Electronic Configuration of Elements

The orbitals of a multi-electron atom are not likely to be quite the same as the hydrogen atom orbitals. For practical purposes, however, the number of orbitals and there shapes in multi-electron cases may be taken to be the same as for the hydrogen orbitals.
    In multi-electron atoms, experimental studies of spectra show that orbitals with the same value of n but different l values have different energies.
    The 3S orbital is lower energy then 3P orbitals which again are of lower energy than 3d orbitals. However, orbitals belonging to a particular type ( P or d or f ) will be of equal energy (degenerate) in an atom or an ion.
    For example, the three P orbitals or the five d orbitals originating from the same n will be degenerate. The separation of orbitals of a major energy level into sub-levels is primarily due to the interaction among the many electrons.
    This interaction leads to the following relative order of the energies of each type of orbitals
1S2S〈2P3S〈3P4S〈3d 〈4P5S〈4d〈5P〈6S〈4f〈5d 6P〈7S〈5f〈6d
Admittedly it is often difficult for the readers to remember the orbital energy diagram. A trivial but distinctly more convenient way is to make giving as,
Electronic Configuration of first thirty six Elements with Pictorial Representation.
Orbital Occupancy Order.
The different orbitals originating from the same principal quantum number n are written in the horizontal lines. Now inclined parallel lines are drawn through the orbitals according to the above picture. Filling up the different orbitals by electrons will follow these lines.
Example:
An element which contains 36 electrons the electronic configuration of this element according to the above diagram is,
1S22S22P63S23P64S23d104P6

The Aufbau or Building Up Principle:

The question that arises now how many electrons can be accommodated per orbital. The answer to this follows from Pauli's Exclusion Principle. 
  • Pauli's Exclusion Principle:
No two electrons in an atom can have the same four quantum numbers.
This principle tells us that in each orbital maximum of two electrons can be allowed. The two electrons have the same three quantum numbers namely the same n, same l, and the same ml. Any conflict with the Pauli Principle can now be avoided if one of the electrons has the spin quantum numbers is (+1/2) and (-1/2).
An alternative statement of Pauli's Exclusion Principle is, 
No more than two electrons can be placed in one and the same orbital. 
When two electrons with opposite spins exist in an orbital, the electrons are said to be paired. These two electrons per orbital are given the maximum accommodation of electrons in an atom.

Capacities Of Electronic Levels:

The total number of electrons for a particular n is given by 2n².
Principle Quantum Number Azimuthal Quantum Number(l) Total No. of Electrons
n l=0 l=1 l=2 l=3 l=4 2n2
n =1 2 2
n=2 2 6 8
n=3 2 6 10 18
n=4 2 6 10 14 32
n=5 2 6 10 14 18 50
To determine the electronic configuration of elements the procedure is to feed electrons in different orbitals obeying certain rules.
The Aufbau or Building up Principle is based on the following Rules: 
  1. Electrons are fed into orbitals in order of increasing energy (increasing n ) until all the electrons have been accommodated.
  2. Electrons will tend to maintain maximum spin. So long orbitals of similar energy are available for occupation electrons will prefer to remain unpaired.
  3. In other words, electrons tend to avoid the same orbital, that is, hate to share space. This rule is known as Hund's Rule of maximum spin Multiplicity.
  4. Spin pairing can occur only when vacant orbitals of similar energy are not available for occupation, and when the next available vacant orbital is of higher energy.
Electronic Configuration of Elements:
  • Electronic Configuration of Elements H to Ne:
Hydrogen (Atomic Number 1) has its only one electron in the 1S orbital and this electronic configuration represented below. the bar on the pictorial representation indicates the orbital and the arrow a single spinning electron in the orbital. In helium (Atomic Number 2) the second electron occupies the 1S orbital since the next 2S orbital is much higher energy.
Obeying Pauli principle the configuration 1S2 represented below.
From the above rule we can represent the electronic configuration from H to Ne:
Hydrogen(H) 1S1
1S
Helium(He) 1S2
1S
Lithium(Li) 1S22S1
1S 2S
Beryllium(Be) 1S22S2
1S 2S
Boron(B) 1S22S12P1


1S 2S 2Px 2Py 2Pz
Carbon(C) 1S22S12P2

1S 2S 2Px 2Py 2Pz
Nitrogen(N) 1S22S12P3
1S 2S 2Px 2Py 2Pz
Oxygen(O) 1S22S12P4
1S 2S 2Px 2Py 2Pz
Fluorine(F) 1S22S12P5
1S 2S 2Px 2Py 2Pz
Neon(Ne) 1S22S12P6
1S 2S 2Px 2Py 2Pz
  • Electronic Configuration of Elements Na to Ar:
After neon, the next available orbital is 3S being followed by 3P. The orbitals are then progressively filled by electrons. 
Thus the electronic configuration of Na to Ar given below.
Sodium(Na) 1S22S22P63S1
1S 2S 2Px 2Py 2Pz 3S
Magnesium(Mg) 1S22S22P63S2
1S 2S 2Px 2Py 2Pz 3S
Aluminum (Al) 1S22S22P63S23P1


