Demonstrations of Chemical Reactions

resourcesforteachingchemistry/maindishes/entrees

Home

A first search the literature resulted in a great many potential chemical demonstrations but on closer inspection many of these are variants of one another, disguised by different catchy titles. First thing was to select the best and then in naming the demo, I sometimes kept the most popular name, but more often than not came up with a phrase that suggested the chemical principle being demonstrated. 

Lots of demos, students take notes.  Try to avoid dangerous demos and sometimes selected simplicity over smoke & mirrors-whatever illustrates the principle best and not necessarily makes the biggest explosion.  Sometimes the special effects overshadow the principle and while entertaining that is doing a disservice to the student .  Indebted to Flinn, ChemWest, Becker, Shakhashir, Teachers at... as the major source of these demos.

There were so many demos that broken up into different web pages.  here are demos Highlighting Chemical Reactions.  See also Demonstration of Chemical Principles

 

Red, White & Blue (acid/base, precipitate, complexation reactions)

Return to Top

Hydrolysis of Salts (pH study of salts dissolved in water)

Return to Top

Fertilizer & Explosives (Ostwald Oxidation of Ammonia)

Return to Top

An Entropy-driven Endothermic Reaction (Freeze stick to a Board)

Return to Top

Thermite (Powdered Al and Metal Oxide to give Molten Metal)

Return to Top

Instant Carnations (Cobalt Coordination is better)

The Pink Catalyst-Demo--Not the same as carnations!!!!

Introduction

Add a pink cobalt chloride solution to a colorless solution containing potassium sodium tartrate and hydrogen peroxide and watch as a very obvious green-colored complex forms. As the reaction ends, the solution will return to its original pink color—indicating that the cobalt chloride catalyst is not used up in the reaction.

 

Chemical Concepts • Catalysts • Kinetics

 

Materials

Cobalt chloride, CoCl2z6H2O, 0.4 g                                                      Graduated cylinder, 100-mL

Hydrogen peroxide solution, 6%, H2O2, 40 mL                                  Hot plate

Potassium sodium tartrate solution, 0.21 M, 100 mL                         Spatula or scoop

Beaker, 50-mL                                                                                           Stirring rod or magnetic stirrer

Beaker, 500-mL or 1-L                                                                            Thermometer

Distilled water, approximately 25 mL

 

Procedure

Prepare a 0.21 M potassium sodium tartrate solution by dissolving 6 g of KNaC4H4O6z4H2O in 100 mL of distilled water.

1. Using a graduated cylinder, measure out 100 mL of 0.21 M potassium sodium tartrate solution. Pour it into a 500-mL or 1-L beaker.

2. Warm the solution slowly to 70 °C on a hot plate.

3. While waiting for the temperature of the solution to increase, dissolve one scoop (0.4 g) of cobalt chloride in a very small amount of distilled water. Show this solution to the class so that the students can note the pink color of the catalyst.

4. When the temperature of the potassium sodium tartrate solution reaches 70 °C, add 40 mL of 6% hydrogen peroxide and the cobalt chloride catalyst to the 500-mL beaker. Stir continuously.

5. As the solution becomes green, a vigorous reaction occurs. When the reaction subsides, the green color disappears and the pink color of the cobalt chloride returns.

 

Tips

• When timing the reaction, the use of a hot plate-magnetic stirrer greatly aids the demonstration of this reaction. Otherwise, continuous stirring is necessary.

• Begin timing the reaction upon addition of the hydrogen peroxide and cobalt chloride solution. Complete the timing after the vigorous reaction subsides and the original pink color of the cobalt chloride solution has returned.

• The use of a large beaker is important so that the reaction does not froth over.

• 12–15 mL of 0.1 molar cobalt chloride solution can be used instead of solid cobalt chloride.

