© 2026 Nanjing Finechem Holdings Co., Ltd,
All Right Reserved
Ammonium iron(II) sulfate hexahydrate, also known as Mohr’s salt, features in labs and schools for good reason. It became a staple for analytical chemistry once August Mohr introduced it in the 19th century, aiming for a stable ferrous salt that resisted oxidation better than simple iron(II) sulfate. Working in the lab, you see firsthand how much easier it gets to standardize solutions and run redox titrations with a salt that isn’t turning to rust every time you blink. Historical records point out that chemists struggled to keep iron(II) pure as oxygen in the air kept converting it, tweaking analytical results. Mohr’s design, pairing ammonium and sulfate ions around the iron, gave science a reliable path forward.
Powdered or crystalline, you spot this salt by its faint green hue—a mark of hydrated ferrous ions. Six water molecules stick in the structure, giving the compound not just its signature color but making it easier to weigh out and dissolve. It is not something you’d toss around lightly—it dissolves in water, forming a solution full of both ammonium and iron(II) ions, and picks up oxygen faster when the air is damp. The crystals stand up well in a dry atmosphere, but you leave them out in humidity and the iron slowly slips into the next oxidation state, leaving brownish stains as a warning that the magic ratio got lost. This property isn’t just chemical trivia; it determines if your standard solution stays true week after week.
You can bet folks care about the purity of their chemicals. For this hexa-hydrated salt, key tech specs touch on iron(II) content, ammonia smell, sulfate content, water of crystallization, and heavy metal impurities. Some products hit over 99% purity, meeting analytical chemistry’s high demand for sharp end points and exact concentrations. Quality labs label bottles with clear lot numbers, proper storage directions, and safety notes about iron’s toxicity and ammonium’s volatility. You don’t buy random green powder and hope for the best—you check these numbers, because when iron content drops, entire experiments go south in an afternoon.
The process feels almost old-fashioned, yet you see the logic in every step. Start by dissolving iron(II) sulfate and ammonium sulfate in water, then coax the solution at low temperatures to slow any lurking oxidation. Add a dash of sulfuric acid if you want to drive the reaction forward and suppress those sneaky ferric formations. Crystals form by cooling the mixture, and you’ll harvest them by filtration, watching for that green, glistening texture. Preparation shows you that chemistry often favors patience; rush the cooling or skip the acid, and you get brown slush instead of crisp crystals.
In the lab, ammonium iron(II) sulfate hexahydrate acts as a mild reducing agent, trading electrons in redox reactions with common oxidizers like potassium permanganate. It’s one of the best paths to calibrate those tough permanganate titrations—fewer errors, tighter numbers. Under acidic pH, the iron(II) center yields electrons and temporarily forms iron(III), until you add more acid or reduce exposure to air. Folks have gone further, modifying this salt’s structure with other ligands to form bespoke catalysts, or swapping cations for research on molecular magnetism. Each reaction proves the salt isn’t some inert green powder; it’s an active player in many chemical plays.
In catalogues and stocksheets, this compound pops up with plenty of synonyms. Mohr’s salt in teaching contexts; ammonium ferrous sulfate hexahydrate in technical protocols; double salt of ammonium sulfate and iron(II) sulfate for the pedants. Its formula, typically (NH4)2Fe(SO4)2·6H2O, follows it from school kits to bulk industrial barrels, with each synonym painting a slightly different picture yet always linking back to the stability boon that made Mohr’s version endure.
Safety isn’t a handout from a supplier; it’s lessons learned by trial and error. Despite its harmless look, ammonium iron(II) sulfate demands respect. Iron salts bring toxicity risk, especially if kids or animals get exposed, while ammonium ion can release ammonia gas under heat or in contact with strong base, irritating to lungs and eyes. Reliable laboratories don gloves, ventilate work areas, and stash samples in tightly sealed containers, away from sunlight and moisture. Training newcomers outweighs sticking up a warning label. I’ve seen careless hands leave bottles open, only to discover ruined samples and rusty stains later—an expensive lesson in lab housekeeping. Regulatory bodies, from chemical agencies to educational accreditation boards, outline safe storage, disposal, and use practices, all grounded in years of accumulated caution.
