Interacting with the world is a sine qua non condition for reassuring my existence as constantly self-actualising being. I want to make a change on Earth. I cannot afford to let the miracle of life that I have been gifted with become a mercurial liquid that vapourises as soon as it falls into death’s hands.
A bit of love turns the darkest depths of sadness into a sea of light
Who said romance is reserved for fancy restaurants? A true scientist at heart pursues knowledge out of love, out of passion for learning about nature. What better gift to enamour someone with scientific awe than this romantic reaction?
Luminol’s a snitch
Luminol (C8H7N3O2) is a crystalline white solid that exhibits chemiluminescence. While at first glance chemiluminescence might seem practical only to spark our wonder, it actually has a crucial role in forensic investigations of crime scenes.
Whenever the luminol solution (containing other compounds, as explained subsequently) comes into contact with a drop of blood (or a few different things), it emits an alarm in the form of a short-lived but intense blue glow that can be photographically documented. No drop of blood will remain unnoticed. From chemistry, you can run, but you can’t hide!
To make luminol glow, two solutions must be made—one containing luminol and ammonia hydroxide (NH4OH) and another one containing hydrogen peroxide (H2O2). The reaction that takes place is shown below in a diagram I painstakingly drew from scratch using Marvin Demo.
First, luminol reacts with hydroxide groups (–OH) from the ammonia hydroxide in solution to form a luminol dianion in the keto form. This dianion undergoes tautomerization to the enol form (tautomers are structural isomers of a compound that readily interconvert).
Then, the hydrogen peroxide solution is added to the first solution, which contains the enol form tautomer of the luminol dianion. Then, the oxygen (O2) given off by the hydrogen peroxide reacts with the dianion to produce a very unstable cyclic peroxide. The latter, in turn, due to its instability, transforms into 3-aminophthalate, releasing nitrogen gas (N2). However, this 3-aminophthalate is excited and thus releases energy in the form of blue light (425nm) when it returns to its ground state.
Fiat lux! (But without the crime part)
“So what does the blood do?” you ask. Well, there’s something I didn’t tell you: hydrogen peroxide (H2O2) releases oxygen more slowly than molasses drips off your spoon on a January morning (if you’re at high latitudes in the Northern hemisphere, that is). Now that’s the reason you need blood: it’s a catalyst!
The haemoglobin (which contains iron) in blood and urine, the catalase in potatoes, the horseradish peroxidase in (guess what?) horseradishes, all are catalysts that speed up the decomposition of hydrogen peroxide into oxygen by providing an alternative pathway with a lower activation energy.
2H₂O₂(aq) → 2H₂O(l) + O₂(g)
As usual, my friend Máximo and I wanted to make light in the lab, though after some consideration, we decided to make luminol shine without giving the police a reason to pay us a visit. Opting for love over violence, we designed a magical heart that, when immersed in our special solution, would glow—a heart crafted from copper.
Copper cations (Cu2+) are another excellent catalyst for decomposing hydrogen peroxide, and are a tad cleaner to use than blood or urine. If we took a heart-shaped piece of copper and dipped it into our potion, the Cu2+ ions would relatively homogeneously mix into it, creating a stunning visual effect as the solution glows everywhere (pictured below). However, the homogeneous glow would render the heart shape invisible.
But don’t worry, I still have another trick up my sleeve! It’s called EDTA, short for Ethylenediaminetetraacetic acid (now that’s a mouthful!). This acid is a chelating agent, meaning it can form a dative bond with a single metal cation (such as our Cu2+) via two or more of its atoms, effectively sequestering the cation and preventing it from reacting with other substances (such as hydrogen peroxide). This chelation process is pictured below. M can represent any metal cation, but in our case, it’s Cu2+.
We added EDTA to our initial luminol solution, as it would act as a cage that traps rogue copper cations escaping from our heart-shaped wire, which would otherwise catalyse the decomposition of hydrogen peroxide, thereby eclipsing our heart’s glow.
Romance is the solution
Finally, after combining the two solutions, we can use a heart-shaped copper wire to catalyse the reaction and see the light only on its surface.
We also experimented with a regular coiled copper wire to see the effect more prominently. As we stirred the solution, the oxygen gas and copper ions dissolved, causing all the luminol to glow, and thereby achieving a higher overall brightness.
Being ardently enamoured by and fiercely hating something are inextricably linked insofar as their nature of intense devotion and dedication. They are two sides of the same coin. Akin to luminol’s light, this intensity of our character is a beacon that guides us through the dark void of absurdity inherent to our existence. Hence, we constantly face a decision: what will we choose to give meaning and light to every day we wake up? Will it be blood spilt with the hatchet of hate, or a heart pumped by the vehemence of love?
