© 2026 Nanjing Finechem Holdings Co., Ltd,
All Right Reserved
Iron compounds have always played a surprisingly vital role in chemistry labs—sometimes outshined by showier transition metals, but rarely less important. Back in the early days of coordination chemistry, the focus hovered around simple salts like iron(II) chloride, better known in some circles as ferrous chloride. Chemists looking to improve solubility and reactivity soon started turning to ligands like tetrahydrofuran (THF). THF isn’t just another solvent; it actively engages with iron ions, offering more than just a vehicle for transport. In the 20th century, as organometallic work gained steam, combining iron(II) chloride with THF gave researchers access to new reactivity, pushing the field of catalysts and precursors forward in both academia and industrial labs. Iron(II) chloride tetrahydrofuran, though still not as celebrated as its noble metal cousins, has quietly carved out a niche.
In practice, iron(II) chloride tetrahydrofuran usually comes to researchers as a pale green crystalline solid or as a solution in THF. This form brings solubility and stability to a compound known for sensitivity to air and moisture. Lab workers who remember fiddling with stubbornly insoluble salts can appreciate the convenience—dissolving iron(II) chloride directly in THF skips several finicky steps. Iron(II) chloride tetrahydrofuran is not a household name outside specialized labs, but those who deal with reduction chemistry, cross-couplings, or unconventional syntheses will know how handy it can be.
Iron(II) chloride by itself doesn’t win any prizes for stability, especially under humid air. The THF-coordinated version presents a more manageable situation—it handles atmospheric exposure slightly better, with higher solubility in compatible organic solvents. The green color hints at its coordination environment, and this vibrancy is more than cosmetic; color changes during reactions often signal subtle shifts in the iron center. Those signals become lifelines when tracking sensitive process steps in synthetic chemistry. Still, ongoing exposure to air or water steadily degrades both the iron and THF, pushing experienced chemists to work swiftly and maintain a dry atmosphere.
The technical details seem straightforward on paper: purity often lands above 97 percent, and the compound exhibits a molar mass in the ballpark of what one expects for a THF adduct. Labeling varies between manufacturers but usually spells out the number of THF equivalents. This sort of clarity helps avoid confusion, since using the wrong stoichiometry torpedoes many sensitive syntheses. Safety data may stress the importance of dry storage and the toxic potential. It’s worth noting that THF complexes encourage a more uniform delivery of the iron center, especially for lab-scale reactions.
Anyone plunging into the synthesis of iron(II) chloride tetrahydrofuran most likely starts with anhydrous iron(II) chloride and dry THF. Results hinge on keeping everything scrupulously moisture-free. Iron(II) chloride powder, added to cold THF under an inert atmosphere, produces the complex after stirring. Recrystallization sometimes follows, depending on the purity level required. Chemists who have juggled Schlenk lines and argon-flushed glassware know the drill. Contamination by water or oxygen can quickly lead to intractable byproducts.
In the lab, iron(II) chloride tetrahydrofuran opens doors to a slew of reactions—Grignard-type couplings, reductions, and catalytic cycles that shy away from pricier palladium catalysts. THF’s role as a stabilizing ligand cannot be overstated; it delivers reactivity not easily matched by ammonia or simple aqueous solutions. The iron center interacts flexibly with organohalides, sometimes opening up uncharted reactivity lines. Custom modifications, like changing the THF/iron ratio, can tune the compound for unusual applications ranging from polymerization to organometallic synthesis.
Iron(II) chloride tetrahydrofuran shows up under a few alternative names—ferrous chloride-THF complex, or FeCl2(THF)x. Commercial offerings carry variations, too, with some highlighting the number of THF ligands per iron. Keeping track of synonyms matters not only to avoid ordering mix-ups but also when trawling the literature for reaction precedents.
