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Copper nitrate hydrate sits in a special corner of chemical history. Alchemists in old Europe stumbled across blue-green copper compounds as far back as the Middle Ages. Early scientists, out of both curiosity and necessity, experimented by dropping copper scraps or shavings into baths of nitric acid, soon noticing the vivid blue crystals that formed. These early forays set the stage for copper nitrate hydrate to play a role in the evolving world of scientific exploration. By the 19th century, advances in industrial chemistry encouraged more systematic studies. Laboratory people uncovered more practical uses, from dyeing textiles to demonstrating chemical reactivity in school classrooms. In my own college days, professors would pull out a chunky glass bottle, shake out a few blue crystals, dissolve them in water, and let students watch a classic double displacement reaction unfurl. These early experiences in the lab do more than feed curiosity—they anchor copper nitrate hydrate as a bridge between historical science and modern research.
Today, copper nitrate hydrate takes the form of gleaming blue crystals or bluish liquors that dissolve easily in water. It finds its way into laboratories, industrial research facilities, and sometimes public infrastructure projects. There’s a sense of reliability and familiarity to copper nitrate; its color announces itself right away, signaling to even a casual eye the presence of copper. This compound is not some obscure laboratory curiosity. Its use pops up across industries as different as ceramics production, agriculture, metallurgy, and analytical chemistry. Chemists look for its reactivity, pigment manufacturers value its intense color, and researchers depend on the straightforward way it reacts with other substances. From a practical perspective, anyone working with copper nitrate hydrate gets used to taking precautions—respect for its oxidizing nature stays at the forefront of every interaction, a point driven home in academic and workplace safety trainings.
The crystal form of copper nitrate hydrate catches the light, sometimes appearing nearly neon blue when sunlight hits. Its molecular structure is more than a simple curiosity—the hydrate form locks in water molecules, which affects everything from how easily it dissolves to the type of reactions it participates in. Handling the material, I’ve noticed it absorbs moisture from the air, growing sticky on humid days. Copper nitrate hydrate is a strong oxidizer, which means it reacts quickly with reducing agents or combustible materials, a property both useful and hazardous. Its solubility ranks high, and it blends readily with water to tap into a host of applications. Unlike some more stubborn copper salts, copper nitrate hydrate can react strongly with ammonia, bases, and several anionic species to give rise to surprising shades or secondary compounds.
On every package or storage container, you’ll find clear details about concentration, hydration level, and hazards. Commercial suppliers keep an eye on both copper content and nitrate percentage to meet customer demands and regulatory requirements. Labels consistently highlight its oxidizing potential and health hazards, echoing safety standards set by national and international organizations. Having seen how a small lapse—a ripped label or faded warning—can lead to lab mishaps, I appreciate robust color-coded hazardous material symbols. Material safety data sheets follow the shipped products everywhere; every laboratory I’ve worked in treats them not only as a bureaucratic hurdle but as necessary reading. Researchers learn early that close attention to labeling brings a sense of safety and accountability to chemical handling.
Manufacturers lean on time-tested synthesis methods, mainly reacting copper metal or copper oxide with nitric acid. This direct route produces a vivid blue solution, from which hydrated crystals can be coaxed by careful evaporation. In my own experience, trying to rush the evaporation step causes the crystals to trap too little water, leading to variable results and impure product. Clean glassware and precise acid strengths matter. In more advanced settings, filtration and purification refine the final product for analytical grades. Industrial production occasionally steps things up with automated batch reactors and vacuum drying chambers meant to streamline processing and avoid unnecessary exposure. It’s this fusion of precision and scale that keeps copper nitrate hydrate available and affordable for both pedagogical and large-scale industrial use.
Few people forget their earliest run-in with copper nitrate’s chemical reactions. Its pungent smell and tendency to release toxic nitrogen oxides makes chemists wary, but also excited. Add it to sodium hydroxide or potassium carbonate and you’ll coax out turquoise or deep blue precipitates. Drop it into a beaker of iron filings and an unmistakable redox reaction springs to life. These reactive properties get harnessed in many synthetic pathways, enabling everything from the production of finely divided copper powders to the elaboration of complex metal-organic frameworks. It’s no exaggeration to say that the dynamic chemistry—especially its role as an oxidizer—gives copper nitrate hydrate a central place in both educational demonstrations and advanced research experiments. In more technical hands, it can serve as a precursor for copper(II) oxide or copper nanoparticles, opening doors in catalysis and nanomaterials science.