I prefer the latter: the ephemeral nature of existence unavoidably compels the self to live and act in the best way possible, and to my mind, that is through love and passion. Dislike and opposition to specific ideas are an integral part of an authentic being. Hate, though… hate demands an insurmountable devotion that one cannot afford to be stolen from where it rightfully belongs.
Life is brief. Luminol is limited in solution. Why choose hate?
All extractions I do are for mere educational purposes to learn about phytochemistry. No alkaloids were completely isolated, and the solutions containing them were disposed of after seeing their fluorescence. This website DOES NOT encourage alkaloid consumption or extraction, and STRONGLY ADVISES each viewer to consult their local restrictions before doing any experiment.
Personal Introduction
Ever since my grandfather gave me a mushroom book when I was around five years old, I have been very, very, intensely passionate about nature. What started with learning how to differentiate a Boletus from an Agaricus mushroom has grown into knowing a plethora of scientific names and travelling hundreds of kilometres to see an endangered plant. Parallel to this, ever since I realised I had hands, I have been joining atoms with my molecular model kit and inventing molecules too big to fit on an A4 sheet of paper. Here, I share with you one of the many testaments to my passion as a witness to the most fundamental characteristic of science: living in a perpetual cycle of melting and recasting to be at peace with each new truth unveiled from nature.
Extracting a molecule from the seeds of a desert-dwelling plant that’s both fluorescent and psychoactive? Is it a tale? NO! It’s reality.
Today, you will learn a bit more about the incredible romance between plants and chemistry alongside my journey as a paparazzi that adores documenting every one of their secret encounters. Embrace yourself for me to take you on this (quite literally) intoxicating love story. You may want to skip the extraction part, as it is tedious and dull (though it is written to maintain scientific rigour), and proceed to the“The Moment of Truth…” heading.
Peganum harmala, also known as Syrian Rue, is a herbaceous perennial plant from the family Nitrariaceae, native to the Mediterranean region and extending eastward to Central Asia. It inhabits saline, dry, and disturbed areas, even those that have been subject to eutrophication. It has been used to make traditional medicines, dyes, and incense, among others.
Like the caapi vine, Banisteriopsis caapi, Syrian rue contains a relatively high concentration of alkaloids from the β-carboline class, including harmine, harmaline, and harmalol. Some of these alkaloids are Monoamine Oxidase Inhibitors (MAOI) and therefore prevent our body’s enzymes (Monoamine Oxidases) from degrading hallucinogens like N,N-Dimethyltryptamine (DMT), making its effects stronger.
Last year, I was paging through articles on Google Scholar to find a research topic for a school task. I was left utterly dumbfounded when I discovered that harmaline and harmine are fluorescent, and what’s more, that I could easily extract these compounds myself in the lab. My school wasn’t impressed. Little did they know that a few months later, I would still do the extraction anyway.
And here begins the hunt…
Hunt
Is there anything more magical than the existence, the spontaneously arisen existence, of life? Stellar and supernova nucleosynthesis (the fusion of lighter nuclei to yield heavier nuclei) gave rise to almost all of Earth’s elements, including the soil’s nutrients, the air’s carbon, and water’s oxygen and hydrogen. Nuclear fusion (of hydrogen atoms to give helium) in the Sun powers almost all of life on Earth, which definitely includes ours. Plants, like artists of the atom, utilise enzymes to combine the Sun’s energy and Earth’s elements, creating a vast array of chemical masterpieces, such as our beloved harmaline and harmine.
Nuclear fusion reaction that occurs in the Sun. Image by Doctor C (Own work) [CC BY-SA 4.0], via Wikimedia Commons.
Consequently, I wanted to do everything from scratch to really admire the beauty of evolution. First, I visited a scrubland area in southern Madrid, where the plant had been spotted on iNaturalist. I went in shorts (clever me), offering ticks a free buffet and thistles a chance to paint my legs red, all to find not a single plant. With a sliver of hope, I travelled almost 200km to a town named Molina de Aragón in Guadalajara, Spain, where there was an iNaturalist observation of this plant, hoping to get seeds… and guess what?
Site of observation of Peganum harmala, with views of beautiful Molina de Aragón, Guadalajara, Spain.
Not a plant. Again, all my effort scrambling through chest-high vegetation (not pictured) was absolutely futile.
School was ending soon, and hence, my access to the laboratory. I urgently needed the seeds to carry out the extraction. I had no other choice but to buy them online.
Extraction
Grinding and defatting
The first step is grinding the seeds. This was done with a mortar and pestle. Then, they were defatted using pure n-hexane, CH3(CH2)4CH3, as it is a very apolar molecule due to the small electronegativity difference between carbon and hydrogen.