 

Discussion

The solution starts out pink due to the pink color of the cobalt chloride catalyst. The solution turns green forming an intermediate between the catalyst and potassium sodium tartrate. The solution returns to the original pink color of the cobalt chloride solution demonstrating and confirming the fact that a catalyst does not get used up in a chemical reaction. Based on experimental findings, the following reactions can be suggested as taking place in the Co(II)-H2O2-tartaric acid system: On the action of hydrogen peroxide, the cobalt(II)-tartrate complex becomes oxidized to a green, probably binuclear, Co(III)-tartrate compound. This cobalt(III)-tartrate is reduced both by tartaric acid and hydrogen peroxide to Co(II)-tartrate with a concomitant evolution of CO2 and O2, respectively. Since the color of the solution is green throughout the reaction, and most of the cobalt is present as Co(III), then the first step (oxidation) is most likely faster than the reduction of Co(III)-complex. (Toth, 1980). This demonstration also demonstrates kinetics—the effect of temperature on the rate of a chemical reaction. For each 10 °C increase in temperature, the reaction rate will approximately double. The reaction may be timed at various temperatures. Suggested temperatures and their corresponding reaction times are: 50 °C—200 seconds; 60 °C—90 seconds; 70 °C—40 seconds

 

Disposal: Dispose of the final solution according to Flinn Suggested Disposal Method #27d.

.

References

Deroo, Julius, Sci Teach., 1974, 41, 44., Ruda, Paul T., J. Chem. Educ., 1978, 55, 652., Toth, Zoltan, J. Chem. Educ., 1980, 57, 464. Flinn Scientific, Inc.

Return to Top

Miner’s Lamp & The Self-Carving Pumpkin (CaC2 + water forms C2H2)

I am thinking of a fuel that...

We think of fuel these days most as gas (natural gas, butane, propane) or liquified or gasoline and jet fuel that can be pumped.  There is also a way to store as a solid and generate gas through a bit of chemistry.  Union carbide got its start with this product and before electric light camp tocars this ws a very common way to produce lights (more hints?) miners use it and in fact it remains popular today with cavers.  What is it

Return to Top

Underwater Fireworks (C2H2 + Cl2)

Chlorine gas is bubbled up along with acetylene gas through a large graduated cylinder filled with water. Where the bubbles of the two collide, an instantaneous, bright flash of light occurs.

Concepts: In some hydrocarbons, two or even three pairs of electrons can be shared between two adjacent carbon atoms. These multiple sharings are known as double or triple bonds, and the areas where they occur are said to have high electron densities. Hydrocarbons with double or triple bonds are referred to as "unsaturated."

Halogens have seven electrons in their outermost level and need only one more to form a stable octet. This gives them a high electron affinity.

Activation energy (AE) is the energy required by reactant particles so that they might collide with enough force to initiate a reaction. Many reactions require high temperatures before they can begin, but some reactions, like those between halogens and unsaturated hydrocarbons, have low AE.

CaC2 + 2 H2O --> C2H2 + Ca(OH)2

 

C2H2 + Cl2 --> 2 HCl + 2 C (black soot)

Procedure

1.       Cut glass tubing to about 10 cm longer than the height of the graduated cylinder.  Fill the graduate with tap water to within 1-2 cm of the top.

2.       Under the hood, place 100 mL of NaClO solution in the flask and carefully pour in 20 mL of 3 M HCl. (Caution: these two react to form Cl2 especially when the flask is swirled or shaken. Insert the 1-holed stopper assembly.)

3.       Swirl the flask gently until 2-3 bubbles of Cl2 bubble up out of the tube. Bubbles should not be emitted in a steady stream but form only when you swirl the flask. If insufficient Cl2, swirl again. If there is still too little Cl2, add 10 more mL of HCl. Swirl. Only a few bubbles from each swirl are required.

4.       Drop 2 pea-sized pieces of CaC2 into the water; immediately generates acetylene gas.

5.       Swirl the flask and maneuver the glass tube along the bottom of the graduate to cause bubbles of Cl2 to collide with bubbles of C2H2. Turn down the lights.

6.       Tip the cylinder to facilitate the reaction, causing bubbles to travel up the inside surface, increasing the likelihood of bubbles colliding with one another.

7.       Turn on the lights to observe any products. As soon as the reaction is over, add base to the bleach solution, and move the entire apparatus to a hood.