Working with iron(II) chloride tetrahydrofuran should always come with a heads-up about its air and moisture sensitivity. Inhalation or skin contact brings risks, resulting in the usual precautionary gear: gloves, goggles, lab coats. Spills of the dry powder or solutions can stain deeply and pose a health hazard, especially if handled in poorly ventilated spaces. Emptying the jar straight into an open beaker rarely ends well. Emphasizing careful handling and the use of glove boxes or Schlenk techniques keeps mishaps at bay, a lesson every chemist driving a project in an academic or industrial setting learns early. Waste management also takes center stage, because both iron residues and THF are considered hazardous in many regions.
Across many chemical industries and university research groups, iron(II) chloride tetrahydrofuran steps up where greener or iron-based transformations are sought. Organometallic synthesis benefits from the compound’s solubility and ease of integration into reaction setups. It plays an unflashy but essential role in cross-coupling reactions as an iron-based catalyst. Pharmaceutical research leans on it for specific transformations that require mild but effective iron(II) centers, especially for constructing complex carbon frameworks. Polymeric material development has started to incorporate it in newer techniques aimed at sustainable plastics and functionalized polymers—clear signs that the compound’s relevance continues to grow.
R&D groups consistently push the boundaries with iron(II) chloride tetrahydrofuran. Newer catalytic routes emphasize earth-abundant, less toxic metals, so iron fits perfectly. Several teams working on active pharmaceutical ingredients look to swap out palladium for iron, slashing both cost and environmental impact. Ongoing collaborations between academia and industry are dissecting the compound’s behavior in unique reaction media, hoping to unlock greater selectivity or faster kinetics. Research grants have spiked for iron-based methods in hopes of advancing sustainable chemistry, bolstered by mounting regulatory pressure and cost considerations.
Toxicological studies on iron(II) chloride tetrahydrofuran reveal the double-edged nature of the compound. Acute exposure to iron salts can result in gastrointestinal and systemic effects, particularly in high doses, and THF is flagged as a potential carcinogen in long-term tests on rodents. Risk increases if spills or vapors go unnoticed in close quarters. Any laboratory storing and using this compound looks for adherence not just to minimum legal frameworks but also to best practice protocols borrowed from decades of occupational health data. Disposal practices reflect the toxic profile, prioritizing containment over dilution.
Iron(II) chloride tetrahydrofuran isn’t about to replace every transition metal catalyst, but it grows in importance as chemists sidestep precious metals in favor of cheaper and more sustainable options. The continued focus on green chemistry, along with advances in organic catalysis, gives iron(II) complexes valuable new relevance. The next decade will likely see expanded uses in battery development, greener polymer synthesis, and more ambitious cross-coupling protocols. As research teams publish detailed mechanistic analyses, the compound’s former reputation as a bench curiosity will keep receding toward a fixture in forward-thinking chemistry.
Take a peek into any organic chemistry lab, and you’ll probably find a bottle labeled Iron(II) Chloride Tetrahydrofuran. To most, that name seems a mouthful, but for folks handling synthesis, this compound does some pretty important work. A blend of iron(II) chloride and tetrahydrofuran (THF), the reagent has found a place in both academic and industrial labs, especially across pharmaceutical and materials research.
I remember my first semester in graduate research, surrounded by glassware and stacks of procedure notes. Nothing prepares you for the subtle differences a reliable reagent brings to the table. Iron(II) Chloride in tetrahydrofuran shines as a strong reducing agent. Chemists count on it for making organometallic complexes, driving cross-coupling reactions, or even helping with reductions that can’t go by standard routes. Unlike harsher or more toxic metals, iron keeps things both effective and a little easier on the environment.
Plenty of modern pharmaceutical compounds depend on stable forms of organic molecules — but getting the right bonds sometimes means a long chase, batch after batch. Iron(II) Chloride Tetrahydrofuran acts as a necessary shortcut. It clears roadblocks in multi-step syntheses where you want to reduce a functional group without touching other parts of the molecule. The chemistry world keeps scaling up processes like this, since pharmaceutical production often juggles tricky intermediates that need precision and reliability.
