In chemistry, the electron configurations of elements are important to know. The number of electrons in an outer shell is used to determine how reactive an element will be and its oxidation state. This introduction explains what these terms mean and how they are applied within a group of elements.
The why are the noble gases relatively unreactive is a question that has been asked many times. The answer to this question is that they have low electron configurations.
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The electron configurations within the same group of the periodic table are similar because the elements have the same number of valence electrons. The valence electrons are in the outermost orbital and determine how the element reacts chemically. The elements in a group have the same number of valence electrons because they have the same number of orbitals.
The Electron Configuration
The electron configuration is the distribution of electrons of an atom or molecule (or other physical structure) in atomic or molecular orbitals. For example, the electron configuration of the neon atom is 1s2 2s2 2p6, using the notation explained below. This notation can be used to represent any orbital configuration, including monatomic hydrogen (1s1), sodium (3s1), and chlorine (3p5). The general form of an extended electron configuration is as follows:
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p4
The s-Block Elements
The s-block elements are the elements in which the valence electrons are present in the s orbital. The s orbital is the first orbital of the shell. The s-block elements are present on the left side of the periodic table and comprise Groups 1 and 2, which includes H and He, Li, Be, Na, Mg, and so on.
The p-Block Elements
The p-block elements are the elements in groups 13-18 of the periodic table. The p-block is to the right of the main body of the periodic table and below the d-block. The elements in the p-block are called the p-block elements.
The electrons in the outermost energy level of p-block elements are called valence electrons. The number of valence electrons in an element influences its chemical properties. For example, chlorine (Cl), which has 17 valence electrons, is a halogen, and therefore has properties that are similar to other halogens, such as fluorine (F), bromine (Br), and iodine (I).
p-block elements can be divided into two groups: metals and nonmetals. The dividing line between these two groups is not always clear, but generally, if an element has more than four valence electrons, it is a nonmetal. If an element has four valence electrons or fewer, it is a metal.
The d-Block Elements
The d-Block Elements
The elements in the periodic table are organized into groups and periods. The d-block elements, which include the transition metals, are found in groups 3-12 on the periodic table. The first two columns of the periodic table (the groups 1 and 2) are called s-block elements, while the second and third columns (the groups 3 and 4) are called p-block elements. The fourth column of the periodic table is the d-block.
The elements in each block have similar electron configurations. The s-block elements have their outermost electrons in an s orbital, while the p-block elements have their outermost electrons in a p orbital. The d-block elements have their outermost electrons in d orbitals.
The following table shows the electron configurations of some representative elements in each group of the periodic table:
Group 1A: H, Li 1s1
Group 2A: Be, Mg 2s2
Group 3A: B, Al 2p1
Group 4A: C, Si 3s2 3p2
Group 5A: N, P 3s2 3p3
Group 6A: O, S 3s2 3p4
Group 7A: F, Cl 4s1 3d10 4p5
Group 8A: Ne 4s2 3d10 4p6
The f-Block Elements
The electron configurations of the elements within the same group have some similarities. The elements in groups 3-12 are transition metals. They all have unpaired d electrons in their ground states, but they also have different numbers of f electrons. The f-block elements are sometimes called the inner transition metals because their d and f orbitals are very close in energy (unlike the s and p orbitals).
The elements in groups 13-18 are main group elements. They all have filled s orbitals in their ground states, but they also have different numbers of p and d electrons.
The table below shows the electron configurations of the first 20 elements in each group. Notice that the last column lists the element’s symbol.
1s^2 2s^2 2p^6 3s^2 3p^6 4s^2
Na (11), Mg (12), Al (13), Si (14), P (15), S (16), Cl (17), Ar (18)
1s^2 2s^2 2p^6 3s^2 3p^6 4s^23d10
Ca (20), Sr (21), Ba (22), Ra (88)
The Transition Elements
The transition elements are located in the center of the periodic table. They are elements that have incomplete d or f orbitals in their ground state electron configurations. The first row of the transition elements contains elements with incomplete d orbitals, and the second row contains elements with incomplete f orbitals.
The table below shows the electron configurations of the first ten transition elements.
| Element | Configuration |
| ——- | ————- |
| Sc | 1s22s22p63s23p64s23d1 |
| Ti | 1s22s22p63s2
The Inner Transition Elements
The inner transition elements are the elements in Groups 3 through 12 on the periodic table (see below). Like the transition elements, they have partially filled d orbitals in at least one common oxidation state. However, in addition to their d orbitals, they also have f orbitals (which are not shown in the diagrams below). The filling of the f orbitals affects the chemistry of these elements in some important ways.
The first thing to notice about the electron configurations of the inner transition elements is that they follow a pattern. For example, look at the configuration of europium (Eu, atomic number 63). The first two electrons go into the 1s orbital; the next six go into the 2s orbital; and so on. The 14th electron goes into one of the 5d orbitals, and then we start to fill the 4f orbitals. Notice that this is exactly the same pattern we saw for chromium (Cr, atomic number 24), except that we start filling the 4f orbitals after we fill all of the 3d orbitals.
This general pattern continues as we go down Group 3. For each element, we fill all ofthe orbitals up to and includingthe nth orbital, and then we start fillingthe (n+1)th orbital. In other words, for each element in Group 3, we fill allofthe 1s orbitals before we start filling anyofthe 2s orbitals; we fill allofther 2s orbitals before we start filling anyofther 3s orbitals; and so on.
What about Groups 4 through 8? The general pattern continues, but with an important difference. For each element in these groups,we fill alloftheorbitalshigher thanand includingthenthorbitalbeforewe start filling anyoftheorbitalshigher than(n+1)thorbital. In other words, for each elementinGroup 4,wewillfillallofthe2sorbitalsaheadofthe2porbitals;we willfillallofthe3saheadofthe3porbitals;and so on.
The reason for this difference is that as we go across a period from left to right,theacidityofthewaterincreases(Group 16 ismostacidic,Group 1 isleastacidic). Thismeans thateachelementisclearingitsorbitsmorequicklythanitonelevelbelow itufffdinotherwords,.
The Rare Earth Elements
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The table below shows the electron configurations for elements within the same group. As you can see, the elements in each group have different numbers of electrons in their outermost shells.
Group 1A: 2, 8, 18, 32
Group 2A: 2, 8, 18, 32
Group 3A: 2, 8, 18, 32
Group 4A: 2, 8, 18, 32
Group 5A: 2, 8, 18
Group 6A: 2, 8
Group 7A: 2
The “what information is provided by the specific block location of an element” is a question that can be answered by looking at the “electron configurations within the same group of elements”.