Nasty Chemicals PRIMER - BOND ENERGY When we talk about "bond energy" or "bond enthalpy", we're talking really about how stable a chemical is. To do a reaction means breaking chemical bonds and making new ones. Making bonds released energy, breaking bonds absorbs it. We usually measure it in kilojoules per mole. Moles are the measure of "amount", in pure number, and one mole is roughly equal to 6x10^23. It's a big number, because we're discussing the atomic scale. So, if we have a bond which has a really high energy and we do some chemistry and the resulting bonds are very low energy, we've released some energy. We can do some accounting with that. Hydrogen gas (H2) is two hydrogen atoms bound by a bond of 436 kJ/mol. Oxygen gas (O2) is two oxygen atoms bound by two bonds, totalling 498 kJ/mol. A single O-H bond is 464 kJ/mol. We burn some hydrogen in oxygen. 2 H2 + O2 = 2 H2O We break 2x436 for the hydrogen and 498 kJ/mol for the oxygen, so we need to break 1,370 kJ/mol worth of bonds. Chemists will often call this the "activation energy". We've put in that much energy to break the bonds. We then make four O-H bonds (two water molecules, with two bonds each) which releases 4x 464 kJ/mol or 1,856 kJ/mol. We've got a net positive of energy of 1,856 - 1,370 = 486 kJ/mol. As we're releasing energy, this reaction will self-sustain. Hydrogen will burn in oxygen and release energy. PRIMER - ELECTRONEGATIVITY Some atoms attract electrons better than others. In general, atoms toward the top right of the periodic table attract electrons very well. Bonding is sharing electrons, so two oxygens can share two electrons and achieve a full shell of electrons (don't worry too much about shells) each. Full shells are thermodynamically very stable, so the bond releases energy. Some atoms don't actually want an electron, those toward the lower left of the table. Sodium has just one electron in its outer shell: Getting rid of it results in all full shells, again quite stable. Put it next to chlorine, which wants just one more to complete a shell, and sodium will happily be a good friend. Sodium then loses a single negative charge, so becomes positive. Chlorine gains one, so becomes negative. The two opposite charges attract and they pack together: We've formed common table salt, sodium chloride. Getting the sodium to release the electron is a little work (ionization energy) as you have to pull a negative charge away from a positive one, but you get all that energy back, and more, when it joins chloride. This is "ionic" bonding, as two ions are formed. In chemistry, losing an electron is "oxidation" and gaining one is "reduction". In the example above, the chlorine has oxidised sodium and become the chloride anion. Anions are negatively charged, cations are positive. PRIMER - BOND POLARISATION If a strongly electronegative element, like chlorine, meets a strongly electropositive one, like sodium, then the predictable happens. Sodium gets rid of the electron, chlorine gains it. This is purely ionic, there is no covalent (electron sharing in chemical bonds) character. If the "pull" of each atom is equal, like two oxygens, you get a purely covalent (but short) bond. If they're equal but not really all that great, like carbon and hydrogen, it's a fairly long bond. They're happy sharing, there's no pulling going on. A C-H bond is very strong because it isn't ionic in any way: It has nothing to gain by dissociating into ions. A single C-H bond is 439 kJ/mol in methane (CH4), while the O=O double bond is only 498, yet it has two bonds going on. Covalent bonds, when they're strongly polar, such as H-Cl (Cl wants the electron much more than hydrogen does, but hydrogen doesn't want to give it up as much as, say, sodium would) have a partial "ionic character". They behave like they're ionic in some ways. When HCl dissolves in water, it dissociates. The hydrogen and chlorine just go their separate ways. Chlorine becomes solvated and the chloride ion. Hydrogen joins with water to become H3O+, the hydroxonium ion. So technically, hydrogen chloride in water becomes hydroxonium chloride. Conventionally, we call this hydrochloric acid. The H3O+ is easily able to perform "proton donation" and give its hydrogen to something else, so it behaves like it's just a H+ ion. This is the defining characteristic of a simple acid. A simple acid can donate H+. A simple base or alkali can accept them. "Lewis acids" broaden this definition into a Lewis acid being able to accept a "lone pair" (a pair of electrons not