Taking Medicines

Is it just me, or do all chemists guess the structure of a medicine before they take it? After having a guess, I then Google it and look up the synthesis, where possible.

I am currently using Ketoconazole for a skin condition (over-share, sorry). Immediately I knew there was a ketone (keto) and a nitrogen containing heterocycle (azole = five-membered nitrogen ring containing at least one other heteroatom) in there.  Conazoles are a common type of fungicide containing an imidazole or triazole ring so that narrowed it down further for me. I could not predict any thing more so looked up the structure (DOI: 10.1021/jm00194a023). It is actually a lot more complex than I expected.

ChemDraw2

Ketoconazole

Similarly, a few years ago I took Metronidazole and, as usual, I tried to predict the structure from the name. I guessed there was a nitrogen heterocycle, probably an imidazole ring and possibly nitro group. I wasn’t far off this time! When looking up the structure, I was very surprised by the simplicity of the structure and the synthesis (DOI:10.1007/BF00764821).

photo (2)

Metronidazole

So, is it just me? Or do you all try and predict aspects of a drug molecule before you take it? Have you ever guessed close to the true structure? If so, share in the comments section below.

Why do I use twitter?

It has been a while since I have blogged. This mini hiatus has been due to a hectic summer in the lab and the arrival of new graduate students, who need frequent direction.

More and more people in my workplace have found out that I use twitter and are surprised to find out how many followers I have and how long I have been a member. This had led lot of people to then ask me why the heck I use twitter. They seem to think that social media, in particular twitter, is for teenagers to follow the exploits of their favourite boyband. I have found twitter to be an excellent tool as a novice scientist and wish to express my feelings here (thanks to @_byronmiller for the helpful inspiration).

1) During my PhD

I found twitter very useful during my PhD write-up. I could vent to people who had gone through the same laborious PhD process as me, using the hashtag #phdlife. There were people on twitter just like me who could offer excellent advice and ideas from how to deal with the stress to how to improve my thesis writing, for example, I found out about Mendeley because I was exasperated with Endnote and tweeted my feelings. Many followers offered their experiences with various software packages and I ended up with Mendeley, which I still use to this day. Chemdraw are even on twitter now, so if you get stuck drawing a molecule you can just ask for help!

2) Literature searching

I am very capable at literature searching. I have RSS feeds, Google scholar alerts and I read my favourite journal website daily to see if there are any new developments in my research area. I have found, however, that twitter has broadened by chemical knowledge, with editors, academics and students regularly tweeting publications of interest from all areas. I even read computational publications now, due to the sheer number of computational tweeps online (I still don’t understand it all though, sorry).

3) Networking

I have been able to converse with the editors of journals and academics, join in with outreach activities (e.g. ScienceGrrl who recently gave me the opportunity to meet Brian Cox), contribute to magazine stories (e.g. Beth Halford’s Postdoc Pains and Gains piece) and even get free bits of lab kit! It is such an easy place to join in a conversation  and ‘meet’ people. It doesn’t matter if it is only over the internet, the people you converse with are just as important as those contacts who you have met in real life.

4) Job postings

People, including me, post jobs of interest that they have found, using the hashtag #chemjobs. There is even a tweep, disguised as a duck, who dedicates much of his time to job postings (@chemjobber). I used to regularly look at the Nature jobs and jobs.ac.uk websites but now I mostly just use their twitter feeds.

5) Conferences

Going to a conference? Don’t know anyone else going? Just tweet about the conference using a relevant hashtag and you are sure to find drinking buddies. I recently did this on a trip to Kyoto, not only did I find drinking and karaoke buddies, but I also found people who could recommend places to visit. Twitter is also a great place to find out about conferences and which talks to go to at a particular conference.

6) #Realtimechem

Ok, so I am a little biased, but I do love the #realtimechem hashtag. This hashtag is used to show the world what chemists are up to everyday, in and out of the lab. Some people post  awesome pictures and I get to learn so much about the techniques that other chemists use.

So readers, what one reason would you give a person to join twitter, or conversely, why do you think it is pointless to join twitter?  Leave your comments below or tweet me!

p.s. If you are new to twitter and not sure where to begin, I highly recommend Heather’s post on ‘How to Use Twitter’

p.p.s If I haven’t convinced you that Twitter is great, then believe Nature Chemistry.