Surveying the uses gives you a sense for the material’s flexibility. In analytical chemistry, it calibrates solutions and runs iron analyses. Environmental science teams reach for it in water quality checks, where its predictable reactivity helps determine trace contaminants. Artists and historians use it, too—photographers once relied on iron(II) salts for blueprinting in cyanotype printing, and restoration experts sometimes seek it to replicate 19th-century color processes. Biomedical researchers study transport of iron in living systems using labeled forms of the salt. Industrially, its steadfast behavior helps produce dyes, treat water, and manufacture coatings where other iron salts just don’t hold up under the same conditions. Anywhere you need a clean, steady iron(II) donor, Mohr’s salt has a shot at making the short list.
Researchers don’t just use chemicals; they interrogate them. Ongoing studies track how altering water content or mixing the salt with varied acids shifts its reactivity. University labs comb through its crystal structure using X-ray diffraction, seeking clues that could help design even tougher double salts. Environmental chemists ask how iron(II) complexes interact with natural processes, like iron cycling in soils or groundwater. Spin-off research investigates how adding ligands or tweaking ionic ratios leads to new compounds for sensors, battery materials, or even targeted drug delivery—bringing iron chemistry out of the bottle and into emerging tech.
Not all risks show up on the surface. Chronic exposure studies remind us that iron(II) salts in the bloodstream lead to harmful effects, from cell damage to liver overload. Ammonium’s risk spikes if mishandled—it gives off fumes, and at high doses can disrupt cellular balance, especially in kids. Animal studies point out environmental toxicity, particularly to aquatic life if this salt seeps into waterways. Modern waste protocols focus on neutralization and dilution before disposal, with best practices calling for education on spill containment and medical intervention at the first signs of acute poisoning. The legacy of improper disposal shows up in contaminated sites, building the case for continued vigilance.
Looking ahead, new uses beckon. Green chemistry seeks iron(II) complexes that minimize byproducts and lower energy usage in process industries, and Mohr’s salt could anchor those solutions. Battery researchers test its abilities as a redox mediator to boost efficiency in next-gen storage devices. Analytical chemistry’s digital age still wants reliable calibrants for automated titrations—sometimes old favorites beat high-tech alternatives by sheer predictability. In my own experience, seeing this humble double salt anchor fresh projects reminds me that historical stability and modern ambition get along just fine when chemistry gets done right.
Ammonium iron(II) sulfate hexahydrate, sometimes called Mohr’s salt in the lab, stands out for its stability and reliability. It’s one of those chemicals that quietly supports a range of everyday science without grabbing headlines. I’ve handled it in university labs, always noticing how it provides a clear, predictable path for chemical analysis. It resists oxidation even when you leave it on the bench for a reasonable span. That trait makes it valuable in teaching, research, and industry.
Chemistry education would look very different without this compound. In titrations, it acts as a primary standard. The purity and stability mean precise measurements, which my classmates and I relied on through endless classes and lab practicals. When labs teach how to determine iron content in samples, ammonium iron(II) sulfate often stands front and center. It helps students see the clear color change between ferrous and ferric ions. Educators count on it because it’s safe to handle, doesn’t give off dangerous fumes, and keeps results consistent semester after semester.
Water treatment plants depend on accurate determination of iron. The presence of excess iron in water leads to staining and taste issues, not to mention infrastructure damage over time. Many municipal labs turn to ammonium iron(II) sulfate for calibrating instruments and standardizing tests. The ability to trust measurements impacts public health and confidence in drinking water. Any slip in monitoring, and residents immediately notice rusty taps or metallic tastes.
Textile factories and dye houses need to maintain color quality. Small amounts of iron in water can change how dyes bond to fabric. Routine checks with chemical standards iron out those issues before products reach store shelves. Each bottle of dye or roll of fabric reflects hundreds of quality control checks, and ammonium iron(II) sulfate sits in the background of that system.
Researchers often tackle questions about soil health, pollution, and minerals. Soil samples from farm fields or forests head straight into analysis. The iron levels extracted often rely on calibration against precisely weighed amounts of this compound. Any mistake in the reference, and conclusions about crop health, fertilizer needs, or environmental stress go out the window. Reliable, reproducible results allow researchers to chart the impact of environmental practices or policy changes.
No chemical comes without risk. Large spills of ammonium iron(II) sulfate can release ammonium ions, which may disrupt local waterways. While the compound itself isn’t highly toxic, improper disposal builds up salts and nutrients in soil and rivers, nudging natural balances. In classroom and industrial stocks, careful storage keeps the substance out of kids’ reach and away from accidental ingestion or mixing with incompatible chemicals, like strong oxidizers.
One thing I learned early: keep chemicals labelled, and teach proper waste handling. Schools and companies reduce risk with good training and strong safety routines. Larger plants switch to digital inventory and automated dosing to shrink accident rates. Sustaining these habits avoids emergencies and long-term environmental harm.