The seed and n-hexane mixture was stirred and then filtered. The filtrate containing the n-hexane was discarded, and the seeds were left to dry.
Acid-base extraction
The now-defatted seeds were placed in a dilute solution of acetic acid (CH3COOH), the main acid in vinegar, to increase the solubility of the alkaloids in water. The reaction mechanism is shown below in my own drawing of harmine’s protonation. The nitrogen in harmaline and harmine’s pyrrole accepts a proton (H+) with its lone pair of electrons, becoming positively charged and increasing the molecule’s solubility in water, a polar solvent.
Neutral harmine on the left and positively charged harmine on the right. Own drawing.
The mixture was then filtered. The residue (seeds) was discarded, and the filtrate was collected. This part contained all the desired alkaloids in their salt form. Now, using the opposite process, by increasing the pH to around 9 with sodium carbonate (Na2CO3), the acidic salts are converted to our desired harmine and harmaline-free bases (the neutral form of the alkaloid, such as the widely known cocaine freebase!). All of this was monitored using a pH meter, which measures the concentration of protons (H+) in solution using electrical conductivity.
Adjusting the pH of solution using acetic acid and sodium carbonate, keeping track of it with a pH meter. Own video.
Organic solvent extraction
The freebase harmine and harmaline solution was poured into a separatory funnel. Ethyl acetate, which is a less polar solvent than water, was added to get the relatively nonpolar alkaloids to migrate to this organic solvent. The two layers were mixed thoroughly and then left to separate. The aqueous and organic layers were separated and stored separately.
The moment of truth…
Like a kid impatient to open a present on a Christmas morning, I couldn’t wait to see if I really had extracted any alkaloids, and even more, I was anxious to see their fluorescence, which requires ultraviolet (UV) light of around 300nm.
I shone an old UV torch I found hidden in some old drawer on each of the solutions (organic and aqueous), but due to the brightness of the room and the visible light the torch was emitting, I couldn’t see any effect. Fortunately, months before, I had spent countless afternoons on a very special instrument working on my school chemical investigation: the mighty spectrophotometer. This instrument generates light of a discrete wavelength and passes it through a sample, measuring the amount of light absorbed. Water, which is (almost perfectly) transparent, for example, would absorb 0% of light at 500nm (greenish-blue).
I could programme it to pass ultraviolet light (I used 365nm) through my sample, and if I saw ANY LIGHT coming from it, then that’d imply fluorescence, as UV light is invisible to our human eyes. My molecule would be converting UV light to visible light!
I switched on the spectrophotometer. I had to wait 20 very long minutes for it to heat up (it uses a deuterium lamp to generate UV light). As an effort to make the laboratory as dark as possible, I took a poster of a scientific project I had presented at a conference with my friend (see the silver mirror post) and stuck it to a window that the blinds didn’t cover. I closed all the doors and switched off every single light, except the one of hope inside me.
Then, I placed my samples, one cuvette with the organic extract and another with the aqueous extract, inside the spectrophotometer, set the wavelength to 365nm, covered myself and the samples with a lab coat, and…
Fluorescence of harmaline/harmine extraction solution in quartz cuvette. Own image.
IT WORKED!
This minuscule speck of light my sample was re-emitting exhilarated me. This confirmed the presence of harmaline and/or harmine in the solution.
In that moment of ecstasy, I threw away the aqueous solution (which supposedly had no alkaloids) and bottled the organic one to keep it as a souvenir of my experiment. However, after this, I realised that only the aqueous solution fluoresced (probably due to an inadequate pH adjustment and a poor freebase conversion). I had thrown away all the alkaloids I had just spent hours extracting… luckily, I had no intention of using them.
Scientific Explanation
Harmine and harmaline fluorescence under UV light. Image by Coaster420, Public Domain, Wikipedia Commons.
A compound, depending on its chemical structure and therefore its molecular orbitals, may absorb photons of a specific, discrete energy, which is related to its frequency by Planck’s equation:
E=hv Where E is the energy, h is Planck’s constant, and v is the frequency of the light.
When a molecule absorbs a photon, it becomes excited, and the conformation of the molecular orbitals changes (some become antibonding, for example). When fluorescence occurs, the excited electron does not change its spin when it is promoted to higher levels and stays in a singlet state, denoted S, which means that all electrons are paired and hence that the net spin angular momentum is 0. The ground singlet state is designated as S0, and the excited states are labelled as S1, S2, S3, and so on.
