Copper (II) Sulfate 98% 5H2O 500g

Copper (II) Sulfate 98% 5H2O 500g

Copper (II) Sulfate 98% 5H2O 500g

Copper(II) Sulfate, also known as copper sulphate, is an inorganic compound with the chemical formula CuSO₄. It forms hydrates CuSO₄·nH₂O, where n can range from 1 to 7. The pentahydrate (n = 5) is the most common hydrate of copper(II) sulfate. Older names for the pentahydrate include blue vitriol, bluestone, vitriol of copper, and Roman vitriol.

The bright blue pentahydrate CuSO₄·5H₂O is the most commonly encountered form of copper(II) sulfate. 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 light grey powder.

Preparation and occurrence

Copper sulfate is produced industrially by treating copper metal with hot concentrated sulfuric acid or copper 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).

Chemical properties

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 when the final two aqua groups are lost. Complete dehydration occurs when the final 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⁻ → [CuCl₄]²⁻

Chemistry education

Copper sulfate is commonly included in teenager 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 or cupric sulphate. Copper(II) sulfate pentahydrate can easily be produced by crystallization from solution as copper(II) sulfate, which is 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.

{\displaystyle {\ce {Fe{+}CuSO4\rightarrow 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.

Uses

As a fungicide and herbicide

Copper sulfate pentahydrate is used as a fungicide. However, some fungi are capable of adapting to elevated levels of copper ions.

Bordeaux mixture, a suspension of copper(II) sulfate (CuSO₄) and calcium hydroxide (Ca(OH)₂), 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.

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 and zebra mussels from water pipes. Copper ions are highly toxic to fish, however. Most species of algae can be controlled with very low concentrations of copper sulfate.

Analytical reagent

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.

In a flame test, its copper ions emit a deep green light, a much deeper green than the flame test for barium.

Organic synthesis

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.

Rayon production

Reaction with ammonium hydroxide yields Tetraamminecopper(II) sulfate or Schweizer’s reagent which was used to dissolve cellulose in the industrial production of Rayon.

Niche uses

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.

Copper sulfate was once used to kill bromeliads, which serve as mosquito breeding sites. Copper sulfate is used as a molluscicide to treat bilharzia in tropical countries.

Art

In 2008, the artist Roger Hiorns filled an abandoned waterproofed council flat in London with 75,000 liters of copper(II) sulfate water 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.

Etching

Copper(II) sulfate is used to etch zinc or copper plates for intaglio printmaking. It is also used to etch designs into copper for jewelry, such as for Champlevé.

Dyeing

Copper(II) sulfate can be used as a mordant in vegetable dyeing. It often highlights the green tints of the specific dyes.

Electronics

An aqueous solution of copper(II) sulfate is often used as the resistive element in liquid resistors.

Medical and public health issues

Copper(II) 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 sulfates

Anhydrous copper(II) sulfate can be produced by dehydration of the normally available pentahydrate copper sulfate. In nature, it is found as the very rare mineral known as chalcocyanite. The pentahydrate also occurs in nature as chalcanthite. Three other copper sulfates comprise the remaining of these rare minerals: bonattite (trihydrate), boothite (heptahydrate), and the monohydrate compound poitevinite. There are numerous other, more complex, copper(II) sulfate minerals known, with environmentally important basic copper(II) sulfates like langite and posnjakite.

• 

  Anhydrous CuSO₄.

• 

  Space-filling model anhydrous CuSO₄.

• 

  The rare mineral boothite (CuSO₄·7H₂O)

Toxicological effects

Copper(II) sulfate is an irritant. The usual routes by which humans can receive toxic exposure to copper(II) 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(II) sulfate can cause conjunctivitis, inflammation of the eyelid lining, ulceration, and clouding of the cornea.

Upon oral exposure, copper(II) sulfate is moderately toxic. According to studies, the lowest dose of copper(II) sulfate that had a toxic effect on humans is 11 mg/kg. Because of its irritating effect on the gastrointestinal tract, vomiting is automatically initiated in case of the ingestion of copper(II) sulfate. However, if copper(II) sulfate is retained in the stomach, the symptoms can be severe. After 1–12 grams of copper(II) 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 severe cases, injury to the brain, stomach, liver, or kidneys may also occur.

Environmental toxicity

Copper(II) 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,^{[citation needed]} copper(II) sulfate exists mainly in the surface soil and tends to bind organic matter. The more acidic the soil is, the less binding occurs. Study also found that copper(II) sulfate has the potential to catalyze the methylation of halide ions in soil and water to produce methyl chloride (CH₃Cl) and methyl bromide (CH₃Br), two major atmospheric halide carriers capable of catalyzing stratospheric ozone depletion.