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My interest in organometallic chemistry first led me to manganese carbonyl, a compound that didn't make much noise outside research circles when it first appeared in the early 20th century. Manganese itself, discovered in the 18th century, bounced around iron smelters and steel foundries before chemists started looking at its complex with carbon monoxide. Manganese(0) carbonyl — Mn2(CO)10, known to some as dimanganese decacarbonyl — grabbed attention in the 1930s as folks chased after the behavior of metals bonded directly to carbon monoxide. The isolation and structural study of this bright orange crystalline solid gave real insight into how transition metals can bond in neat, symmetric arrangements. Early in its career, this molecule turned heads because it didn’t quite fit the usual rules; its bonding pushed against the boundaries of accepted electronic structures, nudging chemists to expand their understanding of molecules that don’t follow conventional patterns.
Walking into a lab supply store or flipping through a catalog, one might run into a bottle labeled Mn2(CO)10. The substance looks unassuming at first — bright orange, landing somewhere between pumpkin and carrot juice in color. It’s labeled as moisture sensitive and volatile. For researchers, especially organometallic chemists, this complex gives a springboard into manganese chemistry. In practice, folks use this compound not because it’s stable or easy to handle, but because its reactivity unlocks a wide range of possibilities. It serves as a starting point for crafting manganese-based catalysts or for investigating carbonyl cluster chemistry. For people like me who have tried to dissolve Mn2(CO)10 in various organic solvents, it proves soluble in benzene, toluene, or ether, but sitting in water turns it into a smelly mess pretty quickly.
Mn2(CO)10 doesn’t hide its character: it’s a low-melting solid, fuming at moderate temperatures and flashing a warning to gloved hands with its volatility and the whiff of carbon monoxide. Under the microscope, its structure fascinates even the most seasoned researcher. Two manganese atoms form a direct metal-to-metal bond, each surrounded by five CO ligands in an almost perfect symmetry. This arrangement isn’t a mere curiosity — it teaches researchers about the electronic environment in metal clusters and reveals how transition metals bend the rules of electron counting. The molecule demonstrates magnetic and electronic properties that shift as soon as it reacts, which researchers take advantage of in multiple fields. With careful handling and the right dry conditions, the substance retains its integrity long enough to take into practical chemistry.
Anyone opening a bottle of Mn2(CO)10 for the first time notices the strict labeling — warnings about air and moisture, not just because of degradation, but because of the release of toxic gases. The packaging usually features glass ampules or tightly sealed vials, reflecting serious attention to containment. Instructions rarely mince words: work in a fume hood, use gloves and eye shields, and avoid heat that might build up inside reaction mixtures. Without this caution and respect for the material, reactions can go sideways fast, making safety information and risk statements more than just box-ticking requirements. Each batch comes with information on purity, synthesis date, and typically some guidance on storage away from bright light or sources of heat.
Mn2(CO)10 doesn’t fall out of simple reagent mixing — it takes careful chemistry. Over the decades, researchers developed a few reliable routes. The method most often used in teaching and research labs employs the reduction of manganese halides with carbon monoxide under pressure. Early protocols involved manganese chloride or bromide, with sodium amalgam or aluminum as a reducing agent, and carbon monoxide coursing through the solution at pressures above atmospheric. More recent advances cut down the hazards by refining the reductions or using alternative reductants, but the challenge remains: coaxing manganese into the zero oxidation state and persuading it to grab onto ten carbonyl groups without falling apart, disproportionating, or oxidizing to something much safer and less useful.
Mn2(CO)10 loves to react. Hit it with light, and it starts cleaving CO ligands, giving a pathway into a realm of substituted carbonyls — phosphines, alkynes, and even halides can replace those tightly bound COs. Treat it with oxidants or acids, and the manganese core surrenders, sometimes forming manganese(II) or higher oxidation state products. The direct Mn-Mn bond and the symmetry of the complex make it a unique target for studying cluster fragmentation and cluster building. For those invested in making new catalysts, these reactions let you tweak the electronic structure around the metal and, with luck, tune the catalyst’s selectivity or strength. Other chemists swap out ligands to probe the boundaries of coordination chemistry, pushing manganese further into territory once reserved for more celebrated metals like iron or ruthenium.
