Who Art Thou Chemist: Victor Snieckus

This next chemist I would like to introduce is Victor Snieckus of Queen's University in Kingston, ON. http://www.chem.queensu.ca/people/faculty/snieckus/

I have personally benefited from this man's work. Snieckus is most famous for the pioneering work that he and his group have done on directed ortho metalation. 

Aromatic rings are extremely common in a wide variety of important synthetic products. Many natural products that have medicinal properties contain aromatic rings. Acetylsalicylic acid (asprin) has an aromatic ring in it. Important to the research that I am doing: polymers that are used in LEDs also have aromatic rings. This means that it is really important for synthetic chemists (like myself) to be able to rearrange bonds on aromatic rings and attach other things to them. But the problem is that aromatic rings are stupidly stable and don't really like having their bonds broken so you need a really strong base, like butyllithium (BuLi). The next problem is that on an aromatic ring there are six possible sites for making bonds. The trick for synthetic chemists is controlling which site the bonds are made at. This is where directed ortho lithiation comes in. The reaction is directed to the "ortho" position on the aromatic ring.

List of the different positions on an aromatic ring

Example reaction of a directed metallation

 DMG stands for Directing Metallation Group. By having one of these on the aromatic ring, a chemist can be sure that their metallation occurs at the "ortho" position over the meta and para positions. Take a look at the scheme below. Here we see an example of an ortho lithiation. When BuLi is added to the aromatic ring, the DMG directs it to the ortho position and the BuLi removes the hydrogen atom, leaving the carbon atom it was attached to with a negative charge and really reactive, ready to react with the next compound, in this example it reacts with carbon dioxide (CO2). The result is that there is carboxylic acid attached to the aromatic ring. This aromatic ring looks very much like asprin. This reaction is extremely versatile. It is definitely an important tool in the synthetic chemist toolbox. The Snieckus is always coming up with variations and expansions on this interesting and useful reaction. 

I have personally had the opportunity to discuss this reaction with Snieckus and have him give me some advice. He was a visiting speaker at the University of Alberta and I was able to catch him after his lecture to ask him about a directed metallation reaction that I was having problem with. He was able to give me some great suggestions and had a graduate student of his email me a procedure that was part of the students thesis, but hadn't yet been published. He also gave away one of his group t-shirts to a lucky attendant of his lecture and I happened to win since I did lithiation reactions. It was very helpful.

Selected Publications:

Org. Lett., 2010, 12, 2198-2201

Chem. Eur. J. 2010, 16, 8155-8161

 
Org. Lett. 2010, 12, 68-71

Who Art Thou Chemist?

This particular post, or rather, series of posts has been inspired by an article in the New York Times http://www.nytimes.com/2011/08/09/science/09emily.html 
It seems that people do not know who leading scientists are. This is a shame since there are great scientist out there, doing work that can have huge implications on daily life. The further implications is the fear that the public has of science. See The Ethical Chemist for further information. I plan to introduce a few interesting chemists so that maybe my loyal readers learn a little about some cool science.


The first scientist I would like to introduce is Marc Hillmyer of the University of Minnesota. http://www.chem.umn.edu/groups/hillmyer/ I was first introduced some of the work by the Hillmyer group at a conference in Australia. I chose to present his work as part of a seminar that was required for my Ph.D. The work that I was most interested in was his work on miktoarm star terpolymers for multicompartment micelles. So what does this mean? The Hillmyer group makes polymers that have three "arms". One arm of the polymer is a water soluble (the term used is "hydrophillic") polymer, like polyethylene glycol. The two remaining arms are both not water soluble (the term used is "hydrophobic"). But what is really interesting is that the two hydrophobic arms also don't mix: one is a hydrocarbon polymer; the other one is a fluorinated hydrocarbon. Think of this like a teflon frying pan and bacon grease. Neither will mix well with water, but the bacon grease (the hydrocarbon polymer) also won't stick to the teflon (the fluorinated polymer). Now these polymers, when added into water will assemble so that the hydrophillic polymers are on the outside, while the hydrophobic polymers are on the inside. Because the two hydrophobic polymers don't mix, they form two different compartments inside. The result are multicompartment micelles. A neat application of multicompartment micelles would be in drug delivery. Two incompatible drugs could be packaged in each of the different compartments and then delivered to the same target. This is just one particular example of a possible application. But my favourite part of this work is the synthesis. I love how this group was able to join all three of these polymers at a single carbon junction.  This is not trivial. I found the synthesis very elegant. 


