# An Introduction to Electron Configurations

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```					An Introduction to Electron Configurations An electron configuration is a description of electron arrangement within an atom, which indicates both population and location of electrons among the various atomic orbitals. General Rules for Electron Configurations 1. Electrons occupy orbitals of the lowest energy available. 2. There can be a maximum of only two electrons in any given orbital. 3. Building-Up Principle (Aufbau Principle): Electrons are added by successively filling subshells (sublevels) with electrons in a specific order based on increasing energies of the subshells.
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The maximum number of electrons in any s subshell is two. (just one orbital in the s subshell) The maximum number of electrons in any p subshell is six. (because there are three orbitals in p) The maximum number of electrons in any d subshell is ten. (because there are five orbitals in d) The maximum number of electrons in any f subshell is fourteen. (because there are seven orbitals in f)

4. Electrons are added to orbitals one at a time. In other words, each of the three orbitals within the p subshell would get one electron before either of them would receive a second. 5. Ground State: The electron configuration associated with the lowest energy level of the atom is referred to as ground state. Each electron in the atom or ion will be in the lowest energy level possible. Configurations associated with electrons in energy levels other than the lowest are referred to as excited states. Example Configurations 1. Hydrogen has a single electron and therefore has the following configuration. H = 1s1  The large 1 refers to the value of n, the principal quantum number, or principal energy level of the electrons.  The letter s is the subshell/sublevel. In the case of the first energy level, there is only one sublevel. For each successive energy level there is an additional sublevel. For atoms with more electrons, there are 2 subshells in the 2nd energy level, three in the third, and so on.  The superscript 1 after the s refers to the number of electrons in the 1s subshell. 2. Nitrogen has 7 electrons, distributed as shown in the following configuration. N = 1s2 2s2 2p3  So that’s 2 electrons in the 1s subshell, 2 in the 2s subshell, and three in the 2p subshell. The p subshell has three orbitals. Recalling rule #4 above, you would conclude that each of those orbitals would have one electron in it. 3. Fluorine has 9 electrons distributed as follows: F = 1s2 2s2 2p5  So how would the p subshell look for fluorine? 4. Magnesium has 12 electrons distributed as follows: Mg = 1s2 2s2 2p6 3s2 Practice: Write the electron configurations for the following elements Neon Oxygen Phosphorus Calcium Iron

Core Electrons (Noble Gas Core) The terms "core electrons" or "noble gas core" refer to the electrons within the atom which have the same electron configuration as the nearest noble gas (group 8A or 18 on the periodic table) of lower atomic number. The core electrons are the inner electrons which are not directly involved in bonding. Those electrons that are outside the core are called valence electrons. Much more about them soon! The core for Li is He. The core electrons of Li have the identical electron configuration as an atom of He. In addition, Li has one more electron (its valence electron) in the 2s orbital. The core for F is also He. How many valence electrons does it have? The core for Al is Ne. How many valence electrons does it have? The electron configuration for argon is: Ar = 1s2 2s2 2p6 3s2 3p6 The electron configuration for potassium is: K = 1s2 2s2 2p6 3s2 3p6 4s1 Potassium has an argon core plus that 4s1 electron. An abbreviated method for writing electron configurations uses a set of square brackets [ ] around the chemical symbol of the noble gas, followed by the configuration of the atom’s valence electron(s). So, the abbreviated electron configuration for potassium is: K [Ar] 4s1 Chlorine has a neon core. The configuration for chlorine is: Cl 1s2 2s2 2p6 3s2 3p5 or the abbreviated method is: Cl [Ne] 3s2 3p5 Valence Electrons Valence electrons are those electrons in an atom outside the noble-gas core. Valance electrons are the electrons in the outermost principal quantum (energy) level. Valance electrons are the electrons which are almost always the ones involved in reactions and forming chemical bonds. Potassium has a single valence electron, 4s1, which comes from the 4s subshell. Chlorine has a total of seven valence electrons, 3s2 3p5, two from the 3s subshell and five from the 3p subshell. Electron Configurations and the Periodic Table The representative elements (also called main group elements) are the elements in Groups 1 (1A) through 17 (7A), all of which have incompletely filled s or p subshells of the highest principal quantum number. The representative elements all have valence shell configurations of nsanpb, with some choice of a and b. s-block elements Group 1 (1A; the alkali metals) and Group 2 (2A; the alkaline earth metals) are referred to as s-block elements Group 1 elements have a noble gas core plus 1 valence electron with an ns1 configuration. Li [He] 2s1 or K [Ar] 4s1 Group 2 elements have a noble gas core plus 2 valence electrons with an ns2 configuration Be [He] 2s2 or Ca [Ar] 4s2 p-block elements Group 13 (3A) elements through Group 18 (8A) are referred to as p-block elements. Group 13 elements have the general configuration of ns2 np1 B [He] 2s2 2p1 Group 14 elements have the general configuration of ns2 np2 C [He] 2s2 2p2

Group 15 elements have the general configuration of ns2 np3 N [He] 2s2 2p3 Group 16 elements have the general configuration of ns2 np4 O [He] 2s2 2p4 Group 17 elements have the general configuration of ns2 np5 F [He] 2s2 2p5 The p-block elements in the fourth period and beyond will have the noble gas core together with (n1)d10. Br [Ar] 4s2 3d10 4p5 In Group 18 (8A; the noble gases) the p subshell has just been filled. Ar 1s2 2s2 2p6 3s2 3p6

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