Orbital Energy Ordering in Multielectron Atoms:
The orbitals of a multielectron atom are not likely to be quite the same as the hydrogen atom orbitals. For practical purpose, however, the number of orbitals and there shapes in multielectron cases may be taken to be the same as for the hydrogen orbitals. In multielectron atoms experimental studies of spectra shows that orbitals with the same vale of n but different l values have different energies.
The 3S orbital is lower energy then 3P orbitals which again are of lower energy then 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 same n will be degenerate. The separation of orbitals of a major energy level into sublevels is primarily due to the interaction among the many electrons.
Orbital Energy Diagram for Multi  Electron Atom. 
This interaction leads to the following relative order of the energies of each type of orbitals:
1S〈2S〈2P〈3S〈3P〈4S〈3d 
〈4P〈5S〈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 give as,
／  
1S

／


↙

／

／


2S

2P

／


↙

／

／

／


3S

3P

3d

／


↙

／

／

／

／


4S

4P

4d

4f


↙

／

／

／

／


5S

5P

5d

5f


↙

／

／

／


6S

6P

6d


↙

／

↙


7S


↙

The different orbitals originating from the same principle quantum number n are written in the horizontal lines. Now inclined parallel lines are drown through the orbital according to the above picture. Filling up the different orbitals by electrons will follow these lines.
Example:
A element which contain 36 electrons the electronic configuration of this element according to the above diagram is,
No more than two electrons can be placed in one and the same orbital.
1S^{2}2S^{2}2P^{6}3S^{2}3P^{6}4S^{2}3d^{10}4P^{6} 
The Aufbau or Building Up Principle :
The question that arise 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 electron have the same three quantum numbers namely the same n, same l and the same m_{l}. Any conflict with 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 the 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 exists in an orbital, the electrons are said to be paired. These two electrons per orbital given the maximum accommodation of electron 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  2n^{2} 
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 bases on the following Rules:
 Electrons are fed into orbitals in order of increasing energy (increasing n ) until all the electrons have been accommodated.
 Electrons will tend to maintain maximum spin. So long orbitals of similar energy are available for occupation electrons will prefer to remain unpaired.
 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.
 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 occupy 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 represented the electronic configuration from H to Ne:
Hydrogen(H)  1S^{1} 
↑ 
1S 
Helium(He)  1S^{2} 
↑↓ 
1S 
Lithium(Li)  1S^{2}2S^{1} 
↑↓  ↑ 
1S  2S 
Beryllium(Be)  1S^{2}2S^{2} 
↑↓  ↑↓ 
1S  2S 
Boron(B)  1S^{2}2S^{1}2P^{1} 
↑↓  ↑↓  ↑  
1S  2S  2P_{x}  2P_{y}  2P_{z} 
Carbon(C)  1S^{2}2S^{1}2P^{2} 
↑↓  ↑↓  ↑  ↑  
1S  2S  2P_{x}  2P_{y}  2P_{z} 
Nitrogen(N)  1S^{2}2S^{1}2P^{3} 
↑↓  ↑↓  ↑  ↑  ↑ 
1S  2S  2P_{x}  2P_{y}  2P_{z} 
Oxygen(O)  1S^{2}2S^{1}2P^{4} 
↑↓  ↑↓  ↑↓  ↑  ↑ 
1S  2S  2P_{x}  2P_{y}  2P_{z} 
Fluorine(F)  1S^{2}2S^{1}2P^{5} 
↑↓  ↑↓  ↑↓  ↑↓  ↑ 
1S  2S  2P_{x}  2P_{y}  2P_{z} 
Neon(Ne)  1S^{2}2S^{1}2P^{6} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓ 
1S  2S  2P_{x}  2P_{y}  2P_{z} 
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)  1S^{2}2S^{2}2P^{6}3S^{1} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑ 
1S  2S  2P_{x}  2P_{y}  2P_{z}  3S 
Magnesium(Mg)  1S^{2}2S^{2}2P^{6}3S^{2} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓ 
1S  2S  2P_{x}  2P_{y}  2P_{z}  3S 
Aluminium(Al)  1S^{2}2S^{2}2P^{6}3S^{2}3P^{1} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑  