Walking through chemistry departments, I've heard Mn2(CO)10 called by several names. The shorthand, dimanganese decacarbonyl or manganese carbonyl, gets tossed around. Some refer to it as manganese –0– carbonyl, contrasting it from higher oxidation state manganese complexes. Across literature, the CAS number and systematic IUPAC name appear, but most stick to Mn2(CO)10 — concise, recognizable, and easy to say in a pinch. Keeping those names straight matters, especially in research collaborations, where a slip can send someone running after the wrong compound.
Every scientist I know who’s worked with Mn2(CO)10 has a story about gloves, fume hoods, and near misses. This is a compound that demands respect, and ignoring the risks can go badly. The main threat, in my experience, isn’t so much the solid itself, but the carbon monoxide that escapes easily from decomposition or careless handling. Carbon monoxide is colorless, odorless, and dangerous—more than one lab has suffered near disasters from leaks or spills. Protective gear counts for a lot, but vigilance and routine working under well-functioning fume hoods keep everyone safe. Spills call for evacuation, ventilation, and careful cleanup. At the waste stage, specialized disposal protocols kick in to keep the environment from suffering unintended consequences, with no shortcuts accepted. Even with all the regulation, the real safeguard remains knowledge and habit built from the first semester of chemistry through to everyday research work.
Mn2(CO)10 isn’t a shelf-filler; it's a workhorse in laboratories studying catalysis, synthesis, and reaction mechanisms. Its reactions give rise to single-site manganese catalysts or manganese clusters that find their way into organic transformations, hydrogenation reactions, and even advanced materials research. Some teams branch out, pushing manganese chemistry into organometallic medicine research or electrochemical studies, where its reactivity helps uncover the fine points of redox cycles. For teaching, this molecule provides an accessible example for students wrestling with metal cluster chemistry, electron counting, and molecular symmetry. Most of its magic plays out away from the spotlight, in small-scale fundamental research that lays the groundwork for better, more efficient industrial processes further down the line.
In recent years, manganese carbonyl chemistry has become a proving ground for those working on transition metal catalysis. While noble metals like palladium or platinum dominate commercial processes, researchers have been hunting for earth-abundant alternatives that cut down on cost and toxicity. Mn2(CO)10 offers a model system to fine-tune redox catalysis and to design new types of hydrogenation or carbonylation reactions. I’ve watched graduate students wrestle with ligand design, hoping to unlock greater activity or selectivity by tweaking just a single attachment on the manganese center. Teams across the globe publish on related compounds, part of the broader drive to demystify first-row transition metal chemistry and contribute to green chemical practices that move past established, resource-hungry metals.
Working with manganese carbonyl, safety doesn't start and stop at the fume hood. The compound carries studied risks, both for acute exposure and longer-term effects. An inhalation of carbon monoxide can cause serious health problems or death, and the persistent handling of manganese salts can impact the nervous system, something well-documented in occupational medicine. Toxicity studies highlight the importance of containment and careful waste management, ensuring researchers don’t inadvertently introduce manganese or carbon monoxide to shared spaces, water supplies, or the air outside laboratory walls. Environmental toxicology also points out the tricky balance: manganese is essential in trace amounts for living systems but turns toxic quickly at elevated concentrations. This duality guides strict laboratory protocols and waste treatment steps, keeping the wider ecosystem out of harm’s reach.
Looking ahead, I see manganese(0) carbonyl holding a firm spot in research for years to come. As pressure mounts to find new catalytic materials that break free of rare, expensive metals, Mn2(CO)10 offers real potential as a starting point. Green chemistry goals demand catalysts that aren’t just effective but sustainable, which puts manganese in the spotlight as a resource accessible worldwide. New research branches out into different ligand environments, cooperative cluster catalysis, and hybrid materials, trying to pull even more useful reactions from this complex. Advances in spectroscopic techniques and computational modeling promise deeper insights into the dynamics and structure of these compounds, helping guide the next generation of researchers. Working with Mn2(CO)10 still carries risk and challenge, but the possibilities for innovation and discovery continue to make it an essential tool for anyone chasing the next breakthrough in sustainable chemical synthesis.
Science classes and chemistry labs often showcase beakers full of colorful reactions, but the small, dull-looking compounds sometimes push the boundaries of innovation more than anything else. Manganese(0) carbonyl—more formally known as Mn2(CO)10—rarely draws much attention outside of scientific circles, yet this orange-yellow crystalline substance plays a sneaky but powerful role in shaping modern research and industry.
