![]() An electron in the n = 1 n = 1 state of a hydrogen atom is denoted 1 s, where the first digit indicates the shell ( n = 1 ) ( n = 1 ) and the letter indicates the subshell ( s, p, d, f … correspond to l = 0, 1, 2, 3 … ). Table 8.5 Electron States of Atoms Because of Pauli’s exclusion principle, no two electrons in an atom have the same set of four quantum numbers.Įlectrons with the same principal quantum number n are said to be in the same shell, and those that have the same value of l are said to occupy the same subshell. ![]() Consistent with Pauli’s exclusion principle, no two rows of the table have the exact same set of quantum numbers. ![]() Sample sets of quantum numbers for the electrons in an atom are given in Table 8.5. This principle is related to two properties of electrons: All electrons are identical (“when you’ve seen one electron, you’ve seen them all”) and they have half-integral spin ( s = 1 / 2 ). The structure and chemical properties of atoms are explained in part by Pauli’s exclusion principle: No two electrons in an atom can have the same values for all four quantum numbers ( n, l, m, m s ). (The spin quantum number s is the same for all electrons, so it will not be used in this section.) The electric potential U( r) for each electron does not follow the simple −1 / r −1 / r form because of interactions between electrons, but it turns out that we can still label each individual electron state by quantum numbers, ( n, l, m, s, m s ) ( n, l, m, s, m s ). The assumption is valid because the electrons are distributed randomly around the nucleus and produce an average electric field (and potential) that is spherically symmetrical. Assume that each electron moves in a spherically symmetrical electric field produced by the nucleus and all other electrons of the atom. In practice, chemists simplify the notation by using a bracketed noble gas symbol to represent the configuration of the noble gas from the preceding row because all the orbitals in a noble gas are filled.To construct the ground state of a neutral multi-electron atom, imagine starting with a nucleus of charge Ze (that is, a nucleus of atomic number Z) and then adding Z electrons one by one. When we reach neon, with Z = 10, we have filled the 2 p subshell, giving a 1 s 22 s 22 p 6 electron configuration:Īs we continue through the periodic table in this way, writing the electron configurations of larger and larger atoms, it becomes tedious to keep copying the configurations of the filled inner subshells. Similarly, fluorine has the electron configuration 1 s 22 s 22 p 5: Fluorine has one electron pair in the 1 s orbital, one electron pair in the 2 s orbital, and 2 electrons pairs with one unpaired electron in the 2 p orbital. Oxygen has one electron pair in the 1 s orbital, one electron pair in the 2 s pair, and one electron pair, and 2 unpaired electrons in the 2 p orbitals. Because all the 2 p orbitals are degenerate, it doesn’t matter which one has the pair of electrons. One electron must be paired with another in one of the 2 p orbitals, which gives us two unpaired electrons and a 1 s 22 s 22 p 4 electron configuration. The electron configuration of nitrogen is thus 1 s 22 s 22 p 3.Īt oxygen, with Z = 8 and eight electrons, we have no choice.
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