1S  2S  2P_{x}  2P_{y}  2P_{z}  3S  3P_{x}  3P_{y}  3P_{z} 
Silicon(Si)  1S^{2}2S^{2}2P^{6}3S^{2}3P^{2} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑  ↑  
1S  2S  2P_{x}  2P_{y}  2P_{z}  3S  3P_{x}  3P_{y}  3P_{z} 
Phosphorus(P)  1S^{2}2S^{2}2P^{6}3S^{2}3P^{3} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑  ↑  ↑ 
1S  2S  2P_{x}  2P_{y}  2P_{z}  3S  3P_{x}  3P_{y}  3P_{z} 
Sulphur(S)  1S^{2}2S^{2}2P^{6}3S^{2}3P^{4} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑  ↑ 
1S  2S  2P_{x}  2P_{y}  2P_{z}  3S  3P_{x}  3P_{y}  3P_{z} 
Chlorine(Cl)  1S^{2}2S^{2}2P^{6}3S^{2}3P^{5} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑ 
1S  2S  2P_{x}  2P_{y}  2P_{z}  3S  3P_{x}  3P_{y}  3P_{z} 
Argon(Ar)  1S^{2}2S^{2}2P^{6}3S^{2}3P^{6} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓ 
1S  2S  2P_{x}  2P_{y}  2P_{z}  3S  3P_{x}  3P_{y}  3P_{z} 
Electronic Configuration of Elements K and Ca: 
Then 4S orbital, being of lower energy then the 3d, is filled. the elements involves potassium and calcium are represented as,
Potassium(K)  [Ar]^{18}4S^{1} 
↑  
[Ar]^{20}  4S 
Calcium(Ca)  [Ar]^{18}4S^{2} 
↑↓  
[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 orbital with the capacity of ten electrons. From scandium to Zinc these 3d orbitals are filled up.
Presence of partially filled d orbitals generates some some special properties of the elements.
Elements with partially filled d or f orbitals in the elementary state or ionic state are called transition elements.
Scandium(Sc)  [Ar]^{18}4S^{2}3d^{1} 
↑↓  ↑  
[Ar]^{20}  4S  3d  3d  3d  3d  3d 
Titanium(Ti)  [Ar]^{18}4S^{2}3d^{2} 
↑↓  ↑  ↑  
[Ar]^{20}  4S  3d  3d  3d  3d  3d 
Vanadium(V)  [Ar]^{18}4S^{2}3d^{3} 
↑↓  ↑  ↑  ↑  
[Ar]^{20}  4S  3d  3d  3d  3d  3d 
Chromium(Cr)  [Ar]^{18}4S^{2}3d^{4} 
↑↓  ↑  ↑  ↑  ↑  
[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]^{18}4S^{1}3d^{5} 
↑  ↑  ↑  ↑  ↑  ↑  
[Ar]^{20}  4S  3d  3d  3d  3d  3d 
This reordering of electrons is due to an extra stability associated with a half filled subshell.
Manganese(Mn)  [Ar]^{18}4S^{2}3d^{5} 
↑↓  ↑  ↑  ↑  ↑  ↑  
[Ar]^{20}  4S  3d  3d  3d  3d  3d 
Iron(Fe)  [Ar]^{18}4S^{2}3d^{6} 
↑↓  ↑↓  ↑  ↑  ↑  ↑  
[Ar]^{20}  4S  3d  3d  3d  3d  3d 
Cobalt(Co)  [Ar]^{18}4S^{2}3d^{7} 
↑↓  ↑↓  ↑↓  ↑  ↑  ↑  
[Ar]^{20}  4S  3d  3d  3d  3d  3d 
Nickel(Ni)  [Ar]^{18}4S^{2}3d^{8} 
↑↓  ↑↓  ↑↓  ↑↓  ↑  ↑  
[Ar]^{20}  4S  3d  3d  3d  3d  3d 
Copper(Cu)  [Ar]^{18}4S^{2}3d^{9} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑  
[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]^{18}4S^{1}3d^{10} 
↑  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  
[Ar]^{20}  4S  3d  3d  3d  3d  3d 
This reordering of electrons is due to an extra stability associated with a filled subshell.
Zinc(Zn)  [Ar]^{18}4S^{2}3d^{10} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  
[Ar]^{20}  4S  3d  3d  3d  3d  3d 
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 orbital with the capacity of six electrons. From Gallium to Krypton these 4P orbitals are filled up.
Gallium(Ga)  [Ar]^{18}4S^{2}3d^{10}4P^{1} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑  
[Ar]^{18}  4S  3d  3d  3d  3d  3d  4P  4P  4P 
Germanium(Ge)  [Ar]^{18}4S^{2}3d^{10}4P^{2} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑  ↑  
[Ar]^{18}  4S  3d  3d  3d  3d  3d  4P  4P  4P 
Arsenic(As)  [Ar]^{18}4S^{2}3d^{10}4P^{3} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑  ↑  ↑  
[Ar]^{18}  4S  3d  3d  3d  3d  3d  4P  4P  4P 
Selenium(Se)  [Ar^{18}4S^{2}3d^{10}4P^{4} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑  ↑  
[Ar]^{18}  4S  3d  3d  3d  3d  3d  4P  4P  4P 
Bromine(Br)  [Ar]^{18}4S^{2}3d^{10}4P^{5} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑  
[Ar]^{18}  4S  3d  3d  3d  3d  3d  4P  4P  4P 
Krypton(Kr)  [Ar]^{18}4S^{2}3d^{10}4P^{6} 
↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  ↑↓  
[Ar]^{18}  4S  3d  3d  3d  3d  3d  4P  4P  4P 