Copper nitrate hydrate gets referenced by many names, reflecting sources from chemistry classrooms to industrial catalogues. I’ve called it by several: cupric nitrate trihydrate, copper(II) nitrate hydrate, or simply “blue copper”. Whatever the name, the same elemental properties, color, and warnings attach. Recognizable synonyms make searching for safety protocols or research notes a bit easier, but they also reinforce just how deeply this compound has woven into different branches of science and manufacturing. Consistency in labeling counts for more than just protocol—it cuts down on dangerous mix-ups.
Working with copper nitrate hydrate leaves you aware of both its promise and its risks. It stands among the more hazardous copper salts, owing to both its corrosive nature and the toxic vapors—primarily nitrogen oxides—that develop during chemical reactions or accidental spills. Every job site I’ve encountered keeps personal protective equipment at arm’s reach, from gloves and goggles to face shields when larger volumes get involved. Fume hoods earn their keep, especially during heating or in reaction steps involving acids. It doesn’t take more than a single incident—a whiff of nitrogen dioxide, or splashed blue solution on a hand—to drive home the importance of procedural discipline. Waste disposal follows strict protocols, ensuring that run-off doesn’t compromise local waterways or sewage lines. Most chemists internalize guidelines for oxidizer storage: keep it away from combustibles, follow fire risk assessments, and never store near incompatible organics or reducing agents. In research and industry alike, a culture of safety reinforces accountability and protects both personnel and the environment.
Copper nitrate hydrate finds itself at the heart of multiple sectors. In ceramics, it delivers colors that range from pale greens to intense blues on tiles, particularly valued in artisan glazes and enamelwork. Metallurgists use it to etch metals for microstructural analysis. In agriculture, specialized applications harness it for micronutrient solutions in hydroponic systems, though its use carries caveats due to environmental sensitivity. Chemical educators rely on copper nitrate hydrate as a visually engaging teaching tool, introducing concepts in redox chemistry and coordination complexes. On the industrial side, it steps into catalyst preparation for chemical reactors, supporting reactions in everything from organic synthesis to pollution abatement. Copper nitrate hydrate also appears in the synthesis of other specialty chemicals, including copper-based catalysts for fuel cells and electrochemical processes. My own brush with copper nitrate as an oxidizer in organic syntheses made me appreciate its powerful, sometimes unpredictable nature—one small measurement error can throw an entire reaction off-track or raise unexpected hazards.
Recent years keep shining fresh light on copper nitrate hydrate’s capabilities. Researchers pursue ways to incorporate it into new materials, especially in nanotechnology and catalysis. The search for energy-efficient production processes and better pollution controls places this compound at the center of initiatives focusing on green chemistry. Collaboration between industrial partners and universities has produced novel copper-based catalysts that reduce harmful emissions or help break down persistent environmental toxins. In my own experience entering research meetings, the talk often circles back to how tweaks—whether in hydration state or reaction temperature—can drive new functionalities from this seemingly old-fashioned chemical. Rapid progress in analytical chemistry relies in part on copper nitrate’s consistent, predictable reactivity, especially in colorimetric analyses and quantitative assays. Growing interest in sustainable agricultural science prompts new investigation into micro-dosing systems that may reduce leaching and runoff risks.
Handling copper nitrate hydrate pushes home the need to understand toxicity. Prolonged exposure, whether by skin contact or inhalation of dust, carries significant health risks—nausea, liver damage, and in severe cases, long-term neurological effects. Toxicologists have documented that nitrate ions contribute to environmental problems, including eutrophication in aquatic systems. Research teams study how copper ions interact with soil and water microfauna, noting detrimental effects at even low concentrations. Measures to protect laboratory and field workers depend on this evolving body of scientific evidence. Health authorities determine safe exposure limits and waste handling guidelines by tracking data from documented incidents and published studies, responding to new findings with updated regulations. Environmental agencies scrutinize run-off and impact reports, particularly in regions where intensive use has led to detectable traces in rivers and groundwater. These ongoing research efforts don’t only guard health—they also push the industry to adopt better containment, decontamination, and treatment technologies.