Selected Publications by Hillmyer:

Liu, C.; Hillmyer, M. A.; Lodge, T. P. – Evolution of Multicompartment Micelles to Mixed Corona Micelles Using Solvent Mixtures – Langmuir 2008, 24, 12001–12009.

Liu, C.; Hillmyer, M. A.; Lodge, T. P. – Multicompartment Micelles from pH Responsive Miktoarm Star Block Terpolymers – Langmuir 2009, 25, 13718–13725. 

 

Li, Z.; Hillmyer, M. A.; Lodge, T. P. – Morphologies of Multicompartment Micelles Formed by ABC Miktoarm Star Terpolymers – Langmuir 2006, 22, 9409–9417.

Lodge, T. P.; Rasdal, A.; Li, Z.; Hillmyer, M. A. – Simultaneous, Segregated Storage of Two Agents in a Multicompartment Micelle – J. Am. Chem. Soc. 2005, 127, 17608–17609.

 I hope that my readers found this interesting. There is lots of interesting science being done out there. This particular post doesn't even cover all the interesting work that is being done in the Hillmyer group.

The Chemistry of Flame Retardants: Part Two-The Environmental Impact of Brominated Flame Retardants

Brominated flame retardants show some of the complexities of the problems faced with many of the materials that we use in life. Obviously, the effects of fires are terrible. There is severe, acute danger to the ease of ignition and flammability associated with the many materials that our daily life is so dependent on. The incorporation of flame retardants immediately reduces this problem. But the use of flame retardants isn't without its own pitfalls. The second most common flame retardant in commercial use is polybrominated diphenylethers (PBDE). These compounds can have up to ten bromine atoms attached to them, and their use is dependent on the number of bromines that are attached to the diphenylether. In chemistry, a difference of one atom can make massive changes to its chemical behaviour. For example, cyanide is one carbon atom triple bonded to one nitrogen atom, and is extremely poisonous. But the atomospere is 75% N2, which is one nitrogen atom triple bonded to one nitrogen atom, and is completely innocuous. These two compounds differ only be one atom. The three most common PBDEs are deca, penta, and octa (10, 5, and 8 bromine atoms). One of the problems with PBDEs is that they are not chemically bonded to any of the materials that they are incorporated into, they are simply physically mixed in. This means that they can be leached from the material and into the environment. 

To examine the problem of their presence in the environment, we need to look at what characteristics makes the chemicals good flame retardants. The chemicals need to be stable and they need to last long. If they weren't, the chemicals wouldn't stay around long enough in the materials that they are incorporated into and eventually those materials would become easily flammable again. So PBDEs are very stable and will last a long time without degrading. This means that if they leach into the environment they will not break down, but persist for years. The other downside of their ability to leach out of materials is a decrease in flame retardance over time. 

PBDE have been detected in arctic life. This suggests that they can be transported through the environment a long way from where they were initially released into it. This is termed "long-range transport". There is also evidence of "bioaccumulation". The chemical is taken up by organisms low on the food chain, and those organisms are in turn take up by organisms higher up on the food chain. The result is that what was a small amount of chemical in an organism low on the food chain becomes a much larger amount of chemical in organisms higher up on the food chain. This process happens when a particular chemical is not broken down in the digestive system of organisms, but rather stored, in fats usually. Mercury is an example of another chemical that is know to bioaccumulate. Beyond that, more labile PBDEs, like deca-PBDE, will break down into its more persistent cousins, penta- and tetra (four bromines) -PBDE. PBDEs have also been shown to degrade overtime, using heat and light, to toxic chemicals: polybrominated dibenzodoxins and polybrominated dibenzofurans. So even though they are not acutely toxic, PBDEs may prove to have chronic effects.

The evidence of environmental impact of PBDEs have prompted legislation against them. In Canada there is legislation against PBDE under subsection 93(1) of the Canadian Environmental Protection Act, 1999. In the United States, there is no federal legislation, but many states have bans against PBDEs. The European Union has banned the use of PBDEs.

What I find interesting about the case of brominated flame retardants is that it highlights many of the complexities associated with the problems with the chemicals in our life. These chemicals are not good for the environment, and we shouldn't use them; however, the results of not using flame retardants are equally damaging. Solutions are being researched to find effective flame retardants that are not environmentally damaging. It is important to understand that these materials weren't designed to be damaging or done by "evil scientists in labs who don't care about the environment". They were designed to solve a problem. That problem was the flammability of materials. Unfortunately, they also created a problem. Every action will have a reaction. 

