It's Elemental, My Dear Mendeleev

The periodic table: one of the most iconic shapes. It is known from pole to pole, by scientist and non-scientist alike, and is as easily recognisable as the shape of Mickey Mouse. It adorns ties, t-shirts, and coffee mugs. It has found reference in art, and in pop culture. So what gave rise to this shape? Why is that that there are spaces and gaps? Are they placed there by accident or for artistic reasons? Well I am here to say: no, the shape of the periodic table was no accident. Each element has a deliberate and predictable space. Todays blog is about the meticulous observations and the man, who was one vote shy of a Nobel Prize, that resulted in one of the worlds most recognisable images: Dmitri Mendeleev. 

Elements: these are the varying types of atoms that make up the Universe. Each element is a single type of atom. Many different types of elements come together to make molecules. Oxygen is an element-only has oxygen atoms; water is a molecule-has both oxygen and hydrogen atoms. By the middle of the 19th century, many elements were discovered by many scientists. Their atomic masses had been calculated, and many of their properties had been observed. But they had yet to be organised in some meaningful way. Though many had tried, none had succeeded.

Mendeleev, Chair of Chemistry at St. Petersburg University, began by writing the properties of each known element on a card. He then began arranging these cards in various ways, and it was not long before a pattern emerged. By arranging the  elements in short rows according to atomic weight, placing the next row under the previous, columns emerged with elements that shared similar properties. Further, for this arrangement to work, gaps had to be left. Mendeleev believed these "gaps" represented elements yet to be discovered. Mendeleev therefore correctly predicted the elements gallium (discovered in 1875), scandium (discovered in 1879), and germanium (discovered in 1886). These elements, when discovered, easily fit into the gaps left for them. This table also showed that some of the atomic masses, such as those of gold and indium, had been incorrectly calculated. Mendeleev recalculated the masses and more accurate measurements would once again prove him correct. 

In the 140 years that have passed since Mendeleev's first periodic table, the shape has changed greatly. However, the concept remains intact. Those patterns first observed in Mendeleev's table have resulted in the continuous, periodic arrangement of each newly discovered element. 

The modern periodic table consists of 18 groups (columns) compared with Mendeleev's 8. Each element is now arranged by its atomic number, rather than its mass. The atomic number is determined by the number of protons in the nucleus. Element 1 is hydrogen; hydrogen has one proton. Element 6 is my favourite element, carbon; carbon has six protons. Element 101, discovered in 1955, is mendelevium, as an homage to the author of the first periodic table; mendelevium has 101 protons. Each group consists of elements with similar properties. Each period (row) has elements increasing sequentially in atomic number. 

One thing that is easily notable in the table is that there are distinct blocks, coloured differently in the table on the right. The green and red blocks consist of what are termed "the main group elements"; the yellow block is referred to as the "transition metals" and the blue block makes up the "rare earth metals" also called the "lanthanides and actinides". Each block has characteristic traits and give rise to interesting science. A lot of research has been devoted to studying periodic trends. It is important to acknowledge the blocks because they influence trends. For instance, going down a column in the main group block (example: group 14-carbon, silicon, germanium, tin, lead) we see that bond strengths decrease; however, going down the column in transition metals (example: group 8-iron, ruthenium, osmium) shows an increase in bond strength. 

Take a look at group 14 again, we see that silicon falls right below carbon. Ever wonder why science fiction nerds make comments, jokes, references to silicon-based life forms? The answer falls in the periodic table. Because silicon is below carbon, it has similar properties. We are carbon-based life forms. The ability of carbon to catenate (meaning it bonds to itself in long chains of covalent bonds) well is what allows for life on this planet; silicon, being in the same group, also catenates well, theoretically meaning that it could form the backbone of life on a different planet. 

It should be noted that there are some out there who feel that there are better, more accurate arrangements of the periodic table.

Mendeleev's table will forever be the first incarnation, quickly igniting revolution in chemistry.

References:

Petrucci, R. H.; Harwood, W. S.; Herring, F. G. General Chemistry 8th ed. 2002 Prentice-Hall Inc. Upper Saddle River, NJ.

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

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

 

Quasicrystals: A Nobel Story-A Follow Up

Since the Nobel Prize ceremony has taken place, I thought I would follow up with Dan Shechtman's story. I just watched a video of his Nobel lecture and wanted to share it. I think the overall message that he talks about is important to anyone pursuing enlightenment, not just scientist. His first statement: "Be humble". What would be the fun of doing research if we already knew everything there was to know about a subject? At the end of his talk (if you aren't going to watch the whole thing, at least skip ahead to the last five minutes) he talks about having courage and tenacity and belief in yourself. I think that that is the best part of this story. He was willing to stand by his results and what he believed was correct in the face of extreme opposition-he did go up against two time Nobel laureate Linus Pauling (and won!).  He backed up his claims with meticulous, well-planned experiments. Very inspirational.