Every industry counting on iron measurements, color quality, or research accuracy benefits from a well-prepared bottle of ammonium iron(II) sulfate. Through careful handling and practical safety measures, its strengths support science, industry, and safety at every scale.
Ammonium iron(II) sulfate hexahydrate, usually called Mohr’s salt, pops up all the time in school labs, research centers, and even water treatment companies. It looks like greenish-blue crystals and dissolves in water, making it handy for experiments or as a source of iron in chemical reactions. The word “hazardous” gets thrown around a lot these days, so the natural question is: Does working with this stuff mean you’re putting your health on the line?
Having spent time in chemistry classrooms and research spaces, I’ve seen Mohr’s salt used by beginners and seasoned scientists alike. It’s not as notorious as some heavy metals, and nobody rushes to hide it in the back of the chemical cabinet. Safety, though, isn’t about fear. It’s about knowing where the real risks are. Even something as familiar as salt can be dangerous in the wrong dose or setting.
Ammonium iron(II) sulfate doesn’t chase you out of the lab with toxic fumes or explosive reactions. At the same time, it shouldn’t be treated like common table salt. Swallowing large amounts can cause trouble. Iron overload can damage organs, leading to symptoms like stomach pain, vomiting, and faintness. Although the hexahydrate form contains more water and dilutes the iron content, it still doesn’t belong anywhere near a child’s snack. Data from the CDC ranks iron toxicity as a real problem, especially among young children. Even modest amounts can overwhelm the body, which can’t quickly get rid of extra iron.
Dust from ammonium iron(II) sulfate isn’t friendly to the lungs, either. Working with any powder calls for careful handling. I’ve been in plenty of situations where simple mishandling led to sneezing fits and minor skin irritation—nothing life-threatening, but always a reminder to wear gloves and a mask. Its ammonium part brings another layer: if it gets hot enough, it can give off ammonia, which stings eyes and nose.
Waste from chemistry experiments ends up somewhere. Pouring solutions down the drain adds iron and sulfate ions to local water systems. Iron isn’t rare in groundwater, yet, loads of it can mess with aquatic life by changing the chemistry and blocking sunlight. Sulfates can speed up the corrosion of pipes and damage certain plant species. Sustainable lab practices mean treating this waste so water quality stays fit for drinking and for wildlife.
The main lesson? Respect the chemicals you work with. Rules in the lab—gloves on, goggles in place, and chemicals stored in clean, sealed containers—aren’t just for looks. I’ve had students roll their eyes at “another safety lecture,” only to admit later that a little extra care saved them headaches. Waste disposal guidelines aren’t up for debate, either. Following local regulations for chemical disposal protects everyone, from the janitor to the entire neighborhood.
Mohr’s salt might not come with a skull-and-crossbones label, but the same common sense applies to it as to stronger substances: handle with care, store it away from kids and pets, clean up spills right away, and toss leftovers responsibly. Safety grows from ritual, not paranoia—just simple respect for what chemistry can really do.
Anyone who’s worked in a laboratory—student or professional—learns quickly to respect chemicals like ammonium iron(II) sulfate hexahydrate. In the world of chemistry, not every compound demands a locked vault or a hazmat suit, but sloppy storage brings trouble. If this particular salt cakes up, oxidizes, or dissolves into a puddle, a whole experiment can end in frustration and error. So, figuring out good storage isn’t only about checking off a safety box; it’s about making sure work and results stay reliable.
This compound absorbs water from humid air like a sponge left on the counter. That’s a recipe for clumping crystals, unwanted reactions, and wasted product. Oxygen present in the air brings another headache: oxidation. Left exposed, iron(II) ions tend to react and jump up an oxidation state, turning the sample brown and practically useless for titrations and other precise work.
After years spent running classrooms, and more time wrangling bottle after bottle in cramped prep rooms, some habits stand the test of time:
I’ve seen someone pop open an unlabeled jar after a summer break only to find the tell-tale yellow-brown crust of hydrolyzed iron. The bottle spent a month tucked near the sink, soaking up every swell of humidity. The lesson stuck: store it tight, dry, and out of sunlight. Without these simple steps, you risk wasting time and funding on ruined chemicals—never a fun conversation with a supervisor.
Spilling ammonium iron(II) sulfate on your hand isn’t nearly as dramatic as some of the nastier acids or reactive metals, but skin irritation happens. Leaving a container near the edge of a bench invites trouble, and open bags spill everywhere. A messy bench isn’t just embarrassing; it brings health risks into play and makes cleanup a drag. Good storage habits, practiced every day, save more than time—they keep people safe.
