ELECTRON CONFIGURATION: ELECTRON CONFIGURATION

CHEMISTRY
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Electronic Structure of Atoms

Each electron in an atom is described by four different quantum numbers. Three of these quantum numbers (n, l, and m) represent the three dimensions to space in which an electron could be found. A wave function for an electron gives the probability of finding the electron at various points in space. A wave function for an electron in an atom is called an atomic orbital. The fourth quantum number (m_s) refers to a certain magnetic quality called spin.

n-The Principal Quantum Number

The n quantum number relates to the size of the atomic orbital. n can have any positive integer value from 1 to 7. The smaller the n, the lower the energy, the higher the value of n, the higher the energy. In the case of any single-electron atom, or hydrogen atom, n is the only quantum number which determines the energy. The size of an orbital depends on n. The larger the orbital, the larger the value of n. Orbitals of the same quantum state belong the the same shell. To use an analogy for n, why not relate it to the size of a computer, where larger values would represent larger houses.

l-The angular momentum quantum number

l can have any integer value from 0 to 3. This quantum number distinguishes orbitals of a given n value which have different states. Or, the secondary quantum number gives the shape of the orbital so the analogy can be made to the shape of the computer with larger values associated with computers with more components.

M-magnetic quantum number

The third quantum number has to do with the orientation of an orbital in a magnetic field. Because of this, we can relate its values to different directions the computer might be facing.

The final quantum number is the spin quantum number, it describes the spin orientation of an electron.

The electron configuration of an atom is the particular distribution of electrons among available shells. It is described by a notation that lists the subshell symbols, one after another. Each symbol has a subscript on the right giving the number of electrons in that subshell. For example, a configuration of the lithium atom (atomic number 3) with two electrons in the 1s subshell and one electron in the 2s subshell is written 1s²2s¹.

sublevel	orbital	maximum # of electrons
s	1	2
p	3	6
d	5	10
f	7	14

The notation for electron configuration gives the number of electrons in each subshell. The number of electrons in an atom of an element is given by the atomic number of that element.

On the left we have a diagram to show how the orbitals of a subshell are occupied by electrons. On the right there is a diagram for the filling order of electrons in a subshell.

Here are some examples that show how to use the filling order diagram to complete the electron configuration for a certain substance.

Element	# of Electrons in Element	Electron Configuration
He	2	1s²
Li	3	1s²2s¹
Be	4	1s²2s²
O	8	1s²2s²2p⁴
Cl	17	1s²2s²2p⁶3s²3p⁵
K	19	1s²2s²2p⁶3s²3p⁶4s¹

Often times you will be asked to find the electron configuration for something that looks like this:

₅₃I

The 53 denotes the number of electrons in an atom of iodine. You would now proceed to do the electron configuration by looking at the filling order chart.

1s²2s²2p⁶3s²3p⁶4s² 3d¹⁰4p⁶5s²4d¹⁰5p⁵^{......................................................................................................}

Aufbau principle and Madelung rule

The Aufbau principle (from the German Aufbau, "building up, construction") was an important part of Bohr's original concept of electron configuration. It may be stated as:

a maximum of two electrons are put into orbitals in the order of increasing orbital energy: the lowest-energy orbitals are filled before electrons are placed in higher-energy orbitals.

The principle works very well (for the ground states of the atoms) for the first 18 elements, then decreasingly well for the following 100 elements. The modern form of the Aufbau principle describes an order of orbital energies given by Madelung's rule (or Klechkowski's rule). This rule was first stated by Charles Janet in 1929, rediscovered by Erwin Madelung in 1936,and later given a theoretical justification by V.M. Klechkowski

Orbitals are filled in the order of increasing n+l;
Where two orbitals have the same value of n+l, they are filled in order of increasing n.

This gives the following order for filling the orbitals:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p

The Aufbau principle can be applied, in a modified form, to the protons and neutrons in the atomic nucleus, as in the shell model of nuclear physics and nuclear chemistry.