Anyone who’s spent time around a research chemistry lab knows that experiments rarely go quickly. Reactions can crawl along, sometimes being stubborn and unwilling to yield practical results. That’s where compounds like manganese carbonyl step in. Researchers across the globe rely on this compound as a catalyst, especially for organic synthesis. It lowers the activation energy, making reactions between stubborn molecules finally cooperate, speeding up processes that might otherwise take hours or even days.
The real excitement comes from manganese carbonyl’s talent for encouraging the coupling of alkynes and alkenes with other chemicals. Instead of using heavy metals like palladium or platinum—which are expensive and often have environmental baggage—scientists can tap into the less-toxic and cheaper manganese counterpart. This has made manganese carbonyl especially popular in academic labs working on innovative pharmaceutical compounds or greener materials.
Cleaning up pollution has never been simple. Industries and researchers have turned to manganese carbonyl as a co-catalyst for breaking down environmental toxins, particularly in the field of hydroformylation. This process takes otherwise stubborn organic molecules and transforms them into more valuable or less harmful substances. Some studies even mention its use in the early breakthroughs of CO (carbon monoxide) conversion—a pivotal step in controlling emissions from factories and power plants.
Its role doesn’t just stop at the academic bench. Companies looking for substitutes for more hazardous catalysts have experimented with manganese carbonyl systems to cut down on toxic byproducts and to reuse waste streams. Less pollution, safer workers, and more sustainable production models are all real-world outcomes that grow from what seems like a niche scientific curiosity.
Electronics and advanced materials research often depends on preparing very thin films of metals. Scientists call it “CVD”—chemical vapor deposition—and manganese carbonyl steps up here as well. Heating the compound causes it to release manganese atoms in a controlled way, allowing engineers to lay down layers just a few atoms thick. These films get used in specialty electronics, measurement sensors, and even in experimental batteries.
Getting these processes right calls for both precision and a deep understanding of how manganese carbonyl decomposes at high temperatures. Anyone who’s worked in a materials lab can recall how heating the wrong carbonyl compound could release nasty fumes or leave unwanted residues. Mastering this technique can bump projects forward and often means the difference between a working prototype and a failed experiment.
Anyone handling a compound like manganese(0) carbonyl needs solid training. The compound’s volatility and toxicity mean that exposure in unventilated spaces could cause harm. Lab workers need gloves, reliable fume hoods, and a healthy respect for the risks. Stories from long-time researchers highlight how proper safety culture keeps accidents at bay and advances good science.
Whether you’re developing new medicines, creating efficient catalysts, or finding new ways to clean up old messes, manganese(0) carbonyl gives innovators a flexible tool. It stands as one of those small, easily missed ingredients that keep the engines of progress turning, underscoring the value of deep chemical know-how and thoughtful, safe lab practices.
One glance at the formula Mn(CO)5 and the world of organometallic chemistry springs to life. Manganese(0) carbonyl sits right on the list of pivotal coordination compounds, giving scientists and industry some solid chemistry to work with. It's built from a central manganese atom sitting at oxidation state zero, surrounded by five carbon monoxide ligands. This compound isn’t as popular as iron pentacarbonyl or nickel tetracarbonyl, but that doesn’t make it any less interesting, especially when we talk about making new materials or catalyzing reactions.
Manganese(0) carbonyl draws attention because it shows how the balance between metal and ligand electronic properties matters. In research labs, creating Mn(CO)5 takes precision. Manganese powder gets treated with carbon monoxide under high pressure and controlled temperatures. Sticking to these strict conditions ensures the product forms cleanly—otherwise, contamination or decomposition runs rampant, and you lose precious experimental time. I've seen chemists spend weeks troubleshooting one step if they skip a small detail in their synthesis.
The structure itself teaches a lot about symmetry and electronic properties. You get a trigonal bipyramidal geometry, something that ends up in quizzes and discussions for chemistry undergrads. These subtle points in geometry actually matter. They influence how the molecule reacts and the way it serves as a building block for larger frameworks in catalysis or material science.
Mn(CO)5 has made its mark in homogeneous catalysis. Some old textbooks ignore this compound, focusing instead on the bigger players, but overlooking it skips a chance to appreciate how minor players shape chemical research. In pharmaceutical synthesis, using the right organometallic catalyst can mean higher yields and cleaner products. Lab scale reactions using manganese carbonyl complexes offer flexibility that traditional base metals can’t. From my experience working on a synthesis project, switching from a common catalyst to a specialized manganese carbonyl led to better selectivity. Sometimes, a compound like Mn(CO)5 can mean the difference between success and a frustrating month of failed experiments.
