# Electron shell

In chemistry and atomic physics, an electron shell, or a principal energy level, may be thought of as an orbit followed by electrons around an atom's nucleus. The closest shell to the nucleus is called the "1 shell" (also called "K shell"), followed by the "2 shell" (or "L shell"), then the "3 shell" (or "M shell"), and so on farther and farther from the nucleus. The shells correspond with the principal quantum numbers (n = 1, 2, 3, 4 ...) or are labeled alphabetically with letters used in the X-ray notation (K, L, M, …).

Each shell can contain only a fixed number of electrons: The first shell can hold up to two electrons, the second shell can hold up to eight (2 + 6) electrons, the third shell can hold up to 18 (2 + 6 + 10) and so on. The general formula is that the nth shell can in principle hold up to 2(n2) electrons.[1] Since electrons are electrically attracted to the nucleus, an atom's electrons will generally occupy outer shells only if the more inner shells have already been completely filled by other electrons. However, this is not a strict requirement: atoms may have two or even three incomplete outer shells. (See Madelung rule for more details.) For an explanation of why electrons exist in these shells see electron configuration.[2]

The electrons in the outermost occupied shell (or shells) determine the chemical properties of the atom; it is called the valence shell.

Each shell consists of one or more subshells, and each subshell consists of one or more atomic orbitals.

## History

The shell terminology comes from Arnold Sommerfeld's modification of the Bohr model. Sommerfeld retained Bohr's planetary model, but added mildly elliptical orbits (characterized by additional quantum numbers and m) to explain the fine spectroscopic structure of some elements.[3] The multiple electrons with the same principal quantum number (n) had close orbits that formed a "shell" of positive thickness instead of the infinitely thin circular orbit of Bohr's model.

The existence of electron shells was first observed experimentally in Charles Barkla's and Henry Moseley's X-ray absorption studies. Barkla labeled them with the letters K, L, M, N, O, P, and Q.[4] The origin of this terminology was alphabetic. A "J" series was also suspected, though later experiments indicated that the K absorption lines are produced by the innermost electrons. These letters were later found to correspond to the n values 1, 2, 3, etc. They are used in the spectroscopic Siegbahn notation.

The physical chemist Gilbert Lewis was responsible for much of the early development of the theory of the participation of valence shell electrons in chemical bonding. Linus Pauling later generalized and extended the theory while applying insights from quantum mechanics.

## Shells

The electron shells are labeled K, L, M, N, O, P, and Q; or 1, 2, 3, 4, 5, 6, and 7; going from innermost shell outwards. Electrons in outer shells have higher average energy and travel farther from the nucleus than those in inner shells. This makes them more important in determining how the atom reacts chemically and behaves as a conductor, because the pull of the atom's nucleus upon them is weaker and more easily broken. In this way, a given element's reactivity is highly dependent upon its electronic configuration.

## Subshells

Each shell is composed of one or more subshells, which are themselves composed of atomic orbitals. For example, the first (K) shell has one subshell, called 1s; the second (L) shell has two subshells, called 2s and 2p; the third shell has 3s, 3p, and 3d; the fourth shell has 4s, 4p, 4d and 4f; the fifth shell has 5s, 5p, 5d, and 5f and can theoretically hold more in the 5g subshell that is not occupied in the ground-state electron configuration of any known element.[2] The various possible subshells are shown in the following table:

Subshell label Max electrons Shells containing it Historical name
s 0 2 Every shell  sharp
p 1 6 2nd shell and higher  principal
d 2 10 3rd shell and higher  diffuse
f 3 14 4th shell and higher  fundamental
g 4 18 5th shell and higher (theoretically) (next in alphabet after f, excluding j)[5]
• The first column is the "subshell label", a lowercase-letter label for the type of subshell. For example, the "4s subshell" is a subshell of the fourth (N) shell, with the type (s) described in the first row.
• The second column is the azimuthal quantum number (ℓ) of the subshell. The precise definition involves quantum mechanics, but it is a number that characterizes the subshell.
• The third column is the maximum number of electrons that can be put into a subshell of that type. For example, the top row says that each s-type subshell (1s, 2s, etc.) can have at most two electrons in it. In each case the figure is 4 greater than the one above it.
• The fourth column says which shells have a subshell of that type. For example, looking at the top two rows, every shell has an s subshell, while only the second shell and higher have a p subshell (i.e., there is no "1p" subshell).
• The final column gives the historical origin of the labels s, p, d, and f. They come from early studies of atomic spectral lines. The other labels, namely g, h and i, are an alphabetic continuation following the last historically originated label of f.

Although it is commonly stated that all the electrons in a shell have the same energy, this is an approximation. However, the electrons in one subshell do have exactly the same level of energy,[6] with later subshells having more energy per electron than earlier ones. This effect is great enough that the energy ranges associated with shells can overlap (see valence shells and Aufbau principle).

