High School SMILE Meeting
1999-00 -- 05-06 Academic Years
Chemical Reactions

28 September 1999: Ed Scanlon (Morgan Park HS)
showed us how to change "orange juice" into a "strawberry float." (handout) He placed a beaker with a solution in it into the middle of a large tray. It had the color or orange juice. He then quickly poured a colorless liquid into it, and it immediately bubbled up and turned from orange to red! Neat! But how does this work? Why does it happen!? Ed told us the "orange juice" was made by mixing 100 gm baking soda into 200 ml water in a 500 ml beaker. Add 3 good squeezes of dish soap and mix well. Add about one tablespoon of glycerin and stir well. Add methyl orange solution (an acid-base indicator) to give the mixture an orange juice color. This was the "orange juice" that we saw. It was orange in color because the baking soda is basic, giving the solution a pH greater than 7. Methyl orange has an orange color in a basic solution. When Ed added the colorless liquid, he was adding about 100 ml of HCl to the "orange juice." The following reaction occurred:

NaHCO3 + HCl ---> NaCl + H2O + CO2 (gas).

The solution turns from basic to acid, shifting the pH to less than 7, and turning the color of the methyl orange indicator to a pink-red, "strawberry." The CO2 bubbled up as a gas. Soap and glycerin mixtures are used for blowing bubbles that will last, so the bubbles foam up and it looks like a "soda." What a good way to motivate student interest in chemistry! Thanks, Ed!

12 October 1999: Ken Schug [IIT Professor of Chemistry]

  1. He had 2 solutions that he poured into separate glasses and then mixed them together. The clear liquids turned a red color.
  2. He then divided a the mixture into 3 parts and put some of the original solutions into separate glasses. They got darker.
    A + B ---> C (red)
    C + A ---> C
    C + B ---> C
    Since the colors got darker this showed that there were all three compounds in each solution. So the equation would be
    A + B <===> C (equilibrium).
    The specific reaction was
    Fe+++ + SCN- <===> FeSCN+++

23 November 1999: Karlene Joseph [Lane Tech High School]
Materials:

  1. 3 plates with lots of surface area.
  2. 1 plate with water
  3. 1 plate with skim milk
  4. 1 plate with whole milk

We observed as she dropped different colors of food coloring on the plates. She then put a drop of dishwashing detergent (JOYTM) in the middle. We observed the colors dissipate, move, flow, and mix in the plates. The there was more activity with milk than with water.

The Science: The hydrophobic and hydrophilic reaction of the soap in milk and water. They also think the soap acts in milk because it cuts grease, which is a soluble fat in milk. She also introduced us to photosynthesis and its parts by doing a photosynthesis play. We acted as parts in photosynthesis, like H+, electron transport chain, light, photosystems I and 11, primary electron acceptors, thylakoid, stroma, ATP synthase, NADP, and Enzyme. A lot of fun was had in making ATP and NADPH in the light reactions.

12 September 2000: Karlene Joseph (Lane Tech HS)
had us write hidden messages on 8.5 in x 11 in paper by using phenolphthalein as an invisible ink. When we sprayed the paper with an ammonia solution (a base), the phenolphthalein turned red, and our messages appeared. Spraying with vinegar (acetic acid) caused the red to turn clear so the messages would disappear!

Now that she had us "hooked," Karlene did a titration of 15 ml of vinegar with ammonia using phenolphthalein as an indicator, and we found that about 12 ml of ammonia was needed to turn the vinegar to deep pink. Thus, the molarity of the ammonia was 15/12 = 1.25 times that of the vinegar. A wonderful phenomenological approach, Karlene!

10 October 2000: Pat Riley (Lincoln Park HS)
showed us how to investigate antacids as bases. Each group of us received 4 resealable plastic sandwich bags. To each we added these ingredients:

Note:  Cabbage juice is a chemical indicator of pH by its color. We used a hammer to crush one tablet of each of 4 different antacids, and we then added each to a different bag. We squeezed each bag until the color change caused by the antacid was complete, and recorded the approximate color. If a pH chart for cabbage juice were available, one would match the final color to a particular pH. [See the website http://chemistry.about.com/library/weekly/aa012803a.htm] The color changes as the acid (acetic acid in the vinegar) is neutralized to some extent by the antacid. The color is pink at the start (indicating low or acid pH), but was converted to blue, to green, to yellow (most basic pH) by the antacids. We inferred, from the color changes, that there were different amounts of base (sodium bicarbonate) delivered by the different antacid tablets. Cool! Such useful and good ideas to use with our students!

