Principles Illustrated:

-Reduction/oxidation induced combustion -High energy of molecules in a living system -Electron excitation and photon emission

Required Equipment:

-Mortar and pestle -Pyrex dish

Required Chemicals:

-Strontium nitrate (SrNO3) -Potassium chlorate (KClO3)

-Shellac Demonstration Description: A small pile of powder is ignited and immediately bursts into a bright red flame. Demonstration Procedure: For the best results mix together 4 grams of strontium nitrate, 1 gram of potassium chlorate and 1 gram of shellac. All three components should be finely ground SEPARATELY and mixed after grinding. The shellac is especially difficult to grind, but should be broken up into small pieces. This product is fairly easily ignited and produces a very intense red flame as well as a fair amount of smoke. Ignition of the powders can be problematic, but we found a candle attached to a meter stick to be quite effective. An alternative method is to use a pea-sized amount of 2 parts potassium chlorate mixed with 1 part of table sugar and place this in the center of the reaction powder. A drop of sulfuric acid will ignite the potassium chlorate/sugar mixture. Theory: One of the most common and important classes of reactions are reduction/oxidation (redox) reactions. In a redox reaction electrons are transferred from one atom or molecule to another. The atom or molecule that loses the electron is oxidized while the atom or molecule that gains an electron is reduced. Perhaps the most familiar redox reaction is the combustion of hydrocarbons as fuel. The basic unbalanced reaction proceeds as follows:

CnHn + O2 + à CO2+H2O In this reaction the carbon on the hydrocarbon is highly reduced, but upon reacting with oxygen, it is oxidized. It can also be said that it acts as the reducing agent. Oxygen, on the other hand, gains electrons from the reaction. It is known as the oxidizing agent. In a combustion reaction such as the one outlined, a high amount of energy is given off and a characteristic flame is seen. This energy release accounts for the common use of reducing agents as fuels.

Shellac is a resin composed of a complex mixture of aliphatic (long hydrocarbon chain) and alicyclic (cyclic hydrocarbon chain) acids. Derived from the insect Laccifer

lacca, shellac is harvested in India and used primarily in lacquers and varnishes. The main components of shellac are shellolic acid and aleuritic acid, composing 5-8% and 43% of shellac respectively.

H

HOOC

COOH

HO

CH2OH

CH3 Shelloic Acid

HOCH2(CH2)5CHCH(CH2)7COOH

OH

OH Aleuritic Acid

Shellac is a reducing agent. Gasoline, propane, fats, and even sugars are also reducing agents. These products are usually derived from living systems. As mentioned before, shellac is derived from insects, while fats and sugars are derived from animals and plants. Fuels such as gasoline and propane are derived from the remains of prehistoric plants and animals. Organisms generally pump huge amounts of energy into the conversion of small, simple, oxidized building blocks into larger, more complex reduced molecules such as sugars and fats. These compounds are used for energy storage, and the huge amount of energy stored in these molecules can be observed upon the burning of fossil fuels and even in the maintenance of the energy needs of our bodies when we eat. The demonstration outlined in this procedure dramatically shows the large amount of energy released upon the oxidation of shellac.

Three compounds used in this demonstration undergo a change in reduction or oxidation. The shellac is oxidized and the nitrogen and chlorine atoms are reduced. Multiple products are formed and the exact chemical equation is not known, although the following unbalanced equation is a reasonable possibility:

Sr(NO3)2 + KClO3 + shellac (CnHnOn) + O2 à KCl + CO2 + H2O + Sr(OH)2 + N2. The high energies of this reaction result in the excitation of the electrons

whereupon their return to the ground state will result in the emission of light. The red color is observed due to the excitation of the valence electrons in the strontium metal. The photon emitted has a wavelength that stimulates our eyes to see red. Safety Considerations:

FIRE HAZARD SMOKE HAZARD Do not store these mixtures for extended periods of time in any quantity. Disposal: Any remaining product can be rinsed and washed down the drain.

Smoke Screen Principles Illustrated:

-Reduction/oxidation (redox) reaction -High energy of sugar -Electron excitation and photon emission

Required Equipment:

-Mortar and pestle -Pyrex dish -Spatula -Fuse

Required Chemicals:

-Potassium nitrate or sodium nitrate (KClO3/NaClO3) -Sugar (C12H22O11)

Demonstration Description: A small pile of powder violently bursts into flames and releases a copious amount of black smoke. Demonstration Procedure: Use the mortar and pestle to grind the potassium nitrate to a smooth consistency. Gently mix in the sugar using a spatula until the two powders are evenly distributed together. Do not grind the powders together with the pestle. Place the mixed powder in the center of the Pyrex dish, add the fuse and ignite. This reaction generates an intense flame and large amounts of heat and smoke. The best mass ratio for this demonstration is 2 parts sugar mixed with 3 parts potassium nitrate. Five grams of total mixture is ample for a large audience. 10 grams of total mixture produce the maximum amount a smoke that a typical fume hood can adequately remove. Several different fuse options are available. By making a small divot in the potassium nitrate/sugar pile and adding a pea sized amount of potassium chlorate/sugar mixture, you can effectively ignite the whole pile when you touch the chlorate mixture with a small drop of sulfuric acid. Cannon fuse can be purchased commercially and works well for large demonstrations. Powdered sugar works slightly better than granulated sugar because the granulated sugar has a greater tendency to splatter during the reaction. Theory: Potassium nitrate is a very strong oxidizing agent and is commonly used in various forms of gunpowder. The general formula is:

4KNO3 + 2C12H22O11 + 19O2 à 24CO2 + 20H2O + 2N2 + 4KOH This formula is what would occur under ideal conditions but the production of CO, NO, NO2, elemental carbon and other side products is highly probable. The redox changes in a wide variety of elements make this experiment very useful for illustrating reduction/oxidation principles.

Safety Considerations: SMOKE HAZARD

FIRE HAZARD Be sure that the Pyrex dish is wide and shallow as sufficient heat is generated to

crack the Pyrex should the flame directly contact the sides of the container. The force of the reaction and the accompanying flame is sufficient to splatter melted sugar up to several feet from the source of the reaction. Disposal: All of the products can be washed down the drain.

Yellow Fire Principles Illustrated:

-Reduction and oxidation (redox) chemistry -Photon emission

Required Equipment:

-Mortar and pestle -Pyrex dish

Required Chemicals:

-Barium nitrate (BaNO3)

-Potassium chlorate (KClO3)

-Sulfur (S or S8) -Sodium oxalate

(Na2C2O4) -Sodium bicarbonate

(NaHCO3) Demonstration Description: A small pile of powder is ignited and immediately bursts into a bright yellow flame. Demonstration Procedure: For the best results mix together 8 grams of potassium chlorate, 1 gram of sodium oxalate, 1 gram of barium nitrate, 1 gram of sodium bicarbonate and 4 grams of sulfur. The first four components should be finely ground SEPARATELY and mixed after grinding. Add the sulfur as the last step and mix carefully. This product is easily ignited and produces a very intense yellow flame as well as a fair amount of smoke. Theory: These compounds undergo an oxidation/reduction reaction where the sulfur and oxalate are oxidized and the nitrate and chlorate ions are reduced. Multiple products are formed and the exact chemical equation is not known. The majority of the redox chemistry occurs between the main two products, potassium chlorate and sulfur, and the other products are primarily responsible for the color:

4KClO3 + 6S + O2 à 2K2O + 6SO2 + 2Cl2

Various nitrogen oxides are formed from the barium nitrate as well as carbon dioxide from the oxalate and carbonate compounds.

The high energies of this reaction result in the excitation of electrons to higher energy levels. Upon their return to the ground state, light is emitted. The net effect of the various light emissions from these elements will result in the observable yellow color, classically associated with both sodium and sulfur.

Safety Considerations: FIRE HAZARD SMOKE HAZARD- Sulfur dioxide is particularly unpleasant. Do not store these mixtures for extended periods of time in any quantity. Disposal: Any remaining product can be rinsed and washed down the drain. Multiple salts are deposited following the flame and are difficult to remove.