Safety datasheets and chemical suppliers recommend storing it in a cool, dry spot with minimal light. Consulting sources like the Sigma-Aldrich MSDS always makes sense before stocking up. Labs worldwide use similar guidance because it works, not because it sounds official.
Automation has helped with some storage headaches. Humidity monitoring sensors and lockable cabinets keep sensitive salts like ammonium iron(II) sulfate in better shape than ever. Open conversations between lab members, routine checks, and training for new joins help cement these habits so they stick for the long run. Small steps count: one overlooked bottle ruins more than an experiment; it can upend a whole teaching schedule.
I remember working in the chemistry lab during my university days and always reaching for a jar of Ammonium Iron(II) Sulfate Hexahydrate. Folks usually call it Mohr’s salt and it’s one of those chemicals that feels familiar, even comforting, if you’ve handled it before. Mohr’s salt doesn’t look intimidating: the crystals show a pale green color, often catching the light with a hint of blue. You feel the cool weight in your hand. That color signals ferrous iron inside—a hint that makes it so popular for titrations where precision matters.
This salt’s recipe brings together iron, ammonium, and sulfate. Each molecule clings to six water molecules. That hydration locks ferrous iron in place and helps the salt stay stable on a lab shelf. Some other iron(II) compounds, like iron(II) sulfate, turn brown way too quickly as oxygen creeps in and oxidizes them. In contrast, Mohr’s salt doesn’t surrender to the air so easily. That means you can store it for months and still get reliable results out of it. Scientists like things they can count on, and this salt pays off in that way.
Drop Mohr’s salt in water and it dissolves cleanly, giving a clear, light green solution—a straightforward signal that it’s working as it should. No weird clumps. No mysterious floating bits. Chemically, the salt offers two ingredients: ferrous iron ions and ammonium ions, along with sulfate. The ferrous ions love to react with air and form ferric ions, but again, thanks to this salt's structure and water content, that reaction stays slow.
I’ve found this effect matters in everyday lab work. Back in my student days, standardizing potassium permanganate using Mohr’s salt felt like following a trusted recipe. The result always lined up with what theory said. That reliability helped build my confidence as a beginner learning about redox chemistry. Educators across the globe stick with this salt—not just because the books tell them to, but because students get to see predictable chemistry that makes sense, every time.
No strong smell. No fuss in daily use. The hydrated crystals rarely cause skin irritation, unlike some harsher reagents. Spills clean up with water. With a melting point above 100°C (dehydration happens first, actual salt melts higher), it holds up fine unless things go seriously wrong in the lab. That resilience helps keep labs running smoothly. Sometimes, folks forget that safety and reliability go together, especially for young learners or technicians with a busy schedule. Mohr’s salt handles rough hands and a touch of humidity better than many alternatives.
More than a teaching tool, Mohr’s salt shows up in industrial analysis of water—testing for iron and tracking chemical oxygen demand. Stability and clarity win out here. Other iron compounds can complicate readings or add unwanted color. In my experience, labs facing a tight budget or busy workflow lean on Mohr’s salt to balance speed and precision. If a chemical does its job consistently, it earns its place on the shelf for years.
Better packaging—airtight containers and clear labeling—go a long way in keeping this salt in good shape. Training new lab staff to spot contamination or dehydration nips mistakes in the bud. I’ve seen resourceful teachers dry out and refresh old salt with careful heating, stretching supply budgets just a bit further. Sharing these small tricks creates a chain of trust and know-how, something every chemistry teacher can appreciate.
For anyone who’s ever worked in a lab, ammonium iron(II) sulfate hexahydrate probably brings back memories—some good, some more along the lines of “uh-oh.” This compound, often called Mohr’s salt, crops up a lot during titrations and redox chemistry. There’s a reason why the rules for cleaning up spills get drilled into our heads: a single mistake can create real problems for people and the environment.
Picture this: a green crystal pile scatters across the floor. Everyone glances around, hoping someone else will step up. My first spill, years ago, taught me that hesitation wastes time and increases risk. The iron(II) in Mohr’s salt quickly oxidizes to iron(III), and water in the air starts the process. Slip risks jump, and the stuff can stain skin and surfaces long after you think you’ve wiped it all up.