Non-fluorescent molecules, when excited, return to their ground state via nonradiative transition, whereby the energy is released not as light, but usually as heat. On the other hand, fluorescent molecules, despite generally undergoing nonradiative transitions from higher excited states (S2, S3…) to the lowest excited state (S1), dorelease a photon when returning from S1 to their ground state (S0), as can be seen in the subsequent Jablonski diagram. As a side note, phosphorescence, a related but distinct phenomenon, is excluded from this discussion to keep it simple.
Jablonski diagram of fluorescence. The levels of excited states and levels of ground states are represented as subsets of the excited state (S1) and ground state (S0), differently from how I am referring to them in this article. Image by Jacobkhed at Wikipedia Commons.
As our fluorescent molecules (harmine and harmaline) lose some of the energy they absorbed via nonradiative transition, the photon emitted during fluorescence is of lower energy and frequency (and hence of longer wavelength) than the one absorbed. This phenomenon is known as the Stokes shift, which explains why I could see the light my sample was emitting, which falls within the visible range (380nm to 750nm), rather than in the UV range (100nm to 400nm), like the photons that excited my sample in the spectrophotometer.
Stokes shift. Image by CactiStaccingCrane – Own work, CC BY 4.0, Wikipedia Commons.
Conclusion
The minute speck of light that my sample emitted was a testament to something often overlooked in our quotidian lives: the truly incredible predictive power of Science. Owing to the invisible effort of a myriad of scientists and philosophers who laid (and lay) the foundation for our understanding of the world, we can know a fact so seemingly simple yet truly complex at heart (quantum mechanics, molecular biology, neuroscience, chemistry): that a plant’s chemical is psychoactive and will fluoresce under UV light.
Nature, like the peak of a towering mountain hidden by clouds, is perpetually enshrouded by an essence of change. In its never-ending chase, Science unveils fragments of its rugged terrain and shows us that even though the admirer and the admired do not share a synallagmatic contract, sometimes Nature does delight us with a glimmer of hope that we might know something of the intimidating cosmos we inhabit. —Own
Making an explosive gas out of sand, magnesium and drain cleaner
As you’ve probably noticed from my previous articles, I really enjoy making things light up, whether it’s through fluorescence, chemiluminescence, or simply causing them to explode. Today, I’ll share with you how my friend Máximo and I made a gas that does the job for you: it explodes on its own, it’s pyrophoric! Combining a few easily accessible materials, we synthesised silane (SiH4), the analogue of methane (CH4) that ignites and explodes upon contact with air.
Chemistry is cooking
In a beaker, we added a handful of magnesium powder and a few pinches of sand, using a rather unscientific approach. We mixed them up and put them in a boiling tube, covering the mixture with a bit of extra sand to prevent pesky oxygen (O2) from oxidising our magnesium. We then put it on a clamp stand outside (somewhere no one could see us) and placed a Bunsen burner beneath it. We lit up the gas and walked a few metres back.
It got scary. Smoke started seeping out, the sand-magnesium mixture turned black, and we heard glass cracking. In case it exploded, we took a long branch and launched the Bunsen burner far away, then ran to turn off the flame before our school went down in flames. I guess we were simultaneously crazy scientists and pyromaniac baseball players scoring a home run.
After waiting for the threat to subside, we approached the boiling tube, wrapped it in paper, and broke it with a stone, releasing the shiny crystals formed inside.
Magnesium silicide crystals formed after heating up silicon dioxide and magnesium powder. OWN IMAGE.
In this process, amorphous silicon (Si) and then magnesium silicide (Mg2Si) are formed via the following reactions:
1. SiO2(s) + 2Mg(s) → Si(s) + 2MgO(s)
2. 2Mg(s) + Si(s) → Mg2Si(s)
Magnesium silicide is the compound that forms the black-violet crystals pictured above.
We took the magnesium silicide crystals and ground them up in a mortar and pestle to increase their surface area. Then, we prepared a beaker containing a pint of dilute hydrochloric acid (HCl) and placed it on the magnetic stirrer in our darkened fume hood. To make silane, we added the magnesium silicide powder to the hydrochloric acid, producing the following reaction:
Mg2Si(s) + 4HCl(aq) → 2MgCl2(aq) + SiH4(g)
The silane bubbles out of solution, and when it comes in contact with the air… BOOM! This is one of the reactions where silane explodes upon coming in contact with air:
SiH4(g) + 2O2(g) → SiO2(s) + 2H2O(g)
A whirlpool of fireworks in a beaker of drain cleaner
To get the best views of this spectacle, you can watch my video in full screen.
Not only can science satisfy our deeply ingrained human desire to learn about the world around us, but it can also please the finest of tastes. Today we will be synthesising two chemicals that know just right which receptors to hijack in our nose: methyl salicylate and butyl salicylate.
