Copper

Properties

Copper is element 29, a transition metal in group 11 alongside silver and gold, the trio often called the coinage metals. Its ground-state configuration is [Ar] 3d¹⁰ 4s¹: a completely filled 3d subshell with a single electron outside it. That lone 4s electron is loosely held (the first ionization energy is 7.726 eV), and in the solid it joins a shared “sea” of electrons that moves freely through the metal. This is the main reason copper carries electricity and heat so well.

Pure copper is soft, ductile and malleable, so it can be drawn into thin wire or hammered into sheets without cracking. It melts at 1357.77 K (1084.62 °C) and boils at 2835 K (2561.85 °C), and it is fairly dense at 8.933 g/cm³. Its warm reddish color is unusual among metals: most look silvery, but copper absorbs part of the blue and green light that falls on it, so the reflected light is shifted toward red and orange.

Copper’s Pauling electronegativity is 1.9 and its electron affinity is 1.228 eV. Natural copper is a mix of two stable isotopes, copper-63 (69.15%) and copper-65 (30.85%), which averages out to an atomic mass of 63.546. You can check that weighted average yourself with the average atomic mass calculator.

Chemistry and compounds

Copper is a fairly unreactive metal. It sits below hydrogen in the reactivity series, so it does not release hydrogen gas from dilute acids such as hydrochloric acid. Oxidizing acids like nitric acid do dissolve it, giving blue solutions of copper(II) ions and brown or colorless nitrogen oxides. In moist air it slowly tarnishes, first darkening and then, over years, building up a green patina.

Its two common oxidation states are +1 and +2. Many copper(I) compounds are colorless, because a filled 3d¹⁰ shell has no d–d transitions to absorb visible light; red copper(I) oxide (Cu₂O) owes its color to a different process. Copper(II) compounds, with a 3d⁹ configuration, are usually blue or green. The best known is copper(II) sulfate pentahydrate, CuSO₄·5H₂O, which forms bright blue crystals. Heating drives off the water and leaves a white powder that turns blue again when water is added, a simple test for water in the lab.

Several classic classroom reactions involve copper. Adding ammonia to a solution of Cu²⁺ produces the deep blue complex ion [Cu(NH₃)₄]²⁺. An iron nail placed in copper sulfate solution becomes coated with copper, because iron is more reactive and displaces it. Benedict’s and Fehling’s tests for reducing sugars rely on Cu²⁺ being reduced to a brick-red precipitate of Cu₂O. Copper compounds also give a blue-green color in a flame test.

Where it comes from

A small amount of copper occurs as native metal, but most is mined from ores. The most important are sulfide minerals such as chalcopyrite (CuFeS₂) and chalcocite (Cu₂S). Oxidized ores include cuprite (Cu₂O) and the green and blue carbonates malachite and azurite, long prized as ornamental stones and pigments.

Because ores often contain only a small percentage of copper, the rock is crushed and concentrated first, typically by froth flotation. Sulfide concentrates are then smelted and converted to impure “blister” copper, which is refined by electrolysis. In that final step, impure copper anodes dissolve and pure copper plates onto the cathodes, while valuable metals such as silver and gold collect in the sludge beneath the anodes. Some oxide ores are instead leached with acid and the copper recovered directly by electrolysis. The electrolysis calculator shows how the mass deposited depends on current and time.

Copper can be recycled again and again without losing its properties, so scrap is an important source alongside newly mined metal.

Uses

Copper’s excellent electrical conductivity makes it the backbone of modern electrical infrastructure, used extensively in wiring, motors, transformers, generators and electronics, including the fine tracks on printed circuit boards. Its high thermal conductivity suits it to heat exchangers, radiators and cookware.

It’s also a key material in plumbing, valued for its resistance to corrosion. Alloyed with other metals, copper forms bronze (with tin) and brass (with zinc), two alloys that have been essential to toolmaking, hardware, musical instruments and decorative objects for thousands of years. Copper-nickel alloys are common in coins and in marine equipment.

Copper’s natural resistance to bacteria and other microbes has led to its use on frequently touched surfaces, like door handles and fixtures, in settings where reducing the spread of germs matters. Copper compounds also serve as fungicides in agriculture and as wood preservatives.

In the body and the environment

Copper is an essential trace element for humans. It is part of several enzymes, including one that performs the last step of cellular respiration and others involved in iron metabolism and connective tissue. Many mollusks and arthropods carry oxygen with hemocyanin, a copper-containing protein that makes their blood blue when oxygenated.

The body keeps copper levels under tight control. In Wilson’s disease, an inherited disorder, copper builds up in the liver and brain. In the environment, dissolved copper is toxic to many aquatic organisms even at low levels, which is why runoff from mines and antifouling paints is monitored.

History

Copper has no single discoverer. It was one of the earliest metals worked by humans, going back roughly ten thousand years, because it occasionally occurs as pure native metal that early people could shape without smelting. As techniques for extracting it from ore developed, copper use became widespread enough that archaeologists mark a distinct Copper Age preceding the Bronze Age, when copper was alloyed with tin to make a harder metal. Its name and symbol, Cu, trace back to the Latin “cuprum,” a term linked to Cyprus, a major copper-producing island in the ancient world.