#Chemnobel Family Tree

I am a glutton for punishment, so I have decided to compose a new chemistry twitter tree. This tree specifically links chemistry Nobel Laureates with twitter users.

Chemnobel6

There is plenty of space left on the tree. Are you linked to a laureate? If so, how? If you want to be added to the tree, post your connection in the comments section or tweet me. If you need help, then use this Academic Tree as a starting point.

If you spot any errors, do let me know.

The Girl with Zero Hangovers

I have not once in my life had a hangover. I am the envy of all my friends, they are disgusted when I wake up the morning after the night before all perky and bright while they are stuck nursing a stonking headache in bed or puking in the toilet bowl. It is not that I don’t drink, I am more than partial to a gin and tonic or five, a bottle of beer, glass bottle of wine etc, it is just that I have never once felt the after effects of a boozy night out. I have wondered for a long time about why I don’t get them and why others do, so I thought I would do some digging around and find out.

Hangovers are a result of many factors, with dehydration being high on that list due to ethanol (alcohol) being a diuretic which causes increased urine production.  The effects of dehydration include headaches, a dry mouth and tiredness. This can be counteracted by drinking plenty of water but often when we are a little tipsy we forget to drink water, preferring yet another margarita. Alternatively, ethanol can be toxic at higher levels and some say that this toxicity causes damage to the stomach lining which results in the nausea. Headaches can also be the result of alcohol being a vaso-dilator. This means that it causes the blood vessels in the body to dilate which can results in reddening of the face and painful headaches. Alcohol also causes a reduction in blood sugar levels which interferes with the generation of glucose by the liver and this can result in tiredness and dizziness.

Ethanol, however, is not the only reason for your hangovers, acetaldehyde is also thought to play its part. Acetaldehyde is known to form adducts with cell proteins which can lead to cell damage (Eur. J. Clin. Pharmacol. 1991, 40,187-188).

Alcohol is detoxified and eliminated primarily in the liver via a series of oxidative metabolic reactions. The three major steps are:

(1) reversible oxidation of ethanol to acetaldehyde, which is an acute toxin according to the MSDS.

(2) non-reversible metabolism of the toxic acetaldehyde to acetate

(3) breakdown of acetate to water and carbon dioxide

An excellent article by de la Monte describes in detail the issues with alcohol consumption and, therefore, acetaldehyde accumulation. (Oxid. Med. Cell Longev., 2010, 3, 3, 178-185).

The article demonstrates that alcohol dehydrogenase  (ADH) is the main alcohol oxidizing enzyme and breaks down ethanol in the cytoplasm. Cytochrome P450 2E1 is employed by a distinct pathway that is induced by alcohol consumption  and results in acetaldehyde formation. The second step, primarily carried out by mitochondrial aldehyde dehydrogenase (ALDH), is the metabolism of acetaldehyde to acetate. The resulting acetate is unstable and spontaneously breaks down to water and CO2. These oxidative mechanisms can become inundated which can result in toxic acetaldehyde accumulation. Disturbingly, experimental animal models have provided evidence that alcohol is a mutagen and that acetaldehyde is a carcinogen due to direct interaction with DNA and proteins. (H. Seitz, P. Becker, Alcohol Res. Health, 2007, 30, 38-41; D. Tuma, C. Casey, Alcohol. Res. Health, 2003; 27, 285-289)

The reason some us, including me, don’t get hangovers is because of polymorphisms in the ADH and ALDH genes. These can affect the rates of acetaldehyde generation and metabolism, and therefore change how much a person is prone to acetaldehyde toxicity.

Other impurities, such as methanol, can also be the cause of hangovers and these are a result of the alcohol fermentation process. These impurities are called congeners. The greatest amounts of these toxins are found in red wine and dark spirits such as bourbon. White wine and white spirits such as vodka contain less congeners and so in theory should give you less of a hangover. Buying more expensive alcohol should reduce the likelihood of these congeners since they should be distilled more rigorously. Different alcoholic drinks have different congeners and combining the range of impurities can result in a particularly bad hangover the next morning.

There are thought to be many other factors that contribute to hangovers including tannins in red wine and flavourings in dark beers. I mainly drink gin (a clear spirit) and tonic which may be another reason why I don’t get hangovers, although, conversely, I drink red wine in the bucket load which contains tannins and higher levels of congeners.