Esters
Oh, gummy bananas! How can the copy smell better and stronger than the real thing!? One word: chemistry. You might be disappointed to know that with a lab and a few chemicals, one could very easily trick you without setting foot on a banana plantation.
Esters are famous for their strong smells and are commonplace in fragrances and aromas in almost everything you can imagine, from cosmetics to food products. They are a type of compound in the form R-COO-R’, which generally form from the condensation of a carboxylic acid (R-COOH) with an alcohol (R’-OH), giving water (H2O) as a side product. This specific process for forming an ester is known as Fischer esterification. The mechanism by which this process occurs is illustrated below.
Now, depending on the alcohol (R-OH) and carboxylic acid (R-COOH) used, different esters, and therefore different aromas, can be formed, as shown in the table below.
Like archaeologists sifting through layers of soil, my friend Máximo and I painstakingly combed through the safety cabinet in our lab, choosing methanol and n-butanol as our alcohol protagonists. To accompany them in their theatre performance, we decided to use salicylic acid, a key precursor to aspirin, as their partner-in-esterification. This would give us two compounds: butyl salicylate (ascribed to an intense raspberry aroma) and methyl salicylate (famously known as wintergreen).
Methyl salicylate is one of the most well-known esters, as it’s used to aromatise gum, root beer and medicines, amongst many others. It has a strong minty smell and was first extracted from the plant pictured at the beginning of this article, American wintergreen (Gaultheria procumbens), which can be used to make “oil of wintergreen”. Hence, while dressed in a white lab coat and at the dawn of a hot Spanish summer, we got straight to work on creating this refreshing aroma.
We only needed two test tubes and a beaker of hot water. In equal quantities, we added butanol to one test tube and methanol to the other. Then, we added a few grams of salicylic acid and stirred it around. We added a dash of 96% sulphuric acid because otherwise the esterification would have taken ages: lowering the pH of the solution catalyses the reaction, saving precious time. Tempus fugit! We put the test tubes in a beaker of hot water and waited about an hour. Unless you want to experience the dizzying and overwhelmingly wonderful mix of raspberry, mint, sulfuric acid, methanol, and butanol smells, then please do this under a fume hood.
A punch in the face
We took out the test tubes and probed the magic by pouring their contents into a beaker with sodium hydroxide (NaOH) to neutralise the pesky sulfuric acid.
WOW!
Unfortunately, we (still) can’t transmit smells digitally, but believe me, both the delightful wintergreen and raspberry smells were so intensely intoxicating that they gave both of us a (nice?) headache, and even people in other parts of the lab could notice their presence.
Don’t believe that an atmospheric-dwelling and life-giving gas can disinter hues from the underworld!? Just watch pool chlorine and hydrogen peroxide mix!
Today, I’m sharing the story of a straightforward yet breathtaking scientific experiment I conducted, which lies at the interface between chemistry and quantum mechanics: the chemiluminescence of singlet oxygen.
Summarised Reaction
Anything that glows is magical. No matter how omnipresent light is in our lives, it never ceases to fascinate our human curiosity. Consequently, if I combine it with chemistry in a fun reaction, then I’ll always be eager to try it out.
The simplicity of this reaction that gave bright red light enticed me: I only needed trichloroisocyanuric acid (TCCA, its abbreviation, (CONCl)3, its formula) and 30% hydrogen peroxide (H2O2) to produce chemiluminescent singlet oxygen. TCCA can be found in many supermarkets as a pool cleaner. 30% Hydrogen peroxide can only be found in authorised laboratories, due to its danger as a strong oxidant.
Trichloroisocyanuric acid reacts with water to give cyanuric acid ((CNOH)3) and hypochlorous acid (HOCl).
Then, the hypochlorous acid dissociates to give the hypochlorite ion (–OCl), which reacts with hydrogen peroxide to produce singlet oxygen (1O2). Chlorine gas arises from a few other reactions that aren’t relevant to the production of singlet oxygen.
My friend Máximo and I gathered both compounds and devised a perfect setup. As very toxic (used in WW1) chlorine (Cl2) gas is released, a fume hood is sine qua non for living on to learn more chemistry.
We covered the fume hood with posters to make it as dark as possible. We then made a hole in the posters to manoeuvre a beaker containing TCCA, allowing us to drop it into the hydrogen peroxide without having to open the fume hood, which would let ambient light in (and put our lives at risk, which might be more valuable than achieving pitch darkness). We set up the filming equipment and started the reaction. In the first trial, attempting to neutralise the chlorine gas with sodium hydroxide (NaOH), the solution splattered everywhere and left corrosion stains (still present today) on my computer, so we switched to a protected phone camera for filming. Here is a video of our beautiful production of chemiluminescent singlet oxygen.
