Electronic Configuration

Electronic Configuration












PART 2

Learning Outcome









At the end of this topic, student should be



able:




• State and apply Aufbau Principle, Pauli



Exclusion Principle, Hund’s Rule in



filling of electron in orbital of an atom.



• Write the electron configuration and



orbital diagram of and atom.

Hydrogen 1e Lithium 3e






















(1) 1s 1
(2.1) 1s 2s 1
2






Carbon 6e Neon 10e























(2.4) 1s 2s 2p 2 (2.8) 1s 2s 2p 6
2
2
2
2

ORBITAL DIAGRAM









Consist of box or line for each


orbital in a given energy level, group by


sublevel, with an arrow indication an


electron’s presence and its spin.







EXAMPLE:


H atom (ground state)










or




1s 1s




Electron configuration: 1s 1

EXAMPLE:


H atom






or






1s 1s




Electron configuration: 1s 1





The arrow show the direction of electron spin.






↑ is +½ and ↓ is –½ , but these is arbitrary.

MAXIMUM NUMBER OF




ELECTRONS





• Each shell or n contains n subshell




Example : n = 2; ℓ = 0, 1




(two possible subshell): 2s, 2p










• Each subshell of quantum number ℓ


contain (2ℓ + 1) orbitals.




Example: if ℓ = 1, there are three p



orbitals.

Maximum number of electrons in an orbital = 2





EXAMPLE:

2p subshell  3 orbitals  Maximum 6 electrons











2p






Maximum number of electrons that an atom


can have a principle level n  2n 2




EXAMPLE:


n = 2 Maximum electrons = 8










2s 2p

Maximum number of electrons:






s subshell = 2








p subshell = 6








d subshell = 10







f subshell = 14

EXAMPLE 3.2








Write the set of four quantum numbers for an


electron in a 3p orbital.









3p



3p orbital: n = 3 l = 1


Possible m values: –1, 0, +1


Each electron can have m values: –½ , +½


There are six possible ways to designate the electron:




( 3, 1, –1, +½ ) ( 3, 1, –1, –½ )




( 3, 1, 0, +½ ) ( 3, 1, 0, –½ )



( 3, 1, +1, +½ ) ( 3, 1, +1, –½ )

AUFBAU PRINCIPLE

















Electrons in an atom should be filled in the



orbitals in the order of increasing energy.














Aufbau: German word aufbauen which means


“ to build up”.

“Fill up” electrons in lowest energy orbitals (Aufbau principle)

1s






2s 2p







3s 3p 3d






4s 4p 4d 4f






5s 5p 5d 5f







6s 6p 6d






7s 7p




1s 2s 2p 3s 3p 4s 3d 4p 5s 4d

PAULI EXCLUSION PRINCIPLE







No two electrons in an atom can


have the same four quantum numbers.












He atom

1s 2



1 electron (1,0,0, +½ ) (1,0,0, -½ ) (1,0,0, +½ )
st


2 nd electron (1,0,0, +½ ) (1,0,0, -½ ) (1,0,0, -½ )



















Wolgang Pauli (1900-1958). Austrian physicist.

MAJOR CONSEQUENCE OF



PAULI EXCLUSION PRINCIPLE










An atomic orbital and holds a maximum of


two electrons they must have opposing spins.









EXAMPLE:
















2p x 2p y 2p z

2
Electron configuration of C (Z=6) = 1s 2s 2p 2
2



Different ways of distributing two electrons among


three p orbitals:















2p x 2p y 2p z 2p x 2p y 2p z 2p x 2p y 2p z




Which one is correct ?

HUND’S RULE










Electrons initially occupy orbitals singly when


orbitals of identical energy (degenerate


orbitals) are available.








EXAMPLE: C (Z = 6)







2p





correct





number of parallel 2p


spin = 2


Frederick Hund (1896-1997). German physicist.

Exercise 1






Draw the orbital diagram and write the


electronic configuration for the following



element:




i. Carbon e = 6


1s 2s 2p 2
2
2


1s 2s 2p

ii. Magnesium e = 12


2
2
6
1s 2s 2p 3s 2
1s 2s 2p 3s


iii. Argon e = 18








1s 2s 2p 3s 3p

Write the electron configuration of all element


in period 1, 2 and 3 (proton number 1 ~ 18)







Z Electron Z Electron

configuration configuration



2
2
1 Is 1 10 1s 2s 2p 6
2
6
2 1s 2 11 1s 2s 2p 3s 1
2
3 1s 2s 1 12 1s 2s 2p 3s 2
2
6
2
2
2
2
6
2
2
4 1s 2s 2 13 1s 2s 2p 3s 3p 1
2
2
2
2
6
5 1s 2s 2p 1 14 1s 2s 2p 3s 3p 2
2
2
2
2
6
6 1s 2s 2p 2 15 1s 2s 2p 3s 3p 3
2
2
7 1s 2s 2p 3 16 1s 2s 2p 3s 3p 4
2
6
2
2
2
2
2
2
6
8 1s 2s 2p 4 17 1s 2s 2p 3s 3p 5
2
2
2
2
9 1s 2s 2p 5 18 1s 2s 2p 3s 3p 6
2
2
6
2
2

Valence electronic configuration








Example :



1.Carbon – 6e



2
Electronic configuration : 1s 2s 2p 2
2
2
Valence electronic configuration : 2s 2p 2





2. Natrium – 11e



6
2
Electronic configuration : 1s 2s 2p 3s 1
2
Valence electronic configuration : 3s 1

An atom which has all paired electrons is said



to be diamagnetic and is repelled by external


magnetic field.