I hope you have learnt something about your hangovers and maybe how to reduce them with your choice of beverage. Feel free to leave any comments about your hangovers and what booze causes your affliction.

Carbon dioxide: A demonstration of the properties by Dr Jess

As you may know, I have recently helped to organise an open day demonstration at my University and I thought I would take this opportunity to show you what the end result was.

Carbon dioxide comprises of two oxygen atoms bound to a carbon atom. It is colourless, odorless, non-flammable, and slightly acidic. For most of these experiments I used solid carbon dioxide or dry ice. Carbon dioxide only exists as a solid and gas at atmospheric pressure and therefore sublimes between the two states. Dry ice sublimates at −78.5 °C (−109.3 °F) at atmospheric pressure and this extreme cold makes the solid dangerous to handle without protection due to burns caused by freezing. I don’t recommend playing with with dry ice. Please do not repeat any of these experiments at home. These were carried out by trained professionals wearing the correct PPE.

Carbon Dioxide is acidic

Carbon dioxide is acidic in solution. When it is dissolved in water it forms a weak acid called carbonic acid. To prove this, solid carbon dioxide (dry ice/cardice) was added to a solution of amine, in this case monoethanolamine, containing universal indicator. Universal indicator is purple/blue for alkalis and is red/orange for acids. You can clearly see that, on addition of the dry ice, we get a colour change from purple to blue to green to yellow to orange. Pretty cool huh?

Carbon dioxide is more dense than air

As you may know, the density of carbon dioxide is around 1.98 kg/m3 which is about 1.5 times that of air. To demonstrate this, we blew up a balloon (you can fill it with nitrogen also) and placed this on a layer of dry ice, which was subliming to give us CO2 gas. The balloon rested nicely on the layer of CO2 glass and gave us the illusion of a levitating balloon. We also showed this by blowing bubbles onto the CO2 but I forgot to take photos of that. The bubbles bob up and down on the layer of CO2 gas. It is pretty neat and well worth seeing for yourselves.

Some people asked the question, “but don’t we see balloons floating all the time when they are filled with helium?”…this of course is the opposite phenomenon, where helium is less dense than air.

Expansion of carbon dioxide from solid to gas

Dry ice and liquid carbon dioxide (which can be formed under pressure) expand they becomes a gas. Simply put, 44g (1 mole) of liquid CO2 will give approx. 22.4 litres of gaseous CO2. We demonstrated this expansion by filling a nitrile glove with dry ice and heating this up with water. We tied up the glove with a cable tie and watched it expand…behind a blast shield, just in case.

The glove got around 5-6 times this size. This experiment was very much a crowd pleaser and I am pleased to say that the glove didn’t once explode.

Carbon dioxide as a fire extinguisher

You will often have come across carbon dioxide fire extinguishers which are often used to put out electrical fires. The fire extinguisher basically works by replacing any oxygen, which is required for a fire, with carbon dioxide.

For this experiment you can use dry ice but you can also use carbonated drinks such as Coke or lemonade. As you can see, you put a CO2 source in a conical flask and seal it. You then attach a tube which allows CO2 to escape and you point the end of this tube at the ‘fire’. In this case we used a tea light but we also demonstrated that six birthday candles could be blown out.

Capture of carbon dioxide with monoethanolamine (MEA)

Carbon dioxide is a major greenhouse gas, responsible for substantial amount of global warming, currently ca. 380ppm. Post-combustion carbon capture removes CO2 from power station flue gas streams and can potentially be fitted to existing power stations. The process utilises amine solvents such as MEA, which reacts with CO2 to form a carbamate (shown below) or bicarbonate salt, and this holds it in an aqueous solution allowing other unreactive gases (N2, O2) to pass through. On heating to around120 ºC, the carboxylated amine regenerates the CO2 as a gas stream, and the amine solution is recycled.
In this experiment we added a solution of MEA to a pre-filled bottle of gaseous carbon dioxide. This creates a vacuum when the amine ‘captures’ the CO2 and crushes the bottle from the inside. This process is exothermic (gives out heat) and so the bottle gets very warm. Conversely, the release of CO2 is endothermic.
After CO2 capture to give carboxylated MEA, we then showed that CO2 release is possible. As previously mentioned this is usually carried out by simply heating the carboxylated amine solution. In this case, however, we used acid to release the amine and thus the CO2 gas. The carboxylated MEA was removed from the crushed bottle and poured in to a conical flask and acid added. The released CO2 was seen as bubbles which could also successfully blow up a balloon.
I hope you have enjoyed reading about my open day demonstration. Any ideas you have on future experiments to show off would be great and any comments on those shown above would be grand.