 

Materials-

Chlorine Trap (in place of hood):

Safety: Good ventilation required. Use a hood if possible. Cl2 is toxic; C2H2 is flammable. Place a fire extinguisher nearby. CaC2 reacts with any source of water including eyes. Use care. Bring only a small amount of CaC2 to the lab bench.

Disposal:

Observations: Carbon forms at the top of the cylinder. If tap water is used, calcium compounds usually precipitate.

Q&A

Q1. In which phase is the reaction occurring?

A1. The reaction takes place between bubbles in the gaseous phase.

Q2. Describe a major difference between this reaction and the reaction of H2 and O2.

A2. The activation energy of this reaction is so low that a spark is not required at RT.

Q3. What side products are present? Describe any precipitate that forms.

A3. A white precipitate forms from the reaction of calcium ions with the tap water. You may sometimes observe carbon black.

Notes: This demonstration was originally developed by Walter Rohr of Eastchester, New York. Shakhashiri, B. Z., in his "Chemical Demonstrations", Volume 2 (p 227., University of Wisconsin Press, 1985), discusses the reactions in this experiment.

Key Words rate of reaction, kinetics, unsaturated, activation energy

Return to Top

 

Burning Magnesium with Dry Ice or Sand (Mg + CO2 or SiO2)

Part 1

Concept: Periodicity (follow up the Mg and dry ice demo with this one)

What are ways to put out a Mg fire? Previously we saw that CO2 reacts with Mg and feeds the fire-water is also a horrible choice.  How else? (Most students think of sand to smother the fire)

 

Materials

Procedure

 

Discussion

As in carbon dioxide, covalent bonds need to be broken.  SiO2 + 2 Mg →Si + 2MgO

Gray silicon can be seen in the tube.  Since the tube is glass (SiO2) there may a reaction with the tube and it may show signs of melting. (An aside: The gray-black discoloration on crucibles may be silicon formed through high heat and the crucible glaze. For this reason, discolored crucibles are better for the formula of Mg experiment.)

 

Part 2 (Spontaneous Alien Combustion)

 Sometimes there are side reactions in doing chemistry in addition to the main reaction.  This often happens in organic synthesis and a great deal of effort goes into trying to purify the product from unintended side reactions to get better yield.  Our sand reaction is a case in point.  The conditions of very high temperature that were needed to get the reaction going also resulted in some side reactions.  One such side reaction was the production of magnesium silicide, Mg2Si. This compound can be used to make silane, a Si and hydrogen compound analogous to methane.

 

Materials:

Procedure

 

Discussion Magnesium silicide reacts with HCl to liberate silane. 

 

The reaction is:   MgSi2(s) + HCl (aq) → SiH4 + MgCl2 (aq)    (higher silanes such as Si2H6 may also be produced)

 

Many hydrocarbons are used as fuels and react exothermically with oxygen.  Silanes do as well, but their activation energies are so low that at RT they spontaneously combust. (Small flashes of light and tiny “pops” will be observed as the residue is shaken into the HCl.)  The difference in reactivity between silanes and alkanes (such as methane and ethane) can be explained by considering the relevant bond energies and the availability of d-orbitals in silicon.

 

The reaction for the combustion of silane is:  SiH4 (g) + O2 (g) → SiO2 (s) + H2O (l)

 

The low activation energy of silane has some interesting implications for science fiction fans.  One common scenario is the existence of life forms based on silicon instead of carbon.  If these life forms were to come to earth, the silicon compounds in their bodies would quickly oxidize.  Hmmmm, so how about those stories we see in tabloids about spontaneous human combustion—they must be the aliens.  We’ll know for sure when the Tabloids report that the ashes were nothing but piles of wet sand.

