Right. You want to know about krypton difluoride. Fine. Don't expect me to gush. It’s a chemical compound, a rather volatile one at that, composed of krypton and fluorine. Apparently, it was the first compound of krypton anyone bothered to discover. How groundbreaking. At room temperature, it’s this pale, crystalline solid. It doesn’t exactly hang around; it’s quite volatile.
Names
They call it Krypton difluoride by its IUPAC name. Sometimes, just to be overly precise, they’ll refer to it as Krypton fluoride or, more technically, Krypton(II) fluoride. Because, of course, the oxidation state is crucial.
Identifiers
If you’re into that sort of thing, it has a CAS Number of 13773-81-4. There’s a 3D model of it, if you’re inclined to stare at a digital representation. It’s also listed in ChemSpider (75543) and PubChem (CID 83721). The UNII is A91DJL4OJC. The CompTox Dashboard from the EPA has it listed too. For the truly dedicated, there’s the InChI string: InChI=1S/KrF2/c1-3-2. And its SMILES notation is a stark F[Kr]F.
Properties
Its chemical formula is, predictably, F₂Kr. The molar mass clocks in at 121.795 g·mol⁻¹. As for its appearance, it’s described as colorless crystals when it’s solid. Don't expect it to stay that way. If you try to measure its density, you’ll find it’s 3.24 g cm⁻³ in its solid state. As for dissolving it? Don’t bother. It reacts. Violently, I’d imagine.
Structure
The crystal structure is a bit of a mouthful: body-centered tetragonal. The space group is P4₂/mnm, number 136. The lattice constant is given as a = 0.4585 nm and c = 0.5827 nm. But the molecular shape? That’s simpler. It’s linear. Imagine two fluorine atoms, perfectly spaced from a krypton atom. No dipole moment, naturally. A perfect zero D. It’s related to Xenon difluoride, which is, frankly, more interesting.
Chemical Compound
So, Krypton difluoride, or KrF₂, as the chemists insist on calling it, is a compound of krypton and fluorine. It holds the dubious honor of being the first krypton compound ever identified. Thrilling, I know. It's a volatile, colorless solid. The molecule itself is linear, with the krypton and fluorine atoms arranged in a straight line, about 188.9 pm apart. It has a tendency to react with strong Lewis acids, forming salts that contain KrF⁺ and Kr₂F₃⁺ cations.
The energy required to break KrF₂ apart into its constituent atoms is quite low – only about 21.9 kcal/mol. This translates to an average Kr–F bond energy of a mere 11 kcal/mol. That’s the weakest bond energy for any isolable fluoride, and that’s saying something. For comparison, breaking the F–F bond in elemental fluorine requires 36 kcal/mol. This makes KrF₂ a rather potent source of atomic fluorine, which, as you might guess, is extremely reactive and a strong oxidizer. It’s not exactly stable either; at room temperature, it decomposes by about 10% every hour. So, don’t leave it out. The formation of krypton difluoride is an endothermic process, meaning it absorbs energy. Its heat of formation, when in gaseous form, is measured at 14.4 ± 0.8 kcal/mol at 93 °C.
Synthesis
There are quite a few ways to cobble together krypton difluoride. Electrical discharge, photoionization, hot wires, even proton bombardment. If you manage to make it, you can store it at −78 °C, and it should stay put. For a while, at least.
Electrical Discharge
This was the original method, the one that first yielded KrF₂. It was even used in an attempt to create krypton tetrafluoride, though that turned out to be a misidentification. The process involves mixing fluorine (F₂) and krypton (Kr) in a 1:1 or 2:1 ratio, at a pressure of about 40 to 60 torr, and then blasting it with a significant electrical arc. You can get about 0.25 g/h this way, but the yield is notoriously unreliable.
Proton Bombardment
This method can produce up to 1 g/h of KrF₂. You bombard a mixture of Kr and F₂ with a proton beam at around 10 MeV and a temperature of roughly 133 K. It’s fast and can yield a decent amount, but you need a cyclotron or something similar to generate those high-energy protons.
Photochemical
Lucia V. Streng first reported its photochemical synthesis in 1963, with J. Slivnik following suit in 1975. This method uses UV light and, under ideal conditions, can produce 1.22 g/h. The sweet spot for wavelength is between 303–313 nm. Anything harsher is detrimental. Using materials like Pyrex glass, Vycor, or quartz, which block higher-energy UV light, significantly increases the yield. Experiments showed that Pyrex, blocking UV light above 280 nm, yielded the most, while quartz, blocking above 170 nm, yielded the least. It seems higher-energy UV light actively works against KrF₂ production. The best conditions involve solid krypton and liquid fluorine, at 77 K. The major drawback here is handling liquid fluorine, which is... problematic, to say the least.
Hot Wire
This technique involves solid krypton and a hot wire placed a few centimeters away. Fluorine gas is passed over the wire, which is heated to around 680 °C by a substantial electric current. This splits the fluorine gas into radicals, which then react with the solid krypton. Under optimal conditions, this can yield up to 6 g/h. The ideal gap between the wire and krypton is about 1 cm, creating a steep temperature gradient. The downside? The sheer amount of electricity required, and the inherent dangers if not set up correctly.
Structure
There are two known crystalline forms of krypton difluoride: the α-phase and the β-phase. The β-phase is generally found above −80 °C, while the α-phase is more stable at lower temperatures. The unit cell for the α-phase is body-centered tetragonal.
Reactions
KrF₂ is primarily known as a potent oxidizing and fluorinating agent, even more formidable than elemental fluorine due to its weaker Kr–F bond. Its redox potential is a staggering +3.5 V for the KrF₂/Kr couple, making it the most powerful known oxidizer. Some hypothetical compounds, like KrF₄, could theoretically be stronger, and nickel tetrafluoride comes close.
For instance, KrF₂ can oxidize gold to its highest known oxidation state, +5:
7 KrF₂ + 2 Au → 2 KrF + AuF⁻₆ + 5 Kr
The resulting KrF⁺AuF⁻₆ salt is unstable and decomposes at 60 °C into gold(V) fluoride, krypton, and fluorine gases:
KrF⁺AuF⁻₆ → AuF₅ + Kr + F₂
KrF₂ can also directly oxidize xenon to xenon hexafluoride:
3 KrF₂ + Xe → XeF₆ + 3 Kr
It's also used to synthesize the highly reactive BrF⁺₆ cation. KrF₂ reacts with SbF₅ to form the salt KrF⁺SbF⁻₆. The KrF⁺ cation can then oxidize BrF₅ and ClF₅ to BrF⁺₆ and ClF⁺₆, respectively.
KrF₂ can even react with elemental silver to produce AgF₃.
When a crystal of KrF₂ is irradiated with γ-rays at 77 K, the krypton monofluoride radical, KrF•, is formed. This species, identified by its ESR spectrum, is violet and remains stable indefinitely within the crystal lattice at 77 K, but decomposes at 120 K.