More and more, the conversation in chemistry circles has shifted toward sustainability. Green chemistry isn’t just a buzzword — it drives decisions in both research and commercial labs. Iron-based reagents stand out here, given that iron is abundant, relatively safe, and gives minimal toxic waste when compared to heavy metals like palladium or platinum. Let’s not forget, those metals come with costs that go way beyond the reagent jar. They’re expensive, they’re scarce, and cleaning up after them can be a headache.
I’ve seen process technologists gravitate to iron(II) chloride solutions for exactly that reason. Easier compliance with safety laws, lower waste disposal bills, and fewer health concerns go a long way. Where older methods relied on nasty solvents or oxidation catalysts, iron salts in THF work with much milder reaction conditions. Graduates and technicians both appreciate not having to dress for a toxic cloud every time a scale-up rolls around.
No reagent can fix every challenge, but the value of iron(II) chloride tetrahydrofuran keeps climbing as researchers demand better ways to build complex molecules. One persistent issue does pop up — storage and handling. THF attracts moisture, and iron(II) chloride can oxidize if left exposed. That means supply chain management and careful storage practices, both in the warehouse and on the bench.
Seasoned synthetic chemists and students alike need hands-on training in handling air- and moisture-sensitive reagents. Better education, robust containers, and tight partnerships with suppliers can cut down on spoilage and wasted batches. Top universities and pharma companies now bake this awareness directly into their lab protocols. The hope is to turn everyday challenges into quick troubleshooting, not stumbling blocks.
Iron(II) chloride tetrahydrofuran won’t make headlines, but it’s a quiet workhorse. Compounds like this continue to push innovations — from streamlined drug creation right down to more responsible waste management. Researchers, manufacturers, and regulators can learn from the chemists who use these solutions every day. Progress comes from practical, down-to-earth improvement, and this humble reagent proves just how much difference that can make.
Iron(II) chloride tetrahydrofuran refers to a coordination compound where a molecule of iron(II) chloride, also known as FeCl2, forms a complex with the organic solvent tetrahydrofuran, or THF. The standard formula for this complex is FeCl2(THF)n, with n usually being 1, 2, or sometimes more depending on synthesis conditions. In most laboratory syntheses and protocols, n=2 is typical, giving the formula FeCl2(THF)2.
Each Fe2+ ion bonds directly to two chloride ions, and the fifth and sixth coordination sites on the iron center play host to the oxygen atoms from two THF molecules. It creates a roughly octahedral environment around the iron, offering enough stability to allow the otherwise reactive FeCl2 to dissolve in non-aqueous solutions.
In the lab, FeCl2 itself doesn’t dissolve well in most organic solvents. Bringing THF into the picture allows chemists to coax FeCl2 into solution, opening up a whole range of possible synthetic pathways. I’ve seen firsthand the difference this makes. Trying to perform organometallic synthesis with plain FeCl2 powder usually leads to frustrating clumps stuck at the bottom of the flask, refusing to take part in the reaction. Once THF coordinates, the mixture turns into a clear solution, making reactions reliable.
The octahedral geometry isn’t just a structural curiosity; it changes how reactive the iron is. Coordination by THF stabilizes the iron center against unwanted oxidation and hydrolysis, so reactions involving sensitive substrates can be run more safely and with higher yields. That’s not a theoretical benefit. Many iron-catalyzed cross-coupling reactions run best with the THF complex because it stays homogeneous throughout—no solid settling out, no unpredictable side products.
Iron-based catalysts carry promise for more sustainable chemistry thanks to iron’s low toxicity and abundance. Most modern industrial applications still rely on palladium and other precious metals. With the right support from ligands like THF, however, iron catalysts can break into these established markets. That’s an important step not just for academic curiosity, but for everyone who cares about the affordability and environmental footprint of everyday products—pharmaceuticals, dyes, agrochemicals.