 

Adapted from Rob Lewis (Retired from Downers Grove)

Return to Top

Cement & Cooked Eggs (Slaking Lime)

Return to Top

Can Ripper (3Cu2+ + 2Al → 3Cu(s) + 2Al3+)

Return to Top

Flash Paper (cellulose nitrate)

Return to Top

Nitric Acid Acts on Copper (Ira Remsen)

Return to Top

Polishing Silver the Lazy Way (Electrochemistry)

Over time silverware and silver jewelry pick up a black tarnish.  (What silver compound is black?)  From the miniscule amount of H2S in the air (rotten egg smell) the silverware develops a coating of Ag2S.  Most often, people remove this layer of tarnish by scrubbing with a polishing compound which is not only removes some of the silver but is difficult to get into the cracks etc.   A much better way is to use our knowledge of chemistry to dip the silver in a non-toxic electrochemical dip.

First line a suitable dish with aluminum foil.  Then fill with hot water.  Add a tsp of salt and a teaspoon of baking soda.  Drop the silver items in so that they are touching each other and touching the aluminum foil.  Watch the tarnish disappear!  Once clean, remove the silverware, rinse with water and gently buff with a dry cloth.

Return to Top

Start A Fire with a Drop of Water (NH4NO3 + Zn + a little NH4Cl catalyst)

Return to Top

Genie in a Bottle (Decomposition of Peroxide using NaI)

Introduction

When sodium iodide is dropped into a flask containing 30% hydrogen peroxide, a “magical” genie appears in the form of water vapor and oxygen.

 

Chemical Concepts

• Exothermic reaction • Catalysts • Decomposition reactions

 

Materials

30% H2O2, 50 mL                                  Volumetric flask, Pyrex®, 1000-mL

NaI, 4 g                                                                  Filter paper

Graduated cylinder, 50-mL or 100-mL

 

Safety Precautions

 The reaction flask will get extremely hot; use only a Pyrex flask and hold with a towel around it to prevent burns. Do not point the mouth of the flask towards yourself or anyone else. Never tightly close a vessel containing hydrogen peroxide—it may explode.

 

Procedure

1. Wrap 4 g of sodium iodide in a small piece of filter paper or tissue. Staple the filter paper so that no sodium iodide leaks out. 2. Add 50 mL of the 30% hydrogen peroxide solution to a 1000-mL Pyrex volumetric flask. Caution: Wear rubber gloves when handling 30% H2O2. Contact with skin may cause burns.

3. Set the flask on a counter and hold the flask with a thick cloth towel. Drop in one packet of the sodium iodide solid. Point the flask up and in a safe direction away from yourself and your students as the magic genie (water–vapor) emerges from the flask. The flask will get extremely hot. The towel will hide the flask contents as well as protect your hand from the heat produced.

 

Tips

• It is very important that this demonstration be done in a borosilicate (i.e., Pyrex) flask. A flask that is not borosilicate glass can crack from the evolution of heat.

• A large flask (1000-mL) is necessary because a brownish liquid can spurt out at the end of the reaction. A large flask will help prevent this from happening. The brown liquid results from the presence of free iodine produced from the extreme oxidizing ability of the 30% hydrogen peroxide.

• A thick cloth towel will prevent your students from seeing what is happening in the flask, as well as protect you from the heat evolved in the reaction. Another option is to wrap the flask in aluminum foil and decorate it like a “Genie bottle”.

• Manganese (IV) oxide can be substituted for sodium iodide in the demonstration. Both chemicals catalyze the reaction and will cause the release of oxygen from hydrogen peroxide.

• The sodium iodide packet can also be attached to a piece of thread and hung inside the flask. Attach the thread to the outside of the flask with tape or a stopper. Warning: Do not use a solid stopper or cap. If the reaction starts prematurely, the pressure buildup may explode the flask. Use a one- or two-holed stopper and place it loosely on the flask.

 

Discussion

The Magic Genie demonstrates the decomposition of hydrogen peroxide into oxygen gas and water vapor. The decomposition is catalyzed by iodide (I ) which is not changed during the reaction. It is an exothermic reaction and will evolve a lot of heat. The reaction is:

                                                                      I(aq)

                                                   2H2O2(aq)    →        2H2O(g) + O2(g) + Heat Energy

Cleanup and Disposal

Immediately clean up any liquid which may have splattered on the floor. Pour any liquid remaining in the flask down the drain with excess water. Rinse the flask thoroughly with water.

