Chemistry
Electron Shells
Electron shells are energy levels within an atom where electrons are found. These shells are designated by letters (K, L, M, etc.) and can hold a specific number of electrons. The innermost shell, K, can hold up to 2 electrons, while subsequent shells can hold more. The arrangement of electrons in these shells determines an atom's chemical properties.
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11 Key excerpts on "Electron Shells"
eBook - PDFChemistry for Today
General, Organic, and Biochemistry
- Spencer Seager, Michael Slabaugh, Maren Hansen, , Spencer Seager, Spencer Seager, Michael Slabaugh, Maren Hansen(Authors)
- 2021(Publication Date)
- Cengage Learning EMEA(Publisher)
Electrons in the third shell all have an n value of 3, all have an energy higher than the energies of electrons in shells 1 and 2, and also are located farther from the nucleus than those in shells 1 and 2. Each shell is made up of subshells that are designated by a letter representing the group s, p, d, and f. Because all subshells are designated by one of these letters regard- less of the shell in which the subshell is found, a combination of both shell number and subshell letter is used to identify subshells clearly. Thus, a p subshell in shell number 2 is referred to as a 2p subshell. The number of subshells found in a shell is the same as the value of n for the shell. Thus, shell number 2 (n 5 2) contains two subshells. The subshells are the 2p mentioned earlier and 2s. Electrons located in specific subshells are often referred to in terms of the same number and letter as the subshell. For example, we might refer to an atom as having three 2p electrons. All electrons within a specific sub- shell have the same energy. The description of the location and energy of electrons moving around a nucleus is completed in the quantum mechanical model by specifying an orbital. Each subshell con- sists of one or more atomic orbitals, which are specific volumes of space around nuclei in which electrons move. These atomic orbitals must not be confused with the fixed elec- tronic orbits of the original Bohr theory; they are not the same. These volumes of space around nuclei have different shapes, depending on the energy of the electrons they contain shell A location and energy of electrons around a nucleus is designated by a value for n, where n 5 1, 2, 3, etc. subshell A component of a shell that is designated by a letter representing the group s‚ p, d, and f. atomic orbital A volume of space around atomic nuclei in which electrons of the same energy move. Groups of orbitals with the same n value form subshells.
eBook - PDF- Edward J. Neth, Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2019(Publication Date)
- Openstax(Publisher)
The secondary (angular momentum) quantum number, l, can have any integer value from 0 to n – 1. This quantum number describes the shape or type of the orbital. Orbitals with the same principal quantum number and the same l value belong to the same subshell. The magnetic quantum number, m l , with 2l + 1 values ranging from –l to +l, describes the orientation of the orbital in space. In addition, each electron has a spin quantum number, m s , that can be equal to No two electrons in the same atom can have the same set of values for all the four quantum numbers. 3.4 Electronic Structure of Atoms (Electron Configurations) The relative energy of the subshells determine the order in which atomic orbitals are filled (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on). Electron configurations and orbital diagrams can be determined by applying the Pauli exclusion principle (no two electrons can have the same set of four quantum numbers) and Hund’s rule (whenever possible, electrons retain unpaired spins in degenerate orbitals). Electrons in the outermost orbitals, called valence electrons, are responsible for most of the chemical behavior of elements. In the periodic table, elements with analogous valence electron configurations usually occur within the same group. There are some exceptions to the predicted filling order, particularly when half-filled or completely filled orbitals can be formed. The periodic table can be divided into three categories based on the orbital in which the last electron to be added is placed: main group elements (s and p orbitals), transition elements (d orbitals), and inner transition elements (f orbitals). 3.5 Periodic Variations in Element Properties Electron configurations allow us to understand many periodic trends. Covalent radius increases as we move down a group because the n level (orbital size) increases.