So without further ado, Dr. Shechtman

Don't Worry, It's a Dry Cold

When people first move to Edmonton, they quickly begin hearing tales of how cold our winters are. These tales are met with the exasperated shout of disbelief: "It gets HOW cold?!" To which native prairie dwellers, such as myself, reply: "Don't worry, it's a dry cold." So what? Ever wonder why it feels way colder at -10 in Guelph than it does at -20 in Edmonton? It has to do with humidity (or the lack of) and a branch of chemistry referred to as "thermochemistry".

Our story begins with James Joule (1818-1920) and the first law of thermodynamics. Joule, being a brewer by trade, also shows us that chemists are rarely far from ethanol. 

First Law of Thermodynamics: energy can neither be created nor destroyed, only converted from one form to another. Heat is a form of energy. It can be transferred across the boundary between a system and its surroundings. Temperature is the measure of heat. Also important to know it that the direction of heat transfer is always from the thing that has the heat to the thing that doesn't. So leaving your front door open in the winter will not let any cold air in. It is physically impossible; however, you can let a whole lot of heat out. Another important term to know is heat capacity: the amount of heat required to change the temperature of a system by 1 degree. 

It is currently a balmy +1 in Edmonton right now.  The North Saskatchewan river isn't even frozen over, but I can tell you with certainty, I will not be jumping in for a swim. Even though the temperature of that river is actually warmer than the air surrounding it, it still feels a heck of a lot colder. This is because heat transfer is more efficient between a liquid and a solid than it is between a solid and gas. Heat is lost, at a molecular level, by collisions between the warmer body (you) and the colder body (the river or the air). Because of the fact that a liquid is more dense than a gas (especially a gas at cold temperature), there are more opportunities for molecules to collide, meaning more opportunities for heat transfer. The other important point is that the heat capacity of water is really big. Meaning that it takes a lot of heat to warm the water just one degree. If there is a lot of water in the air, it can condense on you, giving you that "wet" feeling. This will lead to more opportunity for heat transfer, making you feel much colder at -10, than if you are living in a place like Edmonton, where it is so dry that your skin begins to feel like an exoskeleton that you are much too big for. (I recommend Moisturel as a moisturiser for anyone looking to combat dry skin.) 

This heat capacity of water isn't all bad though. Ask people in Vancouver. See, because water takes so long to heat, it also takes a really long time to cool. Meaning that in the winter, places near water, like Vancouver, don't get that cold. Of course, one big snow storm and the whole city shuts down because they only own one snow plow, and have no idea how to function in weather that gives the rest of Canada the monicker "The Great White North". You can also use this heat capacity to your advantage when cooking dinner. Want to thaw your frozen meat faster? Stick it in room temperature water. 

So Edmontonians, when that mercury dips, and your skin and hair desiccate to a point beyond all recognition, be thankful for it. After all, it's a dry cold, so it really isn't that bad. Just bundle up.  

References:

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.

Laidler, K. J.; Meiser, J. H.; Sanctuary, B. C. Physical Chemistry 4th ed. 2003 Houghton Mifflin Company, Boston, MA.

Winter Tires: Don't Tread the Snow

Well winter has arrived in Edmonton. It is currently -17, with a windchill that makes it feel like -25 C. Over 15 cm of snow has fallen in 72 hours. The roads have become a delightful mix of ice and snow, making driving difficult. And it is not just here in Edmonton that citizens have been hit with a mound of snow and freezing temperatures. Calgarians are currently praying for their next chinook. So how can chemistry help you survive winter? With the science of winter tires! Why are winter tires mandatory in Quebec? Why are some Albertans lobbying for the same law in this province? Are winter tires that important? Well, anyone I have asked have all stated that they love winter tires and are shocked at the difference it has made. The difference all comes down to glass transition temperature (Tg). 

Take a look around your home. I am sure that you can find numerous examples of different types of plastics. Some are rubbery, some are hard, some are fiberous. These characteristics are going to determine how different polymers (plastics are a type of polymer) are going to be used. Now think of a plastic bucket. The kind that you may have used as a kid to build sandcastles. That thing was indestructible during the summer, but leave it outside in Edmonton right now and drop it, that same bucket would shatter into a million pieces. What we are observing is a change in "state" of the polymer. Now this might sound odd, considering it is still solid, and the states of matter are solid, liquid, and gas. So how can we be seeing a change in state? Enter the glass transition.