Scientific explanation of singlet oxygen chemiluminescence
Gas from the heavens
Oxygen is one of the most essential elements for life, and is the third most abundant element in the universe. At standard temperature and pressure (273.15 K and 105 Pa), oxygen generally exists as the allotrope dioxygen (O2).
The ground (most stable, low energy) state of dioxygen is the spin tripletstate, denoted 3Σg–. The image below shows the molecular orbital diagram for this spin state.
What is characteristic of triplet oxygen is that the two non-bonding electrons of the molecule areeach in a different antibonding π orbital (one in πx* and one in πy*) and have the same spin (represented with arrow direction), conferring extra stability.
As a consequence of the aligned spins, triplet oxygen is paramagnetic (forms internal, induced magnetic fields in the direction of the applied magnetic field), as can be seen in the image below.
Paramagnetic property of triplet oxygen best observed in liquid state. Adapted image. By Bob Burk, work supported by the National Science Foundation under grant numbers: 1246120, 1525057, and 1413739 – [1], frame at 4:26, CC BY 3.0, https://commons.wikimedia.org/w/index.php?curid=57047554
Singlet state
The other spin state oxygen can exist in that is relevant to this experiment is the lowest excitedsinglet state, denoted 1Δg, which is higher in energy than the triplet state, 3Σg–. When oxygen is in its singlet state, it will react much more readily with other singlet molecules than when it’s in the triplet state.
In 1Δg, the two aforementioned non-bonding electrons are in the same antibonding π orbital (either πx* or πy*), and hence have opposite spins, following the Pauli exclusion principle.
When two molecules of singlet oxygen collide, they “deactivate” by exchanging their non-bonding electrons, producing two molecules of triplet oxygen (two molecules of 3O2) and LIGHT, due to the energy difference between singlet and triplet oxygen. In the image caption, there is a link to a video which explains it really well.
Deactivation reaction for two molecules of singlet oxygen. Image is a screenshot taken from a video of Random Experiments Int. – Experiments and Syntheses, which can be found at: https://www.youtube.com/watch?v=XDYAzdEhOGc. Permission was very kindly expressely given by the author.
Above, 1O2 and 3O2 represent singlet and triplet oxygen, respectively. The arrows show the spins of the electrons in the antibonding pi orbitals.
This chemiluminescent process is known as dimol emission. The reason why two molecules of singlet oxygen are needed to collide and exchange electrons is that the non-bonding electrons in each molecule of singlet oxygen CANNOT change their spin to occupy the other antibonding pi orbital, as is the configuration in the triplet state.
Using Plack’s equation, we can deduce the wavelength (λ) of the emitted light knowing the energy released (E), Planck’s constant (h) and the speed of light (c).
The energy difference between triplet oxygen, 3Σg–, and singlet oxygen, 1Δg, is 94290 J mol-1. As two molecules (not moles!) of singlet oxygen are combined to form two molecules of triplet oxygen, we must first divide the energy by Avogadro’s number (6.02·1023 mol-1) and then multiply by 2 giving a total release of 3.13·10-19 J per collision. Multiplying Planck’s constant times the speed of light and dividing by the energy provides the wavelength in metres, so it is multiplied by 109 to convert it to nanometres.
The emitted light, with a wavelength of around 634nm, corresponds to the red part of the visible light spectrum, which is precisely how we observe it in the experiment.
I went up on stage. For a moment of suspense that seemed like ages, I rummaged through my pockets and took out a lighter and a string of nitrocellulose pieces stuck with vaseline. As a surprise for the over 300 people in front of me, I ignited it, causing it to burn fiercely. After the few seconds it took to become ashes, I said: “¡Esta es la vida!” (meaning “This is life!” in Spanish). That’s how I started my Baccalaureate graduation speech on the fleetingness of life.
The nitrocellulose I used was made from scratch by my friend Máximo and me. Until the moment of the graduation speech, absolutely no one but myself knew what I was going to use that nitrocellulose for. However, because of my reputation at school as a chemistry enthusiast, after burning the magic paper on stage, many figured that I had synthesised it.
A towering polymer
Cellulose, the most abundant organic polymer on Earth, is a fundamental building block of green plants’ cell walls, as well as algae and oomycetes. The image below perfectly showcases the synergistic power of polymers. This banana species, Musa ingens, native to montane New Guinea, is herbaceous, i.e. non-woody, and thus its very tall stem is mostly built with cellulose, not lignin.
Usually, cellulose burns relatively slowly, requiring an external oxidiser such as diatomic oxygen present in the air (O2) to combust. However, if we replace the hydrogen (H) in the hydroxyl groups (OH) with a nitro group (R-NO2), we can make a nitrate ester (R-ONO2). The NO2 in the nitrocellulose is highly oxidative and thus accelerates combustion, as now it doesn’t depend on the oxygen present in the air.