## Number of electrons in each shell

Shells and subshells. 1 rectangular triangle (1/2 of a cell) = 1 electron on the level. Red color indicates sublevel s; orange - p; yellow - d; green - f; blue - g; indigo - h; violet - i
Shell
name
Subshell
name
Subshell
max
electrons
Shell
max
electrons
K 1s 2 2
L 2s 2 2 + 6 = 8
2p 6
M 3s 2 2 + 6 + 10
= 18
3p 6
3d 10
N 4s 2 2 + 6 +
10 + 14
= 32
4p 6
4d 10
4f 14
O 5s 2 2 + 6 +
10 + 14 +
18 = 50
5p 6
5d 10
5f 14
5g 18

Each subshell is constrained to hold 4 + 2 electrons at most, namely:

• Each s subshell holds at most 2 electrons
• Each p subshell holds at most 6 electrons
• Each d subshell holds at most 10 electrons
• Each f subshell holds at most 14 electrons
• Each g subshell holds at most 18 electrons

Therefore, the K shell, which contains only an s subshell, can hold up to 2 electrons; the L shell, which contains an s and a p, can hold up to 2 + 6 = 8 electrons, and so forth; in general, the nth shell can hold up to 2n2 electrons.[1]

Although that formula gives the maximum in principle, in fact that maximum is only achieved (by known elements) for the first four shells (K, L, M, N). No known element has more than 32 electrons in any one shell.[7][8] This is because the subshells are filled according to the Aufbau principle. The first elements to have more than 32 electrons in one shell would belong to the g-block of period 8 of the periodic table. These elements would have some electrons in their 5g subshell and thus have more than 32 electrons in the O shell (fifth principal shell).

## Valence shell

The valence shell is the outermost shell of an atom. Valence electrons in non-transition metal elements reside in this shell. Such elements with complete valence shells (noble gases) are the most chemically non-reactive, while those with only one electron in their valence shells (alkali metals) or just missing one electron from having a complete shell (halogens) are the most reactive.[9]

However, this terminology is somewhat misleading in the case of transition metals. In these elements, a valence electron can also be in an inner shell. Thus, the electrons that determine how an atom reacts chemically are those that travel farthest from the nucleus, that is, those with the highest energy, and not necessarily in the valence shell.

## List of elements with electrons per shell

The list below gives the elements arranged by increasing atomic number and shows the number of electrons per shell. At a glance, one can see that subsets of the list show obvious patterns. In particular, the seven elements (in   electric blue) before a noble gas (group 18, in   yellow) higher than helium have the number of electrons in the valence shell in arithmetic progression. (However, this pattern may break down in the seventh period due to relativistic effects.)

Sorting the table by chemical group shows additional patterns, especially with respect to the last two outermost shells. (Elements 57 to 71 belong to the lanthanides, while 89 to 103 are the actinides.)

The list below is primarily consistent with the Aufbau principle. However, there are a number of exceptions to the rule; for example palladium (atomic number 46) has no electrons in the fifth shell, unlike other atoms with lower atomic number. Some entries in the table are uncertain, when experimental data is unavailable. (For example, the elements past 108 have such short half-lives that their electron configurations have not yet been measured.)