1S 2S 2Px 2Py 2Pz 3S 3Px 3Py 3Pz
Silicon(Si) 1S22S22P63S23P2

1S 2S 2Px 2Py 2Pz 3S 3Px 3Py 3Pz
Phosphorus(P) 1S22S22P63S23P3
1S 2S 2Px 2Py 2Pz 3S 3Px 3Py 3Pz
Sulphur(S) 1S22S22P63S23P4
1S 2S 2Px 2Py 2Pz 3S 3Px 3Py 3Pz
Chlorine(Cl) 1S22S22P63S23P5
1S 2S 2Px 2Py 2Pz 3S 3Px 3Py 3Pz
Argon(Ar) 1S22S22P63S23P6
1S 2S 2Px 2Py 2Pz 3S 3Px 3Py 3Pz
  • Electronic Configuration of Elements K and Ca:
Then 4S orbital, being of lower energy then the 3d, is filled. the elements involve potassium and calcium are represented as,
Potassium(K) [Ar]184S1

[Ar]20 4S
Calcium(Ca) [Ar]184S2

[Ar]20 4S
  • Electronic Configuration of Elements Sc and Zn:
In Scandium(21) the twenty-first electron goes to the 3d orbital, the next available orbital of the higher energy. There are five 3d orbitals with the capacity of ten electrons. From scandium to Zinc these 3d orbitals are filled up.
The presence of partially filled d orbitals generates some special properties of the elements. Elements with partially filled or f orbitals in the elementary state or ionic state are called transition elements.
Scandium(Sc) [Ar]184S23d1





[Ar]20 4S 3d 3d 3d 3d 3d
Titanium(Ti) [Ar]184S23d2




[Ar]20 4S 3d 3d 3d 3d 3d
Vanadium(V) [Ar]184S23d3



[Ar]20 4S 3d 3d 3d 3d 3d
Chromium(Cr) [Ar]184S23d4


[Ar]20 4S 3d 3d 3d 3d 3d
In reality, experimental studies on Chromium revel their electronic configuration and better represented as:
Chromium(Cr) [Ar]184S13d5

[Ar]20 4S 3d 3d 3d 3d 3d
This reordering of electrons is due to extra stability associated with a half-filled sub-shell.
Manganese(Mn) [Ar]184S23d5

[Ar]20 4S 3d 3d 3d 3d 3d
Iron(Fe) [Ar]184S23d6

[Ar]20 4S 3d 3d 3d 3d 3d
Cobalt(Co) [Ar]184S23d7

[Ar]20 4S 3d 3d 3d 3d 3d
Nickel(Ni) [Ar]184S23d8

[Ar]20 4S 3d 3d 3d 3d 3d
Copper(Cu) [Ar]184S23d9

[Ar]20 4S 3d 3d 3d 3d 3d
Experimental Studies on Copper also reveals that their electronic configuration and are better represented as:
Copper(Cu) [Ar]184S13d10

[Ar]20 4S 3d 3d 3d 3d 3d
This reordering of electrons is due to extra stability associated with a filled sub-shell.
Zinc(Zn) [Ar]184S23d10

[Ar]20 4S 3d 3d 3d 3d 3d
Note: These 3d orbital are represented as dxy, dxz, dyz, d, dx²-y².

  • Electronic Configuration of Elements Ga and Kr:

In Gallium(31) the thirty-first electron goes to the 4P orbital, the next available orbital of the higher energy. There are three 4P orbitals with the capacity of six electrons. From Gallium to Krypton these 4P orbitals are filled up.
Gallium(Ga) [Ar]184S23d104P1



[Ar]18 4S 3d 3d 3d 3d 3d 4P 4P 4P
Germanium(Ge) [Ar]184S23d104P2


[Ar]18 4S 3d 3d 3d 3d 3d 4P 4P 4P
Arsenic(As) [Ar]184S23d104P3

[Ar]18 4S 3d 3d 3d 3d 3d 4P 4P 4P
Selenium(Se) [Ar184S23d104P4

[Ar]18 4S 3d 3d 3d 3d 3d 4P 4P 4P
Bromine(Br) [Ar]184S23d104P5

[Ar]18 4S 3d 3d 3d 3d 3d 4P 4P 4P
Krypton(Kr) [Ar]184S23d104P6

[Ar]18 4S 3d 3d 3d 3d 3d 4P 4P 4P

Electronic Configuration of Elements: Pauli's exclusion principle, Aufbau principle and Electronic and Electronic Configuration from Hydrogen To Krypton.

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