Production of Zinc Sulfide Principles Illustrated:

-Reduction/oxidation (redox) reaction -Electron excitation and photon emission

Required Equipment:

-One 10-15” large gauge wire -One Pyrex dish -Bunsen burner

Required Chemicals:

-Zinc powder -Sulfur powder

Demonstration Description: A heated wire is added to a small pile of powder. A large flash of light and copious amounts of smoke are produced. Demonstration Procedure: Thoroughly mix 6g of zinc dust with 1g of sulfur powder. Place the combined powder in a small pile in the middle of the Pyrex dish. Heat the wire to red hot in the burner and touch it to the pile of powder. Theory: This is a simple reduction/oxidation reaction. Several products are formed:

Zn + S à ZnS (∆Hof = -206kJ/mole)

2Zn + O2 à 2ZnO (∆Hof = -348.3kJ/mole)

S + O2 à SO2 (∆Hof = -297.04kJ/mole)

All three of these reactions are highly exothermic and generate a large amount of energy. Safety Considerations: FIRE HAZARD SMOKE HAZARD Because the reaction is so violent, much of the product is thrown out of the Pyrex dish. Disposal: The products can be washed down the drain or placed in a waste container.

Disappearing Styrofoam Principles Illustrated:

-Non covalent bonding of nonpolar molecules -Solubility

Required Equipment:

-Large beaker or Pyrex dish Required Chemicals:

-Acetone (C2H6O) -Styrofoam (polystyrene- [C14H16]n)

Demonstration Description: A large amount of Styrofoam is quickly “dissolved” in a small layer of liquid. Demonstration Procedure: Fill the beaker or dish with 1-2cm of acetone. Place the Styrofoam in the acetone and it will steadily melt until it is gone. Large amounts of Styrofoam can be consumed quickly. Any source of Styrofoam will work, although Styrofoam cups form a unique taffy-like substance that can be removed from the acetone and manipulated more easily than the final products of other Styrofoam sources. The Styrofoam will harden quite nicely if the acetone is allowed to completely evaporate. Theory: Styrofoam is composed of polystyrene molecules that are packaged in an ordered structure. Polystyrene is a vinyl polymer of styrene. Structurally, it is a long hydrocarbon chain, with a phenyl group attached to every other carbon atom. The phenyl groups will have hydrophobic interactions and pi-orbital stacking interactions with the phenyl groups of the same strand as well as adjacent strands. Acetone acts as a great solvent to disrupt the non-covalent interactions that allow the individual strands of polystyrene to stack together. It is important to note that acetone does not dissolve, depolymerize or melt Styrofoam; it simply disrupts the networking between adjacent polymers. The structure

of styrene and polystyrene are shown:

CHH2C

Styrene

HC CH2 CHH2C

Polystyrenen

Safety Considerations: VOLATILITY HAZARD

Acetone is carcinogenic and should not be handled without gloves. Acetone fumes can be noxious and this demonstration should be performed in a well-ventilated area. Disposal: Acetone should be disposed of in an organic waste container. The Styrofoam can be discarded after the acetone has evaporated.

Nylon 6-10 Principles Illustrated:

-Polymerization -Hydrogen bonds -Amide bond formation

Required Equipment:

-250ml beaker or large petri dish -Tweezers -One large test tube

Required Chemicals: -1,6-hexanediamine

(H2N(CH2)6NH2) solution -Sebacoyl chloride

(ClCO(CH2)8COCl) solution Demonstration Description: Using a pair of tweezers, a thin sheet of Nylon is pulled from solution and stretched over a test tube. The Nylon is continuously produced by the solution as the test tube is rotated and the Nylon is wound onto the test tube. Demonstration Procedure: To prepare the 1,6-hexanediamine solution, dissolve 3.0g of 1,6-hexanediamine and 1.0g of NaOH in 50ml distilled water. The 1,6-hexanediamine can be dispensed by placing the reagent bottle in hot water until sufficient solid has melted and can be decanted. The melting point is 39-40oC. Several drops of food coloring may be added to this solution to increase the visibility of the solvent interface.

To prepare the Sebacoyl chloride solution, dissolve 1.5ml to 2.0ml Sebacoyl chloride in 50ml hexane.

Wearing gloves, place the 1,6-hexanediamine solution in a 250ml beaker or large petri dish. Slowly pour the Sebacoyl chloride solution as a second layer on top of the 1,6-hexanediamine solution, taking care to minimize agitation at the interface. With forceps, grasp the polymer film that forms at the interface of the two solutions and pull it carefully from the center of the beaker. Wind the polymer thread on a large test tube.

Wash the polymer thoroughly with water or ethanol before handling. Theory: Nylon was developed as a synthetic substitute for silk and is a group of linear polymers with repeating amide linkages along the backbone. These are produced by an amidation of diamines with dibasic acids. Nylon is strong and tough. It resists abrasion, fatigue and impact. Nylon offers excellent chemical resistance when used with organic solvents. However, it has poor resistance to strong mineral acids, oxidizing agents and certain salts.

The word "nylon" is used to represent synthetic polyamides. The various nylons are described by a numbering system that indicates the number of carbon atoms in the monomer chains. Nylons from diamines and dicarboxylic acids are designated by two numbers, the first representing the diamine and the second the dicarboxylic acid. Thus nylon 6-10 is formed by the reaction of 1,6-hexanediamine and Sebacoyl chloride. The equation is:

H2N(CH2)6NH2 + ClC(CH2)8CCl

O O

N(CH2)6 N C (CH2)8C

H H

O O

+ 2HCl

Amide group

Notice the amide groups in the structure. The amide groups are what make nylon such a good fiber. They are strongly attracted to each other. When nylon is spun into fibers, the long chain-like macromolecules line up parallel to each other. The amide groups on adjacent chains then form strong bonds with each other called hydrogen bonds, which hold the adjacent chains together, making nylon yarn strong.

N(CH2)6N C(CH2)8C

H H

O O

N(CH2)6N C(CH2)8C

H H

O O

N(CH2)6N C(CH2)8C

H H

O O

N(CH2)6N C(CH2)8C

H H

O O

N(CH2)6N C(CH2)8C

H H

O O

N(CH2)6N

H H

C(CH2)8C

O O

HydrogenBonds

Safety Considerations:

VOLATILITY HAZARD 1,6-hexanediamine is irritating to the skin, eyes, and respiratory system. Sodium

hydroxide is extremely caustic and can cause severe burns. Sebacoyl chloride is corrosive and irritating to the skin, eyes, and respiratory system. Hexane is extremely flammable. Hexane vapor can irritate the respiratory tract and, in high concentrations, be narcotic. Disposal:

Any remaining reactants should be mixed thoroughly to produce nylon. The solid nylon should be washed before being discarded in a solid waste container. Any remaining liquid should be discarded in a solvent waste container

Silly Puddy Principles Illustrated:

-Polymerization -Elastomers

Required Equipment:

-Spatula -Small beaker -Disposable syringe (>10mL)

Required Chemicals:

-Sodium borate (Na2B4O7-10H2O) -All-purpose white glue (~Elmer’s), non-washable -Food coloring

Demonstration Description: Two solutions are mixed thoroughly to form silly puddy. Demonstration Procedure: A 4% solution of sodium borate is prepared by dissolving 40 grams of sodium borate into 1 liter of water. The solution will need to be heated for >30 minutes to completely dissolve the borate. An autoclave works well with a 40-minute cycle. This should be done in a large Erlenmeyer flask that is 3 to 4 times the volume of the liquid (3-4 L). After the borate solution cools to room temperature several drops of food coloring can be added. The sodium borate solution will last indefinitely and can be prepared well in advance. The 55% solution of glue is made by adding 11mL of glue to 9mL of water. This solution should be prepared just prior to or during the demonstration. The water should be measured into a small beaker and the glue can be measured out and injected directly into the water using a large disposable syringe. Mix the water and glue until the glue is completely dissolved. Add 10mL of 4% sodium borate solution to 20mL of the 55% glue solution. Mix thoroughly with the spatula for several minutes. After several minutes, the polymer can be taken out of the beaker and kneaded to the proper consistency. Theory: If a substance springs back to its original shape after being twisted, pulled, or compressed, it is most likely a type of polymer called an elastomer. The elastomer has elastic properties (i.e., it will recover its original size and shape after being deformed). An example of an elastomer is a rubber band or a car tire. This experiment will create an elastomer.