References:

See part one for the references.

The Chemistry of Flame Retardants: Part One-What is a Flame Retardant?

Here in Northern Alberta, Canada we have had a disaster unprecedented in our province: wild fires. We have been hit with numerous fires that have impacted over 10 000 people in the province. Most notably are the residents of the town of Slave Lake, Alberta. On May 15, 2011 the fires actually entered the town, destroying one third of the town. After being forced to flee their homes, many residents returned to find that they had lost their homes and businesses to the fire. This tragedy caught the eyes of the world and even prompted a stop by the Duke and Duchess of Cambridge during their Canadian visit to meet with those affected by this disaster. So I dedicate this post to those affected by the Northern Alberta fires. Anyone wishing to support the many victims of this disaster can do so by making donations to the Canadian Red Cross. Information on the relief and recovery effort of this disaster is also available on the Canadian Red Cross website: www.redcross.ca

I am part of a team of disaster management volunteers and was deployed on May 15th to assist the victims of the Northern Alberta fires. As I drove north, on my way to High Prairie, Alberta, the landscape became an eerie red colour. The road, the treeline, the areas recovering from fires ten years past, all were covered in a red film. That red film was flame retardant and that image was the inspiration for today's blog entry: chemistry of flame retardants.  

Flame retardants are actually the second most common additive to the polymers that make up the various materials on which our modern western culture has become so reliant: these include polystyrenes, polyesters, epoxy resins, polyethylenes, polyurethanes. Take a look around your home: chances are you are currently using a computer, the circuitry, the wiring, and the casing is comprised of these polymers. Many textiles: curtains, upholstery, and clothing are also comprised of these polymers. Anything in your household that is "plastic" is made of these polymers. Why might adding a flame retardant be so important? These polymers are made of hydrocarbons-they come from petroleum, just like the fuel used in combustion engines. Anything made of a hydrocarbon (a chemical that is rich with hydrogen-carbon bonds) burns really well.  Flame retardants are, therefore, used to prevent or minimise the risk of fire.  The use of flame retardants have been documented as early as 450 BC, when Egyptians used alum (potassium aluminum sulfate hydrate) to reduce the flammability of wood. In the 17th century, Parisian theatre curtains were made "incombustible" by soaking them in a mixture of clay and gypsum. In 1735, the first patent was taken out on fire retardants. 

There are four classes of flame retardants: inorganic, halogenated organic, organophophorous, and nitrogen-based. I don't actually know what flame retardant was sprayed  on the land during the Northern Alberta fires. (A quick Google search suggests that it could be some mixture of ammonia based compounds-take that at face value since there is no verification on its make up and therefore I cannot comment on any environmental impact.) The flame retardants that I will write about today are halogenated organic, specifically brominated flame retardants. In chemistry, the term "organic" refers to chemicals that are rich in the element carbon. Halogens are in the 17th column (the second last) of the periodic table. The elements are, in descending order, fluorine, chlorine, bromine, iodine, and astatine. A halogenated organic compound is one where halogens are bonded to carbons.

How might flame retardants prevent combustion? Combustion is an oxidative (this is why the presence of oxygen in fires is so critical) gas phase reaction. The process of combustion can be broken down into four steps: preheating, volatilisation/decomposition (the reaction takes place in the gas phase-volatilisation is this phase change), combustion, and propagation. A flame retardant will target any one of these steps to prevent combustion. Halogenated flame retardants target the propagation step. In the propagation step, many free radicals are produced, which are what continues the chemical reactions in the burning process. A free radical is a chemical that has an unpaired electron-which makes them really reactive. I stated earlier that halogens were in the second last column of the periodic table; the last column is the noble gas column. (Helium, neon, argon, krypton, xenon, radon.) These gases are considered to be "inert". What separates the halogens from noble gases is one electron. If a halogen can get one electron then they can be as inert (and therefore happy) as a noble gas. If radicals lose one electron, then they have all their electrons paired and are super happy and unreactive too. 

