Electronic Configuration of First 30 Elements: Aufbau and Pauli’s Exclusion Principle, Hund’s Rule of Maximum Multiplicity

If you have 30 elements and task to arrange them according to their similar chemical and physical properties, on the basis of which property you try to arrange them?

Electronic configurations tell us about their outermost electronic configuration (valence electrons), we can easily arrange them to their similar chemical and physical property. 

Table of contents

• What is electronic configuration?
• Aufbau’s principle
• Pauli’s Exclusion Principle
• Hund’s rule for maximum multiplicity
• Exceptional electronic configurations
• Electronic configurations of elements from atomic number 1 to 30 (H to Zn)
• Condensed electronic configuration
• Application of electronic configuration:
• Practice Problems
• Frequently Asked Questions-FAQs

What is electronic configuration?

Electrons are filled in orbitals of an element under a set of rules based on the parameters of energy, indistinguishability and orientation.

Rules for Filling Electrons in Orbitals

• Aufbau's principle
• Pauli’s exclusion principle
• Hund’s rule for maximum multiplicity

Aufbau’s principle

Electrons are filled in various orbitals in order of their increasing energies. Orbitals having the lowest energy are filled first. The order in which orbitals increase in energy is called the orbital sequence.

• Energies of subshells of single-electron species

Note: Energy of single-electron species depends only on the principal quantum number (n).

Order of Energy, 1s < 2s = 2p < 3s = 3p = 3d < 4s = 4p = 4d = 4f < …

• Energies of subshells of multi-electron species

Note: Energy of single-electron species depends on the principal quantum number (n) & azimuthal quantum number (l).

The energy of orbitals depends on (n+l) rule.

Note: if two subshells with same (n+l) value, subshell with lower ‘n’ value has lower energy

Order of Energy, 1s < 2s < 2p < 3s < 3p < 3d < 4s < 4p < 4d < 4f < …

• Memory Map for writing Electronic Configuration

Pauli’s Exclusion Principle

No two electrons in an atom can have the same set of all four quantum numbers.

Hund’s rule of maximum multiplicity 

No electron pairing takes place in the orbitals in a sub-shell until each orbital is occupied by one electron with parallel spin. Exactly half-filled and fully filled orbitals make the atoms more stable, i.e., p³, p⁶, d⁵,d¹⁰,f⁷&f¹⁴ the configuration is most stable. 

Maximum spin multiplicity = 2|S|+1

|S| = Modulus of the maximum spin of an atom

Exceptional electronic configurations:

Electronic configurations of elements from atomic number 1 to 30 (H to Zn)

Condensed electronic configuration

Rules to write condensed electronic configuration or noble gas configuration

• Electrons are filled in valence or outermost shell according to given order

(n-2)f^0-14(n-1)d^0-10ns^0-2np^0-6; 

where, n outermost (valence) shell

 (n-1) penultimate shell

 (n-2) anti-penultimate shell

Example: Write the condensed electronic configuration of flerovium (Fl₁₁₄).

Answer: We know, the last element of period 6 is Rn₈₆.

114-86 = 28, only we have arranged 28 electrons in their valence, penultimate & anti-penultimate shells according to increasing (n+l) rule

This element belongs to period 7 because period 6 is completely filled and the last element of period 7 is

Og₁₁₈. So, n = 7, (n-1) =6 and (n-2) =5

Note: subshells having same (n+l) value, subshell with the lower value of n is preferred

