• Soap, Candles, and Colored Flames

    Structure of Sodium Stearate Micelle

    What is soap?  "Soap" is generally composed of sodium salts of long-chain carboxylic acids.  Soap helps us clean things because it allows nonpolar hydrophobic 'oils' to dissolve in polar water:  the soap molecule itself has a nonpolar alkane end, and a polar carboxylate end.  We used this property of soaps last year to do Ben Franklin's oil slick experiment and estimate the size of a molecule:  stearate molecules formed a single-molecule layer on water with the hydrophilic carboxylate groups touching the water while the alkane chains floated upwards.  We also use soap all the time, to clean things:  soap molecules form spheres in water called 'micelles' which can encapsulate oily, hydrophobic substances and allow them to dissolve in water. 

    Candles with colored flames

    Some of you may have purchased colored-flame candles.  We can try to make our own as part of this project!  As you remember from last year, compounds of lithium, strontium, copper, are known to give bright colors in a flame.  We have plenty of lithium compounds in our inventory:  LiOH, LiNO3, LiCl and so on.  If we could dissolve these compounds in candle wax or gel fuel, we might be able to make a colored-flame candle of our own.  But this won't work, because these are all very polar ionic compounds that do not dissolve in nonpolar candle wax or gel.  How can we get lithium to dissolve in candles?  We can make a 'lithium soap', where instead of making a sodium salt of a long-chain carboxylic acid, we make a lithium salt.  If this works, we can look into other elements like strontium, copper, potassium, and so on.  We can also make lithium propionate, which has a shorter 3-carbon carboxylic acid, and see if this works in candles as well.  Professional colored-flame candles are made with a fuel called trimethyl citrate which burns at a hotter temperature, which we could possibly make, but I'd rather buy:  unfortunately, I haven't seen any for sale. 

    Part I:  Making Soap (Sodium Stearate)

    Hydrolysis of Glycerol Tristearate

    Reaction:  C3H5(C17H35COO)3 + 3 NaOH → C3H5(OH)3 + 3 C17H35COONa

    When planning this first reaction, use Melanie Messer's "Let's Lather Up" activity as a guide. In this reaction, we are using coconut oil as our starting material, which we assume to be almost entirely glycerol tristearate ( C3H5(C17H35COO)3) which has a density of 0.86 g/mL according to Wikipedia.  We react it with sodium hydroxide, NaOH, also known as 'lye'.  The salt (NaCl) solution is used as a polar solvent to help separate the sodium stearate product from the glycerol (C3H5(OH)3) product:  it is not a reactant so there is no need to worry about exactly how many moles of NaCl are present.  Note that 'glycerol' is also often called 'glycerine' and is useful in many personal care products.  For purposes of finding percentage yield, we should focus on the sodium stearate (C17H35COONa) product.  

    As we may have multiple groups on this project, I may assign some groups to using Irene Cesa's "Making Soap" procedure instead.  This procedure will require the soap to sit and 'cure' for a few days in order for the reaction to complete:  be aware that the 'uncured' soap will still contain fat and NaOH, and must be treated with as much caution as any other NaOH mixture.  This procedure is somewhat safer as it does not involve the use of flammable solvents (ethanol, aka 'alcohol'!).

    Safety Note:  Melanie Messer is writing from the 1970s, when people really did not care about safety at all!  It's been a long time since then, and we are far more into safety these days.  You should know that the 8M (8 moles/liter) NaOH solution is extremely corrosive and can cause severe burns or even blindness if it gets in your eyes.  Be very, very careful when adding it to your hot oil/alcohol mixture, as it may start to boil and splatter!  Wear full PPE!  Any contact with NaOH should result in you immediately washing under cold running water for 10-15 minutes:  otherwise, you may find you have a chemical burn in an hour or two!  You will not handle solid NaOH, which is even more corrosive:  I will provide you with the 8M NaOH solution and will supervise your addition of it to your reaction. 

    Safety Note:  People are often advised to wash an NaOH burn with vinegar instead of water, so that acid will neutralize the base.  This is common in popular media, including a certain old movie where a main character makes soap. This seems believable, but it is wrong.  If exposed to NaOH, wash with water.  Lots and lots of water!

    Safety Note:  Melanie is using 1970s-era teaching lab equipment, heating a mixture of flammable oil and alcohol in an open dish over a Bunsen flame.  This is very likely to catch fire and she knows this, so she advises you to keep a pile of wet paper towels handy to extinguish it.  I read this and I think nope nope nope.  So we will do this reaction, not in an open dish, but in an Erlenmeyer flask.  We will not heat it over a Bunsen flame, but will use a temperature-controlled digital hotplate.  We will maybe start off at 90-100 C until all the alcohol solvent has boiled off, and possibly raise the temperature if needed to complete the reaction.  If the reaction ignites, which it might, as it is giving off flammable fumes, the correct way to extinguish it is to place a stopper in the mouth of the Erlenmeyer flask.  Not by using wet paper towels.  And no, we don't blast it with a fire extinguisher either, that would make it worse.  Just put a cork in it. 

    Part II-A:  Making Lithium Soap (Lithium Stearate aka 'Lithium Grease')

    Reaction:  C3H5(C17H35COO)3 + 3 LiOH → C3H5(OH)3 + 3 C17H35COOLi

    We could put sodium stearate in a candle, but the flame would just be your basic yellow sodium flame which kinda looks like a regular flame.  Lithium should give a prettier bright red, so let's make lithium stearate!  This can be purchased as a component of 'lithium grease' for automotive use, but we can make our own.  I will provide you with an 8M solution of LiOH which you can use instead of NaOH.  Otherwise, this reaction should be pretty much the same.  If it isn't, we'll find out.  Be observant!

    Safety Note:  The same safety note from the previous applies:  LiOH is just as corrosive as NaOH.

    Part II-B:  Making Lithium Propionate

    Reaction:  C2H5COOH + LiOH → H2O + C2H5COOLi

    Lithium propionate is similar to lithium stearate, but has a shorter carbon chain.  Will this make it better for a colored-flame candle, or worse?  I don't know yet, so we can find out!  We do have propionic acid (C2H5COOH) on order and will hopefully have it in stock.  Plan this reaction:  start with a solution of LiOH in water which I will provide you with, add an equimolar amount of propionic acid, verify that the pH is neutral by pH paper or a pH meter, and see if we can precipitate the product C2H5COOLi.  It may precipitate on its own.  If not, you can try the usual methods such as boiling down the solution, refrigerating the solution, or saturating it with NaCl to make it more polar and cause the nonpolar product to precipitate.

    Safety Note: See previous.

    Part III:  Try to Make a Colored-Flame Candle

    We can figure this one out:  try to mix your lithium stearate with some molten candle wax, or with gel fuel, and make a lithium colored-flame candle.  We can make a plain candle and a sodium stearate candle too, to compare.  Then we light them and see if we get colored flames.  I'm looking forward to this!

    Commercial colored-flame candles, as I mentioned, use trimethyl citrate as fuel.  I haven't found it for sale yet, and I'm not totally committed to making it.  Perhaps a group working on the Ester Project would care to attempt to synthesize this trimethyl citrate fuel from citric acid (mostly harmless) and methanol (flammable and very toxic)?  We shall see.

    Messer, Melanie B., et al. "Let's Lather Up" Introductory Experimental Chemistry: Teachers Manual. Prentice-Hall, 1977.

    “Stearin.” Wikipedia, Wikimedia Foundation, 25 Jan. 2019, en.wikipedia.org/wiki/Stearin.

    Cesa, Irene G. "Making Soap" Laboratory Experiments for General, Organic & Biological Chemistry. Flinn Scientific, 2015.

  • Copper Aspirinate:  Pharmaceutical and Pigment

    Copper Aspirinate Copper Aspirinate Structure

    Copper (II) acetylsalicylate, known commonly as 'copper aspirinate', is a salt known as a 'chelate'.  A 'chelate' is a compound where molecules known as 'ligands', usually organic, form multiple covalent bonds to a metal ion.  According to the Wikipedia entry, copper aspirinate has been proven effective as a medical treatment for rheumatoid arthritis, although it has not been approved as a prescription medication at this point in time.  It has also been used as a pigment.  We generally use basic copper carbonates for blue pigment in the Paint Project, but we could certainly use a more vivid blue!  We will not use it as a medicine, as our lab is not FDA approved.

    The synthesis is described in the Wikipedia entry for copper aspirinate, and seems easy enough, so let's try it.  Starting materials are copper (II) sulfate and sodium carbonate, both of which are easily purchased as household chemicals, and acetylsalicylic acid which is the active ingredient of aspirin tablets and can easily be purified.  

    Part I:  Synthesis of Copper (II) Acetylsalicylate

    Deprotonation of Acetylsalicylate

    Formation of Copper (II) Acetylsalicylate Complex

    Reaction:  2 C8H7O2COOH + Na2CO3 + CuSO4·5H2O → Cu(C8H7O2COO)2 + Na2SO4 + CO2 + 6 H2O

    We need to start with purified acetylsalicylic acid for this reaction - it is purified from Dollar Tree aspirin tablets.  I have 78 grams in inventory, which should be sufficient for the year, prepared by students from previous years using the 'Purification of Acetylsalicylic Acid' handout below.  When calculating your reaction table, please be aware that the blue crystals sold as commercial copper (II) sulfate are actually a pentahydrate:  this increases the molar mass by about 5 x 18 (why?).  For an initial preparation, let's plan to prepare 5 grams of Cu(C8H7O2COO)2 .  If we find it a useful pigment, we can always make a larger batch afterwards.