Halogenated flame retardants are added to a material. If this material is lit on fire then the carbon-halogen bond will break, releasing the halogen into the gas phase-which is where the combustion is taking place. The halogen will find and capture radicals (the term used to describe this process is "scavenge") and will prevent the fire from continuing. Certain halogens are better at scavenging radicals than others. Iodine is actually the most efficient halogen at radical capture, fluorine is least efficient. But this is not the only property considered. The strength of the carbon-halogen bond is also important to consider. The halogen-carbon bond has to be strong enough not to degrade before the combustion temperature is reached, otherwise the halogen will be lost and therefore unable to capture any radicals. But if the bond is too strong then the halogen will never be released into the gas phase, and also will be unable to capture any radicals. It turns out that the carbon-fluorine bond is too strong and degrades at too high of a temperature. Fluorine is not a good flame retardant choice-it has poor efficiency, and too strong of a bond. Iodine is also a poor choice as a flame retardant. Though it has good radical capture ability, the carbon-iodine bond is too weak, and therefore degrades at too low of a temperature. This leaves chlorine, and bromine. While both are used in flame retardants, bromine is the halogen of choice because of its increased radical capture efficiency over chlorine, and the carbon-bromine bonds break at a lower, more ideal temperature for use as a flame retardant. 

The carbon part of the flame retardant is not part of the "flame retardant ability". It is present to hold the bromine (or chlorine) in a manner that doesn't interfere with the polymer function, but will allow it to be released upon decomposition in the flame. Some of the flame retardants are actually chemically bonded into the polymer, these are things like brominated styrenes, or tetrabromobisphenol-A. Other brominated flame retardants are simply mixed into the polymer and are not chemically bonded to the polymer, these are things like polybrominated diphenylethers, or hexabromocyclodocanes.  

Coming next: The Chemistry of Flame Retardants: Part Two-Environmental Impact of Brominated Flame Retardants.

References:

Mehran, A.; Aria, P.; Sjodin, A.; Bergman, A. Environ. Int. 2003, 29, 683-689.

Hindersinn, R. R. Historical Aspects of polymer fire retardance. In: Nelson, G. L., editor. Fire and Polymers hazard identification and prevention. American Chemical Society Symposium Series, vol. 415. New York: American Chemical Society; 1990.

Weil, E. D.; Levichik, S. J. Fire Sciences, 2004, 22, 25-40.

Weil, E. D.; Levichik, S. J. Fire Sciences, 2004, 22, 339-350.

Clarke, F. B. Fire and Materials, 1999, 23, 109-116.

Houde, M.; Muir, D. C. G.; Tomy, G. T.; Whittle, D. M.; Teixeira, C.; Moore, S. Environ. Sci. Technol. 2008, 42, 3893-3899.

Mansour, S. H.; Asaad, J. N.; Abd-El-Messieh, S. L. J. App. Polymer Science 2006, 102, 1356-1365.

Shaghaghi, S.; Mahdavian, A. R. J. Polymer Research 2006, 13, 413-419.

Abdallah, M. A-E.; Harrad, S.; Covaci, A. Environ. Sci. Technol. 2008, 42, 6855-6861.

Weil, E. D.; Levchik, S. J. Fire Sciences 2006, 24, 137-151.

Mandalakis, M.; Stephanou, E. G.; Horii, Y.; Kannan, K. Environ. Sci. Technol. 2008, 42, 6431-6436.

Bromine Science Environment Forum (BSEF), http://www.bsef.com, accessed 16/11/2008.

Weber, R.; Kuch, B. Environ. Int. 2003, 29, 699-710.

Environment Canada, www.ec.gc.ca, accessed 18/11/2008.

Tange, L.; Drohmann, D. Fire and Materials 2004, 28, 403-410.

Ranganathan, T.; Zilberman, J.; Farris, R. J.; Coughlin, E. B.; Emrick, T. Macromolecules, 2006, 39, 5974-5975.