20 November 2000: Pat Riley [Lincoln Park HS] Ionic bonding versus Covalent Bonding
Materials

We placed a small amount of substance A on the slide in its respective location.  then we repeated this for substances B, C, and D.  We then carefully placed the slide on a hot plate.  We placed a thermometer on a clear part of the glass, and turned on the heat.  We then waited and watched to see which substance melted first:

Ionic compounds are usually hard, brittle, and water soluble, with high melting points.  They can conduct electricity when dissolved in water. Molecular compounds can be soft, hard, or flexible; are usually less water soluble; have lower melting points; and cannot conduct electricity when dissolved in water.  So, we guess that A and C were covalent, whereas B and D were ionic.

Then we performed a solubility test in water for each of the samples

A: Insoluble       B: Soluble            C: Soluble          D:  Soluble
Conductivity test:  we inserted an electrode to see if they conducted electricity:
A:   No           B:    Yes              C:   Yes          D:   No
In actuality we had the following materials:
A:  Carnauba Wax     B:  Table Salt          C:  Sugar          D:  Epsom Salts [MgSO4]

What is going to happen next?!
DON'T MISS IT!

30 January 2001: Therese Donatello (St Edwards School):  Double Displacement Reactions
She made the analogy between displacement reactions and two teams of dancers, male and female, exchanging partners.  In other words male: and female partners in dance ® positive ions: + and negative ions: -  in chemical reactions.

She considered the specific example of softening hard water:

Terry continued with a discussion of chemical nomenclature, discussing the meaning of the suffixes "ide" and "ate", as well as the need to balance the number of positive and negative ions in a compound, so that the net charge is zero (e.g., Fe2O3).


30 January 2001: Zoris Soderberg (Webster School)
started out with a discussion of Bessie Coleman, first US black female pilot.  She went to France to learn flying, and demonstrated that "nothing is too hard for you if you try hard enough".  Then Zoris talked about airplanes, and asked the question as to how planes fly.  The answer lies in the air, in that air is a form of matter, and can apply a force to overcome gravity and to keep the plane in the air.

How can you prove that there is air in the room?  Zoris took a 1 liter plastic bottle, containing nothing except air.  She put the cap on the top lip of the bottle, upside down, and held the bottle with both hands.  The cap danced, because of heating of air inside the bottle with her hands.  The kinetic energy of air molecules was transferred to the cap, causing it to dance as they hit it.

Next she put water in a Styrofoam™ cup, and put a 3" ´ 5" index card over the top of the cup, and inverted it.  Air pressure held the water in the cup, against gravity.

Zoris then mentioned Bernoulli's Principle as an explanation for how air holds up an airplane, but she did not go into the details.

13 February 2001: Pat Riley (Lincoln Park HS) Conservation of Mass in Chemical Reactions
used a large 2-pan balance.

10 April 2001: Ben Stark (IIT Biology)
showed us a model of glucose, an organic compound at the heart of the process of our metabolism.  In the biological world plants take in light energy, carbon dioxide C02, and water H20; and produce glucose C6H12O6 through photosynthesis.  Oxygen is given off as a byproduct.  The net chemical reaction is something like the following

6 C02+ 6 H20 ® C6H12O6 + 6 02

Oxygen and glucose combine within animals or fungi to produce energy, and carbon dioxide is given off as a byproduct.  This Oxygen - Carbon Dioxide Cycle requires both green plants and animals for its continuous operation.  Energy is stored in glucose in the detailed arrangement of the chemical bonds of the atoms of Carbon, Oxygen, and Hydrogen.

25 September 2001: Pat Riley (Lincoln Park HS)
Pat asked us how we might make a reaction --- such as dissolving sugar in water --- go faster.  Suggestions included the following:

heating, stirring, breaking sugar into smaller pieces

As to the effect of breaking sugar into smaller pieces, we suggested a refinement, since the surface area was increased and  the particle size was decreased.  [In this sense the "particle size" refer to the size of individual sugar molecules]. Which of these effects is more important?  She considered these two balls of the same size:

The Styrofoam™ ball has more surface area and should have more "sticking power".  Similarly, the increase in surface area, although it does not change the size of each sugar molecule, is the key parameter for speeding the reaction rate.