Looking forward, copper nitrate hydrate’s story feels far from finished. The rise in advanced materials science guarantees ongoing demand for reliable oxidizers and well-characterized copper sources. Renewable energy research counts on copper-based catalysts, spurring fresh rounds of investment and experimentation. Environmental scientists continue to explore better ways to capture and neutralize copper compounds from effluent streams, aiming to balance copper’s industrial utility against ecological impact. As science keeps chipping away at the boundaries of old knowledge, there’s every reason to expect new uses for copper nitrate hydrate to emerge in medicine, nanofabrication, and perhaps fields still beyond current imagination. Sustained innovation, underpinned by rigorous safety and environmental guardianship, charts the best path forward for this vivid blue compound that traces its roots to the earliest days of laboratory chemistry.
Staring at a jar of copper nitrate hydrate, that striking blue color draws people in. On paper, this crystal falls under the “industrial chemicals” label, often left out of dinner-table conversation. In real life, it plays a bigger role than most people think, tucked away in labs, industrial plants, and even hobbyist kits. My own first brush was in high school chemistry, watching that vivid blue solution fizz and bubble during an experiment with metals. Curiosity followed me from there.
Teachers love using copper nitrate hydrate to help students understand reactions between metals and acids. The dramatic color change plugs right into lessons about transition metals. Growing up, friends and I tried some of those demos ourselves. The vivid reactions never failed to draw out “wow” moments from classmates. These hands-on activities foster real understanding, making chemistry seem practical rather than just abstract symbols on a page. Importing hands-on learning into science curriculums helps students picture these chemicals in action, and not just as facts to memorize.
In industry, copper nitrate hydrate steps into metal finishing work, especially for preparing catalyst materials that drive chemical reactions. Factories often rely on those catalysts for producing plastics, refining petroleum, and making medicines. A friend who works with catalysts says that even tiny changes in the way a catalyst is made can impact how efficiently a large-scale process works. Getting the chemistry right with copper nitrate hydrate matters, especially when companies chase tighter emission standards and process upgrades.
Laboratories frequently turn to copper nitrate hydrate for analytical chemistry. It acts as a source of copper ions in tests, helping chemists identify, separate, and analyze other metals. This process keeps quality checks accurate in everything from drinking water testing to metal quality assurance. My advisor once pointed out that an impure batch could throw off months of careful measurements. Consistency counts, not just trendy new methods or gadgets.
Ever seen a fireworks display with a deep blue burst? There’s a fair chance copper nitrate hydrate played a part. In pyrotechnics, chemists use it to add those brilliant blue hues to firework formulas. Getting that rich color isn’t easy; balancing the chemical mix is almost an art, and even a small slip can dull the fireworks spectacle. Artisans creating ceramic glazes or specialty pigments sometimes add copper nitrate hydrate to create subtle blues and greens as well.
The flip side to practical uses shows up in environmental and safety debates. Like many copper compounds, copper nitrate hydrate can harm aquatic life and disrupt ecosystems if companies release waste irresponsibly. Stories continue to surface about factories fined for not treating their runoff. Overexposure also leads to health problems; years back, I remember an industrial accident caused skin irritation and sickness in several workers. That event drove home how protective equipment and strict handling procedures aren’t just box-checking exercises—they really protect people and the environment.
People in the field—scientists, engineers, workers—continue searching for greener alternatives and better recycling methods. Simple changes, like tightening waste capture in factories or switching to less toxic chemicals in teaching labs, can add up. Regulations get stricter each year, and for good reason. If more companies treat copper nitrate hydrate with the respect it deserves, the benefits can outweigh the risks in real-world settings.
Walking into any high school chemistry lab or serious research room, you learn quickly that some chemicals act up in ways regular folks don’t expect. Copper nitrate hydrate looks innocent: bright blue crystals, not much different from colored rock candy. But, facts tell another story. It absorbs water from the air, stains everything green-blue, and can release dangerous fumes if things heat up or get messy.
This compound stands out for two solid reasons. First, it can damage skin and eyes pretty quickly, and second, it reacts strongly with organic stuff. No one wants an unexpected mini-fire. Scars from simple slip-ups or ruined equipment serve as lasting reminders. Taking storage lightly? The record books and safety reports show why that’s never a good idea.
From years spent keeping chemicals in check, labeling containers isn’t just paperwork. Clear labels, with hazard info, and using containers that shut tight, stop most problems cold. For copper nitrate hydrate, use a bottle with a strong cap, preferably plastic or glass. Metal lids don’t mix well if any moisture shows up—they corrode and leave a mess. This compound loves moisture, soaking it from the air. If the cap fits poorly, it’ll cake or liquefy over time. That’s money lost and safety out the window.