Z Element No. of electrons/shell Group
1 Hydrogen 1 1
2 Helium 2 18
3 Lithium 2, 1 1
4 Beryllium 2, 2 2
5 Boron 2, 3 13
6 Carbon 2, 4 14
7 Nitrogen 2, 5 15
8 Oxygen 2, 6 16
9 Fluorine 2, 7 17
10 Neon 2, 8 18
11 Sodium 2, 8, 1 1
12 Magnesium 2, 8, 2 2
13 Aluminium 2, 8, 3 13
14 Silicon 2, 8, 4 14
15 Phosphorus 2, 8, 5 15
16 Sulfur 2, 8, 6 16
17 Chlorine 2, 8, 7 17
18 Argon 2, 8, 8 18
19 Potassium 2, 8, 8, 1 1
20 Calcium 2, 8, 8, 2 2
21 Scandium 2, 8, 9, 2 3
22 Titanium 2, 8, 10, 2 4
23 Vanadium 2, 8, 11, 2 5
24 Chromium 2, 8, 13, 1 6
25 Manganese 2, 8, 13, 2 7
26 Iron 2, 8, 14, 2 8
27 Cobalt 2, 8, 15, 2 9
28 Nickel 2, 8, 16, 2 10
29 Copper 2, 8, 18, 1 11
30 Zinc 2, 8, 18, 2 12
31 Gallium 2, 8, 18, 3 13
32 Germanium 2, 8, 18, 4 14
33 Arsenic 2, 8, 18, 5 15
34 Selenium 2, 8, 18, 6 16
35 Bromine 2, 8, 18, 7 17
36 Krypton 2, 8, 18, 8 18
37 Rubidium 2, 8, 18, 8, 1 1
38 Strontium 2, 8, 18, 8, 2 2
39 Yttrium 2, 8, 18, 9, 2 3
40 Zirconium 2, 8, 18, 10, 2 4
41 Niobium 2, 8, 18, 12, 1 5
42 Molybdenum 2, 8, 18, 13, 1 6
43 Technetium 2, 8, 18, 13, 2 7
44 Ruthenium 2, 8, 18, 15, 1 8
45 Rhodium 2, 8, 18, 16, 1 9
46 Palladium 2, 8, 18, 18 10
47 Silver 2, 8, 18, 18, 1 11
48 Cadmium 2, 8, 18, 18, 2 12
49 Indium 2, 8, 18, 18, 3 13
50 Tin 2, 8, 18, 18, 4 14
51 Antimony 2, 8, 18, 18, 5 15
52 Tellurium 2, 8, 18, 18, 6 16
53 Iodine 2, 8, 18, 18, 7 17
54 Xenon 2, 8, 18, 18, 8 18
55 Caesium 2, 8, 18, 18, 8, 1 1
56 Barium 2, 8, 18, 18, 8, 2 2
57 Lanthanum 2, 8, 18, 18, 9, 2 3
58 Cerium 2, 8, 18, 19, 9, 2
59 Praseodymium 2, 8, 18, 21, 8, 2
60 Neodymium 2, 8, 18, 22, 8, 2
61 Promethium 2, 8, 18, 23, 8, 2
62 Samarium 2, 8, 18, 24, 8, 2
63 Europium 2, 8, 18, 25, 8, 2
64 Gadolinium 2, 8, 18, 25, 9, 2
65 Terbium 2, 8, 18, 27, 8, 2
66 Dysprosium 2, 8, 18, 28, 8, 2
67 Holmium 2, 8, 18, 29, 8, 2
68 Erbium 2, 8, 18, 30, 8, 2
69 Thulium 2, 8, 18, 31, 8, 2
70 Ytterbium 2, 8, 18, 32, 8, 2
71 Lutetium 2, 8, 18, 32, 9, 2
72 Hafnium 2, 8, 18, 32, 10, 2 4
73 Tantalum 2, 8, 18, 32, 11, 2 5
74 Tungsten 2, 8, 18, 32, 12, 2 6
75 Rhenium 2, 8, 18, 32, 13, 2 7
76 Osmium 2, 8, 18, 32, 14, 2 8
77 Iridium 2, 8, 18, 32, 15, 2 9
78 Platinum 2, 8, 18, 32, 17, 1 10
79 Gold 2, 8, 18, 32, 18, 1 11
80 Mercury 2, 8, 18, 32, 18, 2 12
81 Thallium 2, 8, 18, 32, 18, 3 13
82 Lead 2, 8, 18, 32, 18, 4 14
83 Bismuth 2, 8, 18, 32, 18, 5 15
84 Polonium 2, 8, 18, 32, 18, 6 16
85 Astatine 2, 8, 18, 32, 18, 7 17
86 Radon 2, 8, 18, 32, 18, 8 18
87 Francium 2, 8, 18, 32, 18, 8, 1 1
88 Radium 2, 8, 18, 32, 18, 8, 2 2
89 Actinium 2, 8, 18, 32, 18, 9, 2 3
90 Thorium 2, 8, 18, 32, 18, 10, 2
91 Protactinium 2, 8, 18, 32, 20, 9, 2
92 Uranium 2, 8, 18, 32, 21, 9, 2
93 Neptunium 2, 8, 18, 32, 22, 9, 2
94 Plutonium 2, 8, 18, 32, 24, 8, 2
95 Americium 2, 8, 18, 32, 25, 8, 2
96 Curium 2, 8, 18, 32, 25, 9, 2
97 Berkelium 2, 8, 18, 32, 27, 8, 2
98 Californium 2, 8, 18, 32, 28, 8, 2
99 Einsteinium 2, 8, 18, 32, 29, 8, 2
100 Fermium 2, 8, 18, 32, 30, 8, 2
101 Mendelevium 2, 8, 18, 32, 31, 8, 2
102 Nobelium 2, 8, 18, 32, 32, 8, 2
103 Lawrencium 2, 8, 18, 32, 32, 8, 3
104 Rutherfordium 2, 8, 18, 32, 32, 10, 2 4
105 Dubnium 2, 8, 18, 32, 32, 11, 2 5
106 Seaborgium 2, 8, 18, 32, 32, 12, 2 6
107 Bohrium 2, 8, 18, 32, 32, 13, 2 7
108 Hassium 2, 8, 18, 32, 32, 14, 2 8
109 Meitnerium 2, 8, 18, 32, 32, 15, 2 (?) 9
110 Darmstadtium 2, 8, 18, 32, 32, 16, 2 (?) 10
111 Roentgenium 2, 8, 18, 32, 32, 17, 2 (?) 11
112 Copernicium 2, 8, 18, 32, 32, 18, 2 (?) 12
113 Nihonium 2, 8, 18, 32, 32, 18, 3 (?) 13
114 Flerovium 2, 8, 18, 32, 32, 18, 4 (?) 14
115 Moscovium 2, 8, 18, 32, 32, 18, 5 (?) 15
116 Livermorium 2, 8, 18, 32, 32, 18, 6 (?) 16
117 Tennessine 2, 8, 18, 32, 32, 18, 7 (?) 17
118 Oganesson 2, 8, 18, 32, 32, 18, 8 (?) 18

