Contents:
Quantum Numbers and Atomic
Orbitals
1. Principal Quantum
Number (n)
2. Angular
Momentum (Secondary, Azimunthal) Quantum Number (l)
3. Magnetic Quantum
Number (m_{l})
4. Spin Quantum Number (m_{s})
Table of Allowed Quantum Numbers
Writing Electron Configurations
Properties of Monatomic Ions
References
By solving the Schrödinger equation (Hy = Ey), we obtain a set of mathematical equations, called wave functions (y), which describe the probability of finding electrons at certain energy levels within an atom.
A wave function for an electron in an atom is called an atomic orbital; this atomic orbital describes a region of space in which there is a high probability of finding the electron. Energy changes within an atom are the result of an electron changing from a wave pattern with one energy to a wave pattern with a different energy (usually accompanied by the absorption or emission of a photon of light).
Each electron in an atom is described by four different quantum numbers. The first three (n, l, m_{l}) specify the particular orbital of interest, and the fourth (m_{s}) specifies how many electrons can occupy that orbital.
l | 0 | 1 | 2 | 3 | 4 | 5 | ... |
Letter | s | p | d | f | g | h | ... |
The subshell with n=2 and l=1 is the 2p subshell; if n=3 and l=0, it is the 3s subshell, and so on. The value of l also has a slight effect on the energy of the subshell; the energy of the subshell increases with l (s < p < d < f).
n | l | m_{l} | Number of orbitals |
Orbital Name |
Number of electrons |
1 | 0 | 0 | 1 | 1s | 2 |
2 | 0 | 0 | 1 | 2s | 2 |
1 | -1, 0, +1 | 3 | 2p | 6 | |
3 | 0 | 0 | 1 | 3s | 2 |
1 | -1, 0, +1 | 3 | 3p | 6 | |
2 | -2, -1, 0, +1, +2 | 5 | 3d | 10 | |
4 | 0 | 0 | 1 | 4s | 2 |
1 | -1, 0, +1 | 3 | 4p | 6 | |
2 | -2, -1, 0, +1, +2 | 5 | 4d | 10 | |
3 | -3, -2, -1, 0, +1, +2, +3 | 7 | 4f | 14 |
The distribution of electrons among the orbitals of an atom is called the electron configuration. The electrons are filled in according to a scheme known as the Aufbau principle ("building-up"), which corresponds (for the most part) to increasing energy of the subshells:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f
It is not necessary to memorize this listing, because the order in which the electrons are filled in can be read from the periodic table in the following fashion:
Or, to summarize:
In electron configurations, write in the orbitals that are occupied by electrons, followed by a superscript to indicate how many electrons are in the set of orbitals (e.g., H 1s^{1})
Another way to indicate the placement of electrons is an orbital diagram, in which each orbital is represented by a square (or circle), and the electrons as arrows pointing up or down (indicating the electron spin). When electrons are placed in a set of orbitals of equal energy, they are spread out as much as possible to give as few paired electrons as possible (Hund's rule).
examples will be added at a later date
In a ground state configuration, all of the electrons are in as low an energy level as it is possible for them to be. When an electron absorbs energy, it occupies a higher energy orbital, and is said to be in an excited state.
The electrons in the outermost shell (the ones with the highest value of n) are the most energetic, and are the ones which are exposed to other atoms. This shell is known as the valence shell. The inner, core electrons (inner shell) do not usually play a role in chemical bonding.
Elements with similar properties generally have similar outer shell configurations. For instance, we already know that the alkali metals (Group I) always form ions with a +1 charge; the "extra" s^{1} electron is the one that's lost:
IA | Li | 1s^{2}2s^{1} | Li+ | 1s^{2} |
Na | 1s^{2}2s^{2}2p^{6}3s^{1} | Na+ | 1s^{2}2s^{2}2p^{6} | |
K | 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{1} | K+ | 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6} |
The next shell down is now the outermost shell, which is now full — meaning there is very little tendency to gain or lose more electrons. The ion's electron configuration is the same as the nearest noble gas — the ion is said to be isoelectronic with the nearest noble gas. Atoms "prefer" to have a filled outermost shell because this is more electronically stable.
IIA | Be | 1s^{2}2s^{2} | Be^{2+} | 1s^{2} |
Mg | 1s^{2}2s^{2}2p^{6}3s^{2} | Mg^{2+} | 1s^{2}2s^{2}2p^{6} | |
IIIA | Al | 1s^{2}2s^{2}2p^{6}3s^{2}3p^{1} | Al3+ | 1s^{2}2s^{2}2p^{6} |
IVA | Sn | [Kr]4d^{10}5s^{2}5p^{2} | Sn^{2+} | [Kr]4d^{10}5s^{2} |
Sn^{4+} | [Kr]4d^{10} | |||
Pb | [Xe]4f^{14}5d^{10}6s^{2}6p^{2} | Pb^{2+} | [Xe]4f^{14}5d^{10}6s^{2} | |
Pb^{4+} | [Xe]4f^{14}5d^{10} | |||
VA | Bi | [Xe]4f^{14}5d^{10}6s^{2}6p^{3} | Bi^{3+} | [Xe]4f^{14}5d^{10}6s^{2} |
Bi^{5+} | [Xe]4f^{14}5d^{10} |
IVA | C | 1s^{2}2s^{2}2p^{2} | C4- | 1s^{2}2s^{2}2p^{6} |
VA | N | 1s^{2}2s^{2}2p^{3} | N3- | 1s^{2}2s^{2}2p^{6} |
VIA | O | 1s^{2}2s^{2}2p^{4} | O2- | 1s^{2}2s^{2}2p^{6} |
VIIA | F | 1s^{2}2s^{2}2p^{5} | F- | 1s^{2}2s^{2}2p^{6} |
VIIIA | Ne | 1s^{2}2s^{2}2p^{6} | ||
Ar | 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6} |
B-group | Fe | 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{6}4s^{2} | Fe^{2+} | 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{6} |
Fe^{3+} | 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{5} |
Martin S. Silberberg, Chemistry: The Molecular Nature of Matter and Change, 2nd ed. Boston: McGraw-Hill, 2000, p. 277-284, 293-307.