The elastomeric properties of our silly puddy-like polymer are based on the polymer strand stacking interactions between adjacent polymers. The liquid latex (Elmer's glue) used in this experiment contains multiple hydrocarbon polymers. These hydrocarbon polymers stack together, forming noncovalent interactions as a result of adding sodium borate to the solution:

OB

OB

OB

OB

O

O- O-Na+Na+

Sodium borate will form various types of noncovalent interactions with the latex

hydrocarbons. These interactions can be broken and reformed very rapidly. This allows the linear polymer chains to slide past one another, giving the puddy elastic properties. Safety Considerations: Sodium borate is a weak bleaching agent and hands should be washed after handling the product. Disposal: The silly polymer can be thrown away.

Slime Principles Illustrated:

-Polymerization -Hydrogen bonds -Non-Newtonian fluids

Required Equipment: -One large beaker -Two large Erlenmeyer flasks -One small graduated cylinder

-Food coloring *Actual sizes of the glassware are discussed later

Required Chemicals: -Polyvinyl alcohol (CHOH)n -Sodium borate (Na2B4O7-10H2O) Demonstration Description: Two liquids are poured together to form a gelatinous slime. Demonstration Procedure: A 4% solution of polyvinyl alcohol is prepared by dissolving 40 grams of powdered polyvinyl alcohol into 1 liter of water. The solution will need to be heated for >30 minutes to completely dissolve the polyvinyl alcohol. An autoclave works well with a 40-minute cycle. This should be done in a large Erlenmeyer flask that is 3 to 4 times the volume of the liquid (3-4L). A 4% solution of sodium borate is prepared by dissolving 40 grams of sodium borate into 1 liter of water. Heating the solution on a hot plate will facilitate the dissolution of the sodium borate. After the borate solution cools to room temperature several drops of food coloring can be added. The two liquids polymerize upon contact at room temperature. Measure out 200mL of 4% borate solution and 1L of the 4% polyvinyl alcohol solution. Simultaneously pour the liquids into a large beaker and mix until uniform. Theory: The polyvinyl alcohol is a long chain of repeating CH2CHOH subunits. The –OH dipoles on the borate ion are attracted to the –OH dipoles on the polyvinyl alcohol. These dipoles form a bridging hydrogen bond that can be broken and reformed very rapidly. This allows the linear polymer chains to slide past one another, giving the gel liquid-like properties. Polyvinyl slime is a non-Newtonian fluid. Non-Newtonian fluids expand under stress and this gives the slime several interesting properties worth exploring. Under low stress it flows and stretches. High stress will cause the material to break. It firms up when struck and bounces off hard surfaces. Safety Consideration:

The slime will dry out after two days and is not contact sensitive. Avoid ingestion. Polyvinyl alcohol will begin to mold if stored too long, so only prepare enough for one use. The borate solution should last indefinitely. Disposal: The slime can be thrown away when you are done.

Acid Breath Principles Illustrated:

-Solubility of carbon dioxide -Formation of carbonic acid

Required Equipment:

-Drinking straw -Large beaker -Medicine dropper

Required Chemicals:

-Phenolphthalein solution (C20H14O4) -6M Sodium hydroxide (NaOH)

Demonstration Description: By blowing through a straw, into a beaker full of a pink solution, the solution slowly becomes clear. Demonstration Procedure: To prepare the phenolphthalein solution, add 1 gram of solid phenolphthalein to 60mL of alcohol and dilute to 100mL with water.

Fill the beaker approximately 2/3 full of water. Place one drop of 6M NaOH into the water. Add about five drops of the phenolphthalein while mixing until the water is sufficiently pink. Blow into the solution continuously through the straw. The liquid should become clear in about 30 seconds.

By adding additional drops of NaOH, the experiment can be repeated several times.

The color density can be enhanced by addition of phenolphthalein or NaOH, but be careful not to add too much NaOH or it could take several minutes of blowing to change the color. Adding excessive amounts of phenolphthalein will cause the phenolphthalein to precipitate out of solution.

Theory:

The sodium hydroxide solution in the beaker is basic and raises the pH sufficiently to cause the phenolphthalein indicator to turn pink. Carbon dioxide from your exhaled breath will partially dissolve in water to form carbonic acid. Carbonic acid will neutralize the NaOH and lower the pH. When the solution drops to a pH range from 8.2-10 (depending on the ionic nature of the solution) the phenolphthalein will become clear.

The dissolution of carbon dioxide to form carbonic acid: CO2 + H2O à H2CO3

The neutralization of sodium hydroxide by carbonic acid:

H2CO3 + 2NaOH à Na2CO3 + 2H2O

Safety Consideration: STRONG BASE

Be sure that the straw being used is sufficiently long to prevent splashing the solution into your face. Disposal: The pink solution should be sufficiently neutralized to flush down the drain.

Cabbage Principles Illustrated:

-Acid/base indicators Required Equipment:

-Blender -Strainer -Several 200mL beakers -One large beaker -Large kitchen knife

Required Chemicals:

-One red cabbage -Various household chemicals (Baking soda, Clorox, Ammonia, Vinegar, etc.)

Demonstration Description: The juice from a red cabbage turns from purple to red or blue when various household chemicals are added. Demonstration Procedure: Cut up half of a regular red cabbage and place the pieces in a blender. Add ~1 liter of water and blend. Pass your cabbage puree through a strainer to collect the juice in a large beaker and dilute the juice until it is barely translucent. Pour an equal amount of strained juice into several beakers. Add small amounts of various household chemicals to each beaker. A basic solution will turn the cabbage juice from purple to blue and an acidic solution will turn it from purple to red. At a pH > 13 the solution turns clear, most likely due to the decomposition of the indicator molecule. Theory: All indicators work on the principle of a color change associated with a specific pH range. As pH raises or lowers the indicator molecule will change its level of hydrogen ion association and thereby change its chemical structure. The indicator found in red cabbage is anthocyanin. Many berries, blossoms and flowers also contain variations of anthocyanin. Anthocyanins are responsible for many of the blue, red and purple colors in plants. The anthocyanin molecule is chemically different in an acid as opposed to a base. This difference in chemical composition will also change the color of light that is absorbed by the molecule.

O

O

O

O

O

O

O

H

H

HH

H

H

H

H

H

H

+O

O

O

O

O

O

H

H

HH

H

H

H

H

H

H

OH

H+

H+

Red form in AcidBlue form in Base

Safety Considerations: Care should be taken to avoid direct contact with many household chemicals. Disposal: All of these chemicals can be flushed down the drain SEPARATELY.

Collapsing Can Principles Illustrated:

-Gas Laws Procedure A Required Equipment:

-Hot plate or Bunsen burner -Several soda pop cans -Crucible tongs

Demonstration Description: A soda pop can is inverted into a beaker of water and immediately collapses. Demonstration Procedure:

Add ~20mL of water to an empty soda pop can and bring the water to a steady boil. Before the water completely evaporates, invert the can into cold water with the tongs. Procedure B Required Equipment:

-Hot plate or Bunsen burner -A large metal can with a small opening and an airtight lid -Insulated gloves

Demonstration Description: The lid is placed over the opening of a large can and the can is inverted into a tub of ice water. Within minutes the can collapses. Demonstration Procedure: Fill the large can with water until the water is about 1 cm deep. Bring the water to a steady boil. Before the water completely evaporates, remove the can from the heat source, put the lid tightly on the can and quickly invert the can in a tub of ice water. Theory: As the temperature of the can drops, the pressure will drop proportionally:

P(pressure) x V(volume) = n (moles) x R (constant) x T (temperature) The decrease in pressure on the inside of the can will cause of large pressure difference between the inside of the can and the outside of the can. When this difference becomes sufficiently large, the walls of the can will give in to the external pressure. Safety Considerations: If you are using a large metal can that previously contained a flammable material, be sure to rinse the can thoroughly before performing the experiment. Large amounts of heat and pressure differential are generated in this experiment- BE CAREFUL! Disposal: The cans can be taken to a recycling center.

Geyser Principles Illustrated:

-Gas laws -Density

Required Equipment:

-One 1000mL Florence flask -A one-hole stopper with a glass pipette and rubber tube -One large beaker

Required Chemicals:

-Food coloring -Glycerin (optional)

Demonstration Description: Water from a large beaker rushes up a length of tubing to fill an inverted flask. Demonstration Procedure:

Add 20mL of water to the 1000mL Florence flask. Assemble the tubing by inserting a glass pipette through the stopper and slide the tubing over the external portion of the glass pipette. Insert the stopper into the Florence flask. Prepare your water by adding several drops of food coloring to ~1L of water in a large beaker. Place the Florence flask on a hot plate. When the water in the Florence flask is boiling, insert the free end of the rubber tubing into the colored water and invert the flask.