25 September 2001: Chris Etapa (Gunsaulus School)
Chris began by outlining a presentation to emphasize the importance of  laboratory safety.  In her class she took a raw egg and put it into a petri dish placed on the overhead projector.  Then, she added 2 drops of dilute Hydrochloric Acid [HCl].  As the protein "denatures" it looked dark on the screen.  This graphically illustrates the need to wear safety glasses, since our eyes contain the same types of protein as egg proteins, with the same sensitivity to acids.

Next she presented two mini-labs, which are used to get kids into a "science fair" mode.

  1. Observations on a penny
    We answered a series of questions about a penny from memory only.  Then she repeated the questions while we were looking at the penny.  She had us clean the penny by dipping it into vinegar for a few seconds, coating it with salt, and then rubbing it, so that we could see detail on the coin much better.  This cleaning procedure works very well!  We compared our observations with the answers given from memory, to learn about memories as well as our observational skills.
  2. Testing a hypothesis
    Chris lit an ordinary candle.  She asked us to hypothesize whether the candle burns because of (a) a special chemical in the candle or (b) the air in the room.  She covered the candle with a translucent container, and we saw the flame slowly go out.  Because the candle appears the same as before we covered it, the experiment suggests (but does not prove) hypothesis (b).  Note:  Oxygen O2 in air is necessary for combustion.

24 September 2002: Pat Riley [Lincoln Park HS] Milk of Magnesia
Pat led us in a discussion of the properties of Milk of Magnesia Mg(OH)2, [MOM for short] and its use in treating upset stomach, heartburn, and the like.  Then she began an experiment, based upon the laboratory exercise Upset Tummy? MOM to the Rescue! A colorful Antacid Demonstration [Flynn Scientific: http://www.flinnsci.com/]. She proceeded as follows:

  1. She put 20 ml of Mg(OH)2 into a beaker and added 200-300 ml of ice to slow down any reactions, and diluted to 800 ml with water. She then turned on a stirring motor.
  2. Next she added 4-5 ml of a universal pH indicator [http://www.purchon.com/chemistry/ph.htm], which turned purple.
  3. She handed out color cards to use to translate the "indicator color" into pH. [pH is a quantitative measure of hydrogen ion concentration, on a negative logarithmic scale; see http://www.purchon.com/chemistry/ph.htm.]  The color of the solution indicated that the pH was around 10, showing that the Mg(OH)2 is rather basic.
  4. She then added 2-3 ml of Hydrochloric Acid (HCl), and the stirred solution turned orange, indicating a pH of about 5. Over the next minute or so the colors continued to change sequentially to yellow (pH =6) to green (pH = 7-8) to blue (pH = 9) to purple (pH = 10). We concluded that the Mg(OH)2 was gradually neutralizing the HCl.  By the same mechanism, HCl in our stomachs is neutralized by Mg(OH)2.
  5. She added more HCl to the beaker, and the same cycle occurred.
Discussion:
  1. Mg(OH)2 is 5 times more soluble in hot tap water than cold tap water. In water ions are formed through the reaction
    Mg(OH)2 « Mg+2(aq) + 2 OH-(aq)
    The equilibrium constant for the reaction, determined from the concentrations, is much greater in hot than cold water.  In our experiment most of the Mg(OH) does not dissolve, and thus our MOM solution looks milky.
  2. The hydroxyl ions [from soluble Mg(OH)2] and the hydrogen ions [from HCl] combine to completion in the neutralization reaction
    OH-(aq) + H+(aq) ® H2O(l)
    The color indicator thus makes a gradual transition toward the "red" end of the visible spectrum.
  3. This reaction is a direct consequence of Le Chatelier's Principle:
    When stress (an excess of hydrogen ions) is applied to an equilibrium state, the reaction proceeds in the manner (direction) that will relieve the stress.
  4. The experiment can be continued many times, until all the Mg(OH)2 has dissolved and excess OH-(aq) is neutralized.
A terrific miniteach, Pat!