## References

1. ^ a b Re: Why do electron shells have set limits ? madsci.org, 17 March 1999, Dan Berger, Faculty Chemistry/Science, Bluffton College
2. ^ a b Electron Subshells. Corrosion Source.
3. ^ Donald Sadoway, Introduction to Solid State Chemistry, Lecture 5
4. ^ Barkla, Charles G. (1911). "XXXIX.The spectra of the fluorescent Röntgen radiations". Philosophical Magazine. Series 6. 22 (129): 396. doi:10.1080/14786440908637137. Previously denoted by letters B and A (...). The letters K and L are, however, preferable, as it is highly probable that series of radiations both more absorbable and more penetrating exist.
5. ^ Jue, T. (2009). "Quantum Mechanic Basic to Biophysical Methods". Fundamental Concepts in Biophysics. Berlin: Springer. p. 33. ISBN 1-58829-973-2.
6. ^ The statement that the electrons in one subshell have exactly the same level of energy is true in an isolated atom, where it follows quantum-mechanically from the spherical symmetry of the system. When the atom is part of a molecule, this no longer holds; see, for example, crystal field theory.
7. ^ Orbitals. Chem4Kids. Retrieved on 1 December 2011.
8. ^ Electron & Shell Configuration. Chemistry.patent-invent.com. Retrieved on 1 December 2011.
9. ^ Chemical Reactions. Vision Learning (26 July 2011). Retrieved on 1 December 2011.
Actinide chemistry

Actinide chemistry (or actinoid chemistry) is one of the main branches of nuclear chemistry that investigates the processes and molecular systems of the actinides. The actinides derive their name from the group 3 element actinium. The informal chemical symbol An is used in general discussions of actinide chemistry to refer to any actinide. All but one of the actinides are f-block elements, corresponding to the filling of the 5f electron shell; lawrencium, a d-block element, is also generally considered an actinide. In comparison with the lanthanides, also mostly f-block elements, the actinides show much more variable valence. The actinide series encompasses the 15 metallic chemical elements with atomic numbers from 89 to 103, actinium through lawrencium.

Alkaline earth metal

The alkaline earth metals are six chemical elements in group 2 of the periodic table. They are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). The elements have very similar properties: they are all shiny, silvery-white, somewhat reactive metals at standard temperature and pressure.Structurally, they have in common an outer s- electron shell which is full;

that is, this orbital contains its full complement of two electrons, which these elements readily lose to form cations with charge +2, and an oxidation state of +2.All the discovered alkaline earth metals occur in nature, although radium occurs only through the decay chain of uranium and thorium and not as a primordial element. Experiments have been conducted to attempt the synthesis of element 120, the next potential member of the group, but they have all met with failure.

Atomic radii of the elements (data page)

The atomic radius of a chemical element is the distance from the centre of the nucleus to the outermost shell of the electron. Since the boundary is not a well-defined physical entity, there are various non-equivalent definitions of atomic radius. Depending on the definition, the term may apply only to isolated atoms, or also to atoms in condensed matter, covalently bound in molecules, or in ionized and excited states; and its value may be obtained through experimental measurements, or computed from theoretical models. Under some definitions, the value of the radius may depend on the atom's state and context.Atomic radii vary in a predictable and explicable manner across the periodic table. For instance, the radii generally decrease rightward along each period (row) of the table, from the alkali metals to the noble gases; and increase down each group (column). The radius increases sharply between the noble gas at the end of each period and the alkali metal at the beginning of the next period. These trends of the atomic radii (and of various other chemical and physical properties of the elements) can be explained by the electron shell theory of the atom; they provided important evidence for the development and confirmation of quantum theory.

The atomic radius of a chemical element is a measure of the size of its atoms, usually the mean or typical distance from the center of the nucleus to the boundary of the surrounding shells of electrons. Since the boundary is not a well-defined physical entity, there are various non-equivalent definitions of atomic radius. Three widely used definitions of atomic radius are: Van der Waals radius, ionic radius, and covalent radius.

Depending on the definition, the term may apply only to isolated atoms, or also to atoms in condensed matter, covalently bonding in molecules, or in ionized and excited states; and its value may be obtained through experimental measurements, or computed from theoretical models. The value of the radius may depend on the atom's state and context.Electrons do not have definite orbits, or sharply defined ranges. Rather, their positions must be described as probability distributions that taper off gradually as one moves away from the nucleus, without a sharp cutoff. Moreover, in condensed matter and molecules, the electron clouds of the atoms usually overlap to some extent, and some of the electrons may roam over a large region encompassing two or more atoms.

Under most definitions the radii of isolated neutral atoms range between 30 and 300 pm (trillionths of a meter), or between 0.3 and 3 ångströms. Therefore, the radius of an atom is more than 10,000 times the radius of its nucleus (1–10 fm), and less than 1/1000 of the wavelength of visible light (400–700 nm).

For many purposes, atoms can be modeled as spheres. This is only a crude approximation, but it can provide quantitative explanations and predictions for many phenomena, such as the density of liquids and solids, the diffusion of fluids through molecular sieves, the arrangement of atoms and ions in crystals, and the size and shape of molecules.Atomic radii vary in a predictable and explicable manner across the periodic table. For instance, the radii generally decrease along each period (row) of the table, from the alkali metals to the noble gases; and increase down each group (column). The radius increases sharply between the noble gas at the end of each period and the alkali metal at the beginning of the next period. These trends of the atomic radii (and of various other chemical and physical properties of the elements) can be explained by the electron shell theory of the atom; they provided important evidence for the development and confirmation of quantum theory. The atomic radii decrease across the Periodic Table because as the atomic number increases, the number of protons increases across the period, but the extra electrons are only added to the same quantum shell. Therefore, the effective nuclear charge towards the outermost electrons increases, drawing the outermost electrons closer. As a result, the electron cloud contracts and the atomic radius decreases.