Polymers can have two solid states: they can be glassy; these are hard plastics, like cellphone cases and water bottles; or they can be rubbery; these are flexible plastics, like rubber balls, or tires. The glass transition temperature (Tg) is the temperature at which a polymer switches between the glassy state and the rubbery state. If a polymer is used BELOW its glass transition temperature, it will be glassy or hard. If a polymer is used ABOVE its glass transition temperature, it will be rubbery or flexible. The polycarbonate water bottle on my desk is an example of a plastic that I am using BELOW its glass transition temperature, while the flexible silicone spatula I used to make my breakfast is an example of a plastic I am using ABOVE its glass transition temperature. Going back to the plastic bucket example: in the summer, the bucket is above its Tg, so there is some flexibility to it and therefore, doesn't break easily. In the winter time, the bucket is below its Tg, making it glassy, and more fragile, so it breaks. 

At cold temperatures, rubber tires are also going to go through this change. Rubber tires were such a great advancement (thank you John Boyd Dunlop) in the tire because these air-filled rubber tires absorbed shock, had more contact area with the road surface, and consequently, gave more traction. The more a tire interacts with the road, the more traction a vehicle has. In the snowy, icy winter, we need all the traction we can get. To get a nice, flexible tire that has lots of contact with the road, it needs to be used above its Tg. However, in Canada, our winters are going to push that. Our -40 C winter days are going to bring a regular tire down to, if not crossing, its Tg. This will make it more rigid, and therefore, it will have less contact with the road surface, which will decrease the traction, precisely at a time when drivers want MORE traction. Also, the treads on the tire will become less flexible, allowing for snow to build up in them, further reducing traction.

Winter tires are made of a type of rubber that has a much lower Tg than summer tires or all season tires. This means that even as the mercury drops, the tire will not approach the Tg, and will stay flexible, resulting in more road contact, less snow build up, and MORE TRACTION. More traction means less sliding, smaller stopping distances, and safer driving. Enjoy safer winter driving thanks to the chemistry of polymers and the glass transition temperature. Get yourself some winter tires.    

For further winter survival reading check out a previous entry: Careful of the Icy Patch

Want more on winter survival through chemistry? Be sure to leave your questions and comments.

Quasicrystals: A Nobel Story

The announcements have all come in for the Nobel Prizes. It is the highest honour for a scientist. So who won the 2011 Chemistry Prize, in this International Year of Chemistry? Why Dr. Dan Shechtman, a professor of materials science at Technion-Israel Institute of Technology.

Shechtman's work concentrated on crystallinity, but what made his work so interesting, is that he discovered "quasicrystals", a concept previously held to be impossible. His work was met with a lot of criticism and skepticism. His results were suggested to be artifacts, or misinterpretations of microscopy results. His colleagues would hand him textbooks to study, pointing out that these books would clearly show his results were impossible. Overtime, these results were observed by other scientists, and the intense debate began to swing in Shectman's favour.

So what are these interesting and controversial materials Shechtman discovered? Quasicrystals. In crystallography, a crystal is a material in which the atoms are structured in a particular geometric pattern that repeats itself in three dimensional space. These repetitions are at fixed intervals. This means that there are clear definitions of what is allowed symmetry in the crystals, and the 5 and 10-fold rotational symmetry that Dr. Shechtman observed in his materials is certainly NOT allowed by conventional crystallography. In conventional crystallography, crystals may display rotational symmetries of 2, 3, 4 and 6. The reason that 5 fold periodicity is not allowed is because it cannot be exhibited by the crystal as a whole. While a single pentagon many have 5-fold symmetry, an array of pentagons will lack this symmetry. The previously held definitions of crystallinity included periodicity and yet Shechtman's crystals did not have this periodicity. Moreover, these crystals, that lacked the periodicity, were still ordered crystals. The result of Shechtman's work was a redefinition of what it means to be "crystalline" and the introduction of the concept of "quasicrystals".  

To better visualise this concept of periodicity and ordered structure take a look at the picture on the left. Taking a look at it we can agree that there is structure to the pattern, and it does follow mathematical rules, but the pattern is not regular. You cannot fold any part of this image onto itself. The pattern simply doesn't repeat. 

I think one of the most important lessons that we can take from Dan Shechtman is the importance in standing by your data. Too often we as scientists get biased by what we think "should" happen, or what we want to see happen. It becomes easy to shut down new possibilities and new ideas that can change our perceptions and give us a better understanding of the world. After all, is that not our goal in science? Shechtman stood by his data, ensured that his experiments were meticulously done, and did not capitulate to scientific peer pressure. He was able to win ultimate vindication in the form of a Nobel Prize. Congratulations Dr. Shetchman, your story is most inspirational to this chemist, and I hope to many others: scientist and non-scientist alike.

