Copper(II) Nitrate
Properties
| State | Solid (blue, deliquescent crystals; usually as trihydrate) |
| Color | Blue |
| Solubility | Very soluble in water (381 g/L at 20°C); soluble in ethanol and acetone |
| Melting Point | 114.5°C (trihydrate); 255°C (anhydrous decomposes) |
| Boiling Point | Decomposes at ~170°C (trihydrate) |
About Copper(II) Nitrate
Cu(NO₃)₂ is what most chemists actually reach for when they need a clean source of Cu²⁺ in solution — sulfate carries a complexing counterion, chloride forms [CuClₙ]²⁻ⁿ species at any decent concentration, but nitrate is essentially a spectator ion in water, so you get a clean [Cu(H₂O)₆]²⁺ aquo ion and all of its bright blue d-d absorption around 800 nm. The hydrates are the form you typically weigh out: the trihydrate Cu(NO₃)₂·3H₂O melts around 114.5°C in its own water of crystallization, and the hexahydrate is what crystallizes from cold solution. Heating any of these does something visually striking — the blue crystals lose water, then around 170–200°C the salt decomposes via 2 Cu(NO₃)₂ → 2 CuO + 4 NO₂ + O₂, giving a black solid and a plume of brown NO₂. Anhydrous Cu(NO₃)₂ is much weirder: it sublimes around 150°C in vacuum (rare for an ionic-looking salt) because it's actually molecular in the gas phase, with bidentate η²-nitrate ligands chelating Cu — a structure that wasn't pinned down until the 1960s and was a small surprise at the time. In synthesis, Cu(NO₃)₂ supported on bentonite clay (Laszlo's "Claycop") gives a mild, regioselective nitrating agent for activated arenes that doesn't need fuming HNO₃, and Cu(NO₃)₂/Ac₂O does selective benzylic nitration. In pyrotechnics it gives blue-green flame color from CuO* and CuOH* emission.
Where you'll encounter it
If you've ever seen the classic "blue crystals to black powder with brown smoke" decomposition demo, the starting material was almost always Cu(NO₃)₂·3H₂O scooped onto a watch glass over a hot plate. In a research lab, the trihydrate is the go-to source of Cu²⁺ when you need to avoid sulfate or chloride coordination — for example when running EPR on Cu(II) complexes where extra ligands would mess up the spectrum, or when synthesizing copper carboxylate paddlewheel dimers and you need bare Cu²⁺.
Common Uses
- Non-coordinating Cu²⁺ source for synthesis of copper coordination complexes and MOFs
- Precursor for CuO and Cu/CuO catalyst materials via thermal decomposition under air
- Claycop reagent (Cu(NO₃)₂ on bentonite) for regioselective nitration of activated arenes
- Selective oxidant in Sakurai–Hosomi-type allylations and benzylic C–H nitrooxylation
- Blue-green colorant in pyrotechnic stars via CuOH* and CuCl* emission with chloride donor
- Mordant in textile dyeing for fixing acid dyes onto wool and silk fibers
- Patina solution component for chemically aging copper, brass, and bronze sculptures
- Source of Cu(II) for electroless plating bath formulations on plastics
Safety Information
GHS: H272 (oxidizer), H302 (harmful if swallowed), H315/H319 (skin/eye irritation), H410 (very toxic to aquatic life). DOT class 5.1 oxidizer, UN1479. The oxidizer hazard is real — do not store with paper, sawdust, organic solvents, or reducing agents; spills onto cellulose can ignite spontaneously after drying. Thermal decomposition releases NO₂, which is the actually dangerous part if you heat it in a closed space. NIOSH IDLH for NO₂ is 20 ppm. OSHA PEL for copper dusts 1 mg/m³ as Cu. Skin contact gives the characteristic blue-green stain plus low-grade nitric acid burn from acidic hydrolysis.
This safety summary is for educational reference only and may not be complete. It is not a substitute for Safety Data Sheets (SDS), medical advice, or professional chemical safety guidance. Always consult appropriate SDS and qualified professionals before handling chemicals.