22 October 2002: Gary Guzdziol [Rosenwald School]      Handout:  Why do Candles Burn?
Gary put a bunch of candles and other items on the desk, and then passed out the handout.  He lit a candle and asked us what substances were present.  Our list included wax, tallow, scent, wick, and dye.  He then showed us that the wick alone burns completely and very rapidly when lit -- much more so than the wax by itself, which melts and drips but doesn't continue to burn on its own.  How come? We next observed a burning candle, and observed that the wax was melted at the top end, by the flame.  We decided that the candle burned because liquid wax moves up the wick by capillary action, and vapor is rising from the pool of liquid wax.  To test the idea, Gary blew out the candle, and quickly held a tube vertically with its bottom end at the wick, and then held  another lighted candle at the top end.  We saw an occasional spurt of flame move fairly quickly down the tube and re-light the candle.  For additional information on candles, wax, wicks, and flames, see the General Wax and Candle Company website http://www.generalwax.com/misc--candle-making/s___24.html.

03 December 2002: Carl Martikean [Wallace HS, Gary IN]     CATALYSTS 
Carl wrote these words Jersey, Guernsey, ***Angus on the board and asked us what it was. After a couple of rejected answers, the oldest person present (who had seen this before), said "that looks like a cattle list". Carl said, "that's right!--- and today I am going to use some catalysts". --- this was followed by loud groans! He then demonstrated several ways to catalyze the decomposition of hydrogen peroxide , H2O2, using a more concentrated solution (probably 6% by weight and sold as a bleach in hair product shops) than the 3% solution sold at drugstores. In each case he added some bubble soap solution to the hydrogen peroxide, before adding the catalyst, to produce a foam when oxygen gas is formed (the reaction produces oxygen gas and water). The catalysts used were dry yeast (which contains biological catalysts called enzymes), manganese dioxide [MnO2] (with a surface that can act as a catalyst), and some sodium iodide [NaI] (where the iodide ion [I-] is present), into the catalyst. He used the "glowing splint" method to verify that the gas formed really was oxygen. He then demonstrated another way to produce oxygen by adding some cobalt chloride [CoCl2] solution to household bleach. Finally, he enlisted our help in cleaning some grungy looking pennies by immersing them in acetic acid solution (CH3COOH: vinegar); not much happened until we added a pinch of salt (NaCl: sodium chloride) to act as a catalyst, after which the pennies sparkled like new! [*** Carl surely meant to say Aberdeen Angus, although he should not have ignored the Holstein Frisian breed -- PJ!]  Nice work, Carl

We began some informal scientific discussions, which we will follow up in the future.

10 December 2002: Barbara Pawela [retired]      Chemistry of Fire 
Barbara did a lecture experiment that involved the following interactive activities:

As expected, the candle in the larger jar burned slightly longer, but the jars were too close in size to show the effect dramatically.

Good work, Barbara!

11 March 2003: Tyrethis Penrice [McKinley Academy]      Soft Water vs Hard Water
Tyrethis
called our attention to the following websites on Soft Water:  http://pasture.ecn.purdue.edu/~agenhtml/agen521/epadir/grndwtr/softened.html and http://scifun.chem.wisc.edu/HOMEEXPTS/SOFTWATR.html. Then we performed the experiment described in the second website, making our own hard water by mixing 1 teaspoon [5 ml] of Epsom Salts [MgSO4] to 1 cup [250 ml] of distilled water.  We added several drops of liquid dish detergent to this mixture, as well as to an equal amount of distilled water.  We were not able to produce as much suds with the artificially hardened water, in comparison with distilled water.  Still, the difference was not as great as we had expected.  Why?

    Soft Water VS Hard Water
  1. Soft Water (Rainwater): Forms suds easily. No scum or film present. Presence of Na+.
  2. Hard Water ( Ca++ or Mg++ present): Prevents lathering of soap or shampoo. Produces soap scum on surfaces. Clogs pores and coats hair on skin. Residue serves as breeding ground for bacteria. Formation of a hard scale, which can clog plumbing and build up in water heaters, leading to an increase in utility bills and appliance failure.
  3. Methods of treating hard water:

Good Chemistry for everyday life!  Thanks, Tyrethis!