Another important point—safety. These carbonyls, including manganese's, don’t get handled lightly. Carbon monoxide’s toxicity means every glass flask and steel reaction chamber demands hourly monitoring. There’s a hard lesson buried there: solid scientific advancement never forgets about health and safety. Universities and companies run extra training any time carbonyl compounds get involved.
Building knowledge about compounds like manganese(0) carbonyl isn’t just of academic importance. Its study pushes forward the boundaries of green chemistry, efficient catalysis, and improvements in materials science. People often overlook the way simple formulae like Mn(CO)5 shape the world behind the scenes in product development, environmental technologies, and safer chemical processes.
For future innovation, a better understanding of these organometallics will mean less reliance on expensive metals, easier access to clean energy technologies, and smart new ways to recycle chemicals in closed systems. Sharing what we know, through open research and education, gives everyone a chance to see what’s possible from such a straightforward-looking formula.
Manganese(0) carbonyl sits among those chemicals that demand a good dose of respect in the lab. It’s tempting to treat every reagent the same, but there’s a difference between salt crystals and a volatile, potentially toxic organometallic compound. I still remember my early days in graduate school, when someone stashed a carbonyl complex in the regular flammables cabinet and we all scrambled to rectify the mistake. Moments like that sharpen one’s sense for chemical safety.
Working with metal carbonyls, including the manganese species, never feels routine. It’s not just about fire. Many carbonyls, releasing carbon monoxide gas, carry a real asphyxiation risk. Manganese(0) carbonyl falls into this club: its volatility and decomposition products require both caution and a healthy distrust of shortcuts. A cracked vial or a careless transfer can swiftly ruin a researcher’s day and threaten everyone’s health.
Acute inhalation of carbon monoxide fits nobody’s definition of a good workday. OSHA lists it as a hazardous substance. Prolonged exposure also brings neurotoxic effects tied to manganese, a reminder that the damage isn’t always immediate or obvious.
Storing manganese(0) carbonyl never looked like tossing it in any old drawer. Glass-sealed ampoules stand as the norm. Even though some may feel tempted to reuse vials or cheapen safety, a good glass ampoule with a tight seal keeps moisture, air, and clumsy hands away. Keeping the container upright, well-labeled, and segregated from acids and oxidizing agents makes a big difference. Even a minor leak could release CO or ignite a chain reaction with incompatible materials.
Nobody wants to play hide-and-seek for a missing bottle in the fridge, so clear labeling and strict inventory habits matter. Any storage area for organometallics like this deserves signage, easy access to safety equipment, and enough ventilation to curb vapor concentration. Cold, dry, and dark shelves serve the chemical well. My research group always pushed for flammables-rated storage refrigerators—no source of ignition lurking inside.
Every transfer starts with a thorough inspection—no cracks, clean workspace, and fresh gloves. Fume hood use never feels optional. Even quick weighing or a simple dilution should happen behind glass. I always felt reassured knowing the hood was certified and flowing well.
If spills pop up—and they do—spill kits with activated carbon and easy access to emergency procedures make the difference between a controlled site and a disaster. It sounds basic, but even a splash high on the inside of the hood can surprise an unprepared chemist.
Minimizing the time manganese(0) carbonyl sits uncapped also cuts down on vapor exposure and mishaps. After finishing, prompt cleanup and returning the material to storage keeps problems contained. Double-checking that lids fit snug before leaving the lab reinforces the lesson.
Training makes all the difference. Watching a newcomer fumble with a metal carbonyl brings flashbacks to my own cautious beginnings. Regular safety drills, hands-on demonstrations, and up-to-date safety data sheets help prevent disaster. Rehearsing worst-case scenarios, from accidental inhalation to chemical burns, engrains good habits. Labs that value safety keep incident rates low and build trust among workers.
Sometimes equipment upgrades seem expensive, but in the context of a life-changing accident, they pay off instantly. Continuous air monitoring and smart storage cabinets serve as silent sentinels, catching issues before they balloon into emergencies.
At the end of the day, taking care with manganese(0) carbonyl reflects a wider truth: good science and good safety go hand in hand. Valuing the well-being of each team member—and their loved ones back home—matters more than shaving time off a synthetic route.