Glycerin can be used as a lubricant to facilitate the assembly of the apparatus.

Theory:

The Florence flask is filled with steam by the boiling water. As the steam condenses there will by a ~1000 fold

decrease in volume. This decrease in volume will create a vacuum that will draw in the water from the other beaker to fill the void space.

Another way of explaining this experiment is to consider the Universal Gas Law: P(pressure) x V(volume) = n (moles) x R (constant) x T (temperature)

As the temperature drops, so will the pressure. This drop in pressure allows for the water in the beaker to enter the flask and equalize the pressure by adding more moles to the system.

Calculations:

The density of steam is ~22.4L/mole, so a 1000mL flask will hold around 0.04 moles of steam:

1 mole steam/22.4L steam x 1L/1000mL x 1000mL in flask = 0.04 moles steam When the flask is removed from the heat source the 0.04 moles of steam will

condense to form water. The density of water is ~18mL/mole, so the 0.04 moles of newly formed water will occupy less than a 1mL volume:

1mL water/1gram water x 18 grams/mole (H2O) x 0.04 moles water = 0.7mL water This leaves 999mL of empty space that will be filled by the water drawn in from

the other beaker.

Safety Considerations: The demonstrator should wear oven mittens while doing this experiment because of the high temperature of the flask and tubing.

Sinking Bobbers Principles Illustrated:

-Density and water displacement -Gas laws

Required Equipment: 4 disposable plastic Pasteur pipettes 4 10-32 hex nuts 1 two-liter, clear soda pop bottle Scissors A 25-50mL graduated cylinder Analytical scales A candle and matches Demonstration Description: Four bobbers are floating inside of a two-liter soda pop bottle. As you squeeze the bottle the bobbers drop one by one. Demonstration Procedure: Slide the hex nuts up the Pasteur pipette until they sit just below the bulb. Make sure the hex nuts fit tightly. Cut off the spout of the Pasteur pipette leaving about 1 cm below the nut. Weigh the Pasteur pipette bobbers and label them. Find out how much water each bobber displaces by filling them completely with water and dropping them into a half filled graduated cylinder. The difference in water level of the graduated cylinder before and after inserting the bobbers is the water displacement value. Find out the total volume that each bobber will hold. Fill the bobber completely with water and weigh it. Remove some of the water from each of the bobbers to allow air into the bobber. You need to know how much water you remove from each. This can be done by directly weighing the water removed or weighing the bobber before and after the water is removed. To know how much water to remove involves some guesswork. One to two milliliters of air space is a good range to start with. You want to vary the amount of air in each bobber so that they will sink at different times. Fill your 2 liter bottle up about 3/4ths way and add the bobbers. Cap the bottle and squeeze. It works best if the first two bobbers are fairly easy and the last two are difficult. Theory: Objects stay afloat when the weight of the object is less than the weight of the water being displaced. As you squeeze the pop bottle the air space inside the bobber decreases and there is less water displaced. When the air has been sufficiently compressed the bobber will sink. This is an excellent example of the pressure and volume relationships defined by Boyle’s law. Calculations:

We need to know how much air is in each bobber. Because water density is about 1 gram per milliliter, you can assume that the grams of water removed equal the milliliters of air in the bobber. We also need to know how much water the bobber apparatus itself displaces. This is equal to the difference between the water displacement value and the volume of water the bobber holds. This should be around 1 milliliter. Now sum the air volume and the water displacement of the bobber itself. Subtract the weight of the bobber from that sum. This is the excess air value of the bobber. In order for the bobber to sink, this value must be zero. Divide the excess air value by the total air volume of the bobber. This is the fractional decrease of the volume that must occur for the bobber to sink. To know the pressure being applied to the sinking bobber, multiply the inverse of the fractional decrease in volume by atmospheric pressure. Atmospheric pressure will change day to day, but tends to be less at higher elevations. Find out the local atmospheric pressure using a barometer or the Internet. The pressure you are actually applying to the bobber is equal to the difference between the pressure applied to the sinking bobber and the local atmospheric pressure. You can convert that to pounds per square inch by multiplying atmospheric pressure by 14.7.

Siphon Fountain Principles Illustrated:

-Siphoning -Vacuum pressure

Required Equipment:

-One 2L round bottom flask -Two large beakers -One two-hole stopper -Rubber tubing -Glass pipettes

Required Chemicals:

-Food coloring -Glycerin -Sodium hydroxide [NaOH] (optional) -pH indicator (optional)

Demonstration Description: Water will rise from the upper beaker into the inverted flask and then travel back down to the lower beaker. Demonstration Procedure:

Prepare the stopper by inserting the two glass pipettes into the stopper in opposite directions as shown in the diagram. Use the glycerin as a lubricant. Slide the rubber tubing over the ends of both pipettes. Fill the flask with water and insert the rubber stopper. Fill the upper beaker with water, but do not submerge the tubing in the upper beaker. Invert the flask and hold it above the upper beaker, making sure the tubing to the lower beaker is inside the container. Let most of the water drain from the flask and then submerge the tubing into the upper beaker. The water will rise up to the flask and fall back down to the lower beaker.

The best effect can be generated by breaking off the end of the glass pipette exiting the flask while leaving the glass pipette entering the flask drawn to a point.

For added effect two upper beakers can be used with a third hole in the stopper and more glass tubing. Making one solution basic, such as a

1M sodium hydroxide solution, and making the other a solution of phenolphthalein, or another indicator, the two clear solutions will mix together to form a red one.

Theory: The water moving out of the tube and into the lower beaker will create a vacuum that can be used to draw the water from the upper beaker. As long as the weight of the water in the lower tube is greater than the weight of the water in the upper tube, the fountain will continue to run. The purpose of using the narrow glass pipette on the entrance to the flask is to decrease the cross-sectional area of the inlet tubing. The amount of water entering the flask will be the same as the amount of water exiting the flask, so a narrower cross-sectional area on the entrance tube will increase the rate of water coming in, creating a “shooting” stream. Safety Considerations: STRONG BASE

Be careful when handling glassware. Disposal: The sodium hydroxide should be neutralized with a weak acid and flushed down the drain.

Alchemy Principles Illustrated:

-Reduction/oxidation Required Equipment:

-Hot plate -Tweezers -250mL beaker -Bunsen burner

Required Chemicals:

-5M sodium hydroxide solution -Zinc dust -Several copper pennies -Steel wool

Demonstration Description: A penny is placed into a boiling solution for several minutes. It is removed after turning silver. After holding the penny in a Bunsen burner flame for several seconds, it turns gold. Demonstration Procedure: Using a 250mL beaker, prepare 50mL of 5M sodium hydroxide solution by dissolving 10g of sodium hydroxide pellets in water to a final volume of 50mL. Place this solution on a hot plate and add 0.5 grams of zinc dust. Allow the solution to boil until the zinc dust aggregates into a small clump. Add no more than 10 pennies to the boiling solution and leave them in the solution until they are visibly and thoroughly coated in zinc. Remove the pennies and rinse thoroughly before handling without gloves. Be sure to rinse off any zinc clumps that may have stuck to the pennies. After the pennies have been rinsed and cleaned, hold them over a Bunsen burner with the tweezers for several seconds. The penny will change to a mottled copper-gold color. Remove the penny from the flame and set it in a beaker full of water. It should immediately turn gold upon cooling. The largest problem associated with this demonstration is melting the penny during the final bronzing process. Carefully heat the penny until the first indications of color change begin, immerse the penny in water and reheat only if patches of silver remain. Cleaning the penny beforehand with steel wool will allow for a more homogenous final product. Theory: Zinc dissolves in sodium hydroxide to form sodium zincate (Na2ZnO2) and hydrogen gas:

Zn + 2NaOH à Na2ZnO2 + H2 The zincate ions are then reduced by the copper in the penny to metallic zinc, which coats the penny:

Na2ZnO2 + Cu à Na2CuO2 + Zn

Copper still remains underneath, protected from further reaction by the zinc coat. Zinc is commonly used as a redox protective metal coat. Heating the surface allows the metals to diffuse together to form a brass alloy. Safety Considerations: STRONG BASE

SMOKE HAZARD - Boiling sodium hydroxide will produce choking fumes. Hot metals and the hot plate should only be handled with oven mittens.

