Copper(II) sulfate, also known as cupric sulfate, or copper sulphate, is the inorganic compound with the chemical formula CuSO4(H2O)x, where x can range from 0 to 5. The pentahydrate (x = 5) is the most common form. Older names for this compound include blue vitriol, bluestone, vitriol of copper, and Roman vitriol.
Blue vitriol (pentahydrate)
Bonattite (trihydrate mineral)
Boothite (heptahydrate mineral)
Chalcanthite (pentahydrate mineral)
3D model (JSmol)
|E number||E519 (acidity regulators, ...)|
|RTECS number||GL8800000 (anhydrous)
|Molar mass||159.609 g/mol (anhydrous)
249.685 g/mol (pentahydrate)
|Density||3.60 g/cm3 (anhydrous)
2.286 g/cm3 (pentahydrate)
|Melting point||110 °C (230 °F; 383 K) decomposes (·5H2O)
<560 °C decomposes
|1.055 molal (10 °C)
1.26 molal (20 °C)
1.502 molal (30 °C)
insoluble in ethanol
soluble in methanol
10.4 g/L (18 °C)
insoluble in ethanol, insoluble in acetone
Refractive index (nD)
|Orthorhombic (anhydrous, chalcocyanite), space group Pnma, oP24, a = 0.839 nm, b = 0.669 nm, c = 0.483 nm.
Triclinic (pentahydrate), space group P1, aP22, a = 0.5986 nm, b = 0.6141 nm, c = 1.0736 nm, α = 77.333°, β = 82.267°, γ = 72.567°
|5 J K−1 mol−1|
Std enthalpy of
|Safety data sheet||anhydrous
|Lethal dose or concentration (LD, LC):|
LD50 (median dose)
|300 mg/kg (oral, rat)|
|US health exposure limits (NIOSH):|
|TWA 1 mg/m3 (as Cu)|
|TWA 1 mg/m3 (as Cu)|
IDLH (Immediate danger)
|TWA 100 mg/m3 (as Cu)|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
The pentahydrate (CuSO4·5H2O), the most commonly encountered salt, is bright blue. It exothermically dissolves in water to give the aquo complex [Cu(H2O)6]2+, which has octahedral molecular geometry. The structure of the solid pentahydrate reveals a polymeric structure wherein copper is again octahedral but bound to four water ligands. The Cu(II)(H2O)4 centers are interconnected by sulfate anions to form chains. Anhydrous copper sulfate is a white powder.
Preparation and occurrenceEdit
Copper sulfate is produced industrially by treating copper metal with hot concentrated sulfuric acid or its oxides with dilute sulfuric acid. For laboratory use, copper sulfate is usually purchased. Copper sulfate can also be produced by slowly leaching low grade copper ore in air; bacteria may be used to hasten the process.
Commercial copper sulfate is usually about 98% pure copper sulfate, and may contain traces of water. Anhydrous Copper sulfate is 39.81 percent copper and 60.19 percent sulfate by mass, and in its blue, hydrous form, it is 25.47% copper, 38.47% sulfate (12.82% sulfur) and 36.06% water by mass. Four types of crystal size are provided based on its usage: large crystals (10–40 mm), small crystals (2–10 mm), snow crystals (less than 2 mm), and windswept powder (less than 0.15 mm).
Copper(II) sulfate pentahydrate decomposes before melting. It loses two water molecules upon heating at 63 °C (145 °F), followed by two more at 109 °C (228 °F) and the final water molecule at 200 °C (392 °F). Dehydration proceeds by decomposition of the tetraaquacopper(2+) moiety, two opposing aqua groups are lost to give a diaquacopper(2+) moiety. The second dehydration step occurs with the final two aqua groups are lost. Complete dehydration occurs when the only unbound water molecule is lost. At 650 °C (1,202 °F), copper(II) sulfate decomposes into copper(II) oxide (CuO) and sulfur trioxide (SO3).
Copper sulfate reacts with concentrated hydrochloric acid to give tetrachlorocuprate(II):
- Cu2+ + 4 Cl− → CuCl2−
Copper sulfate is commonly included in children's chemistry sets. It is often used to grow crystals in schools and in copper plating experiments, despite its toxicity. Copper sulfate is often used to demonstrate an exothermic reaction, in which steel wool or magnesium ribbon is placed in an aqueous solution of CuSO4. It is used to demonstrate the principle of mineral hydration. The pentahydrate form, which is blue, is heated, turning the copper sulfate into the anhydrous form which is white, while the water that was present in the pentahydrate form evaporates. When water is then added to the anhydrous compound, it turns back into the pentahydrate form, regaining its blue color, and is known as blue vitriol. Copper(II) sulfate pentahydrate can easily be produced by crystallization from solution as copper(II) sulfate is quite hygroscopic.
In an illustration of a "single metal replacement reaction", iron is submerged in a solution of copper sulfate. Iron reacts producing iron(II) sulfate and copper precipitates.