Manganese(0) carbonyl—known in labs as Mn2(CO)10—behaves as more than just a reagent for chemistry textbooks. Its striking orange-red crystals sometimes spark curiosity on a benchtop, but anyone working with it ought to respect what lies beneath the surface. From experience, lab accidents rarely give second chances, especially with compounds that carry a history of hidden dangers.
Researchers or industrial workers often face the biggest problems from exposure through the air. Heating or mishandling manganese carbonyl leads to fumes—colorless, but far from harmless. These fumes sneak into the lungs, where the carbon monoxide ligands break free and manganese atoms slip into human tissue. Carbon monoxide steals oxygen from blood cells one molecule at a time. Shortness of breath, dizziness, or worse can follow, even at low concentrations.
Beyond carbon monoxide poisoning, manganese pieces lodge themselves throughout the body. The nervous system becomes a target. Over the years, neurologists have reported movement problems and memory loss in those stuck too close to manganese for too long—a problem called manganism. It doesn’t always take years; even short exposures, if concentrated enough, can kickstart symptoms usually seen in much older individuals.
Spills and splashes on skin or in the eyes carry their own set of problems. Manganese carbonyl doesn’t just sit on the surface. Through cuts or even intact skin, metal atoms might travel into muscle and bone, raising internal levels. Eyes sting and burn from contact, and permanent vision damage isn’t out of the question if rinsing gets delayed. Gloves and safety goggles seem basic, but those who skip them one day often regret it for much longer.
Manganese(0) carbonyl doesn’t limit its harm to individuals. Spills contaminate air-handling systems, work benches, even clothes. Anyone around might get exposed, including cleaning staff or other researchers who think the danger has passed. On a wider scale, waste disposal mistakes can trickle down—once in the soil or water, manganese spreads its reach to populations who never stepped in a lab.
Chemicals like manganese carbonyl demand more than statements on safety data sheets. Fume hoods remain the front line of protection—no one should work with this compound in an open room. Respirators, sealed containers, and gloves built for chemical resistance aren’t optional; they work because the risk is real. Training new staff about the unique dangers can’t be skipped in favor of quick lab tours.
Institutions should promote reporting of exposure incidents rather than hiding them out of embarrassment. Early intervention allows medical teams to act before symptoms worsen. Emergency eyewash stations and showers right in the workspace make sense only if everyone knows how and when to use them.
Government agencies like OSHA already have guidelines to limit airborne manganese in workplaces, but regular review of those levels responds to real-world cases and new research. More accurate sensors for carbon monoxide and manganese vapor could warn about leaks before anyone breathes a dangerous dose.
A focus on developing safer alternatives for industrial processes lowers overall dependence on such volatile chemicals. Each choice—down to double-checking seals or properly labeling a bottle—reduces the hidden costs that only show up after someone gets hurt.
Balancing the excitement of handling flashy new compounds with the reality of safety quickly teaches plenty of lessons. Manganese(0) carbonyl grabs the attention of chemists both for its vivid orange color and its capacity as a reactive organometallic compound. Old lab tales float around about unmarked vials flaring up in air, scattering flakes and seeding panic. Those stories stick because they're real. Manganese(0) carbonyl reacts fiercely to both oxygen and moisture in the air. Exposing it just for a moment triggers its breakdown, often producing toxic carbon monoxide in the process.
I’ve seen experienced chemists trip up, underestimating the degree of sensitivity of manganese(0) carbonyl. It doesn’t just lose color; the compound rapidly decomposes, giving off carbon monoxide—a gas virtually impossible to detect without proper monitoring. People sometimes joke that it “disappears” if you look away, but it’s no joke. Everyone working with it learns early the importance of gloveboxes and Schlenk lines, tools specifically made to keep the atmosphere devoid of oxygen and water.
Taking a look at the literature and Material Safety Data Sheets, manganese(0) carbonyl crumbles within seconds in humid air. Airborne moisture and oxygen cause a violent reaction, so leaving it uncapped even briefly means most of the sample vanishes into spent carbonyls, manganese oxides, and—most dangerously—releases carbon monoxide. Curious hands or forgetful moments become real health hazards here.