Levchik, G. F.; Vorobyova, S. A.; Gorbarenko, V. V.; Levchik, S. V.; Weil, E. D. J. Fire Sciences, 2000, 18, 172-182 

The Ethical Chemist

I was just at the 94th Canadian Society for Chemistry Conference in Montreal, QC, Canada. I have to say that the best talks that I attended were in the Chemical Education sessions.  My favourite talks were those on academic integrity and on chemiophobia. Both topics were extremely interesting. The academic integrity section was most interesting because it brought up the questions of "what is academic integrity"? "What is academic misconduct?" We as a chemical society need to take a more active role in teaching academic integrity. Our standard of ethics becomes equivalent to a doctor's Hippocratic Oath. I fully accept the point that was made in these talks by Stacey Brydges from the University of California San Diego and Tricia Bertram Gallant also from the University of California San Diego that we cannot just expect students to read the Student Code of Conduct on their first day of university and know exactly what that means. For example: http://www.uofaweb.ualberta.ca/gfcpolicymanual/content.cfm?ID_page=37633

This would be the link to the Code of Student Behaviour at the University of Alberta. I have been at the U of A for several years and have never once read this document in its entirety. It is boring, full of jargon, and confusing. Who reads this thing? So students say: don't cheat-got it. But that grossly overlooks many of the pressures and grey areas involved in academic misconduct. Let's face it, even the most senior academic can make mistakes that can ultimately cost them their job with regard to less than integrous decisions. http://www.theglobeandmail.com/news/national/prairies/deans-plagiarized-speech-prompts-investigation/article2059535/

This also ties into chemiophobia. What scientist has heard of research being over sensationalised in the media and gotten frustrated at a public who doesn't understand the difference between a cause and a correlation? http://www.phdcomics.com/comics/archive.php?comicid=1174 After discussing this with my sister, it occurred to me that it may be part of our ethics as a chemical society to work on educating the public on science. The medical community has put a lot of resources toward trying to educate the public on medicine. Why not the scientific community?  We owe it to our profession to break down the mysticism and scariness associated with chemistry. We owe it to our profession to educate the public on the scientific procedure. We can't just get mad at the media for not presenting it correctly. Did anyone ever try to educate the media on exactly what it is they are presenting?

I guess that is how I see the purpose of this blog. I am doing my best to educate the public on chemical questions. I wish to break down some of those barriers associated with chemistry and the public and present chemistry in a positive light. I suppose some of doing that means I should stop writing as a pseudonym and start writing as the chemist that I am, so that my credentials on this matter can be verified. So please, send me any chemical questions that you may have, and I will do my best to answer them.

Brenna Brown B.Sc. Honours (University of Alberta 2007)

Ph. D. candidate 

Department of Chemistry, University of Alberta

Infamous Inflammability

Hello loyal chemistry fans! I apologise for my long absence on the blog scene but it has been a busy two months in the world of this graduate student. Today's entry is a word lesson brought on by its confusing nature: inflammability.

I don't know about you, but every time I hear the word "inflammable" I think it means "not flammable". This would be very, VERY wrong.

 

Inflammable: (Merriam-Webster's Online Dictionary) 1) Flammable 2) Easily inflamed, ignited, or angered

Along with the word "inflammable" another important term to understand is "flash point". I am not talking about Canadian show with the Pink Ranger in it. I am referring to the lowest temperature at which a volatile liquid can vapourise, and therefore ignite, in air. This is really important to understand because it explains why you don't talk on a cell phone at a gas station or smoke while pouring out diethyl ether. An inflammable liquid is one that is considered to have a flash point below 37.8 degrees C (as stated in the Alberta Fire Code).

The flash point of gasoline is -43 degrees C. This is why here in Edmonton in the middle of winter we are still able to drive our cars. (It may be of interest to those readers who are not from the Canadian prairies to know that the coldest temperature recorded in Edmonton was -49.4 degrees C-and that does NOT include the windchill.) The point I am trying to make here is that knowledge of the flash point is crucial to the function of a combustion engine. It is also important to know when you are filling up engine with said fuel. Because the average temperatures that gasoline is stored at is well above its flash point, it doesn't take that much to volatilise to an ignitable mixture. And it so happens that electric charge from a cell phone may be enough to spark, and therefore set fire to, the gasoline you are trying to get into your car. Diethyl ether has a flash point of -45 degrees C. It is possible for static build up in the solvent bottle to cause this particular solvent to burst into flames.


If you don't believe me that static charge could actually set fire to something, let me tell you a story. The scariest lab moment for me occurred a few weeks ago when I was weighing out ruthenium dioxide hydrate. This is a pretty innocuous substance. (The "hydrate" part means that it has water in it and therefore should not be a fire threat.) I grabbed a flask and washed it out with acetone and water-not a problem because the solvents I planned to use in this reaction were acetone and water. When I added the ruthenium dioxide to the flask a ball of fire came spurting out the top. (No worries, it quickly burnt itself out, but the three seconds of flame left me a little jumpy the rest of the day.) THERE IS NO CHEMICAL REASON FOR THIS. Check the MSDS: ruthenium dioxide hydrate is a really safe chemical (all things considered). So what happened? My hypothesis: the flash point of acetone is -20 degrees C. There was enough heat created by the dissolution of the ruthenium dioxide hydrate in the small amount of water left in the flask that it was able to ignite the small amount of acetone vapours still remaining in the flask. 