involved in bonding) from a Lewis base, which is able to donate them. PRIMER - SIMPLE ACID AND ALKALI, AQUEOUS SOLUTION So if we have our HCl in water (really H3O+ and Cl-) and add in some sodium hydroxide, we then get Na+ and OH-, both solvated. Solvation is the surrounding of the species (e.g. Na+) with water molecules. They'll get around it with their negative (oxygen) ends pointing inwards (the lone pairs do it, look up what they are if you're curious) and allow it to "be dissolved". However, we have two species now which are incompatible: H3O+ and OH-. This is a simple dissociation of water. The dissociation of water happens constantly, and just reforms, so it's an equilibrium (more on these later in the primer) where the forward reaction and backward reaction happen at the same time. As we've added OH- with H3O+, we just get two H2Os. If the OH- and H3O are gone, what's left? Na+ and Cl-. This has become an aqueous solution of salt. PRIMER - EQUILIBRIA AND LE CHATELIER'S PRINCIPLE Some reactions have a reverse reaction. If we knock nitrogen apart in the Born-Haber process to make ammonia, it will just reform. So we need to "push" the reaction by taking away one of the components. This is Le Chatelier's principle, that an equilibrium will act against whatever disturbs it. If we take away one side of the equilibrium, the process will make more of it. This is easily explained: There's not as much left to do the reverse reaction, so we affect the rate of the reverse reaction, thereby enhancing the net rate to be more forward reaction. Helium Hydride You read that right. While the phrase "There are no compounds of helium" is mostly true, there is an ion. It can be prepared by ionising the dihelium cation (which also exists!) and then exposing it to molecular hydrogen. It also naturally occurs when a tritium atom decays in a HT molecule. It has a proton affinity of 177.8 kJ/mol, making HeH+ the strongest known protonic acid. It has an estimated pKa of -63. It will protonate oxygen, ammonia, water, carbon dioxide and simple organics, although they're blown apart by the energy released. More complex ions, up to He6H+ are known, where the helium clusters around the hydrogen. FOOF / O2F2 FOOF combines it all. The peroxide bond (that's a single O-O bond) is well known to be nasty as it's under a lot of strain from the two oxygens pulling at it. We combine that with fluorine on each end, which is more electronative than oxygen is, which increases the O-O bond strain even more, and has its own enormous bond strain. FOOF doesn't want to exist. Fluorine doesn't want to share electrons with oxygen and oxygen doesn't want to share them with fluorine. When it does, conventionally, it's in OF2 (like water, but with fluorine instead of hydrogen). Strained bonds like this are easy to break, so they have a very low bond energy. Oxygen difluoride (OF2) is fantastically strong at oxidising things - It can oxidise xenon! FOOF is, it's just even worse. FOOF, technically O2F2, has to be kept below -160C or it decomposes into O2 and F2, explosively. It reacts which practically everything. It'll burn ice. It'll burn sand. It'll burn through lab benches, grad students, lab tiled floors, classrooms below labs. Nitrogen N2 gas is not really that harmful. It's so inert that it plain doesn't want to react with anything. It's used as "protective atmosphere" in food, for example. The two atoms are bonded with a triple bond which has an enormous bond enthalpy: When it forms, it releases a lot of energy, so it takes just as much to break it. Breaking apart a nitrogen molecule takes 945 kJ/mol - So forming it also releases this amount of energy, and it's a lot. This means that forming nitrogen in a reaction can be extremely hazardous. Many explosives work on the fact that nitrogens in the same molecule are very unstable with respect to becoming nitrogen gas. When they do so, noise and motion often accompanies the reaction. Fluoroantimonic acid A mixture of the FH2+ ion and the SbF6- ion. It is the strongest acid known, a superacid, and can protonate even hydrocarbons. That ungodly fluoronium ion just REALLY wants to get rid of that extra proton and become HF (which isn't particularly friendly). It is 10^16 (10 quadrillion) times stronger than 100% sulphuric acid. It's so strong an acid that it cannot be stored in glass and cannot be solvated in water, it attacks both. It cannot even be stored in laboratory grade plastic containers, as it will eat through them. Only the strength of the C-F bond in PTFE can