27 January 2004: Patricia Riley [Lincoln Park HS, Chemistry]        Types of Chemical Reactions
Pat
set the objectives as (1) performing reactions and recording observations, and (2) identifying reaction types from observations.  She provided us with the following materials, and we worked in groups of 4:

Candle   Beaker Magnesium Metal Strip: Mg
Test tubes Steel wool Hydrogen Chloride Solution: HCl
Sodium Chloride: NaCl Matches  Hydrogen Peroxide Solution: H2O 
Copper Metal strip: Cu   Wood splints Silver Nitrate Solution:  AgN03
Lump of sulfur: S8   Manganese (II) Dioxide: MnO2   
    Procedure [Precaution: wear eye protection when doing these experiments under the supervisiou of a knowledgable person]:
  1. Take a Candle. Record observations. Now, ignite the candle with a match. Record more observations. Hold beaker upside down over the flame for about 30 seconds. Record observation. Finally, extinguish the candle.
  2. Use the steel wool to shine up the copper metal strip.  Use a paper towel to wipe the copper strip clean.  Study the lump of sulfur and the clean copper strip and record your observations.  Now rub the clean copper strip with the lump of sulfur, and record your observations.
  3. Study the hydrogen chloride solution and the magnesium metal and record your observations.  Place 20 drops of hydrogen chloride into a clean test tube.  Drop ONE (1) small piece of magnesium metal into the test tube and record all observations. Wash out the test tube with lots of water.
  4. Study the silver nitrate and sodium chloride solutions.  Record your observations.  Place 20 drops of silver nitrate and 20 drops of sodium chloride into a clean test tube, and swirl the test tube to mix.  Record all your observations and then rinse out the test tube with lots of water.
  5. Study the hydrogen peroxide and manganese (II) dioxide and record your observations. Place 20 drops of hydrogen peroxide into a clean, dry test tube. Using a wooden splint, add a few grains of manganese (II) dioxide to the test tube. Record all your observations, and then rinse out the test tube with lots of water.
    Observations and Commentary:
  1. Combustion:  2CH2 + 3 O2 ® 2H20 + 2 CO2. Note that water is produced and condenses as a fog inside the beaker.
  2. Synthesis: Cu + S ® Cu S.  Note that the copper turns dark where the sulfur is rubbed against it, since CuS  [copper sulfide] is produced 
  3. Single Displacement: Mg + 2 HCl ® MgCl2 + H2.  Bubbles are formed in the test tube, since hydrogen is produced.
  4. Double Displacement: AgN03 + NaCl ® NaNO3 + AgCl.  The NaNO3 [sodium nitrate] is produced and remains in solution, whereas the insoluble AgCl: silver chloride forms as a white precipitate.
  5. Decomposition: 2H2O2 ® 2H20 + O2.  Note that MnO2 [manganese dioxide] serves as a catalyst for the reaction, but does not itself change.  Bubbles indicate that oxygen gas is produced.
Beautiful Phenomenological Chemistry! Excellent, Pat!

08 February 2005: Ron Tuinstra  [Iliana Christian HS, chemistry]       Honing Observation Skills with an Exercise Using Catalysts
Ron
started with two 250 mL Erlenmeyer flasks, and each contained about 50 mL of a clear liquid (it looked like water). Then Ron heated a Platinum wire red hot with a lighter, and placed it in the mouth of the first Erlenmeyer flask.  We watched as the wire rapidly lost its red glow as it cooled. Nothing unusual here! He did the same with the second flask, and  -- amazing to see --  the wire continued to glow! The first flask contained water, but the second contained methanol. The Platinum wire catalyzed the decomposition of the methanol vapors, releasing oxygen. The heat released from this reaction keeps the wire hot (glowing).  We turned out the lights in the room to see this clearly, and it was very dramatic!

Secondly, Ron placed a 500 mL Florence flask on the table.  It was stoppered and covered with Aluminum foil, so that we could not see inside.  As we watched, Ron removed the stopper, and in about 30 seconds a plume of steam boiled out the top of the flask!   How Come!? Ron had placed about 50 mL of 30% hydrogen peroxide in the flask.  A gram of manganese dioxide had been suspended by a string above the hydrogen peroxide.. The stopper held the string in place until Ron removed it, causing the manganese dioxide to drop into the hydrogen peroxide.  Manganese dioxide  acts as a catalyst for the decomposition of the hydrogen peroxide. The heat given off as a result of this reaction causes a vigorous boiling, which results in a dramatic plume of steam coming out the top of the flask, typically in about 30 seconds. The manganese dioxide can be recovered from the water (which remains after the decomposition) and reused!