Nitroxylation on paper
The process of converting cellulose to nitrocellulose, often wrongly called nitration, is denominated nitroxylation, as what is being formed is not a nitro compound (R-NO2), but a nitrate ester (R-ONO2). The mechanism can be seen below.
Mechanism of the nitroxylation of cellulose. Image by the Royal Society of Chemistry, which can be found at https://doi.org/10.1039/D3RA05457H.
To carry out the pictured reaction, nitric acid (HNO3) is used as a “nitrating agent”, and sulphuric acid (H2SO4) as a catalyst, being an excellent proton (H+) donor.
Nitroxylation of paper
After understanding the chemical reactions involved and reading the procedure and safety precautions, my friend and I went to the lab in a flash.
In two beakers, we added some 96% sulphuric acid, which is so viscous you could trick someone into thinking it’s mineral oil. Then, to experiment, instead of using nitric acid directly, we added potassium nitrate (KNO3) that would react with the sulphuric acid to make nitric acid and potassium sulphate (K2SO4). In one beaker we added folded paper towels and in the other, pure cotton. We left it to soak and nitroxylate for around 2 hours.
My friend, Máximo, on the left, and I on the right, with our two nitrocellulose trials in the background.
Afterwards, using a glass rod, we took both the paper towel and cotton out and put them in an alkaline solution of sodium hydroxide (NaOH), to neutralise any remaining acid. The two beakers still containing sulphuric and nitric acid were also neutralised with sodium hydroxide until a pH of 7 was reached (tested with pH strips). Following this, the solutions, now only containing innocuous salts, were washed down the drain. It is of paramount importance to make sure chemical waste is properly disposed of to avoid the slightest harm to the environment. How stupid would it be to destroy our beautiful World, the source of our awe!
The two unappetising solutions of sulphuric and nitric acid where nitrocellulose was in the making.
The now-nitrocellulose was then washed and placed in a food dehydrator to dry.
Nitrocellulose placed on a food dehydrator tray. Own image.
Watch it burn
After doing my last exam, the afternoon before my graduation day, I tested the nitrocellulose, and it worked wonders for the effect I wanted! The cotton one got more nitroxylated, probably due to a higher initial cellulose purity.
Heating sugar, two house cleaners, and silver nitrate to make a mirror
In March 2023, due to our reputation for being science enthusiasts and always messing around in the lab, our school chose us (my friend Máximo and I) to conduct an experiment for its inaugural scientific congress. A few months prior, we had performed the Tollens’ test for fun (as usual) to create a silver mirror on a round-bottomed flask. Hence, for our scientific project, we decided to investigate how the reaction temperature affects the deposition of silver.
The experiment is very straightforward. First, a few drops of a dilute sodium hydroxide solution (NaOH), used as a household cleaner, are added to a 0.1 mol dm-3 aqueous solution of silver nitrate (AgNO3). This converts the silver nitrate to silver(I) oxide (Ag2O) precipitate via the following reaction:
Then, an adequate amount of ammonia solution (NH3) (also used as a household cleaner!) is added to convert the silver(I) oxide into the diamminesilver(I) coordination complex ([Ag(NH3)2]+), the main component of Tollen’s reagent. This happens through the following reaction:
The flasks containing the diamminesilver(I) solutions were placed in a water bath and monitored using a thermometer until the desired temperature was reached.
Now comes the fun part: to make the silver mirror, we only need to add a reducing sugar (glucose). Yes, you heard just right, the solution only requires some sweetening to make the flask shine! We used glucose, but any reducing sugar (containing aldehyde groups) will do.
The reaction that takes place is what Tollens’ test consists of: an aldehyde (R-CHO) is oxidised to a carboxylic acid (R-COOH), and the diamminosilver(I) complex is reduced to elemental silver (Ag), which deposits on the surface of the flask.
After performing the experiment, we produced a graph illustrating the relationship between the mass of deposited silver and the reaction temperature. We found that at higher temperatures, more silver was deposited, probably due to a faster reaction rate (we only allowed the reaction to proceed for a limited time). We then created a poster and a presentation and presented our results at the scientific congress.
Chemistry doesn’t only bond atoms, it also bonds people!
With a wealth of additional scientific knowledge, two years after presenting it, we stood next to our poster in admiration, reviving all our lab adventures together. We took a picture so that in a few years, we can look back and admire how far we’ve come on our journey as pursuers of nature’s secrets.
The cheapest catalyst on the market: a Euro 2-cent to oxidise acetone!