    Wikipedia warns us that we should plan on an excess of acetylsalicylic acid (C8H7O2COOH) in order to prevent side reactions between leftover carbonate (CO3-2) and copper ion (Cu+2), which will contaminate our product with copper carbonates.  For this reason, when preparing your reaction table, use 2.5 molar equivalents of acetylsalicylic acid (C8H7O2COOH) instead of the 2.0 you would use from the balanced equation!

    As described in the Wikipedia entry, this is a 2-step reaction.  Initially you start with a solution of sodium carbonate.  I would start by weighing out the calculated amount of sodium carbonate and then adding enough water to easily dissolve it.  This reaction will not work without water as a solvent!  Add acetylsalicylic acid to this solution, with stirring.  Filter to remove unreacted acetylsalicylic acid. 

    You should make a solution of your copper sulfate (CuSO4·5H2O) as well, using as much water as is necessary to dissolve it:  it is helpful to stir and warm (do NOT boil) during this process. 

    At this point, you can mix the two solutions and the product should precipitate.  Does it make a difference whether you add the acetylsalicylate solution to the copper sulfate solution, or whether you add the copper sulfate solution to the acetysalicylate solution?  This is something you can find out by trying it.  I've never done this reaction so I don't know the answer yet.

    Collect the product, wash with a small portion of ice water, and allow it to dry overnight before weighing it.  When calculating a %yield for this reaction, you should do it with reference to sodium carbonate or copper sulfate - you are using an excess of acetylsalicylic acid here, so some will remain unreacted.

    Safety Note:  Copper (II) sulfate is an irritant and is toxic!  Sodium carbonate is a corrosive base, capable of causing irritation or even burns.  Full PPE must be worn during this experiment!  Leftover waste containing copper (II) ion should be saved for appropriate disposal:  if it's blue, basically, it does not go down the drain!

    Part II:  Synthesis of other D-Block Metal Aspirinates

    Compounds of d-block metals (aka "transition metals") are often colorful due to unpaired electrons in d-orbitals, which can be excited to higher energy levels by various wavelengths of visible light.  Now that you've made copper (II) acetylsalicylate, you can make other analogs by replacing copper (II) with other d-block metal ions.  

    We can do this by replacing the copper sulfate (CuSO4·5H2O) with other transition metal compounds I have in inventory.  I have iron (II) sulfate heptahydrate (FeSO4·7H2O) in stock, and am curious about the color and properties of iron (II) acetylsalicylate (Fe(C8H7O2COO)2).  I also have manganese (II) sulfate monohydrate (FeSO4·H2O), so we could make manganese (II) acetylsalicylate (Mn(C8H7O2COO)2).  As you may remember, iron forms both Fe+2 and Fe+3 ions:  I have iron (III) chloride in stock as well, so perhaps we could try to make iron (III) acetylsalicylate (Fe(C8H7O2COO)3) - you will need to write a new balanced equation here as the stoichiometry is different.  I have cobalt (II) chloride in inventory as well:  as this compound is more toxic than the iron or manganese compounds, we might make cobalt acetylsalicylate on a much smaller scale.  Cobalt compounds can be toxic, but often do show very vivid colors. 

    I do not expect any one group of students to make all these compounds:  definitely it's reasonable for you to make a second acetylsalicylate from this menu, though.


    “Copper Aspirinate.” Wikipedia, Wikimedia Foundation, 18 Aug. 2019, en.wikipedia.org/wiki/Copper_aspirinate.

    “Chelation.” Wikipedia, Wikimedia Foundation, 15 Aug. 2019, en.wikipedia.org/wiki/Chelation.

    Slotsky, Morris M. "Purification of Acetylsalicylic Acid", 2016, Diman Regional Vocational Technical High School. Honors Chemistry I Course Materials.

  • Bananas, Apples, Pineapples, Peaches, and Blackberries:  Ester Synthesis Project

    Ester functional group Mechanism of acid-catalyzed esterification

    Esters are a class of organic compound which often have distinct odors (and flavors, but we don't eat anything we make!)  You may have already synthesized an ester, ethyl salicylate, from aspirin, but if not, it's okay.  Esters are generally formed from a carboxylic acid (R-COOH) and an alcohol (R'-OH).  This reaction is a double displacement where the -OH of the acid and the -H of the alcohol become water, H2O.  Organic chemists refer to this type of reaction as a 'condensation'.  This reaction will not happen without a catalyst.  Either acid or base catalyst may be used, but we generally use sulfuric acid (H2SO4).  Sulfuric acid is a good catalyst for esterifications because it tends to soak up the water product and drive the reaction forward.

    Safety note:  Concentrated sulfuric acid is extremely hazardous, a single drop can produce a burn with a lifelong scar and can permanently blind an eye instantly.  You will never handle it: I will either mix it with your reaction myself, or provide you a premade mixture of sulfuric acid.  The mixture is still hazardous, but not as much as concentrated sulfuric acid.  Gloves, goggles, and an apron are a must!

    Part I:  Synthesis of isopentyl acetate (Commercially sold as "Banana Oil")

     Isoamyl acetate synthesis

    Reaction:  C5H12O + C2H4O2 → C7H14O2 + H2O

    Please base your lab procedure on 'Synthesis of Isopentyl Acetate'.  We currently do have glassware for a 'reflux' setup.  We do not have an 'electric flask heater' but we do have a 'sand bath' setup which we have used successfully in the past.  We do have separatory funnels as well.  We may skip the 'distillation' step here.

    Once you've got a workable procedure, we can try it!

    Part II:  Synthesis of other esters.

    Reaction:  You will be figuring this out!

    We will select at least one ester to make as described by Kennedy's "Infographic".  We hope to have other alcohols available - ethanol, methanol, and butanol, as well as isopentanol.  We also have salicylic acid (made from aspirin by previous students) available, and hope to have propionic acid as well.  

    We will not be making ethyl acetate, as I have about a gallon of it already and it doesn't smell that nice (it's used in nail polish removers).  Also, let's never make pentyl salicylate again, we did once and it was annoying (smelled like flowers in tiny amounts, but very unpleasant in larger quantitites).  Methyl salicylate and ethyl salicylate are nice wintergreen odors but we have made them in previous years.  Really I'm interested in new esters!  Kennedy warns us that butyl salicylate has a 'strong' blackberry odor, so we would want to do this on a small scale if we chose it.

    How do you plan your synthesis?  In general, you want to consider both the reaction setup and the ease of workup.  You generally want a liquid solution for the reaction - if you are using a solid acid, you want enough alcohol to completely dissolve it.  However, if you can use excess acid, this may prove easier to remove in workup by washing with sodium bicarbonate solution.  Excess methanol and ethanol should be removed by distillation before workup in the separatory funnel, as they will make it difficult to separate the reaction into 2 layers.  


    Synthesis of Isopentyl Acetate. Department of Chemistry, Chulalongkorn University, 26 June 2013, www.chemistry.sc.chula.ac.th/bsac/Org Chem Lab_2012/Exp.8[1].pdf. 

    “Isoamyl Alcohol.” Wikipedia, Wikimedia Foundation, 7 Mar. 2019, en.wikipedia.org/wiki/Isoamyl_alcohol. (note:  "isopentyl alcohol" is also known as "isoamyl alcohol")

    “Acetic Acid.” Wikipedia, Wikimedia Foundation, 26 July 2019, en.wikipedia.org/wiki/Acetic_acid.

    “Infographic: Table of Esters and Their Smells.” James Kennedy, 12 Dec. 2013, jameskennedymonash.wordpress.com/2013/12/13/infographic-table-of-esters-and-their-smells/.

  • Manganese Violet Pigment Project - A Quest for New Colors

    Manganese Violet Pigment

    Manganese violet is a pigment we use every year for the 'Make Your Own Paint' activity.  We found an optimized procedure giving high-quality pigment, thanks to the Gold Alpacas and the Magnesium Monkeys.  This year, we will repeat this, and attempt to prepare new pigments by replacing the ammonium ion (NH4+) with other s-block metal ions.  There are multiple paths you can take through this project!  Every group is not expected to do everything.  

    Part I-A: Preparation of Manganese Violet (Classic Method)

    NH4MnP2O7 "Manganese Violet" sample

    Reaction:  NH4H2PO4 + MnO2 + H3PO4  → NH4MnP2O7 + 2½ H2O + ¼O2 (hypothetical, reaction is not completely understood)

     As described by Lee and Browne, monoammonium phosphate (NH4H2PO4 ), manganese dioxide (MnO2), and phosphoric acid (H3PO4 ) are heated in a ceramic dish together.  Once the reaction is complete, the mixture is boiled in water to remove soluble impurities, collected on a filter, and dried. 

    Here's the tricky part:  the Lee and Browne paper has amounts which are wrong.  The Gold Alpacas and Magnesium Monkeys found the correct amounts to obtain good pigment:  we ended up using 2.5 grams of MnO2 (0.029 moles) with 10.6 grams (0.092 moles) of NH4H2PO4 and 12.5 grams of 85% H3PO4.  Hopefully you can reproduce their results.  A link to the Manganese Monkeys final presentation is provided below.