I’ve seen graduate students frustrated by the lack of clear instructions for preparing such iron complexes. Many recipes get passed around by word of mouth. To avoid mistakes, labs should provide direct access to reliable, peer-reviewed protocols and analytical data—NMR spectra, X-ray crystal structures, even photographs of the product. It’s not enough to have a chemical formula on paper. Chemistry thrives in environments where accurate, experience-based information circulates freely.
Iron(II) chloride THF complexes also help foster safer, more predictable experimentation. They reduce the risks associated with working with anhydrous iron salts and offer a stepping stone to more complex iron-based catalysts and materials. By solving a simple solubility and stability problem, FeCl2(THF)2 supports a much wider range of chemical innovation.
With careful training, open data, and shared laboratory experience, more chemists gain the confidence to adapt and experiment with iron-based complexes. THF-ligated iron(II) chloride is just one example of how incremental improvements in coordination chemistry ripple through science, industry, and sustainable practice.
Dealing with Iron(II) Chloride Tetrahydrofuran in the lab takes more than just reading a data sheet. This compound doesn't play nice with water or oxygen in the air. Let it mix freely with either, and you end up with not just a useless mess, but potential hazards and wasted money. Anyone who’s spent hours reordering chemicals after a poorly stored batch knows that frustration.
I learned early on that sloppy storage costs more than just the price of the reagent. Open up a bottle that’s taken on moisture and you get a distinct stench, crusted lumps, and unreliable results. It can mess with experiments and compromise the safety of the lab crew. Laboratories that don’t take this seriously end up with more spills and cleanup than actual science.
Iron(II) Chloride Tetrahydrofuran gets cranky when exposed to air. Oxygen can start to oxidize the iron, which not only changes its chemical properties but can turn a simple procedure into a guessing game. I’ve seen researchers grind their experiments to a halt because of reagents gone bad from a little negligence.
The first rule is simple—seal it tight and keep it dry. Always use containers with air-tight seals, preferably with screw caps or septa that can withstand regular puncturing with needles if you’re working under argon or nitrogen. Avoid glassware that has chipped rims or faulty threads because leaks are the enemy. The presence of moisture, even in small amounts, ruins the mix and triggers unwanted reactions.
A dry, cool cabinet works well. Find one away from direct sunlight, with temperature controlled to stay well below room temperature. Refrigerators dedicated to chemical storage keep the environment stable and safe from humidity, dust, or accidental knocks from non-chemists searching for snacks.
Use a desiccator or dry box if you've ever had a problem with humidity in the lab. I’ve kept many jars dry even in a tropical environment, all thanks to an old desiccation unit packed with fresh silica gel. Add a label with the date of opening, so you know when things start to go stale.
Gloves, lab coats, and goggles aren't just for show. Spilling Iron(II) Chloride Tetrahydrofuran can cause skin and eye irritation. Inhaling dust or vapors won’t do your lungs any favors. Work with this chemical in a fume hood; keeping the work area ventilated cuts down on exposure to vapors that drift off during transfer.
It pays to transfer the substance using syringes or pipettes under inert gas—you avoid splashing air and moisture in. I’ve worked with colleagues who used to rush this step, only to lose half their precious reagent to oxidation or spills. Slow and steady wins here.
Keep cleanup supplies nearby, too. A spill kit and plenty of absorbent pads save a lot of trouble if a beaker gets jostled or the bottle tips over. Make disposal easy with clearly marked containers for solvent waste.
Iron(II) Chloride Tetrahydrofuran rewards careful, methodical storage. The little steps—double-checking seals, working under gas, labeling dates—build a culture of safety and reliability. These habits pay back every time an experiment yields consistent results and nobody dashes to the eye wash in a panic.
Good storage and handling are less about fear and more about respecting the work, time, and money invested in science. Keep the basics solid, and the rest gets easier.
Working with Iron(II) Chloride Tetrahydrofuran isn’t quite the same as mixing sugar in your morning coffee. This chemical comes packed with hazards, many that sneak up fast if you aren’t thoughtful about your approach. Splashes can irritate the skin, vapors can set your eyes and lungs on edge, and chemical spills have a way of turning a good day upside down. Having spent time at benches in academic and industrial settings, I’ve seen people underestimate substances like this because they look harmless right until they’re not.