Wonderful!!

22 February 2005: Carol Giles [Collins HS]            The Heat Treat
Carol gave us small packets [rectangular, about 4 cm ´ 7 cm ´ 0.3 cm] with a paper-like covering, which are called "Heat Treats" (http://www.amazon.com/Heat-Treat-Hand-Warmer-Pair/dp/B0007XFCIK). These are commercially available, very portable hand and foot warmers. Carol asked us to guess what was going on with the packets to make them work as warmers, but without opening the packet. The packets seemed to have a powder inside. When we crumpled a packet to mix the contents, or just held a packet tightly in our hands, the packet got warm.

Pat: We started an exothermic chemical reaction involving the contents of the packet. Perhaps the crumpling or just holding the packet in the hand provided enough activation energy to start the reaction. The website didn't have any information on what the chemicals were, so we opened a pouch. It was full of a blackish powder that almost had the consistency of potting soil. Carol did know what was in it:

But we still could not figure out the reaction that had occurred. Carol's class made their own from the ingredients and it worked fairly well, though not as well as the commercially available ones.

Now, who can find the reaction?  Thanks, Carol.

20 September 2005: Pat Riley (Lincoln Park HS, chemistry)         Rate-Determining Steps in Chemical Reactions
Pat had a beaker containing about 800 ml of a green liquid (water with green dye) for visualizing how sequential chemical reactions occur, and the concept of a "rate-determining step".  Pat attached three funnels to a ring stand, in a vertical stack.  These three funnels had differing sizes, spout lengths, and spout diameters. She poured the liquid into the top funnel. As it passed through the top funnel,  it went into and through the middle funnel and then through the bottom funnel, and was collected in a beaker underneath. Pat set up the apparatus so that the funnel with the most narrow spout was at the middle position. We saw that liquid accumulated in that funnel and slowly passed through it -- this seemed to be the rate-determining step. We measured the time elapsed from when we starting filling the top funnel until the liquid all landed in the collecting beaker.  Why does the spout diameter make so much difference?  Does it matter where we place the funnel with the narrow spout? Here are the data (averaged over several observers) taken for various locations of the narrow spout:

Narrow Spout Location  Time
Top 32.0 sec
Middle 31.0 sec
Bottom 31.5 sec
The order of the funnels did not seem to matter; the overall time of about 31.5 seconds being set by the rate limiting step -- passing through the funnel with the narrow spout! Bill Shanks suggested that we measure the time for the liquid to pass through the narrow funnel alone. We obtained 31.5 seconds again!  For additional information on rate-determining steps in sequential chemical reactions, see http://www.chem.neu.edu/Courses/1382Budil/ComplexChemicalKinetics.htmQuite nice, and most convincing!  Thanks, Pat.

07 February 2006: Pat Riley (Lincoln Park HS)            Covalent Bonding for the Inclusion Class
Pat
brought a tactile activity that she uses in her "inclusion class" which is a combination of special ed students and students with social problems. Pat is trying to teach them about chemical bonding in a way that avoids the math and other abstract concepts with which they have difficulty. Pat's handout sheet explains this in detail. It shows squares containing the symbols for various elements and a pattern of asterisks around the symbol which depict the valence electrons and their paired or unpaired arrangements. Each square then represents an atom of that element. Squares are placed together so that unpaired electrons match to make a pair, and squares are added until all electrons in the pattern are in pairs. For example, a pattern (compound) with only H and S atoms, which uses the fewest atoms overall, would have one S square and two H squares, each H square adjacent to an "unpaired" electron in the S square, and would be H2S.

      ___    _____    ___ 
| | | ** | | |
|H *|==|* S *|==|* H|
|___| |__**_| |___|

H2S Hydrogen Sulfide [covalent bonds as indicated by " == "]
Although this is a low-functioning group, Pat is able to impart some aspects of electronegativity, and the order of electronegativity among atoms, how atoms come together to form compounds (as shown above), various energy states (the compounds being in a more stable state than the individual atoms), and how compounds are named. Pat

Nice stuff!  Thanks, Pat.