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Copper(II) hydroxide

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Copper(II) hydroxide
Copper(II) hydroxide
General
Systematic name Copper(II) hydroxide
Other names Cupric hydroxide
Molecular formula Cu(OH)2
Molar mass 97.561 g/mol
Appearance Blue or blue-green solid
CAS number 20427-59-2
Properties
Density and phase 3.37 g/cm3, solid
Solubility in water insoluble
Solubility in ethanol insoluble
Melting point n/a, decomposes into CuO
Boiling point n/a
Hazards
MSDS External MSDS
Main hazards Skin, Eye, & Respiratory Irritant
NFPA 704
NFPA 704
safety square
NFPA 704 four-colored diamondHealth 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformFlammability (red): no hazard codeInstability (yellow): no hazard codeSpecial hazards (white): no code
2
Flash point Non-flammable
R/S statement R: R36 R37 R38
S: S26
Related compounds
Other anions CuSO4, CuCl2, CuO,
Cu(NO3)2, CuCO3
Other cations NaOH, KOH, Mg(OH)2,
Ca(OH)2, Ni(OH)2, Al(OH)3
Except where noted otherwise, data are given for
materials in their standard state (at 25 °C, 100 kPa)
Infobox disclaimer and references

Copper(II) hydroxide (chemical formula Cu(OH)2) is the hydroxide salt of the metal copper, with an atomic mass of approximately 97.561 atomic mass units. The typical color of copper hydroxide is blue. Some forms of copper(II) hydroxide are sold as "stabilized" copper hydroxide, quite likely a mixture of copper(II) carbonate and hydroxide. These are often greener in color.

History

Copper(II) hydroxide has been known to man since copper smelting began around 5000 BCE although the alchemists were probably the first to manufacture it.[1] This was easily done by mixing solutions of lye and blue vitriol, both chemicals which were known in antiquity.

It was produced on an industrial scale during the 17th and 18th centuries for use in pigments such as blue verditer and Bremen green.[2] These pigments were used in ceramics and painting.[3]

Chemical Properties

Synthesis

Copper(II) hydroxide can be produced by adding a small amount of sodium hydroxide to a dilute solution of copper(II) sulphate (CuSO4 · 5H2O). Treat the experiment with great care; sodium hydroxide is very corrosive and should be handled using rubber gloves and protective goggles. The precipitate produced in this manner, however, often contains an appreciable amount of sodium hydroxide impurity and a purer product can be attained if ammonium chloride is added to the solution beforehand.

Alternatively, copper hydroxide is readily made by electrolysis of water with a copper anode (+). Scrap copper water piping is a good source of copper for this method. Use a low voltage DC supply (4.5 to 12V) capable of supplying 1A, hammer flat two short lengths of copper piping and connect them to the power supply, place the electrodes on opposite sides of a jar ensuring that they cannot touch and fill the jar with tap water. Switch on the power supply, hydrogen bubbles should immediately start rising from the cathode (-), if not add a pinch of sodium bicarbonate (NaHCO3) to the water and stir. The bicarbonate allows the water to conduct electricity better without introducing chloride ions as when using table salt. Turquoise copper hydroxide flakes should soon be visible in the water. The gas released by this electrolysis cell at the cathode (-) is hydrogen; it will readily disperse but if allowed to accumulate can form an explosive mixture.

"Copper in moist air slowly acquires a dull green coating. The green material is a 1:1 mole mixture of Cu(OH)2 and CuCO3."[4]

2Cu(s) + H2O(g) + CO2(g) ---> Cu(OH)2(s) + CuCO3(s)

This is the patina that forms on bronze and other copper alloy statues such as the Statue of Liberty.

Reactions

Moist samples of copper(II) hydroxide slowly turn black due to the formation of copper(II) oxide.[5] When it is dry, however, copper(II) hydroxide does not decompose unless it is heated to 185°C.[6]

Copper(II) hydroxide reacts with a solution of ammonia to form a beautiful dark blue solution consisting of the Cu(NH3)4+2 complex, but the hydroxide is reformed when the solution is diluted with water. Copper(II) hydroxide in ammonia solution, known as Schweizer's reagent, possesses the interesting ability to dissolve cellulose.

Since copper(II) hydroxide is mildly amphoteric, it dissolves slightly in concentrated alkali, forming Cu(OH)4-2.[7]

Use as an Organic Reagent

Copper(II) hydroxide has a rather specialized role in organic synthesis. Oftentimes, when it is utilized for this purpose, it is prepared in situ by mixing a soluble copper(II) salt and potassium hydroxide.

It is sometimes used in the synthesis of aryl amines. For example, copper(II) hydroxide catalyzes the reaction of ethylenediamine with 1-bromoanthraquinone or 1-amino-4-bromoanthraquinone to form 1-((2-aminoethyl)amino)anthraquinone or 1-amino-4-((2-aminoethyl)amino)anthraquinone, respectively.

Copper(II) hydroxide also converts acid hydrazides to carboxylic acids at room temperature. This is especially useful in synthesizing carboxylic acids with other fragile functional groups. The published yields are generally excellent as is the case with the production of benzoic acid and octanoic acid.

Natural Occurrence

Copper(II) hydroxide is found in several different copper minerals, most notably azurite, malachite, antlerite, and brochantite. Azurite (2CuCO3 • Cu(OH)2 ) and malachite (CuCO3 • Cu(OH)2) are carbonates while antlerite (CuSO4 • 2Cu(OH)2) and brochantite (CuSO4 • 3Cu(OH)2) are sulfates. Copper(II) hydroxide is rarely found as an uncombined mineral because it slowly reacts with carbon dioxide from the atmosphere to form a basic copper(II) carbonate.

Uses

Copper(II) hydroxide has been used as an alternative to the Bordeaux mixture, a fungicide and nematacide.[8] Nowadays, it is disfavored because of environmental contamination problems. Copper(II) hydroxide is also occasionally used as ceramic colorant.

Precautions

Copper(II) hydroxide is a skin, eye and respiratory irritant. Always wear safety glasses when handling copper hydroxide. In case of contact with eyes, rinse immediately with plenty of water and seek medical advice.

Notes

  1. ^ Richard Cowen, Essays on Geology, History, and People, Chapter 3: "Fire and Metals: Copper".
  2. ^ Tony Johansen, Historic Artist's Pigments. PaintMaking.com. 2006.
  3. ^ Blue verditer. Natural Pigments. 2007.
  4. ^ Masterson, W. L., & Hurley, C. N. (2004). Chemistry: Principles and Reactions, 5th Ed. Thomson Learning, Inc. (p 331)"
  5. ^ Watts, Henry (1872). A Dictionary of Chemistry and the Allied Branches of Other Sciences, Vol 2. Longmans, Green, and Co. (p 69).
  6. ^ Copper (II) hydroxide. Ceramic Materials Database. 2003.
  7. ^ Pauling, Linus (1970). General Chemistry. Dover Publications, Inc. (p 702).
  8. ^ Bordeaux Mixture. UC IPM online. 2007.

References

  1. Roscoe, H. E., & Schorlemmer, C. (1879). A Treatise on Chemistry 2nd Ed, Vol 2, Part 2. MacMillan & Co. (p 498).
  2. Paquette, Leo A. (1995). Encyclopedia of Reagents for Organic Synthesis, 8 Volume Set. Wiley. ISBN 0-4719-3623-5 .
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