withstand this hellish acid. Reactions with it have to be done in pure hydrofluoric (HF) liquid, which is terrible stuff itself. Sulphur dioxide can also be used to solvate fluoroantimonic acid, where having HF around isn't desirable (such as always). The Hamett acidity function assigns a "H0" value to acids, the lower the stronger. Concentrated sulphuric acid scores a respectaible -12.0 while fluorosulphuric acid gets that up to -15.1 (it's a logarithmic scale). Even the terror of "magic acid", a 1:1 mixture of fluorosulfuric acid and antimony pentafluoride, only gets it to -19.2. Chemists thought the strongest an acid could be was around -20 or so until fluoroantimonic acid hit -31.3 in 1990. Hydrogen fluoride HF is a really tiny molecule, really little more than the size of a fluorine atom (hydrogen is tiny), and a stable, content, full-shell fluorine atom at that. A full shell is slightly smaller than a semi-completed one. HF boils at about 20C and will attack almost anything. The hydrogen bit of it isn't really all that fussed about keeping hold of the electron, because fluorine wants it so bad. The fluorine side of the molecule has a partial negative charge, leaving the hydrogen part with a partial positive charge. While the bond is covalent (sharing), it has quite a bit of ionic character. If one side of a molecule is more charged than the other, we call the molecule "polar". Water is also polar, for the same reasons. Polar molecules dissolve well in water, so HF does. HF passes straight through your skin, dissolves in your blood, acidifies it (it's not a great acid, but it still is one), the potential to form fluoride ions messes with enzymes, it can react with calcium in your bones to form soluble calcium fluoride (thereby dissolving your bones). It causes huge tissue damage which isn't immediately obvious, isn't often fatal, but is immensely painful and debilitating. The call "HF LEAK" in a lab is a priority-1 get-the-fuck-out-now. Water H2O is a bent molecule with the two hydrogens sitting on one side of the larger oxygen. Oxygen wants electrons (it's more electronegative) more than hydrogen, so it becomes slightly negatively charged, and the hydrogen slightly positively charged - This means the molecule is polar, it has a positive end and a negative end. In a liquid, it can surround charged ions, interfacing them with the bulk liquid, and what we call "dissolution". Ions like to be solvated, as there's a small energy loss in it. Dropping salt in water lowers the temperature of the water for just this reason. It can also do the same with other small polar molecules, such as ethanol, methanol, HCl and HF. Water's neat trick is that it has a slight dissociation constant. Water performs the following reaction: 2H2O <=> H3O+ + OH- The hydroxonium ion is the most common proton donator of simple acids, while the hydroxide ion is the most common proton acceptor of simple bases, so water can be both an acid and a base at the same time! The reaction is strongly biased toward the reverse reaction, so very little H3O+ exists at any one time in water. Complicating matters even more is water's tendency to hydrogen bond. The lone pairs on the oxygen atom can form a weak bonds to a hydrogen on a nearby water molecule. These bonds are constantly forming and breaking as water molecules slide past each other, however they do enforce a crystal lattice (it's a liquid crystal) and as water cools, they become more stable. This means that water is most dense at 4C, before it solidifies, as between 4C and 0C, water molecules are slipping out and joining the crystal structure, which is less densely packed. This also means that ice, solid water, is less dense than the liquid form is. Very few materials do this, usually the liquid is less dense than the solid, and the solid will not float on the liquid. Water is very strange compared to other small polar solvents. Hexanitrohexaazaisowurtzitane / HNIW / CL-20 What a name. The chemists among you will be shifting uncomfortably in their seats at reading it. Six nitro... six azo... oh my, that's a lot of nitrogen. Nitro and azo groups contain a lot of nitrogen, and nitrogens love nothing more than to release a lot of energy and become nitrogen gas, nestled comfortably together with that powerful, very high enthalpy triple bond. This one has the somewhat dangerous configuration of having nitro groups attached to azo groups. It takes the brave step of joining nitrogen to nitrogen in a single bond, which is tickling the nitrogen's tail. This