    It may give a better pigment if the ingredients are ground together in a glass mortar before being heated in the ceramic dish. 

    Safety Note:  H3PO4 can cause irritation or burns, and MnO2 will permanently stain anything it touches!

    Part I-B: Preparation of Potassium Manganese(III) Pyrophosphate ('Kyleigh Purple')

    Reaction:  2 K2HPO4 + 4 MnO2 + 6 H3PO4 → 4 KMnP2O7 + 10 H2O + O2 (hypothetical!)

    Sample of KMnP2O7

    This pigment was synthesized last year by various students, and is named after the first student who had an attempt at it.  As you see, it is more of a 'purple' and less of a violet.  This may be useful in the Paint Project, as it can hopefully be mixed with white and yellow pigments to obtain a proper pink color.  Every year, students ask for a better pink pigment!  I spent some time over the summer working on this pigment, and developed a procedure for the synthesis.  This pigment was described in the German patent of Sander and Mansmann below, although we do not prepare it using the same method.  

    The procedure I used involved starting with 0.869 grams of MnO2 (0.0100 mol), which was cooked in a crucible over a Bunsen flame for about 30 minutes to convert it to Mn2O3 as described in Part II-A below.  To this I added 6 molar equivalents (0.0600 mol) of K2HPO4, potassium hydrogen phosphate, and 8.5 equivalents (0.0850 mol) of 85% phosphoric acid H3PO4.  While I did not grind these ingredients together in a glass mortar and pestle, we may consider doing this so that they will be mixed more throughly.  As the phosphoric acid is only 85%, you need to first convert moles to grams and then divide by 0.85 to find the mass of 85% solution.  The mixture was heated to 120 C in a crucible on a hotplate in order to bring it to a melt, and was then heated at 300 C with the lid on the crucible for a full 12 hours.  It does take 12 hours for the reaction to complete:  Mn2O3 residue will contaminate the pigment otherwise.  The product mixture was then stirred in about 150 mL of hot water to remove phosphoric acid impurities, and was collected on filter paper as a fine powder.

    Hopefully you will be able to reproduce this result.  We may consider increasing the scale of this reaction, as this pigment could be a valuable addition to Paint Day!

    Safety Note:  H3PO4 can cause irritation or burns, and MnO2 will permanently stain anything it touches!

    Part II-A: Preparation of Manganese Violet (Potential Improvement)

    Reaction 1:  4 MnO2 →  2 Mn2O3 + O2
    Reaction 2:  
    2 NH4H2PO4 + Mn2O3 + 2 H3PO4 → 2 NH4MnP2O7 + 5 H2O

    It is thought that MnO2 decomposes to Mn2O3 and Ogas under the conditions of the reaction, and the Mn2O3 actually forms the product.  All we have to do to convert MnO2 to Mn2O3 is heat it in a crucible over a Bunsen burner.  You can use the original 2.5 grams specified in the procedure in part I-A.  This should lose mass during the decomposition:  you should be able to determine from mass measurements whether your conversion of MnO2 to Mn2O3 was successful.  You can then proceed with the procedure as in part I-A, using Mn2O3 instead of MnO2 .We are interested in whether the product will have a more vibrant color or otherwise show a different appearance.

    Part II-B: Preparing Related Analog Compounds (XMnP2O7)


       X2CO3 + 2 H3PO4 → 2 XH2PO4 + CO2 + H2O      ['X' is a Group I metal]
       XOH + H3PO4 → XH2PO4 + H2O      

       XH2PO4 + MnO2 + H3PO4  → XMnP2O7 + 2½ H2O + ¼O2 (hypothetical)

    Last year, we successfully prepared CsMnP2O7 and KMnP2O7, which are similar compounds containing the ions Cs+ and K+ instead of NH4+.  Neither one was as brilliantly colored as regular Manganese Violet.   CsMnP2O7 was darker and bluer in hue, while KMnP2O7 is slightly darker and also more reddish/purple.  We found that, to get a good bright color for these pigments, we had to use a much smaller amount of Mn:  whereas for Manganese Violet we can use about a 3:1 ratio of moles Mn/moles NH4+, we needed more like a 10:1 ratio for moles Mn/moles K.

    We may try for the sodium (Na) or lithium (Li) analogs this year.  We do have LiOH in stock, although this is a highly corrosive substance which must be treated with care - I might give it to you as a dilute solution.  I may be able to order some Li2CO3, as it is safer to handle.

    I suggest that you do not try to precalculate the molar quantities when preparing LiH2PO4 or NaH2PO4:  start with a solution of LiOH, Li2CO3, Na2CO3, and add in H3PO4 with stirring until the mixture measures acidic by pH paper.  Then evaporate and crystallize to obtain your product, which you can use instead of NH4H2PO4 in your pigment synthesis.


    Lee, J. D., and L. S. Browne. “The Nature and Properties of Manganese Violet.” Journal of the Chemical Society A: Inorganic, Physical, Theoretical, 1968, p. 559., doi:10.1039/j19680000559.

    “Magnesium Monkeys- Honors Seniors Project.” 2017-2018 Honors Presentations, Diman Regional Vocational Technical High School, 15 May 2017, docs.google.com/presentation/d/1l1VlbXMD1nsCObiPP6TfWvg3Ti7ybgI0bi135-Voo_4/edit?usp=sharing.

    Sander, Hans, and Manfred Mansmann. Use of Potassium Manganese Pyrophosphate for the Production of Coloured Metallic Coatings, and Coloured Metallic Coating Materials Thus Prepared. 11 Mar. 1994.  Patent #DE4314268A1

  • Anodizing Aluminum

    Oxide layer in anodizing colorful anodized aluminum pieces

    Anodizing is a process for protecting and coloring aluminum objects.  Aluminum metal naturally forms a film of oxide, Al2O3, which protects it from corrosion.  However, this film is fairly thin.  In anoziding, aluminum metal is used as the anode (positive electrode) in an electrolysis cell:  this oxidizes the aluminum and forms a thicker film of Al2O3 containing cellular structure, or 'pores'. These pores can then be filled with inorganic pigments or organic dyes and sealed by boiling in water. This will give the aluminum piece a colorful protective coating! In the past, Diman has sent out items for anodizing, but I am hoping that Honors Chemistry II can figure out how we can do anodizing in-house.

    This year, I want to focus on the procedure given by Melanie Messer for making a 'gold' colored coating.  The gold coloring is produced by taking the aluminum anode out of the electrolysis cell, treating it with a solution of iron (III) chloride (FeCl3) and ammonium oxalate ((NH4)2C2O4), and then sealing it in boiling water.  Thankfully, I have iron (III) chloride in stock.  However, I do not have any ammonium oxalate.  Therefore, we need to synthesize it as a first step so that we can prepare the coloring solution.  In the second part of the project, we will make an anodization cell.  

    Part I:  Synthesis of Ammonium Oxalate 

    Reaction:  2 NH3 + H2C2O4 → (NH4)2C2O4

    Safety Note:  Oxalic acid (H2C2O4) is a toxic, corrosive solid.  Ingesting it can cause kidney stones, kidney failure, and even death - don't eat or drink it!  It can also cause irritation or even burns on contact.  You will not receive pure oxalic acid - I will supply it to you as 0.5 M (0.5 moles/Liter) solution in water.  Oxalic acid should not be heated, as it can give off irritating or even toxic fumes!  Please consult the SDS.

    Safety Note:  Ammonia (NH3) is a highly irritating and malodorous gas.  It will be supplied as aqueous solution approximately 0.5 M - 1 M (0.5-1 moles/Liter) solution, as purchased at Dollar Tree or similar vendors.  Ammonia solution should not be heated, or it will give off irritating fumes.  Please consult the SDS for this compound.

    Safety Note:  The ammonium oxalate product itself is toxic, and can give off toxic fumes if heated.  Please consult the SDS for this as well!

    The common theme of these safety notes is that we do not risk overheating either the reactants or the products of this reaction!

    I suggest you plan to synthesize about 0.05 to 0.1 moles of ammonium oxalate for your first attempt (how many grams is this?  Put it in your table!).  Start with my 0.5 M stock of oxalic acid.  Use an Erlenmeyer flask in a tub of icewater to keep it cool.  Slowly add household ammonia (NH3) solution with stirring (a magnetic stirbar would be great here) and monitor the pH - do this addition until the solution is no longer acidic, but neutral.  At this point, we should have a solution of ammonium oxalate ((NH4)2C2O4).  

    We don't want to risk boiling this solution dry, but fortunately we can precipitate our product another way.  Adding ethanol (ethyl alcohol) or isopropanol (isopropyl alcohol) will reduce the polarity of the solution, lowering the solubility of ammonium oxalate.  I'd start by diluting the mixture with a roughly equal amount of alcohol.  The solution should become cloudy as a precipitate forms.  You can refrigerate the solution overnight to completely precipitate the ammonium oxalate product.  This can be collected on filter paper, left overnight to dry, and weighed to calculate a percent yield for the reaction.