Disposal: The sodium hydroxide should be neutralized with a weak acid and flushed down the drain. The residual zinc should be rinsed and disposed of in a solid waste container.

Colorful Cloth Principles Illustrated:

-Lewis acid/base -Coordination chemistry

Required Equipment:

-Three large beakers -A white cloth

Required Chemicals: -0.02M potassium ferrocyanide (K4Fe(CN)6-3H2O) -0.002M potassium thiocyanate (KSCN) -0.05M ferric chloride (FeCl3-6H2O) Demonstration Description: The white cloth will change colors from yellow to red-orange to blue. Demonstration Procedure: Fill each beaker about half full with one of the solutions above. Submerge the white cloth into the ferric chloride solution and it will turn a faint yellow color. Transfer it to the potassium thiocyanate solution and it will turn red-orange. Finally, transfer it to the potassium ferrocyanide solution where it will turn blue. A 200mL solution of each compound can be made using the following procedures:

-Dissolve 1.7grams of potassium ferrocyanide (K4Fe(CN)6-3H2O) in 200mL of water. -Dissolve 0.04grams of potassium thiocyanate (KSCN) in 200mL of water. -Dissolve 2.1grams of ferric chloride (FeCl3-6H2O) in 200mL of water.

It may be necessary to heat the solutions to fully dissolve the salts. Theory: This experiment illustrates the unique properties of transition metals and their ability to act as Lewis acids. The original yellow ferric chloride solution contains six water molecules ligated to the Iron(III) center:

Fe

OH2

OH2

OH2H2O

H2O OH2

3+

This is possible because the positively charged Iron(III) will accept the non-bonding electrons of the oxygen in each of the water molecules into one of its empty orbitals. Thiocyanate can displace one of the water molecules in the ligation sphere to form the following red-orange compound:

Fe

SCN

OH2

OH2H2O

H2O OH2

2+

The addition of potassium ferrocyanide will cause a metathesis reaction with the following equation:

4[Fe(H2O)5SCN]2+ + 3K4Fe(CN)64- à Fe4[Fe(CN)6]3 (blue solid) + 4KSCN + 20H2O

An entertaining variation of this experiment is to dissolve ~0.01g of the potassium thiocyanate in 25mL of warm glycerol. The ferric chloride solution can be generously applied to your hand or forearm. Now, using a dull knife dipped in the cooled potassium thiocyanated-glycerol solution, drag the blade slowly across the area that the ferric chloride solution was applied allowing the glycerol to run off of the blade. If done correctly, it will look like you’ve sliced your arm open as the glycerol turns crimson red and runs down your arm. Some practice is required, but the student response has proved to be outstanding. Safety Considerations: Potassium ferrocyanide must not be mixed with hot or concentrated acids. Hydrogen cyanide gas can be produced and is lethal. Topical application of any of these chemicals carries inherent risk. You should wash thoroughly after contact with these chemicals. The thiocyanate has been shown to be harmful at ~10,000 times the concentration used in the glycerol solution and only when administered orally or intravenously. Disposal: Solutions should be flushed down the drain with water.

Crystal Lattice Principles Illustrated:

-Crystal lattice structure of salts

Required Equipment:

-Mortar and pestle Required Chemicals:

-Lead nitrate (Pb(NO3)2) -Potassium iodide (KI)

Demonstration Description: Two white powders are placed in a mortar. Upon grinding the powders turn yellow. Demonstration Procedure:

Add 5 grams of lead nitrate to 5 grams of potassium iodide in a mortar. Grind the solids together.

Theory: Solid salts exist in a crystalline lattice. This lattice allows the maximal interaction of cations and anions. This lattice structure will isolate a large number of ions from exposure to other chemicals. In aqueous solution, the water molecules often dissolve the crystal lattice and the salts are exposed and will readily react. Without the aid of water to break down the crystal lattice, the solids will react much slower. This is a simple metathesis reaction to illustrate the effect of the crystal lattice on the reaction rate. As the powders are ground together, the crystalline lattice of each salt is broken down and more individual ions are exposed. The increased exposure of ions accelerates the ability of the salts to react. The yellow color is caused by the formation of lead iodide:

Pb(NO3)2 + 2KI à PbI2 + 2KNO3

Safety Considerations: SMOKE HAZARD- Lead salts are toxic and care should be taken to avoid breathing in the powdered lead product. Disposal: The lead iodide product will need to be disposed of properly in a solid waste container.

Malonic Acid Oscillator Principles Illustrated:

- Belousoz-Zhabotinsky oscillating chemical reactions Required Equipment:

-One 1-liter Erlenmeyer flask -Magnetic stir plate -Magnetic stir bar

Required Chemicals:

-Concentrated sulfuric acid (H2SO4) -Malonic acid (C3H4O4) -Potassium bromate (KBrO3) -Manganese(II) sulfate monohydrate (MnSO4-H2O)

Demonstration Description: A solution repeatedly changes from orange to colorless every 20 seconds for about 10 minutes. Demonstration Procedure:

Add 750mL of water to the Erlenmeyer flask. While the water is stirring with a stir bar on a magnetic stir plate, slowly add 75mL of sulfuric acid. Allow the acid solution to cool to room temperature.

Add each of the following solids separately and in the following order. Allow each to dissolve before adding the next:

1. 9 g of malonic acid 2. 8 g of potassium bromate 3. 1.8 g of manganese(II) sulfate monohydrate The solution will require about 1 minute to begin oscillating. The acid solution can be prepared long before the experiment is conducted.

Theory: Chemical reactions proceed by moving to a lower state of energy. Oscillation reactions seem to contradict this principle because the reaction appears to continuously cycle back to the same point. Such is not the case. The oscillation is a side product of the system’s gradual loss of energy. There are two common threads found in all oscillation reactions:

1. The energy releasing reaction of the oscillator can choose from multiple, different pathways and the reaction periodically switches between those pathways.

2. The product of one pathway is the reactant in another pathway. As the system follows one energy-releasing pathway it generates a certain molecule. When a large amount of that molecule has been generated, the system switches to the pathway that consumes that molecule.

The specific type of reaction illustrated here is known as the Belousoz-Zhabotinsky reaction. The chemistry of the chemical reactions is complex and involves multiple intermediates. A simplified explanation follows.

The net equation for this reaction is: 4KBrO3 + 3C3H4O4 à 9CO2 + 4KBr + 6H2O

The direct reaction between malonic acid and potassium bromate is very slow. However, in the presence of a strong acid, potassium bromate will react along a different pathway once a certain amount of potassium bromide has been produced (reaction #1):

3KBrO3 + 5KBr + 18H+ + 10Mn2+ à 4Br2 + 9H2O + 8K+ + 10Mn3+ The generation of diatomic bromine (Br2) accounts for the orange color seen in the oscillation. This reaction is limited by the availability of both potassium bromide and reduced manganese(II). Once the bromide ion has been depleted, this reaction will cease. The large production of manganese(III) will initiate reaction #2:

Br2 + C3H4O4 + 4Mn3+ + 2H2O à 2Br- + 6H+ + 4Mn2+ + 2CO2 + HCO2H (2) During this reaction the diatomic bromine (Br2) is reduced to bromide and the orange color disappears in favor of a colorless solution, characteristic of the bromide ion. Once the manganese(II) levels have reached a certain level, reaction #2 stops and reaction #1 is initiated by the increase in bromide and the cycle repeats. While many of the side reactions and reactions generating intermediates have been omitted, it can be seen that this reaction will eventually end as the potassium bromate and malonic acid levels drop to concentrations insufficient to initiate the reactions. The final state of bromine will be the bromide ion and the solution will eventually remain colorless. Safety Considerations: STRONG ACID Bromates are very reactive oxidants and dry potassium bromate can ignite powdered malonic acid. Malonic acid powder is irritating to the skin, eyes and mucous membranes. Disposal: The reaction mixture can be neutralized and flushed down the drain.

