Chemistry
Electron Configuration
Electron configuration refers to the arrangement of electrons within an atom, which is represented by a series of numbers and letters. It describes the distribution of electrons in the energy levels and sublevels of an atom. The electron configuration of an atom determines its chemical properties and behavior.
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12 Key excerpts on "Electron Configuration"
eBook - PDF- Ageetha Vanamudan(Author)
- 2023(Publication Date)
- Delve Publishing(Publisher)
Usually, a system with one electron is associated its Electron Configuration. Knowledge on Electron Configuration is very important for different fields such as chemistry and physics (Congrains et al., 2013). It enables one properly to understand the structure of the periodic table of elements. Knowledge of Electron Configuration is also useful in describing the chemical bonds holding atoms together. When applied to bulk materials, knowledge of Electron Configuration is used in explaining the peculiar properties of semiconductors and lasers. Figure 6.1: The Electron Configuration explains the distribution of electrons in an atom or molecule. Source: By Patricia.fidi - This W3C-unspecified vector image was created with Inkscape., Public Domain, https://commons.wikimedia.org/w/index. php?curid=1239876 Electronic Configuration and Periodic Table 97 The concept of Electron Configuration was introduced in the Bohr model of the atom. It makes use of terms such as shells and subshells. Advances have been made in understanding the quantum-mechanical nature of electrons that have been used in further explaining the Electron Configuration of atoms. The term electron shell as used in Electron Configuration is used to refer to the set of allowed states that share the same principal quantum number that can be occupied by an electron. The quantum number usually denoted by the letter n is used to refer to the number before the letter in the orbital description. Usually, an atom’s nth electron shell is able to accommodate 2n 2 electrons. For example, the first shell can occupy 2 electrons, the second shell can 8 electrons while the third shell carries 18 electrons. It is noted that there is a factor of two that arises due to the fact that the allowed states are double (Mendoza & Núñez, 2009). This is attributed to the electron spin- each atomic orbital admits up to two, otherwise identical electrons with opposite spin.
eBook - PDF- Leo J. Malone, Theodore O. Dolter(Authors)
- 2012(Publication Date)
- Wiley(Publisher)
Each shell holds 2n 2 electrons and has n different subshells. The electrons in any atom can be assigned to a given shell and subshell. This is known as the ele- ment’s Electron Configuration. Electrons fill subshells according to the Aufbau principle, which simply means that the lowest energy subshells fill first. As we proceed through the Electron Configurations of the elements, one fact makes itself apparent. Atoms of elements in C H A P T E R S U M M A R Y vertical columns or groups have the same outermost subshell configuration. The next higher shell is indicat- ed as one goes down the table. Since the periodic table and Electron Configuration are interrelated, we now put the periodic table to work in writing electron configura- tions. We see that various groups can be identified by their specific configurations as well as their properties. By using orbital diagrams, we can expand the representation of Electron Configuration to include assignment of electrons to orbitals. Two other rules are required to do this successfully. The Pauli exclusion principle relates to the spin of electrons in the same orbital, and Hund's rule relates to the electron distribu- tion in separate orbitals of the same subshell. The periodic table tells us even more. There are general trends in the radii of atoms. In general, atomic radii decrease up and to the right on the table. The radius of an atom is related to the shell of the outer- most electrons and to the number of electrons in the outermost subshell. The higher the energy of the shell and the fewer electrons in the outermost subshell, the larger the atom. Thus, atoms to the lower left are the largest atoms. The opposite trend occurs if we compare the energy required to remove the outermost electron, which is known as the element’s ionization energy (I.E.). The same factors that affect the size of an atom affect the ionization energy. That is, in general, larger atoms have lower ionization energies than smaller atoms.