eBook - PDF- William R. Robinson, Edward J. Neth, Paul Flowers, Klaus Theopold, Richard Langley(Authors)
- 2016(Publication Date)
- Openstax(Publisher)
This allows us to determine which orbitals are occupied by electrons in each atom. The specific arrangement of electrons in orbitals of an atom determines many of the chemical properties of that atom. Orbital Energies and Atomic Structure The energy of atomic orbitals increases as the principal quantum number, n, increases. In any atom with two or more electrons, the repulsion between the electrons makes energies of subshells with different values of l differ so that the energy of the orbitals increases within a shell in the order s < p < d < f. Figure 3.25 depicts how these two trends in increasing energy relate. The 1s orbital at the bottom of the diagram is the orbital with electrons of lowest energy. The energy increases as we move up to the 2s and then 2p, 3s, and 3p orbitals, showing that the increasing n value has more influence on energy than the increasing l value for small atoms. However, this pattern does not hold for larger atoms. The 3d orbital is higher in energy than the 4s orbital. Such overlaps continue to occur frequently as we move up the chart. 148 Chapter 3 | Electronic Structure and Periodic Properties of Elements This OpenStax book is available for free at http://cnx.org/content/col12012/1.7 Figure 3.25 Generalized energy-level diagram for atomic orbitals in an atom with two or more electrons (not to scale). Electrons in successive atoms on the periodic table tend to fill low-energy orbitals first. Thus, many students find it confusing that, for example, the 5p orbitals fill immediately after the 4d, and immediately before the 6s. The filling order is based on observed experimental results, and has been confirmed by theoretical calculations. As the principal quantum number, n, increases, the size of the orbital increases and the electrons spend more time farther from the nucleus. Thus, the attraction to the nucleus is weaker and the energy associated with the orbital is higher (less stabilized).
eBook - PDF- Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2019(Publication Date)
- Openstax(Publisher)
that all electrons have −1 charges, but nuclei have +Z charges). This phenomenon is called shielding and will be discussed in more detail in the next section. Electrons in orbitals that experience more shielding are less stabilized and thus higher in energy. For small orbitals (1s through 3p), the increase in energy due to n is more significant than the increase due to l; however, for larger orbitals the two trends are comparable and cannot be simply predicted. We will discuss methods for remembering the observed order. The arrangement of electrons in the orbitals of an atom is called the electron configuration of the atom. We describe an electron configuration with a symbol that contains three pieces of information ( Figure 6.25): 1. The number of the principal quantum shell, n, 2. The letter that designates the orbital type (the subshell, l), and 3. A superscript number that designates the number of electrons in that particular subshell. For example, the notation 2p 4 (read "two–p–four") indicates four electrons in a p subshell (l = 1) with a principal quantum number (n) of 2. The notation 3d 8 (read "three–d–eight") indicates eight electrons in the d subshell (i.e., l = 2) of the principal shell for which n = 3. FIGURE 6.25 The diagram of an electron configuration specifies the subshell (n and l value, with letter symbol) and superscript number of electrons. The Aufbau Principle To determine the electron configuration for any particular atom, we can “build” the structures in the order of atomic numbers. Beginning with hydrogen, and continuing across the periods of the periodic table, we add one proton at a time to the nucleus and one electron to the proper subshell until we have described the electron configurations of all the elements. This procedure is called the Aufbau principle, from the German word Aufbau (“to build up”).
eBook - PDF- David R. Klein(Author)
- 2016(Publication Date)
- Wiley(Publisher)
Beyond this region, the remaining 5–10% of the electron density tapers off but never ends. In fact, if we want to consider the region of space that con- tains 100% of the electron density, we must consider the entire universe. In summary, we must think of an orbital as a region of space that can be occupied by electron density. An occupied orbital must be treated as a cloud of electron density. This region of space is called an atomic orbital (AO), because it is a region of space defined with respect to the nucleus of a single atom. Examples of atomic orbitals are the s, p, d, and f orbitals that were discussed in your general chemistry textbook. Phases of Atomic Orbitals Our discussion of electrons and orbitals has been based on the premise that electrons have wavelike properties. As a result, it will be necessary to explore some of the characteristics of simple waves in order to understand some of the characteristics of orbitals. Consider a wave that moves across the surface of a lake (Figure 1.6). The wavefunction (ψ) math- ematically describes the wave, and the value of the wavefunction is dependent on location. Locations FIGURE 1.6 Phases of a wave moving across the surface of a lake. Node ψ = 0 Average level of lake ψ is (–) ψ is (–) ψ is (+) ψ is (+) 14 CHAPTER 1 A Review of General Chemistry As we move across the periodic table, starting with hydrogen, each element has one more electron than the element before it (Figure 1.9). The order in which the orbitals are filled by electrons is determined by just three simple principles: 1. The Aufbau principle. The lowest energy orbital is filled first. 2. The Pauli exclusion principle. Each orbital can accommodate a maximum of two electrons that have opposite spin. To understand what “spin” means, we can imagine an electron spinning in space (although this is an oversimplified explanation of the term “spin”).
eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
Table 2.5 .Table 2.5 Properties of the Electron SubshellsSubshell l Number Orbitals (2l + 1) Maximum Electrons 2(2l + 1) Shape Nodal Planes (l ) s 0 1 2 Spherical 0 p 1 3 6 Dumb bell (2 lobes) 1 d 2 5 10 Double-dumb bell (4 lobes) 2 f 3 7 14 Complex (6–8 lobes) 3 Each subshell then contains orbitals to which the electrons are confined. The number of orbitals in each subshell is equal to 2l + 1. So, the s subshell has 2(0) + 1 = 1 orbital, the p subshell has 2(1) + 1 = 3 orbitals, and so on. These are also outlined in Table 2.5 . The orbitals within a subshell have a characteristic shape specific to that subshell shown in Fig. 2.8 . This shape is defined mathematically by the azimuthal quantum number and represents the probability of locating the electron within the area of each orbital. The orbitals in each of the s subshells are spherical in shape and increase in diameter as n increases. The p orbitals all have a “dumb bell” shape with a nodal plane (a region of zero probability) passing through the nucleus. The number of nodal planes in an orbital is equal to the value of l for the subshell. So, the orbital in an s subshell has no nodal planes (l = 0), the three orbitals in the p subshell (l = 1) have one nodal plane, all five orbitals in the d subshell (l = 2) have two nodal planes, and all seven orbitals in the f subshell (l
eBook - PDF- Morris Hein, Susan Arena, Cary Willard(Authors)
- 2016(Publication Date)
- Wiley(Publisher)
10.3 Energy Levels of Electrons • Modern atomic theory predicts that • Electrons are found in discrete principal energy levels (n = 1, 2, 3 . . . ). • Energy levels contain sublevels. 1s Number of sublevels 2 p 2s 3d 3p 3s 4 f 4d 4p 4s • Two electrons fit into each orbital but must have opposite spin to do so. 10.4 Atomic Structures of the First 18 Elements • Guidelines for writing electron configurations: • Not more than two electrons per orbital • Electrons fill lowest energy levels first: • s < p < d < f for a given value of n • Orbitals on a given sublevel are each filled with a single electron before pairing of electrons begins to occur • For the representative elements, only electrons in the outermost energy level (valence electrons) are involved in bonding. KEY TERMS wavelength frequency speed photons KEY TERMS line spectrum quanta ground state orbital KEY TERMS principal energy levels sublevels spin Pauli exclusion principle KEY TERMS electron configuration orbital diagram valence electrons Paired Exercises 213 10.5 Electron Structures and the Periodic Table • Elements in horizontal rows on the periodic table contain elements whose valence electrons (s and p) are generally on the same energy level as the number of the row. • Elements that are chemically similar are arranged in columns (groups) on the periodic table. • The valence electron configurations of elements in a group or family are the same, but they are located in different principal energy levels. KEY TERMS period groups or families (of elements) representative elements transition elements 1. Define the terms wavelength and frequency. What symbols do chemists generally use to represent these quantities? 2. What is the wavelength range for visible light? Which has a longer wavelength, red light or blue light? 3. What is the name given to a packet of energy? 4. What is an orbital? 5. What is meant when we say the electron structure of an atom is in its ground state? 6.
eBook - PDFQuantum Chemistry
A Unified Approach
- David B Cook(Author)
- 2012(Publication Date)
- ICP(Publisher)
It tends to encourage the mistaken impression that: The electrons occupy the AOs in much the same way as rabbits occupy hutches; that the orbitals exist ‘out there’ in the world just waiting to be occupied by elec-trons whereas, as we explain in Appendix C, this is not the case. But, so long as these limitations are kept in mind, this terminology is always useful. A Feature of the Energy Levels The fact that the energy levels of hydrogen and, in fact, other atoms are described by a formula like equation (2.1), which makes the energy depend inversely on the square of an integer, means that, as the allowed energy goes up, the levels get closer and closer together : (1 / 1 2 − 1 / 2 2 ) = 0 . 25 is larger than (1 / 2 2 − 1 / 3 2 ) = 0 . 13 ˙ 8 and so on. This has some consequences for chemistry, as we shall see later. 2.1 Atomic Electronic Structure 39 2.1.2 Many-electron atoms The outlook for the calculation of the electronic structure of many-electron atoms does not look too optimistic. The Schr¨ odinger equation is not soluble, so we cannot get expressions like equation (2.1) for their ener-gies. Think about what is involved in calculating the energies and electron distributions in, for example, the oxygen atom with a nucleus (charge +8) and eight negatively-charged electrons: • The electrons are strongly attracted to the nucleus; they have charges of opposite sign. • But the electrons all repel one another because they have identical charges. Each electron’s distribution depends on the distribution of all the others. This looks like a Catch-22. How can we calculate the distribution of any electron if, to do so, we need the distribution of all the others, and the distribution of these others, in turn, depends on the distribution of the electron we started with? Fortunately, there is a solution based on the idea of self-consistency .