    Part II:  Preparing "Gold" Coloring Solution

    Melanie Messer advises us to prepare a 1 M (1 mol/Liter) solution of iron (III) chloride, a 1 M solution of ammonium oxalate, and mix equal volumes of each.  You can also just dissolve the iron (III) chloride and ammonium oxalate together to make a final concentration of 0.5 M (0.5 moles/Liter) of each.  Figure out how much ammonium oxalate you have available and decide how much iron (III) chloride you will need:  I should have enough for any reasonable need.  This solution can be stored for later use.  

    Safety Note:  This solution is toxic due to the ammonium oxalate (see SDS again).  It will also stain anything it touches permanently due to the iron (III) chloride.  Be very, very careful with iron (III) chloride.  I spilled a tiny amount of this chemical once, while etching a circuit board.  My wife has forever forbidden me to do chemistry at home again.  Wear aprons when handling it.  Do not wear your interview suit or prom dress to this lab, either.

    Part III:  Anodization of Aluminum Sheet

    Follow the procedure given by Melanie Messer below.  Remember that the piece of aluminum you wish to anodize should be attached to the positive terminal of the power supply or 6 Volt battery (whichever we have handy).  I will provide you with the 2 M (2 moles/Liter) sulfuric acid solution used for the electrolysis apparatus - while sulfuric acid is sold commercially as 18 M concentration, this is extremely dangerous and students will never handle it in my lab.  You can handle the 2M with the usual precautions.

    I am not entirely convinced that this setup is totally safe:  if it heats up too much, there is a chance that the aluminum and sulfuric acid will violently react in an exothermic manner.  We should have ice handy in case we need to cool the reaction down!  We should also probably keep the electrolysis setup in a tub of baking soda, so that any spills can easily be neutralized.

    Once the aluminum piece has been anodized in the cell, it should be taken out, immersed in the gold coloring solution for 15 seconds, rinsed in cold water, and then placed in boiling water for 1-2 minutes to seal in the color as described in the procedure.

    Safety Note:  Please look up SDS for 'battery acid' sulfuric acid, roughly 2M concentration.  This is corrosive and can cause burns, or even blindness if splashed in your eyes.  Wear PPE!  Do not wear your best clothes, or when you get them out of the laundry there may be holes where the acid touched them.

    Part IV:  Further Directions

    We might do a more visually impressive aluminum object at some point.  Or we might look into other anodization dye mixtures.  Do your own research!


    Messer, Melanie B., et al. "Color Me Gold" Introductory Experimental Chemistry: Teachers Manual. Prentice-Hall, 1977 pp 403-405.

  • Benzoin Synthesis:  From Almonds to Zombies.

    Benzoin Powder Benzoin Chemical Structure

    I have always wanted to make benzoin, but never got around to it.  This is the year!  What is benzoin?  It is an organic compound, off-white in color, with a distinctive 'camphor-like' odor.  It was first observed as a byproduct of almond essential oil production.  Today it is used as a pharmaceutical synthesis intermediate, as a means of removing oxygen in industrial processes, and other applications.  Why do I want you to make it?  I read a science fiction novel once, "Demon" by John Varley, where it turned out to be a key ingredient in a potion that killed zombies.  Hopefully we'd have a little on hand, if there was ever a zombie situation.  

    Please note that more than one substance is sold as 'benzoin'.  A natural tree resin is also called 'benzoin' or 'benjamin' but does not contain any of the organic compound 'benzoin'.  Confusing!  

    We will start this synthesis from benzaldehyde (C6H5CHO), a substance described in the wikipedia reference below.  Benzaldehyde is found in commercial almond flavoring, but only at a 5-10% concentration in a water-alcohol solvent mixture.  Our first step will be to extract the benzaldehyde from the almond flavoring by reacting it with sodium bisulfite (NaHSO3, aka 'tree stump remover') to precipitate an intermediate compound sodium alpha-hydroxybenzylsulfonate (NaC7H7O4S).   In our second step, we will convert the intermediate back to benzaldehyde by reaction with sodium carbonate (Na2CO3, 'soda ash').  This process allows us to purify the benzaldehyde from the flavoring, according to the reaction scheme below. 

    Reversible Conversion of Benzaldehyde to Sodium alpha-hydroxybenzylsulfonate


    Finally, we will follow my old college organic lab manual by Lehman where benzaldehyde is reacted with sodium hydroxide (NaOH) using vitamin B1 (thiamine hydrochloride) as a catalyst, to produce benzoin.  

    Two equivalents of benzaldehyde condense when treated with NaOH and thiamine to form one equivalent of benzoin.

    Part I:  Isolating Benzaldehyde from Almond Flavoring as Sodium Alpha-hydroxybenzylsulfonate

    Reaction:  C6H5CHO + NaHSO3 → NaC7H7O4S

    In a trial run, I used a 2 ounce (59 mL) bottle of artificial almond extract.  If you find this scale convenient, we may use it, but we may scale up if we have more material available.

    The almond flavoring was stirred in a beaker and table salt (NaCl) was added until the solution was fully saturated.  This increases the polarity of the solution and makes it easier to extract with nonpolar solvent.  The nonpolar solvent I used was ethyl acetate, 10 mL.  Shake the two layers together in a separatory funnel until they are well mixed, and then allow the layers to separate.  Collect the top layer, which is an ethyl acetate solution of benzaldehyde and other nonpolar components of almond flavoring.

    Vogel gives a recipe for bisulfite solution in the reference below (see 'sodium bisulphite test' and 'purification of commercial hexanone' for details).  I used 5.2 g of sodium bisulfite (NaHSO3 ) dissolved in about 9 mL of water, added 6.5 mL of ethyl alcohol, and then added water slowly with stirring until the sodium bisulfite dissolved again.  The idea is to have a solvent mixture which will dissolve the sodium bisulfite, but will cause any less-polar reaction products to precipitate.  Work out the moles of sodium bisulfite in your reaction table, and estimate that the artificial almond is about 0.05 grams/mL benzaldehyde.  This should be a large excess of sodium bisulfite.  If you choose to make a larger amount of reagant, perhaps twice, that would probably be okay.

    I mixed the Vogel solution with the ethyl acetate solution in an stoppered Erlenmeyer with stirring bar heated in a water bath at about 60 C.  The mixture was a clear solution.  If no precipitate is visible after about 10 minutes of heating, you may add ethanol drop by drop until the solution becomes very slightly cloudy.  Heat for about an hour and then place solution in freezer.  The next day, collect the solid precipitate on filter paper and allow it to dry overnight.  This should be your sodium alpha-hydroxybenzylsulfonate.

    We may save a portion of this for analysis using facilities at UMass-Dartmouth.

    Part II:  Converting Sodium Alpha-hydroxybenzylsulfonate back into Benzaldehyde

    Reaction:  2 NaC7H7O4S + Na2CO3 → 2 C6H5CHO + 2 Na2SO3 + H2O + CO2

    This is where it gets difficult.  The sodium alpha-hydroxybenzylsulfonate can be stirred with saturated sodium carbonate solution to regenerate the benzaldehyde.  Use the smallest amount of sodium carbonate solution you can which will fully react all the starting material, add it slowly.  You should smell a strong aroma of benzaldehyde (almond), and the solution will have a cloudy appearance.  However, the benzaldehyde does not form an easily separated layer.  I tried and failed!

    Therefore we will have to do another solvent extraction.  Extract your solution with the smallest amount of ethyl acetate (or other low-boiling nonpolar solvent, as safety/availability provide) necessary to remove the cloudiness from the water layer.  Afterwards, we must distill off the solvent so that the benzaldehyde will be left behind.  Monitor the vapor temperature during your distillation and stop when the temperature rises more than a couple degrees C from the boiling point of ethyl acetate.  Your benzaldehyde product may also be analyzed using facilities at UMass-Dartmouth before we carry it onwards.

    Part III:  Benzaldehyde to Benzoin, the 'Benzoin Condensation'

    Reaction:  2 C6H5CHO → C14H12O2

    Follow the procedure in Lehman below when planning this step.  We may have to adjust the scale:  Lehman uses about 15 grams of benzaldehyde, while we are unlikely to produce that much by our efforts.  On the other hand, some benzaldehyde might be ordered by then.  It's basically cool, dump, stir, heat, cool, collect crystals, I'm sure we can work it out.  I never got to do this in sophomore organic chemistry, we skipped this lab, and so I never got to make any benzoin.  But with any luck, you will!

    Safety Note:  This reaction uses sodium hydroxide solution, which is a strong base.  I will prepare the solution for you, you will not handle solid NaOH pellets.  Even in 3M solution, this is capable of causing severe irritation, burns, and even blindness.  Wear full PPE and take precautions.

    Part III-B:  Can we skip a step?  Sodium alpha-hydroxybenzylsulfonate to benzoin?  Or can we skip all the steps?

    Reaction:  You can figure it out!

    I have noticed that the reaction in Lehman takes place in the presence of base and water.  It is possible that the sodium alpha-hydroxybenzylsulfonate will convert to benzaldehyde under the conditions of the reaction, and then go on to form benzoin.  When planning this, we can just substitute the sodium alpha-hydroxybenzylsulfonate for benzaldehyde, mole for mole.  This would be really nice if it worked!

    Or, maybe, we can just skip most of the hard work and just take some of the almond flavoring, which is already a mixture of water and alcohol and benzaldehyde, react with NaOH and vitamin B1, and see if benzoin crystallizes out of the system?