eBook - PDFChemistry
Principles and Reactions
- William Masterton, Cecile Hurley(Authors)
- 2020(Publication Date)
- Cengage Learning EMEA(Publisher)
In this section, you will learn how to predict the Electron Configurations of atoms of elements. There are a couple of different ways of doing this, which we consider in turn. It should be emphasized that, throughout this discussion, we refer to isolated gaseous atoms in the ground state. (In excited states, one or more electrons are promoted to a higher energy level.) Electron Configuration from Sublevel Energies Electron Configurations are readily obtained if the order of filling sublevels is known. Electrons enter the available sublevels in order of increasing sublevel energy. Ordinarily, a sublevel is filled to capacity before the next one starts to fill. The relative energies of different sublevels can be obtained from experiment. Fig-ure 6.13 is a plot of these energies for atoms through the n 5 4 principal level. From Figure 6.13 it is possible to predict the Electron Configurations of atoms of elements with atomic numbers 1 through 36. Because an s sublevel can hold only two electrons, the 1s is filled at helium (1s 2 ). With lithium ( Z 5 3), the third 2 e 2 1s 2 n 5 1 ℓ 5 0 1s 2s 3s Figure 6.11 s orbitals . The relative sizes of the 90% contours (see Figure 6.9b) are shown for the 1s, 2s, and 3s orbitals. Copyright 2016 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. 6-5 Electron Configurations in Atoms 139 electron has to enter a new sublevel: This is the 2s, the lowest sublevel of the second principal energy level. Lithium has one electron in this sublevel (1s 2 2s 1 ). ▼ With beryllium ( Z 5 4), the 2s sublevel is filled (1s 2 2s 2 ).
eBook - PDFChemistry
The Molecular Nature of Matter
- Neil D. Jespersen, Alison Hyslop(Authors)
- 2014(Publication Date)
- Wiley(Publisher)
When two atoms come together to form new chemical bonds, it must be due to the interactions of the electrons furthest from the nucleus. It also makes sense that electrons buried deep within the atomic structure most likely do not contribute significantly to the chemical properties of an atom. These electrons furthest from the nucleus are the electrons that we are most interested in. Each group within the periodic table has elements with similar chemical and physical properties that vary regularly within the group. We are now ready to understand the rea- son for these similarities in terms of the electronic structures of atoms. Chemical Properties and Valence Shell Configurations When we consider the chemical reactions of an atom (particularly one of the representa- tive elements), our attention is usually focused on the distribution of electrons in the outer shell of the atom (those electrons where n is the largest). This is because the outer electrons, those in the outer shell, are the ones that are exposed to other atoms when the atoms react. (Later we will see that d electrons are also very important in reactions involv- ing the transition elements.) It seems reasonable, therefore, that elements with similar properties should have similar outer shell Electron Configurations. For example, let’s look at the alkali metals of Group 1A. Going by our rules, we obtain the following abbreviated Electron Configurations and complete Electron Configurations. Li 3 He 4 2s 1 or 1s 2 2s 1 Na 3 Ne 4 3s 1 or 1s 2 2s 2 2p 6 3s 1 K 3 Ar 4 4s 1 or 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 Rb 3 Kr 4 5s 1 or 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 5s 1 Cs 3 Xe 4 6s 1 or 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 5s 2 5p 6 6s 1 The abbreviated Electron Configuration clearly illustrates the similarities of the outermost Electron Configurations in these elements. Each of these elements has only one outer shell electron that is in an s subshell (shown in bold, red, type).
eBook - PDF- Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2019(Publication Date)
- Openstax(Publisher)
288 6 • Electronic Structure and Periodic Properties of Elements Access for free at openstax.org FIGURE 6.27 This partial periodic table shows Electron Configurations for the valence subshells of atoms. By “building up” from hydrogen, this table can be used to determine the Electron Configuration for atoms of most elements in the periodic table. (Electron Configurations of the lanthanides and actinides are not accurately predicted by this simple approach. See Figure 6.29 We will now construct the ground-state Electron Configuration and orbital diagram for a selection of atoms in the first and second periods of the periodic table. Orbital diagrams are pictorial representations of the Electron Configuration, showing the individual orbitals and the pairing arrangement of electrons. We start with a single hydrogen atom (atomic number 1), which consists of one proton and one electron. Referring to Figure 6.26 or Figure 6.27, we would expect to find the electron in the 1s orbital. By convention, the value is usually filled first. The Electron Configuration and the orbital diagram are: Following hydrogen is the noble gas helium, which has an atomic number of 2. The helium atom contains two protons and two electrons. The first electron has the same four quantum numbers as the hydrogen atom electron (n = 1, l = 0, m l = 0, ). The second electron also goes into the 1s orbital and fills that orbital. The second electron has the same n, l, and m l quantum numbers, but must have the opposite spin quantum number, This is in accord with the Pauli exclusion principle: No two electrons in the same atom can have the same set of four quantum numbers. For orbital diagrams, this means two arrows go in each box (representing two electrons in each orbital) and the arrows must point in opposite directions (representing paired spins). The Electron Configuration and orbital diagram of helium are: The n = 1 shell is completely filled in a helium atom.