Example : Be – 4e


2
Electronic configuration : 1s 2s 2






An atom which has one or more unpaired


electrons is said to be paramagnetic and is


attracted to en external magnetic field


Example : C – 6e


2
2
Electronic configuration : 1s 2s 2p 2

ELECTRON CONFIGURATION OF



THE TRANSITION ELEMENTS




Transition element  partially filled d orbitals.






By applying Aufbau principle:




4s orbital is filled before 3d orbitals.







EXAMPLE: Titanium, Ti (Z = 22)












1s 2s 2p 3s 3p 4s 3d






2
6
2
6
Electron configuration: 1s 2s 2p 3s 3p 4s 3d 2
2
2

4s is more stable than 3d orbital  in terms of energy.


However, as soon as the electron occupy the 3d

orbitals, the 4s electrons are repelled to higher energy.







EXAMPLE: Titanium, Ti (Z = 22)







2
2
6
6
2
Electron configuration: 1s 2s 2p 3s 3p 4s 3d 2
2
(by applying Aubau principle)



2
2
6
2
Electron configuration: 1s 2s 2p 3s 3p 3d 4s 2
6
2
(due to energy level after repulsion)





Note: Both configurations can be used!

Example







Draw the orbital diagram and write the



electronic configuration for;




a) Potassium (Z=19)










1s 2s 2p 3s 3p 4s


2
2
6
2
6
Electron configuration: 1s 2s 2p 3s 3p 4s 1
b) Scandium (Z=21)










1s 2s 2p 3s 3p 4s 3d


2
6
2
2
Electron configuration: 1s 2s 2p 3s 3p 4s 3d 1
6
2

Writing the electronic configuration






Electronic configuration of Neon, Z = 10










1s 2s 2p


Electron configuration: 1s 2s 2p 6
2
2


Electronic configuration of Aluminum, Z = 13










1s 2s 2p 3s 3p


2
2
Electron configuration: 1s 2s 2p 3s 3p 1
2
6
or

[Ne] 3s 3p 1
2

Example







Write a complete set of quantum numbers for


each of the electrons in boron (B: Z = 5).










1s 2s 2p


1s 2s 2p 1
2
2


(n,l,m,s) = (1,0,0,+½) (1,0,0,-½) represent orbital 1s





(2,0,0,+½) (2,0,0,-½) represent orbital 2s




(2,1,-1,+½)or (2,1,0,+½) or (2,1,1,+½)
represent orbital 2p
(2,1,-1,-½) or (2,1,0,-½) or (2,1,1,-½)

Electronic Configuration of Cations and


Anions






1)Noble gas configuration


Group 1, 2 and 13 elements donate valence


electrons to form cations with noble gas


configurations


Example:



2
2
Na : 1s 2s 2p 3s 1
6
2
6
2
Na + : 1s 2s 2p (isoelectronic with Ne)


Ca : 1s 2s 2p 3s 3p 4s 2
2
6
2
2
6
2
6
2
Ca 2+ : 1s 2s 2p 3s 3p (isoelectronic with Ar)
6
2

 Group 15, 16 and 17 elements accept


electrons to form anions with noble gas


configurations


Example:


2
2
O : 1s 2s 2p 4

O 2 : 1s 2s 2p 6
2
2

(isoelectronic with neon)






2
Cl : 1s 2s 2p 3s 3p 5
6
2
2
6
2
Cl  : 1s 2s 2p 3s 3p 6
2
2
(isoelectronic with Ar)

2) Stability of the fully-filled orbitals






 d block elements donate electrons from


4s orbitals to form cations.




Example:



Zn : 1s 2s 2p 3s 3p 4s 3d 10
6
2
6
2
2
2
6
2
2
Zn 2+ : 1s 2s 2p 3s 3p 3d 10
6
2

3) Stability of the half-filled orbitals






 d block element can also donate electrons to

achieve the stability of half-filled orbitals



Example:


2
6
2
2
2
Mn : 1s 2s 2p 3s 3p 4s 3d 5
6
Mn 2+ : 1s 2s 2p 3s 3p 3d 5
2
2
6
6
2
(stability of half-filled 3d orbital )

6
2
2
2
Fe : 1s 2s 2p 3s 3p 4s 3d 6
6
2
6
2
3+
Fe : 1s 2s 2p 3s 3p 3d 5
6
2
2
(stability of half-filled 3d orbital)