    “Benzaldehyde.” Wikipedia, Wikimedia Foundation, 15 June 2019, en.wikipedia.org/wiki/Benzaldehyde.

     “Benzoin (Organic Compound).” Wikipedia, Wikimedia Foundation, 20 July 2019, en.wikipedia.org/wiki/Benzoin_(organic_compound).

    Vogel, Arthur I. A Textbook of Practical Organic Chemistry: Incl. Qualitative Organic Analysis. 3rd ed., Longman, 1974.

    “Sodium Bisulfite.” Wikipedia, Wikimedia Foundation, 23 June 2019, en.wikipedia.org/wiki/Sodium_bisulfite.

    Lehman, John W. Operational Organic Chemistry: a Laboratory Course. Allyn & Bacon, 1988. (pp. 494-503)

    “Benzoin Condensation.” Wikipedia, Wikimedia Foundation, 28 June 2019, en.wikipedia.org/wiki/Benzoin_condensation.

    Varley, John. Demon. Berkley Books, 1984.

  • Electrotyping:  Using Electricity to Copy a 3D Object in Solid Copper

    Electrotyping is a process for copying a 3D object into copper.  How is it done?  A mold is made from the object, using wax or a similar substance.  The mold is coated with a conductive layer :  powdered graphite, powdered copper metal, or perhaps conductive ink.  
    Copper is then deposited onto the mold by electrochemistry.  A solution containing Cu+2 ions is used as the electrolyte.  The mold is immersed into the solution and is connected to the negative terminal of a power supply.  A copper "anode" is also immersed into the solution, and is connected to the positive terminal of a power supply.  When the current is turned on, copper loses electrons at the anode and goes into solution as Cu+2 ions.  These ions are then attracted to the negative electrode (cathode) and travel through the solution until they reach it.  Once there, they pick up electrons and become copper metal again.  This fills the mold with copper metal.
    This technique was once routinely used to duplicate printing plates, before modern technologies for printing made it obsolete.  It is still used by museums to duplicate works of art, ancient coins, and such items.  

    Part I:  Let's make a copper quarter!
    I have literally always wanted a quarter of solid copper.  Or perhaps, a Kennedy half-dollar, or an old-time silver dollar, any big coin which would look wrong if it was made of copper.  I would be happy to provide the coin, or you can duplicate a coin or similar object of your own.  Making a 2-sided mold might be beyond our capabilities but we can just go for a single-sided coin to start with.  We can consider trying for 2 sides later.

    Plan a procedure for doing this, taking into account available equipment here at school.  I can answer any questions you need about what we have or could get.  "700 Science Experiments for Everyone" contains information on how to make the electrolyte solution and how to set up the experiment using a 3-volt battery and a rheostat.  We will be using a modern solid-state power supply which provides variable DC voltage, so plan accordingly.
    Duplicate coins made by professionals are shown below.
    Ancient coins duplicated by electrotyping
    Part II:  Cross-disciplinary Collaboration (3D Printing in Copper)

    We have multiple 3D printer installations here at Diman.  They can only 'print' in certain types of plastic.  Investigate current 3D printing projects and find a reason (or at least, an excuse) to duplicate a 3D-printed item in copper.

    Alternatively, look into using electrotyping to produce useful items for some shop:  electronics, perhaps?  Plumbing?

  • Hot Ice:  Synthesis and Characterization of Sodium Acetate Trihydrate

    Important Announcement:  We have done this project in previous years.  It has been brought to my attention that it might be of interest to Physics as well.  Therefore I ask that groups working on this project plan to coordinate with Physics when planning and doing Parts II/III.  


    Hot Ice Crystallization


    Note:  This video assumes 80% acetic acid AKA "vinegar concentrate", not the usual 5-6% strength.  It is much more hazardous and gives off unpleasant fumes.  I do not currently have any 80% vinegar in stock.  I currently have a stock of 30% acetic acid solution, so plan for that.

    Part I:  Synthesis of Sodium Acetate Trihydrate (NaCH3COO.3H2O)

    Reaction:  NaHCO3 + CH3COOH + 2 H2O → NaCH3COO.3H2O + CO2

    Write a procedure for synthesis of sodium acetate trihydrate from baking soda (NaHCO3) and 30% acetic acid solution (i.e. 300 grams of acetic acid per liter) based on the references given here and any other sources you have found.  Be sure to cite your sources!

    If you are leaving a solution of sodium acetate out to evaporate and crystallize, please cover it with cheesecloth.  Otherwise you may find it full of dead fruit flies in the morning.  Whoever said 'you catch more flies with honey than with vinegar' never actually did the experiment:  fruit flies love vinegar.  

    Despite the video, I do not allow you to handle chemicals with your bare hands in the laboratory.  That can just end badly in so many ways.

    Part II:  Supercooling and Crystallization of Sodium Acetate Trihydrate (NaCH3COO.3H2O)

    You will require pure, crystalline sodium acetate trihydrate for this.  Take some of it and heat it until it melts - this can be done by immersing a test tube into hot water.  The molten sodium acetate trihydrate can be set aside to cool.  If the glassware and the sodium acetate trihydrate itself are both perfectly clean, the melt will actually stay liquid at room temperature.  However, it is actually below its freezing point!  We call this a supercooled liquid.

    If you pour your molten sodium acetate trihydrate into a beaker containing a single crystal of solid sodium acetate trihydrate to act as a 'seed crystal', it should all crystallize.  Feel the beaker:  does it get hot?  Does it get cold?  Why?  Is freezing exothermic or endothermic?

    Does the volume of the sodium acetate change when it solidifies?  Design an experiment to measure this.  Or work out some other experiment of interest (coordinating with Physics!)

    Part III:  Heat Storage of Sodium Acetate Trihydrate (NaCH3COO.3H2O)

    Weigh out a known amount of sodium acetate trihydrate.  If your synthesized supplies are insufficient, I may be able to provide you with some more.  Melt it and allow it to cool to a supercooled liquid.  Be sure to use a very clean container and avoid any dust or debris, which may cause your supercooled liquid to crystallize prematurely.

    Use a thermometer to measure the temperature of the liquid before and after it crystallizes.  You may add a small 'seed crystal' of solid to start the crystallization if necessary.

    Calculate the heat energy given off:  the change in temperature, ΔT, multiplied by the specific heat capacity (C) will give you the energy given off by crystallization in units of J/mol.  You can multiply this value by the molar mass (in g/mol) to obtain the energy in J/gram.

    Compare this with values from the "Chalk Dust Magazine" article and the Wikipedia entry.  Is it comparable?

    Compare this with the energy density of lithium batteries.  Is this a good means of energy storage?



    “Creating Hot Ice.” Chalkdust, 25 Aug. 2017, http://chalkdustmagazine.com/blog/creating-hot-ice/.

    “Sodium Acetate.” Wikipedia, Wikimedia Foundation, 6 Dec. 2018, https://en.wikipedia.org/wiki/Sodium_acetate.

    Home Science. “How to Make Hot Ice at Home - Amazing Science Experiment.” YouTube, 18 Oct. 2016, https://www.youtube.com/watch?v=pzHiVGeevZE.

  • Raspberry Ketone Project
    Raspberry Ketone Structure raspberry fruit
    Raspberry ketone is a compound of many names - it is also known as 'frambinone', 'rheosmine', or even 'p-Hydroxybenzyl acetone', '4-(p-Hydroxyphenyl)-2-butanone', 'oxyphenylon', or 'rasketone'.  As the Wikipedia article states, it is used in perfumes, cosmetics, and flavorings.  It has also been marketed as a weight loss supplement, but this is not proven medically effective.  

