The Chemistry of Copper

Physical Properties

Atomic Number: 29
Atomic Weight: 63.546
Electron Configuration: 1s22s22p63s23p64s13d10
Crystal Structure: Face-Centered Cubic
Density: 8.96 g/cm3
Melting Point: 1083 oC
Important Oxidation States: +1, +2
Standard Reduction Potential: +0.34 volts (Cu2+ + 2e-=Cu)


Properties of Elemental Copper

Copper is the next to the last member of the first-row transition metals. Ir is one of the two exceptions to the writing of electron configurations. One would expect copper to be an s2d9 ion. Instead, one lectron is "borrowed" fromthe 4s orbital to completely fill the 3d orbitals. Copper metal is therefore paramagnetic due to the unpaired electron in the 4s orbital. Copper has the lustor generally associated with metals but exhibits a characteristic red color. Copper is an excellent electrical conductor and is used to make electrical wires. Copper is commonly alloyed wit zinc to make brass and with tin to form bronze.

Copper has a positive reduction potential, meaning that it is fairly easily reduced. Since reversing an electrochemical reaction changes the sign of the potential, copper metal is not easily oxidized and is fairly unreactive. For excample, copper metal will not react with hydrochloric acid like many of the more reactive metals.However, it will react with concentrated nitric aicd (Figure 2). In this reaction the copper is oxidized to the Cu2+ ion, and the nitrate ion is reduced to nitrogen dioxide, which is the brown gas.

Because of this unreactivity, copper can be found in nature in the elemental form and it has been widely used since ancient times.Today most copper is obtained from sulfide ores. Some common copper minerals are cuprite (Cu2O, Figure 3), Chalcocite (Cu2S), and Chalcopyrite (CuFeS2, Figure 4).


The Chemistry of Copper(I)

The two most common ions formed by copper are the +1 (cuprous) and +2 (cupric).The cuprous ion is the less stable of the two oxidation states and is easily oxidized. Therefore, any work done with copper(I) compounds must be carried out in an inert atmosphere. Since electrons are lost first out of the 4s orbital, the electron configuration of the cuprous ion is 4s03d10 and the it is therefore diamagnetic.As a consequence of the d10 configuration, one would not expect the cuprous ion to be colored,and the compounds CuCl, CuBr, and CuI are indeed white in color. Copper(I) oxide is red and color and is sometimes used as a pigment.In this case, the color arises from a different mechanism. Many copper minerals contain copper in the +1 oxidation state.For example, cuprite is copper(I) oxide.


The Chemistry of Copper(II)

Removal of a second electron results in the a 4s03d9 configuration; therefore the cupric ion is paramagnetic. The cupric ion tends to exhibit square planar or distorted octahedral geometries. In the distorted octahedral geometry, four of the ligands form a square planar geometry and the remaining two bonds are elongated.

Most copper compounds are blue or green in color. One common compound is copper sulfate pentahydrate, known as blue vitriol (Figure 5). Another example is copper chloride dihydrate, which is bluish green in color (Figure 6). The greenish color is due to the coordination of the chloride ion, probably proudcing some tetrachlorocuprate(II).

When dissolved in water most copper compounds form the hexaaquocopper(II) ion is, which is pale blue in color.A wide range of colors can be produced when different ligands are added. For example, the addition of concentrated ammonia ammonia results results in the formation of [Cu(H2O)4(H2O)]2]2+, which blue-violet complex ion that is more intensely colored than the original solution (Figure 7). A similar reaction occurs with ethylenediamine. Addition of the nitrite ion results in the formation of a green complex ion. Last, addition of hydrochloric acid results in the formation of yellow complex ion.
Figure 1. Copper Metal
Figure 2. Copper metal reacting with nitric acid
Figure 3. Cuprite, Cu2O
Photo courtesy of R. Weller, Cochise College
Figure 4. Chalcopyrite, CuFeS2
Photo courtesy of R. Weller, Cochise College
Figure 5. Copper sulfate pentahydrate
Figure 6. Copper chloride dihydrate
Figure 7. [Co(H2O)6]2+ and [Co(NH3)4(H2O)2]2+