Alginates, more than just seaweed.

So anyone that knows me well will know I am bit of a foodie. I love eating at fabulous restaurants and gobbling down weird and wonderful things from scorpions (yes really) to fermented eggs (ick!). Having worked at a Michelin starred restaurant, I have seen molecular gastronomy first hand but lately I have seen more and more of my lab equipment in the kitchen, most recently I have seen chefs using rotary evaporators (!). I could write about so many different molecular gastronomy toys but the one that has got me most interested is the use of alginate gels in the kitchen. I have seen an increase lately in the use of sodium alginate by chefs, in particular to make artificial caviar or beads that supposedly give a taste explosion in the mouth. I really did not know much about how they were made and was embarrassed slightly when a friend of mine asked how they worked, so here is a brief history of alginates, particularly for use in molecular gastronomy.

In 1881, Stanford (E. C. C. Stanford, Chem. News 245-257, 1883) discovered a colloid (a substance that is microscopically dispersed evenly throughout another substance, much like milk) made from brown algae that he named algin. Stanford continued to investigate this substance and found that alkali salts, such as sodium and potassium alginate, gave viscous, aqueous solutions at very low algin concentrations. The algins were precipitated from solution by the addition of metal ions such as those of calcium and aluminum.1

Alginic acid is a linear polymer based on two monomeric units, β-D-mannuronic acid (M) and α-L-guluronic acid (G). The alginate polymer is formed by these monomers at the C-1 and C-4 positions. An alginate molecule is basically a co-polymer and the proportion of M and G blocks varies depending on the seaweed source with the properties of the alginate source being greatly dependent on the G to M ratio. Most alginate is currently extracted from just three of the 265 reported genera of the marine brown algae (Phaeophyceae). Macrocystis is the major genus used and is harvested off the west coast of the USA while in Europe we use Laminaria and Ascophyllum. These plants are mostly harvested naturally although large-scale cultivation does take place in China.


β-D-mannuronic acid 

α-L-guluronic acid


Alginates are used in food because they are excellent thickening, stabilising, and gelling agents and because, unlike other gels such as agar, alginate generally forms thermo-stable gels between 0 and 100 °C. Most alginate used in foods is in the form of sodium alginate. In order to form a gel, sodium alginate needs to come into contact with divalent ions such as calcium (Ca2+). As soon as sodium alginate is added to a solution of calcium chloride, a gel forms as the sodium ions are exchanged with calcium ions and the polymer becomes cross-linked. The longer the alginate is in contact with the calcium chloride solution, the more rigid the gel will become, as more cross-links with the calcium ions can be formed. Also, depending on the concentration of calcium ions, the gels are either thermo-reversible (low concentrations) or not (high concentrations).3

Many chefs are now using this process to make alginate beads, which they call spherification. This video shows how the alginate beads are made. For direct spherification, sodium alginate is added to the food that is being spherified and the droplets of food are then dropped into a calcium bath.Alternatively, for reverse spherification, sodium alginate is added to the bath in which calcium rich food is spherified. Spherification  was first introduced to the culinary world by the chefs at El Bulli and it is worth reading their story here.

Alginates have many uses other than in the kitchen, one of the most important being in medical applications. Alginates are used in wound dressing materials for the treatment of acute or chronic wounds. Calcium alginate is insoluble in water and can be woven into various textiles and bandages.4 The bandage is removed much more easily than other bandages, such as those made from cellulose, because calcium alginate can be dissolved in a simple salt solution. Alginates are also used in the treatment of cystic fibrosis, wherein bacterial biofilms formed from alginate gels are secreted by P. aeruginosa.5 Alginates are also used widely in the drug delivery applications.6,

I hope you have learnt something today about alginates and how they can be used for many different applications. I look forward to trying some spherified foods in the future, and when I do, I will update you further.

