Analyzing the Copper Atom

Analyzing the Copper Atom

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An atom has three main particles—the proton, neutron, and electron. The first two particles are contained in the center of the atom, which is called the nucleus. Electrons, on the other hand, are outside the nucleus in levels called energy levels. Electrons in higher energy levels are located farther from the nucleus. These energy levels are orbitals—regions of space in which the electrons are most likely to lie. Electrons do not lie in orbits—definite paths that chart which way an object goes and where it is. Since the Heisenberg Uncertainty Principle states that it is impossible to know both the direction and position of an electron at the same time, plotting an orbit for an electron is impossible (“Electron Structure Discussion”).

All stable copper isotopes—atoms of the same element but with different masses--need to be considered before coming up with an average atomic mass for the element. There are two such isotopes, copper-63 and copper-65 (“Isotopes of Copper”). Copper has an atomic number of 29. This means it has 29 protons. All stable isotopes are electrically neutral. Therefore, there must be the same number of electrons as protons. In the case of copper, there are 29 electrons.

However, to account for the fact that they are isotopes, they have different numbers of neutrons. Protons have a relative mass (on the carbon-12 scale) of about one, and electrons 1/1836 (almost no mass). Neutrons, with a mass also of about one, account for the difference in masses of different isotopes (“Electron Structure Discussion”). Therefore, copper-63 has 34 neutrons, and copper-65 has 36 neutrons. The natural abundance of copper-63 is 69.17%, and the abundance of copper-65 is 30.83% (“Isotopes of Copper”).

As for the electrons themselves, they completely fill the first, second, and third shells. In addition, one electron enters the fourth shell. Within the subshells, they are distributed in the electron configuration 1s2 2s2 2p6 3s2 3p6 4s1 3d10. The “s,” “p,” and “d” are labels of types of subshells. S subshells can hold up to two electrons, p subshells can hold up to six, and d subshells up to ten (“Electron Structure Discussion”). So, the electron configuration indicates that the s subshell in the first, second, and third shells are completely filled. Also, the p subshells in shells two and three and the d subshell in shell three are completely filled. However, the s subshell in shell four is only half-filled.

But, only two electrons can fit in an orbital (“Electron Structure Discussion).

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This means that s subshells have one orbital, p subshells have three orbitals, and d subshells have five. So, for copper, all the orbitals in the 1s, 2s, 2p, 3s, 3p, and 3d subshells are completely filled, and the orbital in the 4s subshell is only half-filled.

However, copper’s electron configuration is a little bit interesting. Given that nickel’s configuration is 1s2 2s2 2p6 3s2 3p6 4s2 3d8, it would make sense to simply add one more electron to the 3d subshell, giving the next element, copper, the configuration 1s2 2s2 2p6 3s2 3p6 4s2 3d9. However, that configuration is incorrect. Instead, in addition to adding the one electron to the 3d subshell, one of the 4s electrons enters the 3d subshell in order to give copper the configuration 1s2 2s2 2p6 3s2 3p6 4s1 3d10.

As for magnetic character, there are two possibilities—paramagnetic, and diamagnetic. The former means that the substance is attracted to a magnetic field. This is the result of unpaired electrons. The latter says that the substance is repelled by a magnetic field. This is the result of having no unpaired electrons. Therefore, since there is an unpaired electron in the 4s subshell, copper is a paramagnetic element (“The Chemistry of Copper”).
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