The periodic table, which is arranged in accordance with the properties of the elements, can therefore be used to predict the ground state electron configurations of atoms.Īn electron configuration is written symbolically to provide three pieces of information: the principal quantum number (shell number), n a letter that designates the subshell ( s, p, d, etc.) a superscript showing the number of electrons in that particular subshell. It determines many physical and chemical properties of that atom. The specific arrangement of electrons in atomic orbitals is called the electron configuration of the atom. Thus, the orbital diagram for the ground state of carbon is Thus option 3 is lowest in energy and therefore represents the ground state of a carbon atom options 1 and 2 represent excited states.Īccording to Hund’s rule, the lowest energy configuration has the maximum number of unpaired electrons with parallel spin within a set of degenerate orbitals. Electrons having parallel spins cannot occupy the same space (the same orbital), so repulsions between them must be smaller than if they had opposite spins. All three possibilities are valid based on quantum numbers and the Pauli principle, but only one is lowest in energy. These two electrons could (1) pair in a single 2 p orbital or (2) occupy separate orbitals but with opposite spin or (3) occupy separate orbitals but with parallel spin. Note that the three 2 p orbitals are degenerate, so the electron can occupy any one of them.Ī carbon atom has six electrons, so there are two electrons in the 2 p subshell. Because m s can only have two values, +½ or -½, no more than two electrons can occupy the same orbital.īy applying the Pauli exclusion principle, the arrangement of electrons in any multi-electron atom can be determined by recognizing that the ground state of an atom has all its electrons in orbitals with the lowest energies possible.Īs illustrated by the boron example, an orbital diagram includes all orbitals in all subshells within a partially occupied shell, even if some orbitals are unoccupied.If two electrons share the same orbital (have the same n, ℓ, and m ℓ), then their spin quantum numbers m s must have different values we say the two electrons have opposite spin.Each electron in an atom must have a different set of values for the four quantum numbers.The Austrian physicist Wolfgang Pauli formulated what is now called the Pauli exclusion principle: Orbitals within the same subshell (for example 2 p x, 2 p y, and 2 p z) all have the same energy orbitals that have the same energy are said to be degenerate. Thus s-subshell electrons have lower energy than p-subshell electrons, which are lower than d-subshell electrons, and so forth. For the same value of n (the same shell), as ℓ increases the energy also increases. Therefore orbital energy depends on both n and ℓ quantum numbers. However, in a He atom, which has two electrons, electron-electron repulsions between the electrons raise energy levels significantly compared to He +, and a He atom is not as stable as we might have expected.įor atoms with many electrons, the effect of electron-electron repulsions differs for different subshells. These repulsions affect electron energies.įor example, the energy levels in a He + ion (which, like H, has a single electron) are significantly lower than in a H atom because of the stronger Coulomb’s law attraction between the one electron and the 2+ charge of the He nucleus. However, when there are two or more electrons in an atom there are repulsive forces between the electrons as well as attractive forces between electrons and the nucleus. The ideas already developed about quantum numbers, orbitals, and sizes and shapes of electron-density distributions apply to all atoms. We will ask you to refer back to what you have written when you complete this section. Also note any aspects of electron configurations that puzzle you. In your course notebook, make a heading for electron configurations and write down what you remember about electron configurations for atoms from your previous experience.
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