Patriotic Colors Principles Illustrated:

-Color absorption -pH indicators

Required Equipment:

-4 small beakers Required Chemicals:

-Isopropanol (C3H8O)

-Phenolphthalein (C20H14O4)

-Lead nitrate (Pb(NO3)2)

-Copper(II) chloride (CuCl2-2H2O)

-Concentrated ammonium hydroxide (NH4OH)

Demonstration Description: A liquid from one beaker is added to three different beakers. The solutions in each receptacle beaker will change to red, white or blue. Demonstration Procedure: Prepare at least 50mL of the following solutions:

-Dissolve 0.5grams of phenolphthalein into 50mL of isopropanol (70% or greater). A commercially available wintergreen version of isopropanol works exceptionally well because it is green and the color change is more dramatic. -Prepare a 0.066M solution of lead nitrate by adding 1.10 grams of lead nitrate to 50mL of distilled water. -Prepare a 0.012M solution of copper(II) chloride by adding 0.10 grams of copper chloride to 50mL of water. Anhydrous or hydrated copper(II) sulfate will also work in similar proportions using the same copper(II) concentration. -Add 5mL of concentrated ammonium hydroxide to 25mL of water. The ammonium hydroxide can be substituted with 5M NaOH but the blue color is not as dark.

Set out the three first three solutions in any order you choose and add ~5mL of diluted ammonium hydroxide to each. It may be necessary to heat the solutions to dissolve the salt. Theory:

The “red” reaction observed with the phenolphthalein solution is simply a pH indicator reaction as the pH is increased above ~10. The “white” reaction forms a white precipitate, lead hydroxide:

2NH4OH(aq) + Pb(NO3)2(aq) à Pb(OH)2(s) + 2NH4NO3(aq). The “blue” reaction will form an ammonia ligated copper complex, Cu(NH3)4Cl2 according to the following equation:

[Cu(H2O)2]Cl2 + 4NH4OH à [Cu(NH3)4]Cl2(blue) + 4H2O.

Some mention of color absorbance is also applicable. The following color wheel illustrates this general principle. Whatever light color is absorbed by a solution, that solution will appear to be the color on the opposite side of the color wheel. In the case of Cu(NH3)4

2+, the molecules will absorb orange light. As the orange light is absorbed, the light passing through appears to be blue. A similar rationale can be applied to the phenolphthalein solution to determine that it absorbs green light. Safety Considerations: STRONG BASE Lead is also known to be toxic and direct contact should be avoided. Disposal: The lead compound will need to be disposed of in a waste container. All other compounds can be flushed down the drain.

Carbon Snake Principles Illustrated:

-Dehydration of hydrocarbons by sulfuric acid -Expansion of gases

Required Equipment:

-Hot plate or Bunsen burner -Evaporating dish -Pipette

Required Chemicals:

-Concentrated sulfuric acid (H2SO4) -para-Nitroaniline (C6H6N2O2)

Demonstration Description: A small amount of yellow powder is placed in an evaporating dish, melted and boiled. After several minutes a large mass of black carbon rapidly grows out of the evaporating dish. Demonstration Procedure :

Add three grams of p-nitroaniline to an evaporating dish. Add one milliliter of concentrated sulfuric acid to the p-nitroaniline. Warm the contents of the evaporating dish continuously until the snake begins to grow. Quickly remove the evaporating dish from the heat source and allow the snake to finish growing.

Just prior to the formation of the snake, a large amount of gas is released. If you can remove the evaporating dish from the heat source during this time the snake tends to grow faster and achieves a greater length.

Para-nitroaniline can be substituted with para-nitroacetanilide.

Theory: Sulfuric acid is an excellent dehydrating reagent and will drive the following reaction:

2C6H6N2O2 + O2 à 12C(graphite) + 6H2O + 2N2 Sulfuric acid will complex water according to the following reaction:

H2SO4 + nH2O à H2SO4-nH2O The heat source is necessary for the oxidation of graphite by atmospheric oxygen to form carbon dioxide:

C(graphite) + O2 à CO2

The production of carbon dioxide gas and nitrogen gas (N2) will cause a gas expansion of the melted p-nitroaniline and lead to the growth of the carbon snake. Sulfur dioxide gas

is also produced as deleterious side reaction and has a characteristic odor: 2C(graphite) + 2H2SO4 + O2 à 2CO2 + 2SO2 + 2H2O Due to the aromatic stability of p-nitroaniline, the reactants must be heated sufficiently to initiate the reaction. Safety Considerations: SMOKE HAZARD- only a moderate amount of noxious fumes are produced, but care should be taken to perform this experiment in a well ventilated area. FIRE HAZARD Disposal: The carbon snake can be thrown away with no further treatment.

Contact Explosive Principles Illustrated:

-Instability of chemical compounds -Soundwave generation

Required Equipment:

-Filter flask and filter paper -Mortar and pestel

Required Chemicals:

-Iodine crystals -Concentrated ammonium hydroxide (NH4OH)

Demonstration Description: A small dark pile is lightly touched with a long wooden or metal rod. Upon contact it explodes loudly and releases a purple gas Demonstration Procedure: Place 10 grams of iodine crystals in a 100mL beaker. In a fume hood or well-ventilated area, add 40mL of concentrated ammonium hydroxide and stir until precipitation ceases. Filter the solution and catch the precipitated crystals on the filter paper. DO NOT LET IT DRY. Quickly move the moist filter paper with the precipitate to a safe location for detonation. When the crystals are dry, touch them with any object to cause the detonation. A Buchner funnel with a side arm vacuum-flask works well for filtration if a vacuum source is available. Leaving the crystals on vacuum will dry them out quickly and care should be taken to disconnect the vacuum as soon as possible. The vacuuming should last only seconds. One popular variation is to detonate the crystals with a feather that is attached to a long wooden pole. Do not use a glass-stirring rod to detonate the crystals. The crystals can be stored, while moist, in an airtight container for a short period of time. Theory: The iodine crystals react with ammonium hydroxide to form a nitrogen triiodide-ammonia complex: 8NH3-NI3. Under dry conditions, the complex will rapidly decompose:

8NH3-NI3à 5N2(g) + 6NH4I(s) + 9I2(g) This explosive is not very powerful but it is extremely sensitive and very loud.

The reason the explosion generates so much noise is due to the rapid expansion of the gases formed during the decomposition of the nitrogen triiodide-ammonia complex. This gas expansion compresses the air surrounding the crystals and creates a small shockwave similar to what occurs at the muzzle of a firearm when it is discharged.

Safety Considerations: SMOKE HAZARD STRONG BASE

Nitrogen triiodide should be handled sparingly and only when it is obviously wet. Those near the explosion should wear ear protection. The iodine gas released during detonation and the ammonia fumes from the original ammonium hydroxide solution are toxic gases and proper accommodations should be made. Disposal: The detonation should eliminate most of the compound. Wash the area with ethanol. Any unreacted nitrogen triiodide should be soaked for 8~10 hours in ethanol and subsequently washed down the drain.

Endothermic Principles Illustrated:

-Spontaneous endothermic reactions -Gibbs free energy equation

Required Equipment:

-A small beaker -A small piece of wood

Required Chemicals:

-Barium hydroxide octahydrate (Ba(OH)2-8H2O) -Ammonium thiocyanate (NH4SCN)

Demonstration Description: Two powders are mixed together in a beaker situated on a damp piece of wood. After the powders melt together the beaker can be picked up, lifting the wood, which has been frozen to it. Demonstration Procedure:

Add 15 grams of barium hydroxide octahydrate and 5 grams of ammonium thiocyanate to a small beaker. Briefly mix the powders with a spatula and place the beaker on a small, damp, piece of wood. After ~30 seconds, lift the beaker up and the wood will come with it.

This experiment should be practiced beforehand to make sure the wood is not to big and to gauge the time it takes for the beaker to freeze to the particular piece of wood. One of the reagent bottles may also be placed on the wood and lifted with the beaker.

Theory: A reaction is spontaneous when it can occur without external input of energy as a driving force. In order for a reaction to occur spontaneously it must be energetically favorable for that reaction to occur. In order for the reaction to be energetically favorable, the overall free energy of the system must decrease as the reaction proceeds. This energetic “favorableness” is determined by the Gibbs free energy equation:

∆G = ∆H - T∆S, where ∆G/H/S = ∆G/H/Sfinal -∆G/H/Sinitial If ∆G is negative, then the system has lost energy (∆Gfinal < ∆Ginitial) and the reaction is spontaneous. In order for ∆G to be negative, either ∆H must be negative (∆Hfinal < ∆Hinitial), ∆S must be positive (∆Sfinal > ∆Sinitial) or both ∆H is negative and ∆S is positive.