Auger effect

The Auger effect is a physical phenomenon in which the filling of an inner-shell vacancy of an atom is accompanied by the emission of an electron from the same atom. When a core electron is removed, leaving a vacancy, an electron from a higher energy level may fall into the vacancy, resulting in a release of energy. Although most often this energy is released in the form of an emitted photon, the energy can also be transferred to another electron, which is ejected from the atom; this second ejected electron is called an Auger electron. The effect was first discovered by Lise Meitner in 1922; Pierre Victor Auger independently discovered the effect shortly after and is credited with the discovery in most of the scientific community.Upon ejection, the kinetic energy of the Auger electron corresponds to the difference between the energy of the initial electronic transition into the vacancy and the ionization energy for the electron shell from which the Auger electron was ejected. These energy levels depend on the type of atom and the chemical environment in which the atom was located.

Auger electron spectroscopy involves the emission of Auger electrons by bombarding a sample with either X-rays or energetic electrons and measures the intensity of Auger electrons that result as a function of the Auger electron energy. The resulting spectra can be used to determine the identity of the emitting atoms and some information about their environment.

Auger recombination is a similar Auger effect which occurs in semiconductors. An electron and electron hole (electron-hole pair) can recombine giving up their energy to an electron in the conduction band, increasing its energy. The reverse effect is known as impact ionization.

The Auger effect can impact biological molecules such as DNA. Following the K-shell ionization of the component atoms of DNA, Auger electrons are ejected leading to damage of its sugar-phosphate backbone.

Bent molecular geometry

In chemistry, the term "bent" can be applied to certain molecules to describe their molecular geometry. Certain atoms, such as oxygen, will almost always set their two (or more) covalent bonds in non-collinear directions due to their electron configuration. Water (H2O) is an example of a bent molecule, as well as its analogues. The bond angle between the two hydrogen atoms is approximately 104.45°. Nonlinear geometry is commonly observed for other triatomic molecules and ions containing only main group elements, prominent examples being nitrogen dioxide (NO2), sulfur dichloride (SCl2), and methylene (CH2).

This geometry is almost always consistent with VSEPR theory, which usually explains non-collinearity of atoms with a presence of lone pairs. There are several variants of bending, where the most common is AX2E2 where two covalent bonds and two lone pairs of the central atom (A) form a complete 8-electron shell. They have central angles from 104° to 109.5°, where the latter is consistent with a simplistic theory which predicts the tetrahedral symmetry of four sp3 hybridised orbitals. The most common actual angles are 105°, 107°, and 109°: they vary because of the different properties of the peripheral atoms (X).

Other cases also experience orbital hybridisation, but in different degrees. AX2E1 molecules, such as SnCl2, have only one lone pair and the central angle about 120° (the centre and two vertices of an equilateral triangle). They have three sp2 orbitals. There exist also sd-hybridised AX2 compounds of transition metals without lone pairs: they have the central angle about 90° and are also classified as bent.

Electron capture

Electron capture (K-electron capture, also K-capture, or L-electron capture, L-capture) is a process in which the proton-rich nucleus of an electrically neutral atom absorbs an inner atomic electron, usually from the K or L electron shell. This process thereby changes a nuclear proton to a neutron and simultaneously causes the emission of an electron neutrino.

p + e− → n + νeSince this single emitted neutrino carries the entire decay energy, it has this single characteristic energy. Similarly, the momentum of the neutrino emission causes the daughter atom to recoil with a single characteristic momentum.

The resulting daughter nuclide, if it is in an excited state, then transitions to its ground state. Usually, a gamma ray is emitted during this transition, but nuclear de-excitation may also take place by internal conversion.

Following capture of an inner electron from the atom, an outer electron replaces the electron that was captured and one or more characteristic X-ray photons is emitted in this process. Electron capture sometimes also results in the Auger effect, where an electron is ejected from the atom's electron shell due to interactions between the atom's electrons in the process of seeking a lower energy electron state.

Following electron capture, the atomic number is reduced by one, the neutron number is increased by one, and there is no change in mass number. Simple electron capture by itself results in a neutral atom, since the loss of the electron in the electron shell is balanced by a loss of positive nuclear charge. However, a positive atomic ion may result from further Auger electron emission.

Electron capture is an example of weak interaction, one of the four fundamental forces.

Electron capture is the primary decay mode for isotopes with a relative superabundance of protons in the nucleus, but with insufficient energy difference between the isotope and its prospective daughter (the isobar with one less positive charge) for the nuclide to decay by emitting a positron. Electron capture is always an alternative decay mode for radioactive isotopes that do not have sufficient energy to decay by positron emission.

Electron capture is sometimes included as a type of beta decay, because the basic nuclear process, mediated by the weak force, is the same. In nuclear physics, beta decay is a type of radioactive decay in which a beta ray (fast energetic electron or positron) and a neutrino are emitted from an atomic nucleus.