eBook - PDF- Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2015(Publication Date)
- Openstax(Publisher)
There is no simple method to predict the exceptions for atoms where the magnitude of the repulsions between electrons is greater than the small differences in energy between subshells. Electron Configurations and the Periodic Table As described earlier, the periodic table arranges atoms based on increasing atomic number so that elements with the same chemical properties recur periodically. When their Electron Configurations are added to the table (Figure 6.30), we also see a periodic recurrence of similar Electron Configurations in the outer shells of these elements. Because they are in the outer shells of an atom, valence electrons play the most important role in chemical reactions. The outer electrons have the highest energy of the electrons in an atom and are more easily lost or shared than the core electrons. Valence electrons are also the determining factor in some physical properties of the elements. Elements in any one group (or column) have the same number of valence electrons; the alkali metals lithium and sodium each have only one valence electron, the alkaline earth metals beryllium and magnesium each have two, and the halogens fluorine and chlorine each have seven valence electrons. The similarity in chemical properties among Chapter 6 | Electronic Structure and Periodic Properties of Elements 317 elements of the same group occurs because they have the same number of valence electrons. It is the loss, gain, or sharing of valence electrons that defines how elements react. It is important to remember that the periodic table was developed on the basis of the chemical behavior of the elements, well before any idea of their atomic structure was available. Now we can understand why the periodic table has the arrangement it has—the arrangement puts elements whose atoms have the same number of valence electrons in the same group.
eBook - PDFFoundations of Chemistry
An Introductory Course for Science Students
- Philippa B. Cranwell, Elizabeth M. Page(Authors)
- 2021(Publication Date)
- Wiley(Publisher)
The electrons are forced to have opposite spins in helium. The Electron Configuration for the atom is a shorthand that describes the occupancy of the orbitals. For hydrogen, the Electron Configuration is 1s 1 . For helium, the electron configura-tion is 1s 2 . The superscripted number represents the number of electrons in the 1s orbital for each element. Once the 1s orbital is full, the next electron at lithium ( Z = 3) must enter the second energy level, as this is the next lowest in energy. Within this energy level, the 2s orbital is of lower energy than the 2p orbitals. Thus the electron occupies the 2s orbital. The Electron Configuration of lithium is 1s 2 2s 1 . The next element, beryllium ( Z = 4), has four electrons, and the fourth elec-tron must pair up with the electron in the 2s orbital, as this is of lower energy than the 2p orbitals. The Electron Configuration is 1s 2 2s 2 . The two electrons in the 2s orbital have opposite spins. At boron ( Z = 5), the 2s orbital is full, and so the next electron must occupy a 2p orbital. All are empty and of the same energy, and so we arbitrarily place the electron in the 2p x orbital, although it could equally occupy 2p y or 2p z . The elec-tron configuration for boron is 1s 2 2s 2 2p 1 . The next two electrons at carbon and nitrogen enter the other empty 2p orbi-tals. Nitrogen has the Electron Configuration 1s 2 2s 2 2p 3 . As all 2p orbitals are half-filled at nitrogen, the next electron of oxygen must pair with another p elec-tron to give one fully occupied and two half-occupied 2p orbitals: 1s 2 2s 2 2p 4 . This can be visualised more easily by using the representation with electrons in boxes, as in Figure 1.10 for oxygen. The next electrons complete the remaining two 2p orbitals so that at neon ( Z = 10), we have a filled second shell of electrons: 1s 2 2s 2 2p 6 . We will see that this is a very stable arrangement of electrons and has significant consequences for the chemical reactivity of the element.