This isn’t just a chemical oddity. Research teams around the world want to use organometallic carbonyls in catalysis and materials science. They see potential for new energy technologies and catalysts. These ambitions run hard against the basic handling risk. You can’t scale up or commercialize compounds that react with the air faster than you can open a bottle. It’s a bottleneck for innovation in organometallics.
Anyone weighing the use of manganese(0) carbonyl needs to rethink their lab habits. Use gloveboxes filled with inert gases like argon or nitrogen—nothing else keeps this stuff intact. Make sure all glassware is dry and that transfer tools don’t carry over a hint of humidity. Carbon monoxide alarms count as essential kit in any research area. Getting trained in emergency protocols for spills keeps everyone safer, and teaching that respect for volatile chemicals shapes habits permanently. I took part in a small lab incident, and seeing firsthand the scramble to clear the air, set up fume hoods, and account for missing sample taught stronger lessons than any seminar.
The future belongs to teams who figure out new ways to stabilize unstable organometallics. Encapsulation, customized storage vessels, and better transfer techniques could help. Expect more robust sealants and smarter warning systems. Until new methods arrive, respect for air- and moisture-sensitive chemicals stays non-negotiable. Shortcuts or wishful thinking about stability invite stories nobody wants to tell.
| Names | |
| Preferred IUPAC name | carbonylmanganese |
| Other names |
Manganese carbonyl Manganese pentacarbonyl Pentacarbonylmanganese |
| Pronunciation | /ˈmæŋɡəniːz ˈzɪərəʊ ˈkɑːbənɪl/ |
| Identifiers | |
| CAS Number | 12065-55-3 |
| Beilstein Reference | 3581446 |
| ChEBI | CHEBI:30597 |
| ChEMBL | CHEMBL2106469 |
| ChemSpider | 143454 |
| DrugBank | DB11545 |
| ECHA InfoCard | 100.007.864 |
| EC Number | 235-166-0 |
| Gmelin Reference | 714 |
| KEGG | C06189 |
| MeSH | D008353 |
| PubChem CID | 6818343 |
| RTECS number | OP2275000 |
| UNII | 3M7G63B639 |
| UN number | UN2813 |
| Properties | |
| Chemical formula | Mn(CO)5 |
| Molar mass | Mn(CO)5: Mn = 54.94, C = 12.01 × 5, O = 16.00 × 5 ⇒ 54.94 + 60.05 + 80.00 = **194.99 g/mol** |
| Appearance | silver-white crystals |
| Odor | sweet odor |
| Density | 1.83 g/cm³ |
| Solubility in water | insoluble |
| log P | -1.6 |
| Vapor pressure | 0.0004 mmHg (25°C) |
| Acidity (pKa) | -5.1 |
| Basicity (pKb) | 11.06 |
| Magnetic susceptibility (χ) | Paramagnetic |
| Refractive index (nD) | 1.527 |
| Viscosity | Viscous liquid |
| Dipole moment | 0.00 D |
| Thermochemistry | |
| Std molar entropy (S⦵298) | 188.8 J·mol⁻¹·K⁻¹ |
| Std enthalpy of formation (ΔfH⦵298) | -36.6 kJ/mol |
| Std enthalpy of combustion (ΔcH⦵298) | -297 kJ mol⁻¹ |
| Pharmacology | |
| ATC code | V03AW33 |
| Hazards | |
| GHS labelling | GHS02, GHS07, GHS08 |
| Pictograms | GHS02,GHS06 |
| Signal word | Danger |
| Hazard statements | H250, H260, H301, H331, H373, H410 |
| Precautionary statements | H260, H301 + H331, H373, P210, P222, P260, P264, P270, P271, P301 + P310, P304 + P340, P308 + P311, P370 + P378, P403, P405, P501 |
| NFPA 704 (fire diamond) | 2-3-0 |
| Flash point | Flash point: -18 °C |
| Autoignition temperature | 140 °C (284 °F; 413 K) |
| Lethal dose or concentration | LD50 Oral Rat 46 mg/kg |
| LD50 (median dose) | LD50 (median dose): Oral-rat LD50: 45 mg/kg |
| NIOSH | PC83000 |
| PEL (Permissible) | 5 mg/m3 |
| REL (Recommended) | 2000 ppm |
| IDLH (Immediate danger) | IDLH: 35 mg/m³ |
| Related compounds | |
| Related compounds |
Tricarbonyl(methylcyclopentadienyl)manganese Manganese(I) chloride Manganese(III) acetate Manganese(III) oxide |