Electron capture is sometimes called inverse beta decay, though this term usually refers to the interaction of an electron antineutrino with a proton.If the energy difference between the parent atom and the daughter atom is less than 1.022 MeV, positron emission is forbidden as not enough decay energy is available to allow it, and thus electron capture is the sole decay mode. For example, rubidium-83 (37 protons, 46 neutrons) will decay to krypton-83 (36 protons, 47 neutrons) solely by electron capture (the energy difference, or decay energy, is about 0.9 MeV).

Electron configuration

In atomic physics and quantum chemistry, 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.

Electronic configurations describe each electron as moving independently in an orbital, in an average field created by all other orbitals. Mathematically, configurations are described by Slater determinants or configuration state functions.

According to the laws of quantum mechanics, for systems with only one electron, a level of energy is associated with each electron configuration and in certain conditions, electrons are able to move from one configuration to another by the emission or absorption of a quantum of energy, in the form of a photon.

Knowledge of the electron configuration of different atoms is useful in understanding the structure of the periodic table of elements. This is also useful for describing the chemical bonds that hold atoms together. In bulk materials, this same idea helps explain the peculiar properties of lasers and semiconductors.

Energy level

A quantum mechanical system or particle that is bound—that is, confined spatially—can only take on certain discrete values of energy. This contrasts with classical particles, which can have any energy. These discrete values are called energy levels. The term is commonly used for the energy levels of electrons in atoms, ions, or molecules, which are bound by the electric field of the nucleus, but can also refer to energy levels of nuclei or vibrational or rotational energy levels in molecules. The energy spectrum of a system with such discrete energy levels is said to be quantized.

In chemistry and atomic physics, an electron shell, or a principal energy level, may be thought of as an orbit followed by electrons around an atom's nucleus. The closest shell to the nucleus is called the "1 shell" (also called "K shell"), followed by the "2 shell" (or "L shell"), then the "3 shell" (or "M shell"), and so on farther and farther from the nucleus. The shells correspond with the principal quantum numbers (n = 1, 2, 3, 4 ...) or are labeled alphabetically with letters used in the X-ray notation (K, L, M, …).

Each shell can contain only a fixed number of electrons: The first shell can hold up to two electrons, the second shell can hold up to eight (2 + 6) electrons, the third shell can hold up to 18 (2 + 6 + 10) and so on. The general formula is that the nth shell can in principle hold up to 2(n2) electrons. Since electrons are electrically attracted to the nucleus, an atom's electrons will generally occupy outer shells only if the more inner shells have already been completely filled by other electrons. However, this is not a strict requirement: atoms may have two or even three incomplete outer shells. (See Madelung rule for more details.) For an explanation of why electrons exist in these shells see electron configuration.If the potential energy is set to zero at infinite distance from the atomic nucleus or molecule, the usual convention, then bound electron states have negative potential energy.

If an atom, ion, or molecule is at the lowest possible energy level, it and its electrons are said to be in the ground state. If it is at a higher energy level, it is said to be excited, or any electrons that have higher energy than the ground state are excited. If more than one quantum mechanical state is at the same energy, the energy levels are "degenerate". They are then called degenerate energy levels.

Helium-4

Helium-4 (42He or 4He) is a non-radioactive isotope of the element helium. It is by far the most abundant of the two naturally occurring isotopes of helium, making up about 99.99986% of the helium on Earth. Its nucleus is identical to an alpha particle, and consists of two protons and two neutrons.

Alpha decay of heavy elements in the Earth's crust is the source of most naturally occurring helium-4 on Earth. While it is also produced by nuclear fusion in stars, most helium-4 in the Sun and in the universe is thought to have been produced by the Big Bang, and is referred to as "primordial helium". However, primordial helium-4 is largely absent from the Earth, having escaped during the high-temperature phase of Earth's formation. Radioactive decay from other elements is the source of most of the helium-4 found on Earth, produced after the planet cooled and solidified.

Helium-4 makes up about one quarter of the ordinary matter in the universe by mass, with almost all of the rest being hydrogen.

When liquid helium-4 is cooled to below 2.17 kelvins (–271.17 °C), it becomes a superfluid, with properties that are very unlike those of an ordinary liquid. For example, if superfluid helium-4 is kept in an open vessel, a thin film will climb up the sides of the vessel and overflow. In this state and situation, it is called a "Rollin film". This strange behavior is a result of the Clausius–Clapeyron relation and cannot be explained by the current model of classical mechanics, nor by nuclear or electrical models – it can only be understood as a quantum-mechanical phenomenon. The total spin of the helium-4 nucleus is an integer (zero), and therefore it is a boson (as are neutral atoms of helium-4). The superfluid behavior is now understood to be a manifestation of Bose–Einstein condensation, which occurs only with collections of bosons.

It is theorized that, at 0.2 K and 50 atm, solid helium-4 may be a superglass (an amorphous solid exhibiting superfluidity).

Helium-4 also exists on the Moon and—as on Earth—it is the most abundant helium isotope.

Inert gas

An inert gas is a gas which does not undergo chemical reactions under a set of given conditions. The noble gases often do not react with many

substances, and were historically referred to as the inert gases. Inert gases are used generally to avoid unwanted chemical reactions degrading a sample. These undesirable chemical reactions are often oxidation and hydrolysis reactions with the oxygen and moisture in air. The term inert gas is context-dependent because several of the noble gases can be made to react under certain conditions.