References:

Chemical and Engineering News, October 10th, 2011

West, A. R. Basic Solid State Chemistry 2nd Ed. 1999 John Wiley and Sons Inc. West Sussex, England.

The Chemistry of Flame Retardants: A Follow Up

Here is an interesting follow up to a previous entry on flame retardants. In the September 5th issue of Chemical and Engineering News an article was published on "greener flame retardants" http://pubs.acs.org/cen/news/89/i36/8936news8.html

Researchers have developed a two different types of coatings. In one case, researchers are using a thin coating of two polymers, poly(sodium phosphate) and poly(allyamine), in alternating layers on cotton. Material that was coated was shown only to char on exposure to flame, where uncoated material completely disintegrates. 

The second coating is one made of clay, or more scientifically, montmorillonite, that is alternated with chitosan-a polysaccharide that makes up the exoskeleton of insects. This was then coated onto a polyurethane foam and the result of the flame test showed that the material only charred.  

So what we see here are strides that are being made to find flame retardants that are as effective as halogenated flame retardants, but do not posses the environmental impact of halogenated flame retardants. 

To quote Nobel laureate Richard Smalley: "Save the world, become a scientist."

The Wizard of Os

This post has been inspired by a reaction that I am too terrified to try. The reaction is called an oxidative cleavage and is done using osmium tetroxide, or OsO4. I was once referred to as "osmiphobic" by one of my committee members during my candidacy exam. So lets take a look at this element called Osmium and then discuss why I am so afraid to work with OsO4. 

Osmium is element 76 on the periodic table. It is found in the transition metal section (the middle part) at the bottom of the same column as Ruthenium (Ru), element 44, and Iron (Fe), element 26. Osmium has an atomic mass of 190.23 g/mol. What is so neat about this metal? Well, unlike most metals, Os isn't grey (or silver). It is actually VERY faint blue colour. Other metals that are not grey or silver are gold (Au), element 79, and copper (Cu) element 29. 

Osmium is actually the hardest pure metal on the Brinell scale. http://en.wikipedia.org/wiki/Hardnesses_of_the_elements_%28data_page%29 Osmium is not the hardest element (carbon: in its diamond allotrope) or even the hardest material, these are mixtures of various elements. But in its pure form, it is the hardest metal. And metals make up the majority of the periodic table. Its hardness has made osmium useful for writing implements and is therefore found in the nibs of fountain pens. It was also used in the tips of needles for phonographs, which are (were) subject to significant ware. 

Interestingly, compared with other metals in this area of the periodic table, gold, rhenium, platinum, that are oxidation resistant, osmium will oxidise slowly in air, forming OsO4. Which brings us to this compound I am freaked out about working with. 

OsO4 is a solid material, but it is actually quite volatile. It will sublime (a process whereby a solid is converted directly to a gas without proceeding through a liquid state)  at room temperature. This is HIGHLY TOXIC! As a matter of fact, a small, sealed ampule of OsO4 obtained from Sigman-Aldrich has "highly toxic" or "very toxic" written on the label in no less than four places, with the accompanying "skull and cross bones" symbol of death. Have a look at the Materials Safety Data Sheet (MSDS) for this compound: http://www.2spi.com/catalog/msds/msds02595.html it is NOT pleasant. It has an extremely low LD50 (the amount necessary to cause the death of half of a sample set population-usually conducted on mice, rats, or guinea pigs). It is rumoured to have a smell something akin to that of ozone; however, if inhaled at levels well below that required to register a smell, this compound can cause pulmonary edema, resulting in death, so I am not anxious to confirm the smell of this compound.

Now I have worked with some pretty nasty chemicals: lead tetracetate, sodium cyanide, sodium azide, dimethyl aminopyridine (DMAP) the list goes on and on. Why I feel this compound is so different is because of this volatility. In its vapour phase, it is much easier to penetrate the safety barriers set up. For example: one should not wear typical safety glasses, but seal goggles because of the ease in which a vapour gets past safety glasses and the fact that one of the main target organs for this compound is the eye. It is much harder to protect against a vapour. 