Every time someone heads to the bench with open-toed shoes or skips their goggles, I wince. Long lab coats, closed shoes, chemical-resistant gloves, and protective goggles take the sting out of surprise splashes or drips. Cotton clothing helps since it doesn’t melt onto skin the way synthetics can if things go off the rails. Gloves deserve special mention—choose ones that actually fend off both Iron(II) Chloride and the solvent tetrahydrofuran, since some gloves crumble with time or become brittle after exposure.
Tetrahydrofuran’s vapor packs a punch; inhaling it for even a short time brings headaches and dizziness. I’ve watched colleagues relax in the fume hood, trusting its shield, and forgetting that if you open a bottle outside its reach, the air can fill with hazardous fumes. Flame sources should stay far away since the solvent burns easily. Any heating step or open flame nearby increases the risk of something catching. Fire blankets and extinguishers have their place in the lab, but far better to steer clear of risky setups to begin with.
People forget just how unpredictable chemicals become if stored near incompatible substances. Iron(II) Chloride with moisture or strong oxidizers can create a toxic mess. Don’t stock it on just any shelf. Dry, cool spots that stay out of direct sunlight help. Containers need clear labeling and solid seals since leaks or evaporating solvent creates a new problem before anyone notices.
I’ve seen panic freeze even seasoned chemists when a bottle tips or a beaker shatters. Small spills respond well to neutralizers and absorbent materials made for chemical clean-up. Large releases push the lab into emergency mode. Every researcher should know right where the eyewash station and safety shower stand, not just in theory, but as muscle memory. Quick action—strip contaminated clothing, rinse skin, and get outside air—can prevent a burn from turning into a hospital stay.
Many forget the value of good preparation. Training isn’t box-checking or a formality. No one becomes careless on purpose, but accidents start creeping in when safety talks feel optional. Labs where people support each other and call out shortcuts before corners get cut keep accidents at bay. Sharing best practices and talking openly about close calls turns every mistake into a lesson everyone learns from, not just the one person involved.
Working with chemicals like Iron(II) Chloride Tetrahydrofuran doesn’t have to turn risky. Habits forged from genuine respect for what these substances can do keep days uneventful, and that’s how people go home with all ten fingers.
Solubility questions look straightforward, until you have to mix the compounds for real. Iron(II) chloride itself causes trouble since it reacts fast with moisture and air. Tetrahydrofuran, or THF, usually acts as a good solvent for a lot of inorganic compounds, thanks to its polar and aprotic nature, but iron(II) chloride doesn't always play along.
Put a pinch of iron(II) chloride into water, and you get rapid dissolving. A green solution forms, sometimes with a brown tint if there’s too much contact with oxygen. Many undergraduate assignments ask about water as a solvent, which gives a clear answer. Water works, fast and completely.
Switch to THF, and the scenario changes. Based on practical chemistry work — and what's documented in chemical literature — iron(II) chloride tetrahydrofuran forms as a clathrate or complex. Instead of breaking down smoothly into THF, it tends to stay as a poorly soluble crystalline solid. Stirring and heating offer little help. You feel like you’re waiting for a sugar cube to dissolve in cold oil. Some partial solvation occurs, but not enough for most reactions.
In the last decade, organometallic chemists started using THF to try and stabilize iron(II) chloride in reactions needing soft donors. THF can bind as a ligand, providing some stabilization, but hardly improves solubility. The crystals can grow slightly sticky if left overnight, cluing you in to partial solvation, but full dissolution slips out of reach. Attempts at filtration, or warming, usually just lead to decomposition. Any exposure to air gives you rust-colored iron(III) oxides.