    Because natural raspberries only contain 1-4 parts per million of raspberry ketone, it is generally produced synthetically.  By chemistry!
    Part I:  "Aldol Condensation" of 4-hydroxybenzaldehyde with acetone
    Base-catalyzed reaction of acetone with 4-hydroxybenzaldehyde  
    Reaction:  C7H6O2 + C3H6O → C10H10O2 + H2O. 
    A "condensation" is what organic chemists call a double displacement where one product is water.
    Last year, we successfully obtained good results in this reaction using the procedure described by Smith (1-18).  It is important to mix throughly after adding each reactant - do not simply mix at the end!  The reaction will not work that way.  We did scale up the reaction by a factor of 4, and did the reaction in an Erlenmeyer flask rather than a screw-capped test tube.  We were able to verify the purity of the product using facilities at UMass-Dartmouth.  We may do this reaction on a larger scale this year, in order to have materials for multiple tries at Part II.
    Safety Note:  This reaction requires the use of NaOH or KOH base.  These are strong bases and can cause severe burns!   
    Part II:  Hydrogenation of 4-hydroxybenzalideneacetone to raspberry ketone (AKA "Frambinone" or "Rheosmine")
     Aldol condensation followed by 'transfer' hydrogenation
    Note:  I would write it NH4HCO2, with the ammonium cation first, 'ammonium formate'.  Why is it useful?  It's basically solid hydrogen storage: it decomposes under the reaction conditions to form NH3, CO2, and H2 gases, but is much safer and easier to handle than H2 gas.  The H2 adds across the double bond under the influence of the Pd catalyst to give a single bond.
    Reaction:  C10H10O2 + NH4HCO2 → C10H12O2 + NH3 + CO2
    We followed the 'Step 2' procedure as described in 'Synthesis of Frambinone' last year.  Where the procedure uses dichloromethane as a solvent, we replaced it with ethyl acetate - it is much less toxic, and works well.  Note that while dichloromethane sinks in water and will form a layer below it, ethyl acetate floats up to the top.  We scaled up the reaction by a factor of 2, to use 200 mg of starting material (C10H10O2 from Part I), and the purity of our product was verified using facilities at UMass-Dartmouth.
    However, we have serious concerns about the safety of this reaction and would like to revisit the procedure this year!  The procedure in 'Synthesis of Frambinone' instructs us to swirl together the solid Pd/C catalyst, C10H10O2 intermediate, and NH4HCO2 , before addition of the methanol solvent.  We observed the mixture of solids became warm:  likely due to reaction between the NH4HCOand oxygen in the air.  This could present a fire hazard if we attempted to scale this reaction up from milligrams to a gram scale.  Therefore, we would like to develop a safer procedure.  
    An old mentor of mine has suggested the following order of addition:  swirl reaction flask in an ice/water bath, add Pd/C catalyst first, followed by methanol.  Warning - adding methanol first and then the Pd/C is likely to result in ignition of the methanol, as the catalyst particles fall through the vapor coming off the methanol.  Next we should add the ammonium formate NH4HCO2, continuing to swirl in the ice bath, and finally our C10H10O2 intermediate.  If the reaction does ignite, it should be immediately capped to exclude oxygen.  After the addition, the flask should be refluxed as described in 'Synthesis of Frambinone'.
    Safety Note:  This reaction is performed in methanol as a solvent.  Methanol is toxic and flammable.  Ingestion of even the small amount used in this reaction can cause permanent blindness or death.  Do not ingest methanol under any circumstances.  Any exposure to methanol must be reported immediately to your instructor, and to the school nurse.
    Part III:  Scaleup of hydrogenation of 4-hydroxybenzalideneacetone to raspberry ketone 
    If we can solve the exothermic issues in Part II, we may wish to scale up the reaction to a 500 mg size.  This should yield enough product that it can be recrystallized from hot water, as in Part I, and obtained as crystals in good purity.

  • "Crystal Microphone": Rochelle Salt and Piezoelectricity

    Crystals of Rochelle Salt, KNaC4H4O6·4H2O, used to be commonly used in microphone and small loudspeakers due to their property of 'piezoelectricity'.  Piezoelectricity means that voltage applied to a crystal will cause small size changes which can produce sound, or alternatively, vibrations applied to a crystal will produce electrical signals.  Today other piezoelectric materials are more common, but rochelle salt is easier to make. 
    Last year, we made a couple Rochelle Salt speakers, and found that they could successfully reproduce high-pitched sounds above 2 kHz or so.  We had to drive them with about 300 Volts to get full volume, and used an audio amplifier connected through a step-up transformer.  We still have this equipment available.  However, this year, we would like to try for a microphone!

    Rochelle Salt Crystals
    Part I:  Synthesis of Rochelle Salt
    Reaction:  KHC4H4O6 + NaHCO3 + 3 H2O → KNaC4H4O6·4H2O + CO2
    As described in the "Forsaken Technology" article below, Rochelle salt can be made from common grocery store items.  The ingredients here are sodium bicarbonate (NaHCO3) and cream of tartar (potassium hydrogen tartrate, KHC4H4O6).  
    Supermarket "cream of tartar" may not be pure (potassium hydrogen tartrate, KHC4H4O6. To purify it, recrystallize it from boiling water as described in Reference 1.  Be aware that boiling water can cause very severe burns!  This would be a great step for your first project lab day.  If I can obtain a purer grade of potassium hydrogen tartate, skip this step.  Don't worry about the sodium bicarbonate (baking soda), as it is generally sold as a highly pure substance.
    Once you have made or obtained purified potassium hydrogen tartrate, plan the synthesis as described in "Forsaken Technology".  Be sure to work out a full reaction table with grams and moles and such.
    Part II:  Growing crystals of Rochelle Salt
    Once you have obtained solid crystals of Rochelle Salt, you may be disappointed that they aren't larger.  We will want large crystals to make 'singing crystals', with as few imperfections as possible.  As described in "Singing Crystals", you will need to use an existing crystal of your Rochelle salt as a 'seed crystal' - you will hang it using fishing line into a saturated solution of Rochelle salt and allow slow evaporation.  It may take days or weeks to grow a proper crystal of Rochelle salt:  we need to find an area for this experiment where it will not be excessively disturbed.  "Singing Crystals" suggested a warm, temperature controlled enclosure - we may consider finding or making one.
    Part III:  Make a Microphone
    Details of contruction may be found in "Singing Crystals" and you are free to do your own research as well.  The basic setup requires that the crystal be connected to electrodes - aluminum foil may be clamped to it, or something similar.  The crystal might also be connected to a paper loudspeaker cone to help vibrations of the crystal couple mechanically to air for efficiency.

    We can connect the crystal to an audio amplifier or oscilloscope, and see if it works as a microphone.  I can probably come up with a guitar amp or something.  But there might be a proper preamplifier setup in the electronics shop?  
    Luminol:  The 'Volcano' and the 'Oscillating Glow Stick'
    CSI - Luminol Reaction
    "Luminol", known to organic chemists as 3-aminophtalic hydrazide, can be oxidized by hydrogen peroxide in the presence of a catalyst to produce a molecule in an 'excited' state.  The excited electrons will return to their ground state with emission of photons.  That's why it glows!  This poster is from Compound Chemistry and is part of an article referenced below.
    In this project, we will generally be scaling down the size of the reactions we're working with.  This is because I only have about 1.5 grams of luminol, and it is expensive.  It also can be toxic so keeping quantities small is a good idea!
    Part I:  The Luminol Volcano

    In this experiment, luminol reacts with hydrogen peroxide using sodium carbonate as a base and copper (II) sulfate as a catalyst.  Light should be emitted, as well as some offgassing of oxygen bubbles.  The above video is from MEL science, which has a webpage on this experiment that is referenced below.  MEL science sells a kit for this reaction!  However, we do not have this kit, and as we have all the needed ingredients in stock we can do without it.
    I suggest that, when planning this experiment, you should follow the procedure given in the Edith Cowan University lab manual referenced below, "Demonstration 7:  Chemiluminescence Reaction".  They are working on a very large scale - a 1 Liter reaction volume!  In order to reduce safety hazards and expense, I would like your first attempt to be scaled down substantially.  Scaling down by a factor of 5 would be a good idea.  I do have 12% H2Oin stock, which can probably be substituted for the 10% they mention to worry about.  If you were worried, you might decrease the amount and add more water to compensate.  I also suggest we use distilled water for this reaction. 
    When doing a gram/mole table for this reaction, you can assume that '%' means grams per 100 mL.  A 12% solution would have 12 grams of H2Oper 100 mL, or 120 grams/Liter.

    Ideally you'll be able to capture pictures and/or video for your final presentation:  you may have to work in the darkroom for some of this.
    Safety Note:  I find the procedure given in "Demonstration 7" where the liquids are poured into a beaker from a height to be dangerous and unnecessary.  Students have gotten splash burns by doing just that - pouring chemicals from a height.  Please do not do this!
    Safety Note:  While 3% drugstore H2O2 isn't very dangerous, stronger solutions like the 12% I have in the lab can cause irritation, blisters, and even burns.  The blisters are usually temporary, but can be quite painful.  And it could be really bad to get it in your eyes!  Sodium carbonate is corrosive and can cause irritation or possibly burns.  Copper sulfate is an irritant, and can be toxic if ingested in sufficient quantity.  Please review SDS sheets for these substances.
    Safety Note:  Luminol is toxic!  Do not ingest it.  Please look up SDS sheets for it as well.
    Part II:  Variations on the Volcano!

    We should try at least one experiment based on the Volcano demonstration, maybe a couple!  Things to think about:  
    What happens if you use 3% peroxide instead of 12%?  I think the reaction will glow less intensely, but for a longer period of time.  What is the limiting reactant here?  Luminol, or peroxide?  The moles will help you figure this out!  As long as luminol is limiting, you'll get the same total number of light photons out, but the percent of peroxide will probably affect the reaction rate greatly.
    If you leave out the copper sulfate (CuSO4), there will be no reaction, because copper is the critical catalyst for this reaction.  It seems that blood is a catalyst for this reaction - hence the use by forensics/CSI types to detect blood residues.  If you ask Ms. Cooney nicely, she might loan you some 'simulated blood' which you could try to detect with your luminol reaction.  
    Part III:  The Oscillating Luminol Reaction
    As Sattar and Epstein describe in the linked article below, a solution containing H2O2, KSCN, CuSO4, NaOH together will undergo oscillations in concentrations of various substances.  This is visible as a repeated change in reaction color between yellow and colorless.  This is not a very spectacular reaction, but can be made visually exciting by addition of luminol.  The yellow species in solution is thought to be a complex ion of formula Cu(O2H)(OH)2-2. This complex will catalyze the reaction of luminol with hydrogen peroxide, yielding repeated pulses of light.
    Please construct a procedure to reproduce the oscillating luminol experiment of 'NurdRage' as shown in the following video:

    Again, pictures and/or video are highly desirable!
    Safety Note:  This reaction is complicated and I worry about the potassium thiocyanate (KSCN) becoming oxidized to the much more toxic potassium cyanide (KCN).  This is really not likely to happen, according to the references that NurdRage was kind enough to link from his video description, but I think just the possibility of this happening should encourage us to be really, really safe here.  It possibly could happen if you forgot the NaOH.  Also the luminol wouldn't dissolve.  So basically, do NOT forget the NaOH!  Without it, the luminol won't dissolve anyway.  Also I'm going to ask that we scale his reaction down by a factor of 10.  As weighing out 3.7 mg of copper sulfate might be impossible, I suggest you go ahead and dissolve 37 mg of it in 100 mL of water, use 10 mL of the solution, and save the other 90 mL for later.

