what is the klmn atomic number
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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 n th shell can in principle hold up to 2( n 2) 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 ℓ 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 but the 5f subshell, although partially occupied in actinides, is not filled in any element occurring naturally. [2] The various possible subshells are shown in the following table:
Subshell label ℓ Max electrons Shells containing it Historical name
s 0 2 Every shell s harp
p 1 6 2nd shell and higher p rincipal
d 2 10 3rd shell and higher d iffuse
f 3 14 4th shell and higher f undamental
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
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 ℓ 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 but the 5f subshell, although partially occupied in actinides, is not filled in any element occurring naturally. [2] The various possible subshells are shown in the following table:
Subshell label ℓ Max electrons Shells containing it Historical name
s 0 2 Every shell s harp
p 1 6 2nd shell and higher p rincipal
d 2 10 3rd shell and higher d iffuse
f 3 14 4th shell and higher f undamental
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
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