The moral of the story: that sign at the gas station about not talking on your cell phone or having a smoke is not there just because the owners feel that talking on your phone while using their services is rude or they don't like the smell of cigarette smoke. It is there because of your safety and that of others.

Today's Blog is Brought to You by the Letter N and the Number 7

I recently was part of a conference where young girls in grade 6 are brought to the University to take part in science experiments. The lab that I personally teach with them is "Cryogenics". This lab is a lot of fun and involves investigating dry ice and liquid nitrogen. So today's blog is dedicated to the young lady who asked me about the discoverer of liquid nitrogen.

Nitrogen was discovered by Joseph Priestley (1733-1804) between 1773-1780. In this time he also discovered ammonia (a nitrogen based gas), carbon monoxide, sulfur dioxide, silicon tetrafluoride, and his most famous discovery: oxygen. On the periodic table, nitrogen is element 7, and is symbolised by the letter "N". Its boiling point is -196 C.

Nitrogen exists as a diatomic molecule, N2, making up over 75% of the Earth's atmosphere. The two atoms are held together by a triple bond that has a bond energy of over 800 kJ/mol. (This is A LOT, making this bond very favourable!)  In the form N2, nitrogen is completely inert; however, other molecules that contain nitrogen are not inert. Many explosive compounds contain nitrogen, the most notable being TNT: TriNitroToluene. Anything that has a high nitrogen content risks becoming explosive due to the extreme favourability in the formation of N2. 

The Haber Process: NH3 is ammonia and is an important fertilizer for crops (plants need nitrogen). However, very few plants actually have the ability to draw nitrogen from the air and convert it into a usable form. They require a symbiotic relationship with microorganisms in their roots to do this. This is called "nitrogen fixing". One of the most important discoveries in human history is the Haber Process.  While N2 may be inert, ammonia is not and in 1908 Fritz Haber worked out the necessary conditions for converting nitrogen from the air into ammonia. This process became instrumental to Germany's World War I effort because it allowed for the production of inexpensive explosives. Of course today this process is still used to generate ammonia and is celebrated for the feeding 1/3 of today's population with ammonia-fertilised crops. (Interesting Historical Side Note: Haber also worked with chlorine gas, and was instrumental in developing the practice of using this poisonous gas by German troops in the trenches during World War I. In 1933 this Jewish scientist, who had contributed to the German war effort considerably in WWI, was driven from his work by the Nazi regime.)

Back to nitrogen: liquid nitrogen was first obtained in 1895 with the invention of the liquifaction of air. This process also allowed for the separation of the components of air by distillation. Nitrogen is the most volatile (lowest boiling), boiling at -196 C. This is followed by oxygen at -186 C. And lastly, argon at -183 C. Liquid nitrogen is often used in flash freezing. Anything that has a high water content will freeze extremely quickly. In the lab I ran, I froze flowers, tomatoes, bananas, and latex gloves which the girls then smashed, much to their delight. (My favourite are always flowers as they shatter like light bulbs.) Being inert, liquid nitrogen is relatively safe and easy to work with. I actually freaked out some of the girls by showing that you can pour liquid nitrogen over the hand of a person (I had another lab helper not one of the kids do this) and it will not freeze them because the temperature of a person's hand at +37 C causes the liquid nitrogen to immediately evapourate into gas before touching the skin. Not unlike dropping water on a hot skillet. The big danger is with its ability to liquefy oxygen. This is a very reactive compound, to the point of scary! My favourite thing to do with liquid nitrogen is to make ice cream. If you ever have the chance to have liquid nitrogen ice cream, take it!

 
References:

Gray, T. The Elements 2009 Black Dog & Leventhal Publishers Inc., New York, NY.

Petrucci, R. H.; Harwood, W. S.; Herring, F. G. General Chemistry 2002 Prentice Hall Inc., Upper Saddle River, NJ.

Balchin, J. Quantum Leaps: 100 Scientists Who Changed the World 2010 Arcturus Publishin Limited, London.