Shortcomings of the Aufbau principle

The Aufbau principle rests on a fundamental postulate that the order of orbital energies is fixed, both for a given element and between different elements: neither of these is true (although they are approximately true enough for the principle to be useful). It considers atomic orbitals as "boxes" of fixed energy into which can be placed two electrons and no more. However the energy of an electron "in" an atomic orbital depends on the energies of all the other electrons of the atom (or ion, or molecule, etc.). There are no "one-electron solutions" for systems of more than one electron, only a set of many-electron solutions which cannot be calculated exactlyHartree–Fock method). (although there are mathematical approximations available, such as the
The fact that the Aufbau principle is based on an approximation can be seen from the fact that there is an almost-fixed filling order at all, that, within a given shell, the s-orbital is always filled before the p-orbitals. In a hydrogen-like atom, which only has one electron, the s-orbital and the p-orbitals of the same shell have exactly the same energy, to a very good approximation in the absence of external electromagnetic fields. (However, in a real hydrogen atom, the energy levels are slightly split by the magnetic field of the nucleus, and by the quantum electrodynamic Lamb shift). effects of the

Ionization of the transition metals

The naive application of the Aufbau principle leads to a well-known paradox (or apparent paradox) in the basic chemistry of the transition metals. Potassium and calcium appear in the periodic table before the transition metals, and have electron configurations [Ar] 4s¹ and [Ar] 4s² respectively, i.e. the 4s-orbital is filled before the 3d-orbital. This is in line with Madelung's rule, as the 4s-orbital has n+l = 4 (n = 4, l = 0) while the 3d-orbital has n+l = 5 (n = 3, l = 2). However, chromium and copper have electron configurations [Ar] 3d⁵ 4s¹ and [Ar] 3d¹⁰ 4s¹ respectively, i.e. one electron has passed from the 4s-orbital to a 3d-orbital to generate a half-filled or filled subshell. In this case, the usual explanation is that "half-filled or completely-filled subshells are particularly stable arrangements of electrons".
The apparent paradox arises when electrons are removed from the transition metal atoms to form ions. The first electrons to be ionized come not from the 3d-orbital, as one would expect if it were "higher in energy", but from the 4s-orbital. The same is true when chemical compounds are formed. Chromium hexacarbonyl can be described as a chromium atom (not ion, it is in the oxidation state 0) surrounded by six carbon monoxide ligands: it is diamagnetic, and the electron configuration of the central chromium atom is described as 3d⁶, i.e. the electron which was in the 4s-orbital in the free atom has passed into a 3d-orbital on forming the compound. This interchange of electrons between 4s and 3d is universal among the first series of the transition metals.
The phenomenon is only paradoxical if it is assumed that the energies of atomic orbitals are fixed and unaffected by the presence of electrons in other orbitals. If that were the case, the 3d-orbital would have the same energy as the 3p-orbital, as it does in hydrogen, yet it clearly doesn't. There is no special reason why the Fe²⁺ ion should have the same electron configuration as the chromium atom, given that iron has two more protons in its nucleus than chromium and that the chemistry of the two species is very different. When care is taken to compare "like with like", the paradox disappears.

Other exceptions to Madelung's rule

There are several more exceptions to Madelung's rule among the heavier elements, and it is more and more difficult to resort to simple explanations such as the stability of half-filled subshells. It is possible to predict most of the exceptions by Hartree–Fock calculations,effects of Special Relativity on the energies of the atomic orbitals, as the inner-shell electrons are moving at speeds approaching the speed of light. In general, these relativistic effects tend to decrease the energy of the s-orbitals in relation to the other atomic orbitals. which are an approximate method for taking account of the effect of the other electrons on orbital energies. ELECTROQUIZ
Quiz 1 - Atomic number from full electronic configurationQuiz 2 - Number of outer shell electrons from full electronic configuration Quiz 3 - Chemical symbol from full electronic configuration Quiz 4 - Name from full electronic configuration Quiz 5 - Full electron configuration from name Quiz 6 - Atomic number from electrons in boxes notationQuiz 7 - Number of outer shell electrons from electrons in boxes notation Quiz 8 - Chemical symbol from electrons in boxes notation Quiz 9 - Name from electrons in boxes notation Quiz 10 - Electrons in boxes notation from name GAMES CORNER ............................................^{ELECTRON CONFIGURATION "PAIR GAMES"}............................................ ELECTRON CONFIGURATION "ORDER GAMES" ............................................

ELECTRON CONFIGURATION SONG

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i HEART... ELECTRON CONFIGURATION
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