1         A. B. Steiner and W. H. McNeely, Ind. Eng. Chem., 1951, 43, 2073–2077.

2         P. Gacesa, Carbohyd. Polym., 1988, 8, 161-182.

3         A. S. Waldman, L. Schechinger, G. Govindarajoo, J. S. Nowick, and L. H. Pignolet, J. Chem. Educ., 1998, 75, 1430-1431.

4         C. . Knill, J. . Kennedy, J. Mistry, M. Miraftab, G. Smart, M. . Groocock, and H. . Williams, Carbohyd. Polym., 2004, 55, 65-76.

5         S. N. Pawar and K. J. Edgar, Biomaterials, 2012, 33, 3279-305.

6         S. A. Abukalaf, A. Badwan, A. Abumalooh, and O. Jawan, Drug. Dev. Ind. Pharm., 1985, 11, 239-256.

7         H. H. Tønnesen and J. Karlsen, Drug. Dev. Ind. Pharm., 2002, 28, 621-30.

Picture from molecularrecipes.com

I am the captain of Team Fluorine

So..I thought it about time I joined in with #ToxicCarnival since I spent 3.5 years of my life playing with oh so scary elemental fluorine for my PhD.

According to gospel Wikipedia “above a concentration of 25 ppm, fluorine causes significant irritation while attacking the eyes, respiratory tract, lungs, liver and kidneys. At a concentration of 100 ppm, human eyes and noses are seriously damaged”. The MSDS of fluorine also states that fluorine gas is corrosive to exposed tissues and to the upper and lower respiratory tracts.  Fluorine penetrates deeply into body tissues and will continue to exert toxic effects unless neutralized.  Workers should have 2.5% calcium gluconate gel on hand before work with fluorine begins. Fluorine also reacts violently and decomposes to hydrofluoric acid (which has previously been described as part of #ToxicCarnival) on contact with moisture.  Fluorine is the most powerful oxidiser known.  It reacts with virtually all inorganic and organic substances.  Fluorine ignites in contact with ammonia, ceramic materials, phosphorus, sulfur, copper wire, acetone and many other organic and inorganic compounds.

As you can tell, it is pretty darn unpleasant. Thankfully, the very pungent odour is detectable at concentrations as low as 20 ppb so you have time to escape, should you come across a fluorine leak.

A little history lesson… Fluorine was isolated successfully over a century ago by Moissan (Ann. Chim. Phys., 1891, 19, 272.) who gained a Nobel Prize in 1906 for his achievement. He produced fluorine by electrolysing a solution of potassium hydrogen difluoride in non-conducting liquid anhydrous HF. The electrolytic cell was constructed from platinum/iridium electrodes in a platinum holder and the apparatus was cooled to -50 °C. Today, fluorine is still manufactured using this electrochemical process.

The first large-scale production of fluorine was actually associated with the Manhattan Project during World War II, where uranium hexafluoride (UF6) was used to allow separation of the 235U and 238U isotopes. The radioactive uranium was used for the construction of the first atomic bombs in 1945 and uranium refining for nuclear energy is still one of the major uses for elemental fluorine.

So..so far you have learnt that fluorine is scary stuff and can be used to make atomic bombs. Now I am going to tell you why we should all love fluorine a little more.

Organo-fluorine compounds are almost non-existent as natural products but these days 20–25 % of pharmaceuticals contain at least one fluorine atom with these drugs treating a huge variety of diseases. One of the earliest synthetic fluorinated drugs was the anti-neoplastic agent 5-fluorouracil, an anti-metabolite first synthesised in 1957 (Nature, 1957, 179, 663-666). It shows high anticancer activity by inhibiting the enzyme thymidylate synthase, thereby preventing the cellular synthesis of thymidine. Since 5-fluorouracil, fluorine substitution is commonly used in med. chem. to improve metabolic stability, bioavailability and protein–ligand interactions amongst other things. An increasing number of related fluorinated anti-tumour agents have now becoming available as cancer treatments, including 5-fluoro-2’- deoxyuridine and its derivatives (Frontiers Biosci., 2004, 9, 2484-2494).

5-Fluorouracil is synthesised by bubbling fluorine through a solution  of uracil in a high di-electric constant solvent. If used correctly and safely fluorine can be a cheap and easy reagent, especially in large scale synthesis. EASY PEASY! Other fluorinating agents (mainly N-F) are seen to be an easier and safer alternative but these reagents can be expensive and wasteful. Using elemental fluorine is really all about knowing how to use it, for example it is best when the reaction is carried out at around -10C as a low concentration mixture in nitrogen. The key is to stop the competing radical reaction and promote the electrophilic process by polarising the F-F bond.