eBook - PDF- Kenneth S. Krane(Author)
- 2020(Publication Date)
- Wiley(Publisher)
For example, the first 18 elec- trons fill the levels up through 3p, and the energy levels (subshells) available to the 19th electron in potassium (Z = 19) or calcium (Z = 20) are well described by Figure 8.1. However, the energy levels appropriate to the 19th electron in a heavy element such as lead (Z = 82) would be very different. In this case, it is more correct to describe the atom in terms of shells—all of the n = 3 states (the M shell) are grouped together, as are all of the n = 4 states (the N shell), and so forth. When we discuss the inner structure of the atom, as in the case of X rays, the ordering of Figure 8.1 is not appropriate, and it is more appro- priate to group the levels by shells, as we do in Section 8.5. The Periodic Table Figure 8.2 shows the periodic table, which is an orderly array of the chemical elements, listed in order of increasing atomic number Z and arranged in such a way that the vertical columns, called groups, contain elements with rather similar physical and chemical properties. In this section, we discuss the way in which the filling of electronic subshells helps us understand the arrangement of the periodic table. In later sections, we examine some of the physical and chemical properties of the elements. In attempting to understand the ordering of subshells and the periodic table, we must follow two rules for filling the electronic subshells: 1. The capacity of each subshell is 2(2l + 1). (This is of course just another way of stating the Pauli exclusion principle.) 2. The electrons occupy the lowest energy states available. To indicate the Electron Configuration of each element, we use a notation in which the identity of the subshell and the number of electrons in it are listed. The identity of the subshell is indicated in the usual way, and the number of electrons in that subshell is indicated by a superscript.
- Eric R Scerri(Author)
- 2008(Publication Date)
- ICP(Publisher)
In a section entitled “A case of reduction”, the authors conclude that the sophisticated periodic table can be explained using atomic theory. The basis for this conclusion is the authors view that the term z can be identified with the number of electrons in any neutral atom. They state that, The necessary link between chemical similarity and equal outer Electron Configuration states that the latter causes the former. (p. 403) It is interesting to contrast the above statement with that of one of the leading authorities on electronic configurations of atoms, No simple relation exists between the Electron Configuration of the ground state of the atom and the chemistry of the element under consideration. (Jorgensen, 1973). Moreover, a reduction of chemistry, or more specifically the periodic table, to quantum mechanics requires far more than a mere approximate explanation of the properties of elements in terms of outer electron con- figurations. After all, quantum mechanics or “atomic theory”, which the authors constantly allude to, is not a qualitative theory dealing in outer- shell electrons. Such explanations are indeed frowned upon by physicists as being of a typically picturesque and naively realistic kind, typical of chemists. Worse still, according to quantum mechanics, the very notion of electron shells or Electron Configurations becomes strictly invalid as mentioned in the introduction.21 Nevertheless, the connection between chemical behavior and electronic configurations can be improved by approaches practiced in computational quantum chemistry. Calculations generally consist in expanding the wave- function of a many-electron atom, for example, as a linear combination of terms representing excited state configurations, in addition to the ground state configuration. A more realistic approximation to chemical behav- ior of atoms is thus achieved through a superposition of numerous, often 99
eBook - PDF- Young, William Vining, Roberta Day, Beatrice Botch(Authors)
- 2017(Publication Date)
- Cengage Learning EMEA(Publisher)
All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. WCN 02-300 Unit 4 Electron Configurations and the Properties of Atoms 92 Notice that elements within a group have the following in common: ● They have the same number of electrons beyond the core electrons (represented by noble gas notation) and similar Electron Configurations. For example, both Li and Na have one electron in addition to the core electrons, and both elements have the general Electron Configuration 3 noble gas 4 ns 1 . ● For the main group elements, the number of electrons beyond the core electrons is equal to the group number (with the exception of He ). The electrons beyond the core electrons are the valence electrons for an element. The valence electrons are the highest-energy electrons and are the electrons least strongly attracted to the nucleus. It is these electrons that are involved in chemical reactions and the formation of chemical bonds. As shown in Interactive Figure 4.3.3, for the Group A elements, the number of valence electrons is equal to the group number of the element. The fact that elements within a group have similar Electron Configurations and the same number of valence electrons suggests that elements within a group have similar properties, something we will investigate later in this unit. The orbital filling order is related to the structure of the periodic table, as shown in Interactive Figure 4.3.4. Interactive Figure 4.3.4 Identify Electron Configuration and valence electrons for elements using the periodic table. 