 

Reference

Stone, Charles, H. J. Chem. Ed., 1944, 21, 300 and Flinn Scientific. 

Return to Top

 

Elephant Toothpaste & Kinetics (Kinetic Studies of Peroxide Decomposition)

 

Chemical Concepts

• Kinetics/Catalysts • Decomposition reactions

• Reaction intermediates • Test for oxygen gas

 

Materials Needed (for each demonstration)

Alconox® detergent, 3–4 g                                                                 Graduated cylinder, 500-mL

Hydrogen peroxide, H2O2, 30%, 10%, 3%, 20 mL of each             Large, plastic demonstration tray, several inches deep

Sodium iodide solution, NaI, 2 M, 4–5 mL                                       Lighter or matches and wood splint

Graduated cylinders, 10-mL and 100-mL, 3 of each                        Spoon or scoop

 

Procedure

Part 1 — Effect of Concentration on the Rate of the Reaction

1. Place three 100-mL graduated cylinders in a large, plastic demonstration tray. Add 20 mL of 30% hydrogen peroxide to the first cylinder, 20 mL of 10% hydrogen peroxide to the second cylinder, and 20 mL of 3% hydrogen peroxide to the third cylinder.

2. Add 1 small scoop (3–4 g) of solid Alconox® detergent to each cylinder and swirl to dissolve the detergent.

3. Measure out 5 mL of 2 M sodium iodide solution in each of three 10-mL graduated cylinders. Ask your students to predict the rate at which each of the peroxide solutions will react with the iodide.

4. Ask for three student volunteers. Make sure the students are wearing chemical splash goggles; warn them to step back as soon as they pour. Have the students simultaneously pour the sodium iodide solution into the three cylinders containing the differing concentrations of hydrogen peroxide. Make observations. White foam erupts from the cylinder with the 30% peroxide the fastest, the 10% peroxide next, and only slowly rises up from the cylinder with 3% peroxide.

Part 2 — Old Foamey—Observing a Reaction Intermediate and Products

1. Place a 500-mL graduated cylinder in a large, plastic demonstration tray. Measure out 20 mL of 30% hydrogen peroxide and add it to the cylinder.

2. Add 1 small scoop (3–4 g) of solid Alconox® detergent to the cylinder and swirl the mixture to dissolve the detergent.

3. Measure out 5 mL of 2 M sodium iodide solution and, quickly but carefully, pour this into the cylinder. In a few seconds, copious amounts of white foam will be produced. Observe closely at the beginning of the reaction. A brown foam is produced at first but then turns white before it erupts out of the cylinder. This is due to the presence of the free iodine produced by the extreme oxidizing ability of the 30% hydrogen peroxide.

4. Notice the steam coming off the foam—this indicates that the decomposition reaction is quite exothermic. Light a wood splint and blow out the flame. Insert the glowing wood splint into the foam. The wood splint will re-ignite in the foam—this indicates that the gas in the foam is pure oxygen. Take the glowing splint out of the foam, re-insert it, and watch it reignite again. This can be repeated numerous times.

 

Discussion

Hydrogen peroxide decomposes to produce oxygen and water according to the decomposition reaction shown below:

                                                2H2O2(aq) 2H2O(l) + O2(g) + Energy

 

The reaction is quite slow unless catalyzed by a substance such as iodide ions, manganese metal, manganese dioxide, ferric ions, and many other substances such as yeast or even blood. A catalyst is a substance that, when added to a reaction mixture, participates in the reaction and speeds it up, but is not itself consumed in the reaction. The iodide ion is used as a catalyst in this demonstration.

 

Return to Top

Alka Seltzer Fizz (Temperature and Reaction Rate)

Return to Top

Kinetics in a Flash (Photon-Initiated Reaction)

Return to Top

The Gummi-Bear Torch (Perchlorate is a strong oxidizer)

Return to Top

Two Inch High 4th of July (Perchlorate meets the Flames Tests)

Return to Top

Microscale Electrolysis (Carbon electrode, CuCl2 soln)

Return to Top

Chemluminescence (Light Sticks & Luminol)

Return to Top