Purified argon and nitrogen gases are most commonly used as inert gases due to their high natural abundance (78.2% N2, 1% Ar in air) and low relative cost.

Unlike noble gases, an inert gas is not necessarily elemental and is often a compound gas. Like the noble gases the tendency for non-reactivity is due to the valence, the outermost electron shell, being complete in all the inert gases. This is a tendency, not a rule, as noble gases and other "inert" gases can react to form compounds.

Internal conversion

Internal conversion is a radioactive decay process wherein an excited nucleus interacts electromagnetically with one of the orbital electrons of the atom. This causes the electron to be emitted (ejected) from the atom. Thus, in an internal conversion process, a high-energy electron is emitted from the radioactive atom, but not from the nucleus. For this reason, the high-speed electrons resulting from internal conversion are not called beta particles, since the latter come from beta decay, where they are newly created in the nuclear decay process.

Internal conversion is possible whenever gamma decay is possible, except in the case where the atom is fully ionised. During internal conversion, the atomic number does not change, and thus (as is the case with gamma decay) no transmutation of one element to another takes place.

Since an electron is lost from the atom, a hole appears in an electron shell which is subsequently filled by other electrons that descend to that empty, lower energy level, and in the process emit characteristic X-ray(s), Auger electron(s), or both. The atom thus emits high-energy electrons and X-ray photons, none of which originate in that nucleus. The atom supplied the energy needed to eject the electron, which in turn caused the latter events and the other emissions.

Since primary electrons from internal conversion carry a fixed (large) part of the characteristic decay energy, they have a discrete energy spectrum, rather than the spread (continuous) spectrum characteristic of beta particles. Whereas the energy spectrum of beta particles plots as a broad hump, the energy spectrum of internally converted electrons plots as a single sharp peak (see example below).

Ionization energy

In physics and chemistry, ionization energy (American English spelling) or ionisation energy (British English spelling), denoted Ei, is the minimum amount of energy required to remove the most loosely bound electron, the valence electron, of an isolated neutral gaseous atom or molecule. It is quantitatively expressed as

X + energy → X+ + e−where X is any atom or molecule capable of ionization, X+ is that atom or molecule with an electron removed, and e− is the removed electron. This is generally an endothermic process. Generally, the closer the outermost electrons are to the nucleus of the atom , the higher the atom's or element's ionization energy.

The sciences of physics and chemistry use different measures of ionization energy. In physics, the unit is the amount of energy required to remove a single electron from a single atom or molecule, expressed as electronvolts. In chemistry, the unit is the amount of energy required for all of the atoms in a mole of substance to lose one electron each: molar ionization energy or enthalpy, expressed as kilojoules per mole (kJ/mol) or kilocalories per mole (kcal/mol).Comparison of Ei of elements in the periodic table reveals two periodic trends:

Ei generally increases as one moves from left to right within a given period (that is, row).

Ei generally decreases as one moves from top to bottom in a given group (that is, column).The latter trend results from the outer electron shell being progressively farther from the nucleus, with the addition of one inner shell per row as one moves down the column.

The nth ionization energy refers to the amount of energy required to remove an electron from the species with a charge of (n-1). For example, the first three ionization energies are defined as follows:

1st ionization energy

X → X+ + e−2nd ionization energy

X+ → X2+ + e−3rd ionization energy

X2+ → X3+ + e−The term ionization potential is an older name for ionization energy, because the oldest method of measuring ionization energy was based on ionizing a sample and accelerating the electron removed using an electrostatic potential. However this term is now considered obsolete.

Some factors affecting the ionization energy include:

Nuclear charge: the greater the magnitude of nuclear charge the more tightly the electrons are held by the nucleus and hence more will be ionization energy.

Number of electron shells: the greater the size of the atom less tightly the electrons are held by the nucleus and ionization energy will be less.

Effective nuclear charge (Zeff): the greater the magnitude of electron shielding and penetration the less tightly the electrons are held by the nucleus, the lower the Zeff of the electron, and hence less will be the ionization energy.

Type of orbital ionized: the atom having a more stable electronic configuration has less tendency to lose electrons and consequently has high ionization energy.

Occupancy of the orbital matters: if the orbital is half or completely filled then it is harder to remove electrons

Kinetic diameter

Kinetic diameter is a measure applied to atoms and molecules that expresses the likelihood that a molecule in a gas will collide with another molecule. It is an indication of the size of the molecule as a target. The kinetic diameter is not the same as atomic diameter defined in terms of the size of the atom's electron shell, which is generally a lot smaller, depending on the exact definition used. Rather, it is the size of the sphere of influence that can lead to a scattering event.