was popularised by the development and "success" of RDX. Given that a nitro group also has two oxygens sitting nearby, this chemical wants to be anything but what it is, and it doesn't care how big a scar on the lab wall it makes to get there. The structure is three fused boat-conformation cyclohexane rings with the fusion points being carbon, but all points between being a C-N-C azo structure (triazine), with the N of that being attached to a dangling nitro group. It's like three molecules of RDX got together in an orgy of explosion. Work by Matzger and Bolton in 2011 showed that this ungodly product was stabilised by adding it as a cocrystal to TNT: Twice as stable in fact. Later, in 2012, another very brave group found it was stabilised by adding it as a cocrystal to HMX. When you have a chemical which is stabilised by adding it to a high explosive, you deserve a high five. Or four, three, however many fingers you've got left. Hexogen / Cyclonite / RDX It's like benzene, except some brave man replaced alternate carbons with nitrogens, then stuck a nitro group to that nitrogen. This "azo-nitro" is so common in explosives chemistry that it needs its own name. Perhaps "Fingerremover" RDX is quite stable and detonates only with a proper detonator. The Royal Air Force delivered lots of RDX to the Germans. Octogen / HMX Take a cyclo-octane. Remove every other carbon and replace it with a nitrogen. Off that nitrogen, hang an extra nitro group. Like HNIW, this is doing the unwise step of bonding nitrogen to nitrogen. It's what happens when you dissolve the high explosive RDX in a strong nitric acid solution, then heating it for six hours. It's chemically similar to HNIW, meaning it really dislikes everything. It was developed as a solid rocket propellant, but works well as an explosive. Octanitrocubane / ONC Cubane really hates existing. It's a cube made of eight carbons with eight hydrogens, bond angles are 90 degrees and immensely strained, the ideal bond angle is 109.5 degrees, from the core of a tetrahedron. Pushing the bonds closer to each other increases electrostatic repulsion between them, increasing their energy potential, or bond enthalpy. "So", some genius asked, holding all six of his fingers of both hands in the air, "what if we replace the hydrogens with nitro groups?" Presumably, his finger count was soon to be reduced further. ONC hates existing. It has an RE factor (how good an explosive it is) of 2.38. Octogen scores barely 1.7 and the index is benchmarked against TNT, which is defined as 1.0. Even the ungodly HNIW scores only 1.80. ONC has the incredible detonation velocity of 10.1 km/s, which is coincidentally the distance travelled in one second to get safely away from a laboratory handling it. Oxygen Difluoride OF2 hates everything. Oxygen is strongly electronegative and fluorine even more so. They hate sharing with each other. In it, oxygen has a formal oxidation number of +2, rather than its usual -2. Heat it with xenon and you get XeF4 and various xenon oxyfluorides. Its molecular configuration is much like water, with fluorine taking the place of the hydrogen. That's where the similarity ends. It is magnificently oxidising. It attacks most metals, forming oxides and fluorides. It'll oxidise glass, and glass is ALREADY oxidised! It'll oxidise its way through the strong hydrocarbon bonds of plastic labware, then oxidise through the lab's tiled floor, classrooms below... Perxenates The perxenate anion is an octohedral thing made of XeO6 with a 4- charge. Its conjugate acid is perxenic acid, which can't actually exist. It attacks its own XeO6 to produce xenon trioxide and oxygen. Like most noble gas compounds, it'll oxidise damn near anything. Krypton Difluoride The strongest stable oxidising agent known (atomic fluorine is stronger, but recombines into F2 molecules almost instantly). It has a standard electrode potential of +3.27 volts. Even hydrogen peroxide is only 2.08 V. Every moment this molecule exists is a thousand years of pain for both it and anyone nearby. Heat it above -78C and it decomposes explosively. It's more powerful an oxisiding agent than elemental fluorine, thanks to the Kr-F bond being ridiculously weak, even weaker than the F-F bond in fluorine. It can oxidise gold to a +5 state, forming gold (V) fluoride as Au2F10 molecules, a bright red solid which decomposes on heating. Xenon Difluoride OF2 will make this and FOOF will make it near instantly, if anyone had the stupidity to try. Guess what? It's a magnificent oxidising agent.