Keeping the bottle away from sunlight and afternoons in hot rooms matters. Store it cool, in a dry cabinet, with a secondary container or spill tray underneath. I’ve seen folks trust old wooden shelves or leave bottles close to sinks and steam. Several ruined labels and mysterious blue puddles later, even the most stubborn chemist learns to move it into ventilated cabinets, away from acids, organics, and anything flammable.
Regulations for labs exist for solid reasons. Putting copper nitrate hydrate on top of a random shelf near food or casual break areas isn’t just foolish, it flouts workplace safety laws. This chemical reacts with reducing agents, so forgetting where you put a bottle or stacking incompatible reagents next to each other sets up trouble. Sulfurs, powdered metals, solvents—keep these far from hydrated copper nitrate salts.
Anyone working with this stuff deserves real training, even before they pour a single gram. Training isn’t a paperwork box to tick. Knowing emergency plans, location of safety showers, eyewash stations, and understanding chemical spill kits lowers the risk for everyone who steps inside that room. I remember the chaos during a small chemical spill; quick action and knowing which neutralizer to use saved us from bigger headaches.
Many accidents come from rushing or skipping steps at the end of the day. Good storage habits include inspecting containers for leaks or crusted-up caps, checking for broken seals, and running regular inventory. If something turns runny, grows crystals outside the bottle, or smells sharp, it’s a sign to call in a supervisor or chemical safety rep. Don’t try cleaning up unknown spills without backup.
Waste copper nitrate hydrate needs a dedicated disposal program. Tossing it in the regular trash does nobody any favors. Use marked waste containers, follow hazardous waste pickup routines, and keep records for audits. Labs that treat waste seriously usually dodge fines and show visitors they care about health and the environment.
Storage habits say a lot about a lab’s culture and values. People look out for each other by double-checking lids, keeping chemicals sorted, and reminding new staff to respect every bottle. Mistakes can happen, but routine, honest conversations about risk and responsibility help everyone head home safe and sound. Copper nitrate hydrate only becomes a problem when people forget it’s not just another pretty crystal in the cabinet.
Copper nitrate hydrate has a striking blue color that’s hard to miss in the lab. It’s tempting to see chemicals as just another toolkit item, but this one comes with more risk than most. It will eat right through skin if spilled and throws off toxic fumes. Breathing in dust or fumes lands you with respiratory trouble. Ingestion brings on stomach pain and could hurt your kidneys. Let’s be clear: taking shortcuts with this compound has real consequences.
Protective gear shrinks those risks. Lab coats aren’t just for looks—they break the direct path between chemical and skin. Nitrile gloves block direct contact, since latex can dissolve when it touches copper nitrate solutions. Safety glasses catch accidental splashes. Any experienced lab tech will agree that eye damage from chemicals is no joke, and even tiny splatters can leave permanent damage. In my years around busy labs, I've seen more than one colleague regret skipping goggles, especially during cleanup.
Respirators matter when dust or vapors could fly. A cheap paper mask does nothing here; a cartridge respirator rated for acid gases keeps you in the clear. If you’re not sure about the right cartridge, check OSHA guidelines—it’s a detail that saves lungs.
Fume hoods aren’t optional if you’re dissolving this salt, heating it, or working with bigger amounts. Fans by an open window just don’t compare. In places without hoods, limit handling to very small quantities and keep the area clear of bystanders. Good air flow protects the whole building.
Keep copper nitrate hydrate stored in tightly sealed containers far from heat sources and direct sunlight. Anything wet or organic nearby ramps up the chance of an accidental reaction, especially since copper compounds can drive other chemicals to burn. Storage shelves need solid labels, since mixing up bottles leads to unsafe combinations. I remember one time in a teaching lab where mislabeled salicylic acid got mistaken for copper nitrate—the mess was a tough lesson in why accurate labeling matters.
Disposal should follow strict local hazardous waste rules. Dumping down the drain contaminates water and threatens wildlife. Many cities run collection days for chemical waste. If you’re working in a larger facility, coordinate with environmental health staff. On my last campus job, missing a disposal deadline filled our storage room with more aging chemicals, making accidents more likely. Relabeling and routine pickups kept things safe.
Spills call for a clear roadmap. Don’t tackle them alone. Granular neutralizer absorbs the solution, but sweeping dry powder risks airborne dust exposure, so avoid it. Eye and skin exposure needs an immediate flush with running water—for at least 15 minutes. Seek a medical check even if it stings less after rinsing. Colleagues must know where emergency showers and eyewash stations are, not just in theory. Fire training helps, too, since copper salts add fuel to a blaze; water or dry chemical extinguishers handle these best.