On top of that, ammonium salts shouldn’t get anywhere near drains unless neutralized and diluted. Down the pipe, and you’ve handed your problem to a waterway. Sulfate ions might not be the most threatening, but iron contamination messes with aquatic life. Leaving even a small spill can encourage rust, rot, and long-term lab headaches.
If you accidentally spill some, don’t just wipe and hope for the best. Grab gloves and goggles, as iron salts can be irritating. A broom and dustpan work for dry crystals, followed by a damp wipe to finish. Liquid or dissolved salt needs more: contain with absorbent material, scoop it into a labeled waste container, and wash down the area with a dilute acid or buffer if directed.
Personal experience taught me not to rely on memory for waste labels. Check local chemical waste guidelines every time—regulations change, and surprises are rare if you pay attention. Chemical splash showers get used less than they should, so if it’s on your skin or in your eyes, rinse right away. Once, an impatient student tried to brush powder off a bench with bare hands. He spent the rest of the day regretting the rash.
Waste ends up in clearly marked containers, never mixed with other chemicals. This salt decomposes a bit in storage, so keep things sealed. I’ve seen teachers pour small amounts into sinks “with lots of water.” That’s not just outdated; it’s lazy. Better practice: collect it all and arrange for a professional hazardous waste pickup. The cost protects more than just your budget—it keeps you on the right side of the law.
Routine matters most. Drill safety steps into daily work, and encourage questions. A shared jar of emergency clean-up powder, clear waste bins, and regular review sessions work. In my own labs, I push students to document every spill, no matter the size. That habit reveals patterns before accidents become serious.
If every workplace paid a little more attention to small chemical spills, we’d see less stress, better health, and less harm to local ecosystems. Sometimes experience sounds like common sense. In chemistry, common sense saves time, money, and lives.
| Names | |
| Preferred IUPAC name | ammonium tetrahydroxy(oxidosulfato)ferrate(II) hexahydrate |
| Other names |
Mohr’s salt Ferrous ammonium sulfate hexahydrate |
| Pronunciation | /əˈmoʊniəm ˈaɪərən ˈsʌlfeɪt ˌhɛksəˈhaɪdreɪt/ |
| Identifiers | |
| CAS Number | 7783-85-9 |
| Beilstein Reference | 87816 |
| ChEBI | CHEBI:63051 |
| ChEMBL | CHEMBL1201562 |
| ChemSpider | 14023 |
| DrugBank | DB14435 |
| ECHA InfoCard | 03b68b57-c89c-4842-81f6-f480e2a3e7ea |
| EC Number | 231-753-5 |
| Gmelin Reference | 13122 |
| KEGG | C01839 |
| MeSH | D019285 |
| PubChem CID | 24857378 |
| RTECS number | WS4250000 |
| UNII | 7Y6B1A9E4F |
| UN number | UN3260 |
| CompTox Dashboard (EPA) | DTXSID0025739 |
| Properties | |
| Chemical formula | (NH4)2Fe(SO4)2·6H2O |
| Molar mass | 392.14 g/mol |
| Appearance | Light green crystals |
| Odor | Odorless |
| Density | 1.86 g/cm³ |
| Solubility in water | Soluble in water |
| log P | -4.5 |
| Acidity (pKa) | ~9.25 |
| Basicity (pKb) | 5.94 |
| Magnetic susceptibility (χ) | +1.6E-4 |
| Dipole moment | 0 D |
| Thermochemistry | |
| Std molar entropy (S⦵298) | 286 J·mol⁻¹·K⁻¹ |
| Std enthalpy of formation (ΔfH⦵298) | -2290 kJ·mol⁻¹ |
| Hazards | |
| Main hazards | Harmful if swallowed. Causes serious eye irritation. May cause respiratory irritation. |
| GHS labelling | GHS05, GHS07 |
| Pictograms | GHS07,GHS09 |
| Signal word | Warning |
| Hazard statements | Hazard statements: "H315, H319, H335 |
| Precautionary statements | P264, P270, P273, P280, P301+P312, P305+P351+P338, P501 |
| Lethal dose or concentration | LD50 (oral, rat): 3250 mg/kg |
| LD50 (median dose) | 3250 mg/kg (Rat, oral) |
| NIOSH | LW8225000 |
| PEL (Permissible) | Not established |
| REL (Recommended) | Not established |
| IDLH (Immediate danger) | Not listed |
| Related compounds | |
| Related compounds |
Ammonium Iron(III) Sulfate Ferrous Sulfate Ammonium Sulfate Potassium Iron(II) Sulfate Iron(II) Sulfate Heptahydrate |