1 s 1 s 2 s 3 s 4 s 5 s 6 s 2 p 3 p 4 p 5 p 6 p 4 f 5 f 7 s 3 d 4 d 5 d 6 d s –block elements p –block elements d –block elements (transition metals) f –block elements: lanthanides (4 f ) and actinides (5 f ) Subshells and the periodic table Copyright 2018 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. WCN 02-300
eBook - PDF- John A. Olmsted, Gregory M. Williams, Robert C. Burk(Authors)
- 2020(Publication Date)
- Wiley(Publisher)
The arrangement of the periodic table provides a simple way to determine the filling order of the elements, as shown in Figure 5.10 and applied in Example 5.2. FIGURE 5.9 The calculated energy-level diagram for neutral atoms with Z between 19 and 30 shows that the 3d and 4 s atomic orbitals have nearly the same energy. E 3p 4s 3d 4p 1s 2p 2s 3s 3p 4s 3d 4p 5s 4d 5p 6s 4 f 5d 6p 7s 5 f 6 7p d 1s 2s 3s 4s 5s 6s 7s 3d 4d 5d 6d 4f f 5 2p 3p 4p 5p 6p Begin s block p block d block f block here to 3 to 2 1 2 3 4 5 6 7 to 4 to 5 to 6 to 7 7p FIGURE 5.10 The periodic table in block form, showing the filling sequence of the atomic orbitals. Filling proceeds from left to right across each row and from the right end of each row to the left end of the succeeding row. 5. 2 Structure of the Periodic Table 213 Valence Electrons The chemical behaviour of an atom is determined by the electrons that are accessible to an approaching chemical reagent. Accessibility, in turn, has a spatial component and an ener- getic component. An electron is accessible spatially when it occupies one of the largest orbit- als of the atom. Electrons on the perimeter of the atom, farthest from the nucleus, are the first ones encountered by an incoming chemical reagent. An electron is accessible energetically when it occupies one of the least-stable occupied orbitals of the atom. Electrons in less stable (higher energy) orbitals are thus more chemically active than electrons in more stable orbitals. Similar electron accessibility generates similar chemical behaviour. For example, iodine has many more electrons than chlorine, but these two elements display similar chemical behaviour, as reflected by their placement in the same group of the periodic table. This is because the chemistry of chlorine and iodine is determined by the number of electrons in their largest and least-stable occupied orbitals: 3s and 3p for chlorine and 5s and 5p for iodine.
eBook - PDFFundamentals of Inorganic Chemistry
An Introductory Text for Degree Studies
- J Barrett, M A Malati(Authors)
- 1997(Publication Date)
- Woodhead Publishing(Publisher)
This is the rationalization of Hund's rules. These may be stated in the following way: In filling a set of degenerate orbitals the number of unpaired electrons is maximized, and such unpaired electrons will possess parallel spins. An alternative, more exact, statement appears in section 3.7. The electronic configuration of the carbon atom is ls 2 2s 2 2p 2 or, if the detailed content of the 2p orbitals is being discussed it may be written as: ls^s^p^p,, 1 , the choice of x and v being merely alphabetical. The 2p configuration is sometimes indicated diagrammatically by entering arrows (representing electrons with a particular spin) into two of three boxes (representing atomic orbitals): Î Î Such diagrams are only approximate descriptions of the electron arrangements in atoms, a full description is dealt with in Section 3.6. By similar arguments it is concluded that the nitrogen atom has a lowest energy electronic configuration of ls 2 2s 2 2p^ 1 2p^, 1 2p z 1 . In the oxygen, fluorine and neon atoms the extra electrons doubly occupy the appropriate number of 2p orbitals since pairing is the lowest energy option - the 3s - 2p gap being greater than any interelectronic repulsion energy involved. The atoms of the elements Li, Be, B, C, N, O, F and Ne form the second period of the Periodic Classification of the Elements. The third period contains the elements Na, Mg, Al, Si, P, S, Cl and Ar, which have core electronic configurations which are that of neon (1 s 2 2s 2 2p 6 ) plus those derived from the regular filling of the 3s and 3p orbitals as have been described for that of the 2s and 2p orbitals. Identical arguments apply and the elements of the third period are arranged under their counterparts in the second period with identical 'outer electronic configurations', except for the change in value of« from 2 to 3. The term 'outer' is a reference to the value of n and is related to the greater diffuseness of orbitals as n increases in value - the orbitals become larger.
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