So what have we learnt?

1) fluorine is toxic and smells bad

2) fluorine can be used to make life-saving drugs (cheaply and easily if the infrastructure is in place)

3) I love fluorine a little too much.

If you want to know more about elemental fluorine in synthesis then check out publications by R. D. Chambers, G. Sandford and S. Rozen. All legends in their own right.

If you want to know more about fluorine, then just get in touch. I may or may not know the answer but I will probably be able to point you in the right direction.

Chemistry at the hairdresser

So, very recently I spent four (yes, really) hours sat in my hairdresser’s chair. After having read all the trash about what all those celebrities were up to and what ridonculous items of clothing I should be wearing in this so called British summer, I got to thinking about chemistry, more specifically, what transformations were actually happening to me and my hair, other than me slowly but surely looking like an alien from the planet Foil.

So what chemical processes have my ~100000 strands of hair gone through?

Hair is mainly keratin, the same protein found in skin and fingernails. The natural colour of hair depends on the ratio and quantities of two proteins: eumelanin and pheomelanin. Eumelanin gives us the brown/black hair shades (my natural colour) whereas phaeomelanin is responsible for the red based colours. Conversely, the absence of either type of melanin protein produces white/gray hair. 1

Hair dyeing by oxidation been practiced for well over 100 years and came from the observation that colourless p-phenylenediamine (shown below) produces a coloured compound when subjected to oxidation, and that this reaction could be used to colour a variety of substrates. The first patent relating to oxidation dyeing of human hair was applied for in 1883 by Monnet (F.P. 158,558).2 More scientifically put, permanent hair colouring involves the in-fibre formation of indo-dyes from colourless precursors by oxidation with hydrogen peroxide, under alkaline conditions. The primary intermediates are p-phenylenediamines (shown below) or p-aminophenols which are easily oxidised by hydrogen peroxide to form p-benzoquinone imines. 3

The use of hydrogen peroxide to develop the colour also allows for bleaching of the natural pigment by one or two shades at the same  time as the synthetic  colour is being formed.

The mechanism of oxidation dyes involves three steps:

  • The first step shows the oxidation of p-phenylenediamine (or similar stuctures, below) to the quinonediimine derivative

  • The second step involves the attack of this quinonediimine on the coupler (with chosen colour properties) by electrophilic aromatic substitution.
  • In the third and final step, the product from the quinonediimine-coupler reaction oxidises to the final hair dye.


Hair can also be dyed by bleaching of the hair’s natural pigments only. Bleaching is simply the removal of colour from hair. Bleaching can be caused by the sun’s ultra violet rays breaking bonds in the pigment molecules but hair bleaching is most commonly achieved by using hydrogen peroxide. Typically, a low volume of peroxide (5-30%) is applied to hair and left until the required amount of colour is stripped from the hair, at which point it is rinsed out. Before the bleach can change the colour of the hair, however, it must first penetrate below the surface of the hair’s cuticle. This is achieved by mixing the peroxide bleach with an alkaline solution, most commonly ammonia. The ammonia swells the hair fibres causing the cuticles to separate and open allowing the bleach to penetrate the cortex of the hair. This cuticle opening effect is also important when colour is being added to or implanted into the hair.

Hydrogen peroxide reacts with the melanin within the hair and in an irreversible reaction, the peroxide oxidises the melanin which renders it colourless. Complete bleaching tends to leave hair a pale yellow colour rather than pure white, however.

A really good review on hair dye in the modern world has been written by Christie and Morel. Well worth a read if you want to know more about dyes.

1            J. F. Corbett, Dyes Pigments, 1999, 41, 127-136.

2            J. F. Corbett, J. Soc. Dyers Colour., 1976, 285-303.

3            C. Incorporated and A. Seminar, J. Soc. Cosmet. Chem., 1984, 310, 297-310.

4            O. J. X. Morel and R. M. Christie, Chem. Rev., 2011, 111, 2537-2561.

5            http://www.chem.shef.ac.uk/chm131-2003/cha02js/dye.html