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 - PDF- Allan Blackman, Steven E. Bottle, Siegbert Schmid, Mauro Mocerino, Uta Wille(Authors)
- 2022(Publication Date)
- Wiley(Publisher)
As the number of electrons in an atom increases, a listing of all quantum numbers quickly becomes tedious. For example, iron, with 26 electrons, would require the specification of 26 sets of four quantum numbers. To save time and space, chemists have devised a shorthand notation to write electron configura- tions. The orbital symbols (1s, 2p, 4d etc.) are followed by superscripts designating how many electrons are in each set of orbitals. The compact configuration for a hydrogen atom is 1s 1 , indicating 1 electron in the 1s orbital. For an iron atom the compact configuration is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 . The third way to represent an electron configuration uses an energy level diagram to designate orbitals. Orbital energy levels are indicated by a horizontal line, and these are arranged vertically in order of increasing energy. Each electron is represented by an arrow and is placed on the appropriate horizontal line. The direction of the arrow indicates the value of m s . By convention, we fill orbitals starting from the left-hand side, and the arrow points upward for m s = + 1 2 and downward for m s = - 1 2 . The configuration of hydrogen can be represented by a single arrow in a 1s orbital. or 1s 1s A neutral helium atom has two electrons. To write the ground-state electron configuration of He, we apply the Aufbau principle. One unique set of quantum numbers is assigned to each electron, moving from the lowest energy orbital upward until all electrons have been assigned. The lowest energy orbital is always 1s (n = 1, l = 0, m l = 0). Both helium electrons can occupy the 1s orbital, provided one of them has m s = + 1 2 and the other has m s = - 1 2 . Below are the three representations of helium’s ground-state electron configuration. n = 1, l = 0, m l = 0, m s = + 1 2 n = 1, l = 0, m l = 0, m s = - 1 2 1s 2 1s The two electrons in this configuration are said to be paired electrons, meaning that they are in the same orbital, with opposing spins.
eBook - PDFChemistry
Principles and Reactions
- William Masterton, Cecile Hurley(Authors)
- 2020(Publication Date)
- Cengage Learning EMEA(Publisher)
As n in-creases, the radius of the orbital becomes larger (Figure 6.11). This means that an electron in a 2s orbital is more likely to be found far out from the nucleus than is a 1s electron. The shapes and orientations of p orbitals are shown in Figure 6.12. Notice that ■ a p orbital consists of two lobes along an axis ( x, y, or z ). Among other things, this means that, in a p orbital, there is zero probability of finding an electron at the origin, that is, at the nucleus of the atom. ■ the three p orbitals in a given sublevel are oriented at right angles to one an-other along the x-, y-, and z -axis. For that reason, the three orbitals are often designated as p x , p y , and p z . Although it is not shown in Figure 6.12, p orbitals, like s orbitals, increase in size as the principal quantum number n increases. Also not shown are the shapes and sizes of d and f orbitals. We will say more about the nature of d orbitals in Chapter 19. 6-5 Electron Configurations in Atoms Given the rules referred to in Section 6-3, it is possible to assign quantum num-bers to each electron in an atom. Beyond that, electrons can be assigned to specific principal levels, sublevels, and orbitals. There are several ways to do this. Perhaps the simplest way to describe the arrangement of electrons in an atom is to give its electron configuration , which shows the number of electrons, indicated by a superscript, in each sublevel. For example, a species with the electron configuration 1s 2 2s 2 2p 5 ▼ has two electrons in the 1s sublevel, two electrons in the 2s sublevel, and five electrons in the 2p sublevel. 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.
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