  • Cinnamon to Strawberry and Spice

    Cinnamaldehyde, as the name suggests, is the main component of cinnamon essential oil, and it is an aldehyde.  As it is a purified essential oil component, it should be handled with the respect due any other chemical:  a bottle of cinnamon flavoring from the grocery story might only have 1% of cinnamaldehyde.  In this project, we can try to turn cinnamon into strawberry-scented ester.  If time permits, we can also attempt to prepare a 'floral spice' flavoring compound.  Note that, as our lab is not food grade and neither are our ingredients, we never taste anything we make.  We may 'waft' the odor, though.
    Part I:  Cannizzaro Reaction of Cinnamaldehyde
     Cannizzaro Reaction of Cinnamaldehyde
    This is an old reaction, recently of new interest as a type of 'Green Chemistry', which converts an aldehyde into equimolar amounts of carboxylic acid and alcohol, as described in the 'Cannizzaro Reaction' entry on Wikipedia. We will be following the solvent-free Cannizzaro method described by Phonchaiya et. al.  However, they used a different aldehyde:  2-chlorobenzaldehyde, whereas we will be using cinnamaldehyde.  
    Last year, students measured out 5 mL of cinnamaldehyde into a mortar with 1.06 g of solid NaOH and students ground it for about 30 minutes.  They covered the mortar with saran wrap to avoid splashing:  cinnamaldehyde is an irritant, and sodium hydroxide is very corrosive and can cause severe burns and even blindness.  They then left the reaction mixture overnight to complete and harden into a gummy solid.  Yield was poor, and it is likely most of the cinnamaldehyde did not react.  This year, I would like us to try a method often used in biochemistry:  the reactants can be mechanically ground together in a centrifuge tube containing a ball bearing, spun on a vortex mixer.  I am also open to student-designed mechanical grinding equipment, if anyone is interested in developing some.
    The reaction product was then dissolved in 30 mL of ethyl acetate and 20 mL of water.  A separatory funnel was used to separate the two layers.  
    The bottom layer was water with sodium cinnamate dissolved in it.  50 m L of 2 M HCl (hydrochloric acid) was added, precipitating cinnamic acid.  This product was collected on a filter and left to dry.  Purity of this product was confirmed by facilities at UMass-Dartmouth.
    The top layer was ethyl acetate containing cinnamyl alcohl and other impurities in solution.  It was washed with 30 mL of water to remove leftover NaOH, followed by 2 portions of 50 mL of NaHSO3 solution (sodium bisulfite) to remove leftover cinnamaldehyde.  This solution was then distilled to remove ethyl acetate solvent, leaving behind solid product.  This product was analyzed by facilities at UMass-Dartmouth, and did contain some cinnamyl alcohol, but it was highly impure.  It is my hope that we can do better this year, and that mechanical grinding of the Cannizzaro reaction will improve our results.
    Part II:  Esterification of Cinnamic Acid to Methyl Cinnamate (Strawberry!)
    Methyl Cinnamate Synthesis
    This is an esterification reaction where a carboxylic acid reacts with an alcohol yielding an ester and water (not shown) - see the above 'Ester Synthesis Project' on this page for background.  You may collaborate with students working on ester synthesis if this will be helpful.  For details of this particular reaction, see 'Thomas Chemistry' below.  We will not be using ether as an extraction solvent:  I have ethyl acetate available instead.
    Part III:  Esterification of Cinnamyl Alcohol to Cinnamyl Acetate (Floral Spice!)
    Esterification producing cinnamyl acetate
    As I have had difficulties obtaining reasonably pure cinnamyl alcohol, we should hopefully only do this after we have verified the purity of our cinnamyl alcohol.  We will have to develop our own esterification procedure for this reaction, as I could not find one - again see above in 'Ester Synthesis Project'. 

  • Lycopene:  Natural Red Pigment from Tomatoes


    Chemical Structure of Lycopene

    I have been struggling to find a good red pigment for the 'paint project' which isn't either toxic or requires toxic materials to synthesize.  We have iron (III) oxide of course, but it's more of an orange than a red, and it often clumps up in the paint medium.  It's better in oil paint, but we don't use that for Paint Day because it might take a month to dry.  It was suggested by a colleague that I have my students extract a natural red, and I think this is something we should try!

    Tomatoes contain a natural red pigment called 'lycopene'.  As you can see from the chemical structure above, it's complicated.  Synthesizing it would be very difficult.  But we do not have to synthesize it, because we can extract it from tomato paste.  We will basically extract the 'red' out of tomato paste using organic solvents.  Unfortunately, this will not be pure lycopene - it will be mixed up with various fats and oils also present in tomatoes.  We can then purify it by 'column chromatography', which works like the paper chromatography you may have already done but allows you to purify larger quantities of substances.  

    I am hoping that we can check the purity & identity of our lycopene using facilities at UMass-Dartmouth although I cannot guarantee it.  This is a more complicated molecule than we will see in any of the other organic chemistry projects!

    For most of these projects, I actually do an extensive writeup.  I really don't need to here, because it turns out we have a kit for doing this exact extraction.  Please see the referenced "Column Chromatography Kit" writeup for details.


    Part I:  Extraction of Tomato Paste Components

    Please consult the section in "Column Chromatography Kit" titled "Extract Preparation and Separation of Tomato Paste".  You should plan to do steps 1-4, which should give you a tomato extract.  However, we might not want to go ahead and do the chromatography just yet! Unlike in the writeup, you will prepare your own 10 mL of 80/20 hexane/acetone mixture.  Please store this in a closed, labeled vial.  Please label the vial using tape and write in pencil, as any stray drops of this solvent will render ink totally illegible.  

    This mixture will have solid residues of insoluble bits of tomato.  Please filter out the solids, and store the liquid extract.  How can you filter a 1 mL mixture?  A big funnel and coffee filter would soak up almost all of it.  You could try a pipette with a small piece of cotton wool!  We should keep this extract in a closed vial in the B119 refrigerator when not using it.  This will not take an entire lab period, but I know from experience that column chromatography is time-consuming and probably will take an entire period.  Therefore, we should prepare as much in advance as we can.


    Part II:  Purification of Lycopene by Chromatography

    tomato pigments on column How does chromatography separate different substances?  It's all about polarity, like you learned in Honors Chem I.  The aluminum oxide in the column is relatively polar, so polar substances in the mixture will tend to stick to it.  You start with a nonpolar solvent (pure hexane), and the least-polar substances will come off the column, or 'elute', more quickly.  You then increase the polarity of the solvent (80/20 mixture of hexane and acetone) to bring the more-polar substances off.  

    The kit writeup is confusing here!  You need to follow the instructions for 'Preparation of the Chromatography Column'.  From there, you should follow steps 6-19 of "Extract Preparation and Separation of Tomato Paste".  You should end up with liquid solutions of 'tomato pigment 1' and 'tomato pigment 2'.  One should be a pink-red color, this is your lycopene!  The other is yellowish and is 'beta carotene', the main pigment found in carrots.

    As the volumes of these solutions will be very small, we can pour them into glass crystallization dishes and allow them to evaporate overnight.  This should yield solid pigment.  Will there be enough quantity to be useful?  We shall see!


    Part III:  Can we Scale this Up? 

    Our column is only capable of purifying very small quantities of lycopene.  It'd be nice if we could develop a procedure capable of isolating larger amounts of lycopene!

    I found a paper titled 'A simple method for the isolation of lycopene from lycopersicon esculentum'.  I have no idea what sort of plant or algae 'lycopersicon esculentum' actually is.   Try to find out if you can.  

    The basic idea is that the pigment is extracted using the moderately-polar solvent ethyl acetate, is evaporated mostly (but not entirely) dry, and then methanol is added.  Methanol is very polar, and should precipitate the less-polar lycopene while keeping beta-carotene and other contaminants in solution.  We should definitely try this, on a small scale.  The paper talks about doing it on a scale of Liters.  That's definitely bigger than we want.  Somewhere in the 10-100 mL range would be better for us.

    Once isolated, we can compare this lycopene with the product from Part II.  And maybe, this way, we can obtain enough for Paint Day!



    "Column Chromatography Kit" CHEM-FAX. Vol. 4847, Flinn Scientific, 2015.