Chemicals like copper nitrate hydrate demand respect grounded in facts, not fear. Sharing incident stories opens eyes more than warning signs do. In every lab I’ve worked with, real safety comes from repetition: gear checks, clear steps, and making sure no one is alone during risky handling. Safe habits come from seeing what happens if you lose focus, and doing it right keeps everyone working another day.
Chemistry classes love this question because it brings together observation, pattern recognition, and a dash of patience. The formula for copper(II) nitrate hydrate tells a story about both copper and water. The main component, copper(II) nitrate, carries the formula Cu(NO3)2. Toss hydration into the mix—since this salt loves hanging onto water molecules—and the result is typically copper(II) nitrate trihydrate: Cu(NO3)2·3H2O.
I spent time in university labs, learning quickly that those little blue crystals never stay dry. That’s because copper nitrate pulls water from the air, clinging to moisture by forming these “hydrates.” Identifying the hydration number is not just nitpicking; it points to how the chemical acts, how it stores, and how it can react in real life.
Hydrates hold water inside their crystal structure. If you left copper(II) nitrate hydrate on a bench overnight during humid weather, the substance would stay sticky and blue thanks to those water molecules locked in. Heating it in a flask would drive that water out, leaving you with anhydrous copper(II) nitrate. This hydrated form isn’t just a footnote for students cramming for a quiz—industry chemists and teachers alike rely on its predictable properties, especially when precision means safety.
The formula’s real value appears in practice. If you’re mixing solutions in a chemistry lab or measuring out chemicals for an experiment, you want accuracy at every step. Hydrated and anhydrous copper nitrate don’t weigh the same. If you think you’re using pure Cu(NO3)2 but use the hydrate instead, your calculations collapse quickly. Experiments fizzle or, worse, result in unpredictable reactions. During my school days, it was easy to overlook how a missing three waters could skew an entire lab result. That oversight magnifies in real industry work—think of water treatment or copper plating, where tight margins matter.
Copper(II) nitrate hydrate isn’t just a pretty sky-blue crystal. This chemical can irritate skin and eyes and causes problems when not handled with respect. It also serves as a source of nitrate ions, which contribute to environmental problems like eutrophication—essentially fertilizer for algae, choking ponds and streams. Knowing the exact composition of what you’re handling keeps exposure within safe limits. In my own lab, it paid to check and double-check the label before uncapping anything blue and crystalline.
Better labeling on bottles helps avoid confusion. Teachers should spend more time drumming in the habit of checking hydrate composition. Digital inventory systems can connect barcodes directly to hazard databases and recommended storage tips. For companies managing chemicals at scale, investing in automated monitoring for humidity and compound freshness means safer, more effective labs. Memory fades, but detailed safety data keeps everyone alert. And if you’re ever stuck with a mystery bottle of “copper nitrate,” clear records and some quick math with the hydrate formula save a lot of grief down the line.
Copper nitrate hydrate shows up in school labs, industrial plants, and even a few specialty art shops. Honestly, most people walk right past it in catalogs and never learn its real story. But there’s a reason teachers run those “wear your goggles” safety lectures: copper nitrate hydrate doesn’t mess around. The material can stain skin blue-green and irritate eyes within seconds. You smell the sharp tang just by cracking the lid. Accidental inhalation rattles the nose and throat, and spills turn tabletops into temporary hazards.
The bigger danger isn’t skin stains. Scientists found that even a small amount of copper nitrate hydrate in dust or liquid form causes real trouble if swallowed or breathed in. The body reacts quickly. Metallic taste, nausea, stomach pain, dizziness—these symptoms don’t take much time to show up. Some folks experience vomiting or diarrhea. The danger runs deeper for workers who touch copper nitrate hydrate every day. Chronic exposure sometimes triggers kidney or liver trouble, and copper builds up in the body over the years if nobody pays attention. The U.S. Centers for Disease Control and Prevention lists way higher risks for kids; their small bodies take on toxic loads faster, leading to permanent damage or even organ failure.