Catalysis is a crucial process for a myriad of chemical reactions, whereby the catalyst, without being consumed, lowers the activation energy required for the reaction to occur. This increases the proportion of molecules with kinetic energy equal to or greater than the activation energy, and therefore considerably speeds up the reaction. Without catalysts (such as enzymes), this article would never have been written, as most probably humanity wouldn’t even exist.
Today, we’ll be catalysing the oxidation of acetone (CH3COCH3) with a Euro 2-cent coin, which is made of copper-coated steel.
The procedure is as straightforward as possible. First, a borosilicate beaker was filled with approximately 20 mL of pure acetone. Then, a 2-cent coin was tied to a copper wire and activated by heating it over a Bunsen burner flame (far away from the acetone). Once the coin was red hot, the wire was tied to a spatula so that the coin hung a few centimetres above the acetone. The magic starts immediately. The reactions involved are the following:
First, a coating of copper oxide (CuO) is formed when the coin comes into contact with air after being heated with the Bunsen burner.
Cu(s) + ½O2(g) → CuO(s)
Then, the copper oxide reacts with the acetone vapours to give ketene, methane, oxygen and copper, regenerating our initial catalyst. Acetone’s oxidation is highly exothermic, causing the coin to glow red-hot and catalyse even more acetone. This process can continue as long as there is sufficient acetone.
The temperature of the coin while it was catalysing the oxidation of acetone. Own image.
Leidenfrost effect
The Leidenfrost effect occurs when a water droplet is close to a solid surface that is hotter than its boiling point (100ºC at sea level), such as our copper coin, causing an insulating vapour to form and preventing the droplet from actually touching the surface.
The surface structure of the oxide layers on this copper wire, which we also used to catalyse the oxidation of acetone, gives rise to beautiful colours.
A powder that liberates a purple cloud on the caress of a feather?
That’s Nitrogen Triiodide (NI3), a contact explosive that releases nitrogen (N2) and a purple cloud of iodine gas (I2) at the slightest disturbance via the following reaction:
As my friend Máximo and I love both explosive and colourful chemicals, we got straight to making this compound when we had time.
The preparation is very straightforward; you only need iodine and a standard cleaning product (ammonia) to make this exotic compound. However, I recommend that you not try it at home unless you’re really longing for a purple wallpaper (apart from the fact that it’s dangerous even in small amounts). First, elemental iodine (I2) is ground in a mortar and pestle to increase its surface area. Then, it is added to a dilute ammonia (NH3) solution. The following reaction takes place:
3I2(s) + NH3(aq) → NI3(s) + 3HI(aq)
Once the reaction had occurred, the mixture of ammonia and nitrogen triiodide was spread on filter paper and left to dry, as nitrogen triiodide is insensitive to contact when wet.
We took the wet filter paper outside and left it on a table to dry. A few minutes later, we got near the incipient explosive to detonate it and watch the beautiful purple cloud form.
What a deception… Despite hearing a small crack, the explosion was far less impressive than we had hoped for, likely due to the infinitesimal quantity we had produced for safety reasons.
You reap what you sow, and you detonate what you synthesise!
Yesterday, we were driving on a narrow and very congested road. Cars in front of us were (luckily) avoiding something, strangely deviating their paths.
When we got closer, we saw it. There was a fledgling wood pigeon on the road, standing on its two feet, but tilting its head to the sky, as if the life in it was evaporating away, as if it had become immobilised after being knocked by a racing driver.
We slowly braked to a stop, putting our warning lights on. Out of an instinct even more animal than the one on the road, I jumped out of the car, I slammed the door behind me and ran on the asphalt until the bird was at my feet. In the few seconds elapsed between leaving the car and getting to it, a river of thoughts flooded my mind: “Is it half-knocked and suffering a terrible pain?” “Is it already dead?” “Am I going to witness its last breath?” “Does it have brain damage?”…
However, very, very fortunately, when I glanced down at it while strangled by adrenaline injected into me by rushing cars on the other lane, I saw its eyes glisten, glisten with life and hope. I was unspeakably relieved.
Usually (yes, unfortunately this is not the only disoriented/dead bird that I’ve found on the road), I pick up the birds and examine them to determine whether they need healing or veterinary attention. However, its ablaze eyes that only blossomed a few weeks ago lit a fire in me, I knew it was full of life and potentiality. I grasped it as if it were my own life that was at play. I scrambled through thistly, chest-high vegetation scorched by the Spanish August sun, I scratched all my legs, I rose my arms aloft and…
It flew!
Yes, it was alive, but not only alive, it was full of life.
When I went back to the car, I saw the driver of the van behind us smile at me and show a thumbs up. Exhilaration circulated through every vein and artery of my body, and a new star of anthropological optimism was born in my constellation of life experiences.