There is a related compound for oxidative cleavage called ruthenium tetroxide, RuO4. Ruthenium is right above osmium on the periodic table. The difference between RuO4 and OsO4 in this reaction is that that RuO4 oxidises to the carboxylic acid-the highest oxidation state of carbon, while OsO4 oxidises only to the aldehyde-the second highest oxidation state of carbon. This is because of a neat little trend of transition metals. Anywhere else on the periodic table, the bonds get weaker as one moves down the table. But in the transition metals have the opposite trend, the bonds get stronger. This is why most catalysts are still made from the more toxic transition metals further down, rather than the less toxic transition metals in the first row, like nickel, copper, and iron. The bonds are too weak and fall apart easy. The metals in the last row are often not used because the bonds are too strong. Leaving the middle row as the "Goldie Locks" row-the bonds are just right. In the case of OsO4 though, those strong bonds work to our advantage by preventing over-oxidation. 

So that is the story on osmium and osmium tetroxide. I am sure one day I will get up the courage to work with this chemical, making sure to practice safe chemical handling techniques when doing so. I have, after all, worked with other equally nasty chemicals. 

Reference:

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

 

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. 

What's in a Chemical? Would a Rose Made of Any Other Smell as Sweet?

Hello followers. I apologise for the long delay since my last entry; as many chemists will understand, I had an exam to prepare for. An exam which is the most important of the Ph.D. path. So I can now get back to entertaining your chemistry questions (please keep them coming). During my hiatus, I also took in a chemistry conference in Hawaii. It was there that the inspiration for today's blog comes from. Here is a picture of a packet of sunscreen. If you notice, it says that it is "chemical free". So what is in this sunscreen? Let's take a look at the back of it: Active ingredients are zinc oxide and titanium dioxide. Those sound like chemicals to me.


Chemistry is the study of matter; therefore, anything that has matter is a chemical. Matter is comprised of atoms. Anything with atoms is a chemical. I am sure you can now see where I am going with this. This sunscreen is a chemical. This makes this packet blatant false advertising. If it isn't matter, then what is it? Energy?


This to me highlights one of the many negative public images that chemistry has and that chemists like me try to fight against. Just because something is a chemical, doesn't make it evil or dangerous! This actually goes hand-in-hand with the idea that just because a chemical is natural it is better. This one always makes me scratch my head as the most toxic things on this earth are nature made not man made. We have yet to even come close. Some chemicals are dangerous. Some cause cancer. Others can cause burns. Some will kill you. But you use chemicals everyday, and some are great. Some fight cancer. Can you imagine life without medicines? Those are chemicals. Imagine life without PLASTIC. Hell, imagine life without matter. Without chemicals there wouldn't be life.


Also, be careful when using things that say "natural, chemical-free". Often what these products are made of are botanicals (chemcials from plants). These are still chemicals and it is very easy to develop allergies to these chemicals, especially if you already have a sensitive skin. (This comes from a personal anecdote and the advice of a doctor.) 


Let's talk about matter for a bit. As I said, matter is comprised of atoms (side note: the concept of the atom was proposed by Democritus of Abera around 300 B.C. long before John Dalton in 1803 A.D.) When we talk about elements, we are talking about a substance that is comprised of only a single type of atom. Copernicium is the newest element: symbol Cn, number 112.  When we take about molecules (I, for example, am a small molecule synthetic chemist) we are taking about compounds comprised of more than one atom type. Example: Hydrogen is an element; it is comprised of only one type of atom. Water is a molecule; it is made of hydrogen atoms and oxygen atoms. Chemical changes or reactions refer to the conversion of one (or more) type of matter into another type of matter. For example: converting water into hydrogen gas and oxygen gas. Remember: matter cannot be created or destroyed, but merely change form.


I mentioned that I am a small molecule synthetic chemist. What this means is that my research is concerned with making small molecules. But what is meant by "small molecule" especially considering we are talking about atoms here. To measure molecules we talk about their atom/molecular weight and this is in the units grams per mole (g/mol). (The concept of the mole is a whole other entry so I will skip over that for now.) To get the molecular weight you simply add the atomic weights of all atoms in the molecule. Hydrogen is 1g/mol; oxygen is 16 g/mol; water, which has two hydrogen atoms and one oxygen atoms, is 18 g/mol. A "small molecule" is generally considered to be one under 700 g/mol. There is another type of molecule called "macromolecules". These are molecules whose molecular weight is well over 1000 g/mol. These are generally polymers. So things like DNA, proteins, and plastics are all macromolecules. You are one big, walking, talking collection of macromolecules!


Hopefully I have dispelled some previously negative views you had about what exactly chemicals are. Keep those chemistry questions coming friends!

References:

Petrucci, R. H.; Harwood, W. S.; Herring, F. G. General Chemistry 8th ed. 2002; Prentice Hall Inc. Upper Saddle River, N. J. 


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