Kinetic diameter is related to the mean free path of molecules in a gas. Mean free path is the average distance that a particle will travel without collision. For a fast moving particle (that is, one moving much faster than the particles it is moving through) the kinetic diameter is given by,

${\displaystyle d^{2}={1 \over \pi ln}}$
where,
d is the kinetic diameter,
r is the kinetic radius, r = d/2,
l is the mean free path, and
n is the number density of particles

However, a more usual situation is that the colliding particle being considered is indistinguishable from the population of particles in general. Here, the Maxwell–Boltzmann distribution of energies must be considered, which leads to the modified expression,

${\displaystyle d^{2}={1 \over {\sqrt {2}}\pi ln}}$
Lone pair

In chemistry, a lone pair refers to a pair of valence electrons that are not shared with another atom and is sometimes called an unshared pair or non-bonding pair. Lone pairs are found in the outermost electron shell of atoms. They can be identified by using a Lewis structure. Electron pairs are therefore considered lone pairs if two electrons are paired but are not used in chemical bonding. Thus, the number of lone pair electrons plus the number of bonding electrons equals the total number of valence electrons around an atom.

Lone pair is a concept used in valence shell electron pair repulsion theory (VSEPR theory) which explains the shapes of molecules. They are also referred to in the chemistry of Lewis acids and bases. However, not all non-bonding pairs of electrons are considered by chemists to be lone pairs. Examples are the transition metals where the non-bonding pairs do not influence molecular geometry and are said to be stereochemically inactive. In molecular orbital theory (fully delocalized or otherwise), the concept of a lone pair is less distinct, but occupied non-bonding orbitals (or mostly nonbonding) are frequently regarded as "lone pairs" as well.

A single lone pair can be found with atoms in the nitrogen group such as nitrogen in ammonia, two lone pairs can be found with atoms in the chalcogen group such as oxygen in water and the halogens can carry three lone pairs such as in hydrogen chloride.

In VSEPR theory the electron pairs on the oxygen atom in water form the vertices of a tetrahedron with the lone pairs on two of the four vertices. The H–O–H bond angle is 104.5°, less than the 109° predicted for a tetrahedral angle, and this can be explained by a repulsive interaction between the lone pairs.Various computational criteria for the presence of lone pairs have been proposed. While electron density ρ(r) itself generally does not provide useful guidance in this regard, the laplacian of the electron density is revealing, and one criterion for the location of the lone pair is where L(r) = –∇2ρ(r) is a local maximum. The minima of the electrostatic potential V(r) is another proposed criterion. Yet another considers the electron localization function (ELF).

Octet rule

The octet rule is a chemical rule of thumb that reflects observation that atoms of main-group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electron configuration as a noble gas. The rule is especially applicable to carbon, nitrogen, oxygen, and the halogens, but also to metals such as sodium or magnesium.

The valence electrons can be counted using a Lewis electron dot diagram as shown at the right for carbon dioxide. The electrons shared by the two atoms in a covalent bond are counted twice, once for each atom. In carbon dioxide each oxygen shares four electrons with the central carbon, two (shown in red) from the oxygen itself and two (shown in black) from the carbon. All four of these electrons are counted in both the carbon octet and the oxygen octet.

Quantum number

Quantum numbers describe values of conserved quantities in the dynamics of a quantum system. In the case of electrons, the quantum numbers can be defined as "the sets of numerical values which give acceptable solutions to the Schrödinger wave equation for the hydrogen atom". An important aspect of quantum mechanics is the quantization of the observable quantities, since quantum numbers are discrete sets of integers or half-integers, although they could approach infinity in some cases. This distinguishes quantum mechanics from classical mechanics where the values that characterize the system such as mass, charge, or momentum, range continuously. Quantum numbers often describe specifically the energy levels of electrons in atoms, but other possibilities include angular momentum, spin, etc. An important family is flavour quantum numbers – internal quantum numbers which determine the type of a particle and its interactions with other particles through the forces. Any quantum system can have one or more quantum numbers; it is thus difficult to list all possible quantum numbers.

Valence electron

In chemistry, a valence electron is an outer shell electron that is associated with an atom, and that can participate in the formation of a chemical bond if the outer shell is not closed; in a single covalent bond, both atoms in the bond contribute one valence electron in order to form a shared pair. The presence of valence electrons can determine the element's chemical properties, such as its valence—whether it may bond with other elements and, if so, how readily and with how many. For a main group element, a valence electron can exist only in the outermost electron shell; in a transition metal, a valence electron can also be in an inner shell.

An atom with a closed shell of valence electrons (corresponding to an electron configuration s2p6) tends to be chemically inert. Atoms with one or two more valence electrons than are needed for a "closed" shell are highly reactive due to the following reasons:

1) It requires relatively low energy (compared to the lattice enthalpy) to remove the extra valence electrons to form a positive ion.

2) Because of their tendency either to gain the missing valence electrons (thereby forming a negative ion), or to share valence electrons (thereby forming a covalent bond).

Similar to an electron in an inner shell, a valence electron has the ability to absorb or release energy in the form of a photon. An energy gain can trigger an electron to move (jump) to an outer shell; this is known as atomic excitation. Or the electron can even break free from its associated atom's valence shell; this is ionization to form a positive ion. When an electron loses energy (thereby causing a photon to be emitted), then it can move to an inner shell which is not fully occupied.

Valence energy levels correspond to the principal quantum numbers (n = 1, 2, 3, 4, 5 ...) or are labeled alphabetically with letters used in the X-ray notation (K, L, M, …).

Quantum numbers
Periodic table
Principles

This page is based on a Wikipedia article written by authors (here).
Text is available under the CC BY-SA 3.0 license; additional terms may apply.
Images, videos and audio are available under their respective licenses.