Copper

Copper

Copper is a ductile metal with very high thermal and electrical conductivity; its symbol is Cu and its atomic number is 29. Pure copper is soft and malleable; a freshly exposed surface has a reddish-orange color. It is used as a conductor of heat and electricity, a building material, and a constituent of various metal alloys. Its compounds are commonly encountered as copper(II) salts, which often impart blue or green colors to minerals, such as turquoise, and have been widely used as pigments. Copper(II) ions are water-soluble, meaning they function at low concentration as bacteriostatic substances, fungicides, and wood preservatives. In sufficient amounts, they are poisonous to higher organisms; at lower concentrations, they are an essential trace nutrient to all higher plant and animal life. In animals, copper is mainly found in the liver, muscles, and bones.

Copper forms a rich variety of compounds with oxidation states +1 and +2, which are often called cuprous and cupric, respectively. It does not react with water, but reacts slowly with atmospheric oxygen, forming a layer of brown-black copper oxide. In contrast to the oxidation of iron by wet air, this oxide layer stops further corrosion. Hydrogen sulfides and sulfides react with copper to form various copper sulfides on the surface. In the latter case, the copper corrodes, as is seen when copper is exposed to air containing sulfur compounds.

Copper(I) oxide

Copper (I) has a red color.

The simplest compounds of copper are binary compounds, i.e. those containing only two elements. The principal compounds are the oxides, sulfides, and halides. Both cuprous and cupric oxides are known. The cuprous halides with chlorine, bromine, and iodine are well known, as are the cupric halides with fluorine, chlorine, and bromine.

Copper, like all metals, forms coordination complexes with ligands. In aqueous solutions, copper(II) exists as [Cu(H2O)₆]²⁺. This complex exhibits the fastest water exchange rate (speed of water ligands attaching and detaching) of any transition-metal-aquo complex. Adding aqueous sodium hydroxide causes the precipitation of light blue solid copper (II) hydroxide. A simplified equation follows:

$Cu^{2+} + 2 OH^{-} \rightarrow Cu(OH)_{2}$

Aqueous ammonia results in the same precipitate. Upon adding excess ammonia, the precipitate dissolves, forming tetraamminecopper (II):

${ Cu({ H }_{ 2 }O) }_{ 4 }{ (OH) }_{ 2 }+4{ NH }_{ 3 }\rightarrow { [{ Cu({ H }_{ 2 }O) }_{ 2 }{ ({ NH }_{ 3 }) }_{ 4 }] }^{ 2+ }+2{ H }_{ 2 }O+2O{ H }^{ - }$

Tetramminecopper (II) sulfate

Copper (II) acquires a deep blue coloration in the presence of ammonia ligands.

Copper forms a rich variety of compounds with oxidation states +1 and +2, which are often called cuprous and cupric, respectively. It does not react with water but reacts slowly with atmospheric oxygen, forming a layer of brown-black copper oxide. In contrast to the oxidation of iron by wet air, this oxide layer stops the further, bulk corrosion. Hydrogen sulfides and sulfides react with copper to form various copper sulfides on the surface. In the latter case, the copper corrodes, as is seen when copper is exposed to air containing sulfur compounds. Oxygen-containing ammonia solutions yield water-soluble complexes with copper, as do oxygen and hydrochloric acid, which form copper chlorides, and acidified hydrogen peroxide, which form copper(II) salts. Copper(II) chloride and copper combine to form copper(I) chloride.

The simplest compounds of copper are binary compounds (i.e., those containing only two elements). The principal compounds are the oxides, sulfides, and halides. Both cuprous and cupric oxides are known. Among the numerous copper sulfides, important examples include copper(I) sulfide and copper(II) sulfide. The cuprous halides with chlorine, bromine, and iodine are well known, as are the cupric halides with fluorine, chlorine, and bromine. Attempts to prepare copper(II) iodide yield cuprous iodide and iodine.

Copper, like all metals, forms coordination complexes with ligands. In aqueous solutions, copper(II) exists as [Cu(H2O)₆]²⁺. This complex exhibits the fastest water exchange rate (speed of water ligands attaching and detaching) of any transition-metal-aquo complex. Adding aqueous sodium hydroxide causes the precipitation of light blue solid copper(II) hydroxide. A simplified equation follows:

$Cu^{2+} + 2 OH^{-} \rightarrow Cu(OH)_{2}$

Aqueous ammonia results in the same precipitate. Upon adding excess ammonia, the precipitate dissolves, forming tetraamminecopper(II):

${ Cu({ H }_{ 2 }O) }^{ 4 }{ (OH) }_{ 2 }+4{ NH }_{ 3 }\rightarrow { [{ Cu({ H }_{ 2 }O) }_{ 2 }{ ({ NH }_{ 3 }) }_{ 4 }] }^{ 2+ }+2{ H }_{ 2 }O+2O{ H }^{ - }$

Many other oxyanions form complexes: these include copper(II) acetate, copper(II) nitrate, and copper(II) carbonate. Copper(II) sulfate forms a blue crystalline pentahydrate, which is the most familiar copper compound in the laboratory. It is used in a fungicide called the Bordeaux mixture. Polyols, compounds containing more than one alcohol functional group, generally interact with cupric salts. For example, copper salts are used to test for reducing sugars. Specifically, using Benedict's reagent and Fehling's solution, the presence of the sugar is signaled by a color change from blue copper(II) to reddish copper(I) oxide. Schweizer's reagent and related complexes with ethylenediamine and other amines dissolve cellulose. Amino acids form very stable chelate complexes with copper(II). Many wet-chemical tests for copper ions exist; one, for example, involving potassium ferrocyanide, which yields a brown precipitate with copper(II) salts.