As the principal quantum number, n, increases, the size of the orbital increases and the electrons spend more time farther from the nucleus. The filling order is based on observed experimental results, and has been confirmed by theoretical calculations. Thus, many students find it confusing that, for example, the 5 p orbitals fill immediately after the 4 d, and immediately before the 6 s. Such overlaps continue to occur frequently as we move up the chart.įigure 6.24 Generalized energy-level diagram for atomic orbitals in an atom with two or more electrons (not to scale).Įlectrons in successive atoms on the periodic table tend to fill low-energy orbitals first. The 3 d orbital is higher in energy than the 4 s orbital.
However, this pattern does not hold for larger atoms. The energy increases as we move up to the 2 s and then 2 p, 3 s, and 3 p orbitals, showing that the increasing n value has more influence on energy than the increasing l value for small atoms. The 1 s orbital at the bottom of the diagram is the orbital with electrons of lowest energy. Figure 6.24 depicts how these two trends in increasing energy relate. In any atom with two or more electrons, the repulsion between the electrons makes energies of subshells with different values of l differ so that the energy of the orbitals increases within a shell in the order s < p < d < f. The energy of atomic orbitals increases as the principal quantum number, n, increases. The specific arrangement of electrons in orbitals of an atom determines many of the chemical properties of that atom. This allows us to determine which orbitals are occupied by electrons in each atom. Having introduced the basics of atomic structure and quantum mechanics, we can use our understanding of quantum numbers to determine how atomic orbitals relate to one another.
Identify and explain exceptions to predicted electron configurations for atoms and ions.Derive the predicted ground-state electron configurations of atoms.By the end of this section, you will be able to:
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