∆H is generally known as the heat or enthalpy of the system. Typically it is the major factor in determining the spontaneity of a reaction, and when heat is given off, ∆H is negative, the reaction is exothermic and proceeds spontaneously. In this demonstration, ∆H is positive, heat is taken in by the system, cooling the surroundings, and the reaction is called endothermic.

In order for such a reaction to occur, the entropy (∆S) must be sufficiently large and positive to offset the effects of a positive ∆H. Entropy is commonly referred to as the

measure of the disorder of a system. When a system becomes more random and disordered, the entropy increases. One way of analyzing the disorder of a system is to compare the number of products to reactants. If the number of products is much larger than the number of reactants, it may be assumed that the system is more disordered and the entropy has increased. Such a system may allow for a spontaneous reaction in spite of an increase in enthalpy (positive ∆H). The two solids mix according to the following equation:

Ba(OH)2-8H2O + 2NH4SCN à Ba(SCN)2 + 2NH3 + 10H2O Note that the powders do not “melt.” The water formed in the reaction will cause the solids to partially dissolve. This highly endothermic reaction will proceed because of the entropy considerations. While the change in enthalpy is clearly positive, the formation of 13 molecules of product from 3 molecules of reactant would suggest a large increase in entropy. This large increase in entropy is sufficiently large to allow the T∆S term to offset the positive ∆H term and the reaction proceeds spontaneously. Safety Considerations: SMOKE HAZARD- Ammonia is produced in small amounts. Barium salts are toxic and should not directly contact the skin. Disposal: The barium salt products are toxic and should be disposed of in an appropriate container.

Kinetic Molecular Theory Principles Illustrated:

-The kinetic molecular theory

Required Equipment: -Two large beakers -Hot plate

Required Chemicals:

-Food coloring -Ice

Demonstration Description: A single drop of food coloring is added to each of two large beakers. The drop diffuses more slowly through the water of one beaker than the other. Demonstration Procedure: Fill the first beaker nearly full with warm tap water and heat it on a hot plate to near boiling. Fill the second beaker approximately 2/3rds full with cool tap water. Fill it the rest of the way with crushed ice. When the ice has just melted in the second beaker, place the beakers side by side on the table. Allow the water in each beaker to settle. Add one drop of red food coloring to the first beaker and one drop of blue food coloring to the second beaker. Allow the class to guess at the relative temperatures of the two beakers. Theory: This experiment is painfully simple, but serves as an excellent demonstration of the kinetic molecular theory. According to the kinetic molecular theory, particles are in constant motion and temperature is related to the average speed of the particles in solution. A higher temperature denotes faster molecular movement. The rate of diffusion of the food coloring is directly proportional to the speed of the water molecules and therefore directly proportional to the temperature of the water. Safety Considerations: The hot plate should be treated with care. Disposal: Everything can be flushed down the drain.

Luminol Principles Illustrated:

-Chemiluminescence Required Equipment:

-Three large beakers (1 liter volume or greater) -Spiral condensing tube apparatus (optional)

Required Chemicals:

-Sodium carbonate, anhydrous (Na2CO3) -Luminol (5-amino-2,3-dihydrophthalazine-1,4-dione- C8H7O2N3) -Sodium bicarbonate (NaHCO3) -Ammonium carbonate monohydrate ((NH4)2CO3-H2O) -Copper(II) sulfate (CuSO4) -Hydrogen peroxide (H2O2)

Demonstration Description: In a darkened room, two liquids are mixed together to produce a blue, luminescent solution. Demonstration Procedure:

Add 500mL of distilled water to a large beaker. Dissolve each the following salts into the water separately and in the following order:

1. 4.0 g of sodium carbonate 2. 0.2 g of luminol 3. 24.0 g of sodium bicarbonate 4. 0.5 g of ammonium carbonate monohydrate 5. 0.4 g of copper(II) sulfate pentahydrate or 0.26g of copper(II) sulfate

anhydrous Dilute the solution to a final volume of 1 liter. In the second large beaker add 5mL of 30% hydrogen peroxide or 50mL of 3% hydrogen peroxide. Dilute with water to a final volume of 1 liter. Turn the lights out and mix the two solutions together into the third large beaker. The solutions should glow for about one minute. Do not prepare the solutions more than 6 hours before the demonstration. An optional method of displaying this reaction is to pour the liquids together through a spiral condensing tube apparatus. Using a ring stand and several clamps, position the condensing tube over a separate beaker. Place a glass funnel in the top of the condensing tube. Pour equal amounts of the two previously prepared solutions into the funnel. Allow the glowing solution to spiral down the condensing tube and exit into the drainage beaker.

Theory: The chemiluminescent properties of luminol are dependent upon a series of chemical reactions. These reactions are not entirely elucidated but a reasonable conjecture is illustrated.

The structure of luminol is:

NH

NH

O

ONH2

Luminol

Because of the multiple carbonate compounds added in this experiment, the pH of the luminol solution is ~9 and the luminol molecule is deprotonated to form:

:NH

:N:-

O

ONH2 The copper(II) cation, when complexed with ammonia from the

ammonium carbonate to form the tetramminecopper(II) complex, will initiate the oxidation reaction by becoming reduced:

:N

:N:-

O

ONH2

:N.

:N:-

O

ONH2

H

Cu(II)

+ Cu(I) + H+

Copper(II) has 9 electrons in its d-orbital. Addition of a single electron and reduction to copper(I) fills the d-orbital providing sufficient stability to justify the action of a metal as an oxidant. The high pH of the solution also stabilizes the generation of a hydrogen ion. Resonance will lead to the relocation of the unpaired electron:

:N.

:N:-

O

ONH2

CN:

N:

O

:O:-

NH2 ..

.

Hydrogen peroxide is added to generate the superoxide radical, which will react with the luminol at this point to bridge the carbonyl carbons:

CN:

N:

O

:O:-

NH2 ..

.

O:-..

..O....

.

CN

N

O

NH2

-O O:

-....O

....

CN:

N:

:O:-

NH2 ..

O

O

:O:-..

The final reaction step will generate a molecule of nitrogen gas and a molecule of 3-amino-phthalate in the excited state:

CN:

N:

:O:-

NH2 ..

O

O

:O:-..

O-

O-

O

ONH2

+ N N

*

During the final step of this reaction, the excited 3-amino-phthalate will release a photon of light (hν) and fall to a lower state of energy:

O-

O-

O

ONH2

*

O-

O-

O

ONH2

+ hν

In a standard, light-emitting chemical reaction, energy is absorbed by the promotion of the electrons to higher energy levels. As the electrons return to the ground energy state, they release that absorbed energy as light photons. In a chemiluminescence reaction, changing the molecular composition will create a new and lower ground state of energy for the electron because this ground state is dependent upon the molecular environment. As that environment changes, so will the energy level of the ground state. The electron will release energy to drop to this new, lower energy ground state and light is emitted without the electron having been excited. Safety Considerations: Hydrogen peroxide is a strong oxidant and should not directly contact the skin. Disposal: All substances used in this demonstration can be flushed down the drain. Before disassembly of the spiral condensing tube apparatus, flush it thoroughly with distilled water.

Magic Eggs Principles Illustrated:

-Density and water displacement -Generation of carbon dioxide from carbonate

Required Equipment:

-Two 1000mL graduated cylinders -Hot plate -Large beaker >1000mL

Required Chemicals:

-2 eggs -Sodium chloride (NaCl) -Hydrochloric acid (HCl) -Sodium Bicarbonate- optional (NaHCO3)

Demonstration Description: Two eggs are placed into separate graduated cylinders. One egg sinks until it reaches the middle of the cylinder where it remains suspended. The other egg sinks to the bottom, eventually rises to the top, returns to the bottom after a slight nudge and continues this cycle indefinitely. Demonstration Procedure: Prepare the saltwater solution by dissolving 200g of salt in 500mL of hot water in a large beaker. Stir the salt water continuously while heating near boiling, until as much salt is dissolved as possible. Allow the solution to cool and decant off around 450mL of salt water into one of the 1000mL graduated cylinders. Carefully pour water down the sides of the cylinder containing the salt solution until it is nearly full. Try not to disturb the saltwater. You should see a faint but visible separation between the salt water and the regular water. Place an egg into the cylinder and it will sink to the top of the saltwater solution. Into a large beaker add 35mL of concentrated hydrochloric acid (12M HCl) to ~1000mL of water. Stir this solution and pour the contents into the second graduated cylinder. Add the egg. It should rapidly sink to the bottom and after 30 seconds it will gradually rise to the top. When it reaches the top, you can bump it with your finger and it will return to the bottom. This cycle can be repeated indefinitely. Adding the correct amount of acid to the water is difficult to generalize for all eggs. Have some acid and water on hand to dilute or concentrate the solution as necessary. It is important to use the salt solution shortly after preparation and not to add the regular water to the graduated cylinder containing the salt solution until just before you plan to perform the demonstration. The acid solution has a final HCl concentration of 0.42M and pH <0.5 An interesting variation is to place the egg in about 200mL of concentrated HCl. The eggshell will completely dissolve in about 5 minutes. You can transfer the egg to a 3M sodium bicarbonate solution to show that the acid has entered the egg. It will bubble in the NaHCO3 solution for a prolonged period of time.