    Roh, Myong-Kyun, et al. “A Simple Method for the Isolation of Lycopene from Lycopersicon Esculentum.” Botanical Sciences, vol 91, no. 2, 2013, pp. 187-192

  • Return of the Bismuth:  Bismuth Metal Project

    Rainbow Bismuth Crystal

    Bismuth is an interesting metal.  It seems an ordinary, dull, low-melting, heavy metal rather like lead or tin.  But unlike its P-Block metal neighbors, bismuth can be grown into beautiful cubic crystals which form a rainbow-colored oxide layer.  It is also, unlike its neighbors lead, tin, antimony, tellurium, polonium, generally nontoxic as the metal and only mildly toxic in its compounds!  It's an element I don't mind having around.  

    Once upon a time, a group of students known as the Niobium Nutrias came to me and requested to do their own Honors Project.  They wanted to isolate bismuth metal from Pepto-Bismol tablets, and to grow crystals of it.  They succeeded.  And it was about the messiest, most painful reaction we had ever done in this lab up until that point.  So we may not do what they did ever again, but we can be inspired by their example!

    Brief Summary of the Niobium Nutria's Procedure:

    You can view the Niobium Nutria's slideshow below.  They were inspired by NileRed's video "How to Extract Bismuth Metal from Pepto-Bismol Tablets", but their procedure was different.  They did obtain a higher percentage yield of bismuth recovery, so their efforts were not in vain.  Their reaction scheme is shown below:

    bismuth reduction scheme

    The first reaction shows bismuth subsalicylate, BiC7H5O4, reacting with hydrochloric acid to yield a solution of BiCl3.  Salicylic acid, C7H6O3, precipitates out due to poor solubility in water.  In the second reaction, the BiCl3 solution is reacted with aluminum foil to precipitate bismuth as a powder.  The powder is finally melted to form an ingot of bismuth metal.  The details are what make this such a pain!

    The Nutrias started by taking 1 box of 40 pepto-bismol tablets (262 mg of bismuth subsalicylate in each) and stirring in 600 mL of 5% HCl over the weekend.  It turned out to be important to do this reaction over a long time, as otherwise the salicylic acid precipitating will clog up filter paper.  The mixture is filtered through coffee filters - this takes time - to obtain a solution of BiCl3. This BiCl3 solution was then reacted with aluminum foil, added piece by piece.  The Nutrias tested the solution by spotting it on clean foil with a glass rod:  a dark streak indicated the presence of more Bi in solution.  The precipitated Bi was obtained as a dark powder and collected on filter paper.

    This is where things got to be a total pain:  the precipitated Bi still had residual Al in it, and was washed with even more 5% HCl to dissolve the remaining Al.  Then it had to be washed with water to remove the HCl.  Then there was leftover salicylic acid in it as well, so it was washed with 91% isopropyl alcohol to remove that.  

    Finally, the Nutrias attempted to melt the powdered Bi to form an ingot, but it reacted with oxygen to form beautiful yellow Bi2O3.  They attempted to do the melt under argon gas, but failed to sufficiently exclude oxygen.  On their final attempt, they took their powdered Bi to Metal Fab and had a student turn a fuel-rich acetylene 'reducing flame' directly onto it, and finally obtained a small ingot of bismuth metal.

    The Nutrias then worked on growing crystals with commercially-purchased bismuth metal, as I was not about to have them scale this process up by a factor of 100.  Just too big!

    How We Could Do this Differently: Electrical Deposition of Bismuth Metal from Solution

    Part I:  Digesting Bismuth Subsalicylate Tablets in Dilute Sulfuric Acid
    2 BiC7H5O4 + 3 H2SO4 → 2 C7H6O3 + 2 H2O + Bi2(SO4)3

    Instead of digesting the pepto-bismol tablets in hydrochloric acid, we need to use sulfuric acid.  Why?  Because if we use hydrochloric acid, we might end up generating chlorine gas during the electrochemistry step later on.  That would be bad!  What we're hoping to obtain is a solution of Bi2(SO4)for the next step.  This salt may only be soluble in a solution with excess acid:  check the pH and make sure that if the solution loses acidity, we add more H2SO4 (perhaps more concentrated for convenience, like 3M).  I suggest we start with 1 M solution of H2SO4 which I will provide for you.  Stir tablets in it overnight or longer, filter.  When you do a reaction table, remember you're trying for excess H2SO4 if possible.  We may want to go smaller-scale than a 40 tablet box in 600 mL acid for our first try!  This is more a proof of concept here, we've never done this reaction before.  Smaller is safer when you try things for the first time.

    Reactions to Test for Bi (We should do at least one):
    Reaction with Al Metal:
    Bi2(SO4)+ 2 Al → Al2(SO4)+ 2 Bi
    Decomposition @ High Temp to Yellow Bi2O3 Pigment:
            Bi2(SO4)→ Bi2O3 (beautiful yellow pigment!) + 3 SO3(gas)
    Alternate Prep of Yellow Bismuth Oxide Pigment:
            3 H2O + Bi2(SO4)3 + 3 Na2CO3 → 2 Bi(OH)3 + 3 Na2SO4 + 3 CO2
            2 Bi(OH)3 → Bi2O3 + 3 H2O

    We will have to decide whether it's worth boiling down the solution to obtain solid Bi2(SO4):  while this would allow us a chance at a percentage yield, there is a good chance it will decompose to bismuth hydroxides or oxides, and we will also have to get it back into solution again before Part II.  This will require more H2SO4.  Perhaps a sample of the solution could be boiled down, and used to estimate the yield of the whole reaction?  If we do obtain solid Bi2(SO4)3, we may consider converting a small amount of it to yellow-colored Bi2O3.  This might be possible by directly heating the Bi2(SO4)3 in a crucible over a Bunsen.  Alternatively, the Bi2(SO4)could be reacted with a base dissolved in water to precipitate Bi(OH)3, which should easily decompose to Bi2O3 upon heating (why?).  Bismuth and Uranium are the only metals which have bright yellow oxides, so this proves the presence of bismuth.

    We should also test the solution for the presence of Bi+3 ion as the Niobium Nutrias did, following NileRed's video:  streak a drop of the solution onto a piece of aluminum foil.  Formation of black precipitated Bi metal powder indicates that there was Bi+3 ion in solution, which underwent single displacement from the aluminum metal.

    Part II:  Electrodeposition ("Electrowinning") of Bismuth Metal from Bismuth Sulfate Solution
       2 Bi2(SO4)3 + 6 H2O → 4 Bi + 3 O2 + 6 H2SO(Electrochemical Cell)

    This reaction will not spontaneously happen on its own, but can be forced to happen by electrical current.  This is like the 'Avogadro's Number' lab we did last year.  Oxygen gas should be evolved at the positive electrode, and bismuth metal will hopefully be deposited on the negative electrode.  There may be some hydrogen gas evolved at the negative electrode as well, potential fire hazard.  Consult an AP student if you can find one, to explain how half-reactions and electrochemistry work here.  You can also read the Wikipedia article on "Electrowinning".

    We should not use metal electrodes here, because there will be side reactions.  The positive electrode, if a metal, will dissolve into the solution as positive ions, just like the zinc did in the 'Avogadro's Number' lab.  We should use carbon rods. I mean, we probably could use platinum, that's pretty inert, but I don't have the budget for platinum right now. Somebody can probably find carbon rods here at Diman.  There may be some in the Welding side of Metal Fabrication.  Or some brave soul could try to scavenge one by cutting open a cheap "Super Heavy Duty" battery from Dollar Tree.  This is dangerous, as batteries contain corrosive substances, but we might end up doing that.  Maybe there are some in the Physics labs?  Scavenger hunt!

    We will do a similar setup to the 'Avogadro's Number' lab, clipping the 2 carbon rods into a beaker of your acidic bismuth sulfate solution and connecting them to a DC power supply.  I am hoping that the bismuth will atually deposit as crystals on the negative carbon rod, but that's mostly a hope.  We may obtain beautiful crystals of bismuth just from this step, but if not ....

    Part III: Crystallization of Bismuth Metal

    Watch NightHawkInLight's video to plan this one.  We found out that this cannot be done in a stainless steel or aluminum container very effectively.  You need to melt the bismuth in a container that is not a strong conductor of heat, or as it cools the bismuth will all solidify and stick to the sides.  We can use a ceramic crucible, perhaps, or a small cast iron pan from Dollar Tree/Five Below/etc.  Ideally, as molten bismuth cools, crystals form floating in the solution and can be fished out with tongs.  What is the melting point of bismuth?  Look it up.  We can do this with a hotplate, it's not that hot.  Still hot enough to cause serious burns, though, precautions and full PPE including aprons must be worn!


    “Niobium Nutrias- Honors Seniors Project.” 2017-2018 Honors Presentations, Diman Regional Vocational Technical High School, 15 May 2017, https://docs.google.com/presentation/d/1c-CvxNaHtKIFt11ZTnuXz9TCmBSUObLGzMsSSK3uDrk/edit?usp=sharing.

    NileRed. “How to Extract Bismuth Metal from Pepto-Bismol Tablets.” YouTube, 29 July 2015, youtu.be/grpSfjUImUs.

    “Electrowinning.” Wikipedia, Wikimedia Foundation, 5 Aug. 2019, en.wikipedia.org/wiki/Electrowinning.

    NightHawkInLight. “How to Make Bismuth Crystals.” YouTube, 29 Mar. 2015, youtu.be/v8KYZHMkTHw.