- Fe + CuSO4 → FeSO4 + Cu
In high school and general chemistry education, copper sulfate is used as electrolyte for galvanic cells, usually as a cathode solution. For example, in a zinc/copper cell, copper ion in copper sulfate solution absorbs electron from zinc and forms metallic copper.
- Cu2+ + 2e− → Cu (cathode) E°cell=0.34V
As a fungicide and herbicideEdit
Bordeaux mixture, a suspension of copper(II) hydroxide Cu(OH)2 and calcium sulfate, is used to control fungus on grapes, melons, and other berries. It is produced by mixing a water solution of copper sulfate and a suspension of slaked lime, a suspension of copper(II) hydroxide Cu(OH)2 and calcium sulfate, which is used to control fungus on grapes, melons, and other berries.
Cheshunt compound, a commercial mixture of copper sulfate and ammonium carbonate (discontinued), is used in horticulture to prevent damping off in seedlings. As a non-agricultural herbicide, is it used to control invasive aquatic plants and the roots of plants situated near water pipes. It is used in swimming pools as an algicide. A dilute solution of copper sulfate is used to treat aquarium fishes for parasitic infections, and is also used to remove snails from aquariums. Copper ions are highly toxic to fish, however. Most species of algae can be controlled with very low concentrations of copper sulfate.
Several chemical tests utilize copper sulfate. It is used in Fehling's solution and Benedict's solution to test for reducing sugars, which reduce the soluble blue copper(II) sulfate to insoluble red copper(I) oxide. Copper(II) sulfate is also used in the Biuret reagent to test for proteins.
Copper sulfate is used to test blood for anemia. The blood is tested by dropping it into a solution of copper sulfate of known specific gravity – blood which contains sufficient hemoglobin sinks rapidly due to its density, whereas blood which does not sink or sinks slowly has insufficient amount of hemoglobin.
Copper sulfate is employed at a limited level in organic synthesis. The anhydrous salt is used as a dehydrating agent for forming and manipulating acetal groups. The hydrated salt can be intimately mingled with potassium permanganate to give an oxidant for the conversion of primary alcohols.
Copper(II) sulfate has attracted many niche applications over the centuries. In industry copper sulfate has multiple applications. In printing it is an additive to book binding pastes and glues to protect paper from insect bites; in building it is used as an additive to concrete to provide water resistance and disinfectant qualities. Copper sulfate can be used as a coloring ingredient in artworks, especially glasses and potteries. Copper sulfate is also used in firework manufacture as a blue coloring agent, but it is not safe to mix copper sulfate with chlorates when mixing firework powders.
It corrects copper deficiencies in the soil and animals and stimulates farm animals' growth.
In decoration, copper sulfate adds color to cement, metals and ceramic. Some batteries, electrodes and wire contain copper sulfate. It is used in printing ink and hair dye and it creates a green color in fireworks.
In 2008, the artist Roger Hiorns filled an abandoned waterproofed council flat in London with 75,000 liters of copper sulfate solution. The solution was left to crystallize for several weeks before the flat was drained, leaving crystal-covered walls, floors and ceilings. The work is titled Seizure. Since 2011, it has been on exhibition at the Yorkshire Sculpture Park.
Medical and public health issuesEdit
Copper sulfate was used in the past as an emetic. It is now considered too toxic for this use. It is still listed as an antidote in the World Health Organization's Anatomical Therapeutic Chemical Classification System.
Other form of copper sulfatesEdit
Anhydrous copper(II) sulfate is a white solid. It can be produced by dehydration of the normally available pentahydrate copper sulfate. In nature, it found as the very rare mineral known as chalcocyanite. The pentahydrate also occurs in nature as chalcanthite. Two other copper sulfates comprise the remaining of these rare minerals: bonattite (trihydrate) and boothite (heptahydrate).
Copper sulfate is an irritant. The usual routes by which humans can receive toxic exposure to copper sulfate are through eye or skin contact, as well as by inhaling powders and dusts. Skin contact may result in itching or eczema. Eye contact with copper sulfate can cause conjunctivitis, inflammation of the eyelid lining, ulceration, and clouding of the cornea.
Upon oral exposure, copper sulfate is moderately toxic. According to studies, the lowest dose of copper sulfate that had a toxic impact on humans is 11 mg/kg. Because of its irritating effect on the gastrointestinal tract, vomiting is automatically triggered in case of the ingestion of copper sulfate. However, if copper sulfate is retained in the stomach, the symptoms can be severe. After 1–12 grams of copper sulfate are swallowed, such poisoning signs may occur as a metallic taste in the mouth, burning pain in the chest, nausea, diarrhea, vomiting, headache, discontinued urination, which leads to yellowing of the skin. In cases of copper sulfate poisoning, injury to the brain, stomach, liver, or kidneys may also occur.
Copper sulfate is highly soluble in water and therefore is easy to distribute in the environment. Copper in the soil may be from industry, motor vehicle, and architectural materials. According to studies, copper sulfate exists mainly in the surface soil and tends to bind organic matter. The more acidic the soil is, the less binding occurs.
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Salts and esters of the sulfate ion