Theory: In the first cylinder, the egg will remain on top of the saturated salt solution because the density of the saltwater is greater than the net density of the egg. In the second cylinder, the hydrochloric acid reacts with the calcium carbonate on the surface of the eggshell:

2HCl + CaCO3 à H2O + CO2 + CaCl2 Carbon dioxide gas is evolved in the reaction and remains attached to the eggshell. When enough carbon dioxide is produced, the net density of the egg will decrease and the egg will float to the surface of the solution. When the egg surfaces, the carbon dioxide bubbles clinging to the sides of the egg will release and the egg will again sink. The reaction will proceed as long as there remains a means of producing carbon dioxide gas. Safety Considerations: STRONG ACID Disposal: The acid solution should be neutralized and flushed down the drain with the salt solution. Do not eat the egg as it will absorb some of the acid.

Peroxide Foam Principles Illustrated:

-Catalysis -Disproportionation reactions

Required Equipment:

-Large graduated cylinder, vacuum flask or volumetric flask Required Chemicals:

-30% Hydrogen peroxide (H2O2) -Dish soap -Potassium iodide (KI)

Demonstration Description: A small amount of powder is added to a clear solution. The powder turns yellow upon contact and quickly generates a profuse amount of foam. Demonstration Procedure: Add 50-100mL of 30% hydrogen peroxide to 1mL of dish soap in a large, narrow flask. Swirl gently until the dish soap is entirely dissolved. Several drops of food coloring may be swirled in as well. Add 2 grams of potassium iodide at one time and allow the reaction to proceed. The key to making this reaction work is to use the right glassware. A 2L volumetric flask works the best, although a 1L graduated cylinder is also acceptable. One variation is to carry out the reaction in a 1L vacuum flask and quickly stopper it as soon as the potassium iodide crystals are added. The foam generated will shoot out of the vacuum port with considerable pressure for several seconds. Theory: This is an excellent illustration of a catalytic reaction where a chemical is consumed to initiate the reaction, but is regenerated as the reaction terminates. The initiation equation is:

2I-(aq) + H2O2(l) +2H+(aq) à I2(g) + 2H2O(l) The regeneration equation is:

I2(g) + H2O2(l) à 2I-(aq) + 2H+(aq) + O2(g) The sum of the two equations shows that the catalyst is not consumed:

2H2O2(l) à 2H2O(l) + O2(g) Because the potassium iodide is a catalyst, only a few grains are theoretically

required to catalyze the reaction. Experience has shown that the reaction will proceed much better when two grams of potassium iodide are added. This is due to the rising foam pushing some of the iodide ions out of the beaker before they can catalyze further reactions. The hydrogen peroxide is this experiment will undergo both oxidation to form oxygen gas and reduction to form water. This type of chemical reaction, where one reactant is both oxidized and reduced, is called a disproportionation reaction.

Safety Considerations: Hydrogen peroxide is a strong oxidant and should be handled with gloves and goggles. The reaction can generate a fair amount of heat, but the end products should be considered safe. Disposal: All of the products are safe to flush down the drain.

The Glowing Pickle Principles Illustrated:

-Electrolysis of water -Conductivity of electrolytes -Electron excitation and photon emission

Required Equipment:

-Large Pickle -Two 3-4 inch large gauge copper wires -Extension cord -Surge protector -A large glass dish

Demonstration Description: A pickle glows as a strong electric current is passed through it. Demonstration Procedure: While the extension cord is unplugged, cut off the extension input (female) end and expose ~1 inch of each of the wires by removing the insulation. Wrap each exposed wire around one end of the large gauge copper wires. Insert the copper wires along the length of the pickle so that they are parallel to each other and the extension cord wires are on opposite sides of the pickle. Now, set the pickle on a non-conducting glass plate and plug the extension cord into a surge protector. The pickle may remain “plugged in” for 20-30 seconds or until it ceases to conduct electricity efficiently. A common problem with this experiment is that enough disturbance is produced by the electric charge passing through the pickle to push the electrodes out of the pickle. Theory: This experiment is an exciting illustration of an electrolytic cell. Water is being electrolyzed according to the following equation:

2H2O à 4H+ + 2O2- à 2H2 + O2 This equation is divided up into two half-reactions where the oxidation occurs at one of the copper wires (the anode):

2O2- à O2 + 4e- Reduction occurs at the other copper wire (the cathode):

4H+ + 4e- à 2H2

Because of the high salt-ion content of the pickle, there exists a sufficient movement of charge through the system and the reaction will proceed until the pickle is dehydrated. The light emitted by the pickle works according to the same principle found in the light emission of a light bulb. As the pickle increases in heat and energy, due to the charge passing through it, the electrons in the various atoms of the pickle are excited and promote to higher energy levels. As these electrons return to their ground state energy level, light is emitted. Safety Considerations: This experiment is extremely dangerous. The demonstrator should wear non-conducting, rubber gloves. If the pickle is left plugged in for too long it can become a serious fire hazard. It is also advisable to perform this experiment under a fume hood to avoid the strong odor of burnt pickle. Disposal: The pickle can be thrown away and all other parts of the experimental apparatus can be saved for future use.

Sugarloaf Principles Illustrated:

-Action of sulfuric acid on hydrocarbons -Expansion of gases -Le Châtelier’s principle

Required Equipment:

-One ~250mL tall form beaker -Stirring rod -100mL graduated cylinder

Required Chemicals:

-Concentrated sulfuric acid (H2SO4) -Granulated table sugar (sucrose- C12H22O11)

Demonstration Description: A mixture of acid and sugar gradually turns black and begins to grow as a solid mass. Demonstration Procedure:

With the sugar already in the beaker, slowly stir in the concentrated sulfuric acid. Remove the stirring rod and allow the acid/sugar paste to react.

Theory: Concentrated sulfuric acid is very hygroscopic and an excellent dehydrating reagent of hydrocarbons. The ability of sulfuric acid to complex water drives the formation of carbon:

C12H22O11 à 12C(graphite) + 11H2O Sulfuric acid can complex water and in doing so will generate a large amount of heat:

H2SO4 + nH2O à H2SO4-nH2O + heat This heat will lead to the oxidation of graphite by atmospheric oxygen to form carbon dioxide:

C(graphite) + O2 à CO2. As the heat of the sulfuric acid hydration melts the sugar, the production of

carbon dioxide gas will cause a gas expansion of the melted sugar and lead to the growth of the carbonaceous sugarloaf. Sulfur dioxide gas is also produced as deleterious side reaction and has a characteristic odor:

2C(graphite) + 2H2SO4 + O2 à 2CO2 + 2SO2 + 2H2O Le Châtelier’s principle is well illustrated by this reaction. The decomposition of

sucrose to form graphite and water is not energetically favorable, but the removal of the water product by the sulfuric acid pushes the reaction to the formation of more products.

Safety Considerations: STRONG ACID SMOKE HAZARD- this reaction produces a large amount of very noxious smoke. This reaction will produce a lot of heat and the product should be allowed to stand for ~10 minutes before handling.

The amount of energy produced in this reaction is sufficient to vaporize small amounts of sulfuric acid and it should be performed in a fume hood or other VERY well ventilated area. Disposal: The remaining carbon mass should be washed thoroughly with water and a dilute sodium bicarbonate solution before discarding.