My own attempts at dissolving iron(II) chloride in THF back at university led mostly to frustration and a mess at the bottom of the flask. Even purchasing the pre-formed "FeCl2·THF" complexes from vendors didn't help much, as they too came as clumpy, light-sensitive powders. Chemists with years of glovebox experience agree: the complex isn't soluble in THF beyond trace amounts.
Popular solvents like ethanol, methanol, and acetone barely impact the issue. Iron(II) chloride stays mostly at the bottom, either unaffected or decomposed. Diethyl ether, highly nonpolar, offers no benefit for iron(II) salts. The only liquids that actually give full solution are water (as discussed) and toxic dimethyl sulfoxide (DMSO), which brings a whole new layer of handling hazards.
Iron(II) chloride’s reluctance in regular organic solvents restricts its use in catalytic reactions or industrial processes. Reliable solvents improve reproducibility, safety, and scaling. Chemists have to work harder, drying everything and often relying on gloveboxes. For anyone working outside a well-funded research lab, these hurdles pile up.
A possible workaround? Chemists sometimes start with soluble precursors or generate the iron(II) chloride in situ, right inside the reaction mix. This means adding reagents directly to the main flask, making the lab workflow simpler and safer. The approach doesn't solve every problem but can save resources and reduce toxic solvent use. Iron(II) chloride’s solubility quirks push professionals to get creative, make safer choices, and keep learning from each experiment.
| Names | |
| Preferred IUPAC name | iron(2+) tetrachloride tetrahydrofuran |
| Other names |
Iron dichloride tetrahydrofuran complex Ferrous chloride tetrahydrofuran complex Dichloroiron(II) tetrahydrofuran complex FeCl2·THF |
| Pronunciation | /ˌaɪərən tuː ˈklɔːraɪd tɛtrəˌhaɪdrəˈfjʊəræn/ |
| Identifiers | |
| CAS Number | 13815-68-8 |
| Beilstein Reference | 3589812 |
| ChEBI | CHEBI:132679 |
| ChEMBL | CHEMBL1230533 |
| ChemSpider | 53701112 |
| DrugBank | DB14608 |
| ECHA InfoCard | 18f66aa3-98e1-4fe1-a9af-8f8bc7b7f981 |
| EC Number | 231-843-4 |
| Gmelin Reference | 82258 |
| KEGG | C06125 |
| MeSH | D017701 |
| PubChem CID | 16211070 |
| RTECS number | NO RTECS number found |
| UNII | RWK4822H7N |
| UN number | UN3260 |
| Properties | |
| Chemical formula | FeCl2·2C4H8O |
| Molar mass | 162.93 g/mol |
| Appearance | Light yellow-green powder |
| Odor | Odorless |
| Density | 1.6 g/cm³ |
| Solubility in water | Soluble |
| log P | -1.6 |
| Magnetic susceptibility (χ) | +2900.0e-6 cm³/mol |
| Refractive index (nD) | 1.439 |
| Dipole moment | 0 D |
| Thermochemistry | |
| Std molar entropy (S⦵298) | 387.8 J·mol⁻¹·K⁻¹ |
| Pharmacology | |
| ATC code | B03AA07 |
| Hazards | |
| GHS labelling | GHS05, GHS07 |
| Pictograms | GHS05,GHS07 |
| Signal word | Danger |
| Hazard statements | H302, H315, H319, H335 |
| Precautionary statements | P210, P273, P280, P305+P351+P338, P337+P313, P370+P378 |
| NFPA 704 (fire diamond) | 1-0-0 |
| Lethal dose or concentration | LD50 (Oral, rat): 450 mg/kg |
| LD50 (median dose) | LD50 (median dose): Oral-rat LD50: 450 mg/kg |
| NIOSH | NL |
| PEL (Permissible) | PEL (Permissible Exposure Limit) for Iron(II) Chloride: 1 mg/m³ (as Fe) |
| REL (Recommended) | Not established |
| Related compounds | |
| Related compounds |
Ligand-stabilized iron(II) chloride Iron(II) chloride Tetrahydrofuran |
