Copper (\(\text{Cu}\)) is a reddish-orange transition metal widely used in electrical wiring, plumbing, and coinage due to its conductivity and malleability. Although often considered stable, copper is not inert and readily engages in various chemical reactions that alter its surface and form new compounds. The metal’s versatility in forming \(+1\) and \(+2\) oxidation states allows it to react vigorously in specific chemical environments, especially when a powerful oxidizing agent is present.
Reactions with Atmospheric Gases and Sulfur
Copper reacts slowly upon exposure to the atmosphere, leading to gradual surface changes. In dry air, copper reacts with oxygen (\(\text{O}_2\)) to form an initial layer of reddish-brown copper(I) oxide (\(\text{Cu}_2\text{O}\)), which eventually oxidizes further to black copper(II) oxide (\(\text{CuO}\)) over time. This thin oxide layer initially causes the metal to darken, protecting the underlying copper from more rapid corrosion.
The familiar green surface layer, known as patina, forms through a slow reaction involving oxygen, water (\(\text{H}_2\text{O}\)), and carbon dioxide (\(\text{CO}_2\)). In unpolluted environments, this protective coating is primarily basic copper carbonate, or malachite (\(\text{Cu}_2\text{CO}_3(\text{OH})_2\)). In urban or industrial areas, the patina often incorporates sulfur compounds, resulting in basic copper sulfates like brochantite (\(\text{Cu}_4\text{SO}_4(\text{OH})_6\)) as the main component. This dense, adherent layer allows copper structures, such as roofing, to last for centuries without significant structural degradation.
Copper also reacts readily with sulfur-containing gases, such as hydrogen sulfide (\(\text{H}_2\text{S}\)). This reaction forms black tarnish, a layer of copper sulfide (\(\text{CuS}\) or \(\text{Cu}_2\text{S}\)), which is frequently observed on copper jewelry or electrical components exposed to industrial air. The sulfide layer is less protective than the oxide or carbonate patina, and its formation can be a significant issue in environments where sulfur-containing compounds are abundant.
Dissolution in Acidic Environments
Copper’s position below hydrogen in the metal activity series dictates that it will not react with non-oxidizing acids, such as dilute hydrochloric acid (\(\text{HCl}\)) or dilute sulfuric acid (\(\text{H}_2\text{SO}_4\)). These acids rely on the reduction of hydrogen ions to hydrogen gas to dissolve the metal. However, the metal readily dissolves in acids that possess strong oxidizing properties.
Nitric acid (\(\text{HNO}_3\)) is a potent oxidizing acid that dissolves copper. When concentrated nitric acid is used, the reaction releases the reddish-brown, toxic nitrogen dioxide gas (\(\text{NO}_2\)), while the copper is oxidized to soluble copper(II) nitrate (\(\text{Cu}(\text{NO}_3)_2\)). Conversely, dilute nitric acid produces the colorless gas nitric oxide (\(\text{NO}\)), which rapidly reacts with air to form \(\text{NO}_2\), and the resulting solution turns a characteristic blue color due to the presence of the copper(II) ions.
Concentrated sulfuric acid will also dissolve copper, but only when heated, as the heat increases the acid’s oxidizing strength. In this high-temperature reaction, the sulfuric acid is reduced, producing sulfur dioxide gas (\(\text{SO}_2\)), alongside the soluble copper(II) sulfate (\(\text{CuSO}_4\)). Even non-oxidizing acids can be made to dissolve copper if an external oxidizing agent is introduced, such as hydrogen peroxide (\(\text{H}_2\text{O}_2\)) or dissolved oxygen. This principle is utilized in the etching of printed circuit boards, where hydrochloric acid is combined with hydrogen peroxide to quickly oxidize the copper metal into soluble copper(II) chloride (\(\text{CuCl}_2\)).
Formation of Coordination Complexes
Beyond its reactions as a solid metal, the copper(II) ion (\(\text{Cu}^{2+}\)) exhibits a strong tendency to act as a Lewis acid, accepting electron pairs from molecules or ions called ligands to form coordination complexes. This interaction dramatically changes the properties of the copper ion in solution.
A particularly striking example involves the reaction of \(\text{Cu}^{2+}\) ions with excess aqueous ammonia (\(\text{NH}_3\)). The pale blue color of the hydrated copper ion transforms instantly into the deep, intense ultramarine blue of the tetraamminecopper(II) ion, \([\text{Cu}(\text{NH}_3)_4]^{2+}\). This complexation is so stable that it is often used to dissolve copper compounds, such as copper(II) hydroxide, that are otherwise insoluble in water.
Copper(II) ions also form complexes with halide ions, such as chloride (\(\text{Cl}^-\)). For instance, adding concentrated hydrochloric acid to a blue copper solution will displace the water ligands, leading to the formation of the tetrachlorocuprate(II) ion, \([\text{CuCl}_4]^{2-}\), which gives the solution a yellow or olive-green color. In contrast, very strong ligands like cyanide ions (\(\text{CN}^-\)) typically reduce the copper(II) ion to the more stable copper(I) state (\(\text{Cu}^+\)) before forming highly stable, colorless complexes, such as the tricyanocuprate(I) ion, \(\text{Cu}(\text{CN})_3^{2-}\).