Electron Configurations Help (page 2)
Introduction to Electron Configurations
Electrons play a big role in the way the entire atom gets along with its friends, family, and neighbors. They don’t stick together on a whim. Usually, atoms of an element will only combine with identical atoms or those of other elements in certain ways. Electrons serve as the glue between the nuclei of two or more atoms.
The word electron was coined by G. Johnstone Stoney in 1891. It was used to describe a unit of electrical charge measured in his experiments where an electrical current was sent through various chemical solutions to test its effect.
A chemical bond is the attachment between atoms within a molecule.
The number of bonds an atom can form with other atoms depends upon the number of electrons it can easily share with its neighbors. The more electrons available to be shared, the more readily an atom will bond. In general, atoms combine in the numbers shown in Table 6.1 .
Elements can be placed in rows and columns of the Periodic Table by knowing something about their properties. The numbers at the top of the chart, Roman numerals I–VIII (or alternatively 1–18), are used to identify groups and chemical properties. The element groups usually give the number of electrons in the outermost orbital of the atoms in each column. Remember, these outer electrons are known as valence electrons. For example, the atoms of the elements in column IV have four electrons available to create bonds. Elements in column II have two free electrons in the outermost orbit around the nucleus.
Just as John Newlands overlapped elements around a cylinder and saw repeating patterns of characteristics, so too do elements in each family change in predictable ways. As you look down the columns of the Periodic Table, these characteristics can be compared. Table 6.2 shows the pattern of a few elements from the alkali metal family.
The electron configuration of an atom describes the specific dispersal of electrons among available subshells.
The capacity of an energy level can be found in the following formula:
2 n s
where n is called the principal quantum number and indicates the energy level. In the first energy level, n = 1; the second energy level n = 2; the third energy level n = 3; and so on. Figure 6.1 shows the electrons in these different orbital subshell energy levels.
Electron subshells or orbitals are written with s, p, d, and f as terms for each known energy level. If you were to write the electron configuration of calcium (Ca), with an atomic number of 20, it would look like the following: 1s 2 2s 2 p 6 3s 2 p 6 4s 2 . It has 2 electrons in the 1s subshell, 8 electrons in the 2sp, 8 electrons in the 3sp, and 2 electrons in the 4s subshell. Another way of writing this is that there are 2 electrons in the n = 1 level, 8 electrons in the n = 2 level, 8 electrons in the n = 3 level, and 2 electrons in the n = 4 level.
The overall capacity of the energy levels then looks like that shown in Table 6.3 .
Orbitals of the s-type are always singular, p-types form orbital sets of 3, d-type orbitals come in sets of 5, and f-type orbitals are written in sets of 7.
The 4s electrons of calcium are found in the outermost orbit to be filled and from this position react with other elements. Valence electrons affect the reactivity of atoms with other elements.
Molecules that share electrons are generally smaller, have lower melting and boiling points, are insoluble in water, and do not conduct electricity. The s and p orbitals of nearby atoms overlap to form a mixed orbital.
A simple example is that of oxygen (O). The atomic number of oxygen is 8. The electron configuration of oxygen is:
1s 2 2s 2 p 4
2 electrons in the 1s subshell, 2 electrons in the 2s subshell, and 4 electrons in the p subshell give a total of 8.
An orbital diagram (shown as circles) is the notation used to show the number of electrons in each subshell. Each subshell is labeled with its subshell notation, s, p, d, or f. An orbital diagram also makes it easy to see the sequence of how subshells are filled. If you use small circles to stand for a subshell, then the orbital diagram can be used to find the orbital configuration of nearly every element. Figure 6.2 shows the order of an orbital filling sequence.
Atoms contain an infinite number of possible electron configurations. The configuration associated with the lowest energy level of the atoms is called the ground state . When an atom’s energy levels go from a lower energy to a higher energy level, the change is sometimes seen as a flash of color or sudden heat.
The Aufbau principle , also called the building-up principle, is used to show an atom’s ground state.
Atoms fill up subshells like eggs in a carton. An atom’s ground state adds electrons in a particular building order.
The standard building order for the description of an atom’s orbital configuration is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, and 5f. An example of an atom’s orbital configuration is illustrated in Figure 6.3 for the element magnesium.
The Aufbau principle shows how energy increases in subshells. By filling the lowest energy subshells first, the ground state is built up. As the energy of the atom increases, the number of subshells filled increases.
When the number of electrons in an atom equals the atomic number (the number of protons), then the atom is neutral. Electron configurations are seen in the Periodic Table under the atomic number and element symbol.
The noble gases are in the ground state with the 1s 2 (He) and p orbitals filled (Ne, Ar, and Kr) so they usually have no interest in reacting with anything else. In Figure 6.4 the Periodic Table is shown with the location of the element orbitals.
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