Copper nitrate hydrate finds ways to escape the lab or processing plant. Leaks or spills wash easily into drains, winding up in lakes and rivers. There, it dissolves and releases copper ions that fish and aquatic insects can’t tolerate. Runoff turns a healthy pond into a graveyard for sensitive life in one season if the concentration climbs. Many countries track copper runoff from factories because toxic levels derail local food webs, starting with snails and heading up the chain. In soil, copper nitrate doesn’t just linger—plants soak it up, affecting roots and blocking nutrients vital to crops or grasslands. Farmers downwind worry they’ll see lower yields or stunted growth if the water table picks up copper over time.
I remember working in a community science center: more than once, young students tried to pour “just a bit” of leftover copper solution down the drain, not realizing how quickly that small amount snowballs. It’s habits like these that justify strict Industrial stormwater permits in the U.S. and across Europe. Factories keep piles of copper nitrate locked up behind warning labels. Workers must glove up, wear face protection, and train for spills as seriously as lab accidents. Nobody wants to risk a rash or eye injury, much less a regulatory penalty for polluting a nearby stream.
The path to safety relies on diligence, education, and infrastructure. For labs and industry, solid safety stations and regular drills prevent panic when something goes wrong. Wastewater filters and spill containment gear block copper leaching into city water. Community groups and schools do best by teaching young scientists what danger symbols mean and what can go wrong with a single sloppy disposal. Some municipalities even pair up with hazardous waste collection drives, so the public can ditch leftover chemicals responsibly instead of dumping them.
Copper nitrate hydrate carries real hazards for health and environmental quality. With informed staff and strong rules, risks drop, but nobody can afford to let their guard down.
| Names | |
| Preferred IUPAC name | copper(2+) dinitrate hydrate |
| Other names |
Cupric nitrate hydrate Copper(II) nitrate hydrate Nitric acid, copper(2+) salt, hydrate |
| Pronunciation | /ˈkɒpər ˈnaɪtreɪt ˈhaɪdreɪt/ |
| Identifiers | |
| CAS Number | 10031-43-3 |
| Beilstein Reference | 358960 |
| ChEBI | CHEBI:63035 |
| ChEMBL | CHEMBL1201572 |
| ChemSpider | 22882 |
| DrugBank | DB11440 |
| ECHA InfoCard | 03b2658979-ea3d-4dc2-b554-4a2f2e88d188 |
| EC Number | 231-841-6 |
| Gmelin Reference | 51697 |
| KEGG | C18636 |
| MeSH | D003059 |
| PubChem CID | 159386 |
| RTECS number | QL5950000 |
| UNII | 27ZB2U855Z |
| UN number | UN1477 |
| Properties | |
| Chemical formula | Cu(NO3)2·xH2O |
| Molar mass | 241.60 g/mol |
| Appearance | Blue to blue-green crystalline solid |
| Odor | Odorless |
| Density | 2.32 g/cm³ |
| Solubility in water | Highly soluble |
| log P | -2.34 |
| Basicity (pKb) | 8.7 |
| Magnetic susceptibility (χ) | +2.0E-6 |
| Refractive index (nD) | 1.682 |
| Viscosity | Viscous mass |
| Dipole moment | 0 D |
| Thermochemistry | |
| Std molar entropy (S⦵298) | 254.0 J·mol⁻¹·K⁻¹ |
| Std enthalpy of formation (ΔfH⦵298) | -447.0 kJ/mol |
| Pharmacology | |
| ATC code | V07AY Copper Nitrate Hydrate |
| Hazards | |
| Main hazards | Oxidizing, harmful if swallowed, causes serious eye irritation, may cause respiratory irritation. |
| GHS labelling | GHS02, GHS07, GHS09 |
| Pictograms | GHS03, GHS07, GHS09 |
| Signal word | Warning |
| Hazard statements | H272, H302, H314, H410 |
| Precautionary statements | P264, P270, P273, P280, P301+P312, P330, P302+P352, P305+P351+P338, P310, P501 |
| NFPA 704 (fire diamond) | 2-3-2-OX |
| Lethal dose or concentration | LD50 oral rat 940 mg/kg |
| LD50 (median dose) | LD50 (oral, rat): 940 mg/kg |
| NIOSH | WL3675000 |
| PEL (Permissible) | 1 mg/m3 |
| REL (Recommended) | 3×10⁻⁵ |
| IDLH (Immediate danger) | IDLH: 100 mg/m3 |
| Related compounds | |
| Related compounds |
Copper(II) nitrate Copper(II) sulfate Copper(II) chloride Copper(II) acetate Copper(I) nitrate |