Condensed electronic configurations of elements

H₁- 1s¹

He₂- 1s²

Li₃- [He] 2s¹

Be₄- [He] 2s²

B₅- [He] 2s²2p¹

C₆- [He] 2s²2p²

N₇- [He] 2s²2p³

O₈- [He] 2s²2p⁴

F₉- [He] 2s²2p⁵

Ne₁₀- [He] 2s²2p⁶

Na₁₁- [Ne] 3s¹

Mg₁₂- [Ne] 3s²

Al₁₃- [Ne] 3s²3p¹

Si₁₄- [Ne] 3s²3p²

P₁₅- [Ne] 3s²3p³

S₁₆- [Ne] 3s²3p⁴

Cl₁₇- [Ne] 3s²3p⁵

Ar₁₈- [Ne] 3s²3p⁶

K₁₉- [Ar] 4s¹

Ca₂₀- [Ar] 4s²

Sc₂₁- [Ar] 3d¹,4s² or [Ar] 4s²,3d¹

Ti₂₂- [Ar] 3d²,4s² or [Ar] 4s²,3d²

V₂₃- [Ar] 3d³,4s² or [Ar] 4s²,3d³

Cr₂₄- [Ar] 3d⁵,4s¹ or [Ar] 4s¹,3d⁵

Mn₂₅- [Ar] 3d⁵,4s² or [Ar] 4s²,3d⁵

Fe₂₆- [Ar] 3d⁶,4s² or [Ar] 4s²,3d⁶

Co₂₇- [Ar] 3d⁷,4s² or [Ar] 4s²,3d⁷

Ni₂₈- [Ar] 3d⁸,4s² or [Ar] 4s²,3d⁸

Cu₂₉- [Ar] 3d¹⁰,4s¹ or [Ar] 4s¹,3d¹⁰

Zn₃₀- [Ar] 3d¹⁰,4s² or [Ar] 4s²,3d¹⁰

Application of electronic configuration:

A. Identification of color of metallic ions compounds
Compounds generally exhibits colours due to excitation and deexcitation of electrons. In compounds of d block elements colors are produced by mainly 2 reasons

• d-d transition
• Charge transfer

The energy of excitation relates to the frequency of light absorbed when an electron from a lower energy d orbital is pushed to a higher energy d orbital. This frequency is usually in the visible range. The colour seen matches the light absorbed's complementary colour.

• Generally, ions having configurations d⁰ & d¹⁰ are colorless and configuration d¹ to d⁹ exhibit colours. (in KMnO₄ , K₂Cr₂O₇ Mn and Cr both have d⁰ configuration but they exhibit colour due to charge transfer, not due to d-d transition.)
• Magnetic behaviour 

Case 1: if unpaired electrons are present then, the ion exhibits paramagnetic behaviour.

Case 2: if all electrons are fully paired then, the ion exhibits paramagnetic behaviour.

Case 3: if half-filled subshells are present, the ion exhibits ferromagnetic behaviour (d⁵- extreme case of paramagnetism).

• Spin magnetic moment

With the help of electronic configuration we can easily find its spin magnetic moment,

n=number of unpaired electron

Example 1: Find the spin magnetic moment of d⁶ configuration?

Answer: d⁶ configuration :

 number of unpaired electron = 4

• Maximum spin multiplicity:

Maximum spin multiplicity = 2|S|+1 

S = total spin 

E.g- for d⁶ configuration 

• Exchange energy

Electrons having the same spin and energy present in degenerate orbitals can exchange their positions and in this exchange process, the energy is released and the released energy is termed exchange energy. 

The higher the number of exchanges in a particular configuration, the stability of the configuration becomes higher.

The exchange energy is the basis for Hund's rule, which allows maximum multiplicity, that is electron pairing is possible only when all the degenerate orbitals contain one electron each.

more the number of exchange ∝ more the stability of configuration

=maximum number of possible exchange

Where n is the total number of electrons having same energy and spin

r = 2 (minimum 2 electrons required for exchange)

E.g- for d⁸ configuration 

d⁸ configuration:

Case 1:

Number of electrons having same energy and spin (n) = 5

Minimum number of electrons required for exchange = 2

Case 2: 

Number of electrons having same energy and spin (n) = 3

Minimum number of electrons required for exchange = 2

Total number of possible exchange = 10 + 3 = 13

• Chemical properties

We can categorize elements having same chemical properties on the basis of their same outer electronic configuration.

E.g: Li₃- 1s²,2s¹

Na₁₁- 1s²,2s²2p⁶,3s¹

K₁₉- 1s²,2s²2p⁶,3s²3p⁶,4s¹

Li, Na, K have same outermost electronic conjugation ns¹

F₉- 1s²,2s²2p⁵ and Cl₁₇- 1s²,2s²2p⁶,3s²3p⁵ have same chemical behaviour both have same outer electronic configuration ns²np⁵

• To predict the group number, period number and block name of elements according to modern periodic table

Rules to identify group number, period number & block name of elements

• Period number: Value of n (principal quantum number) of valence shell or outermost shell decide the number of period of element

Eg. F₉ [He],2s²2p⁵; n=2 - belongs to 2^nd the period of periodic table

Na₁₁ [Ne]3s¹; n=3 - belongs to 3^rd the period of periodic table

• Block name: The last electron enters in which subshell, elements generally belongs to that block

Eg. 

F₉ [He],2s²2p⁵; the last electron enters in p a subshell. So, Fluorine is a p block element

Na₁₁ [Ne]3s¹; the last electron enters in s a subshell. So, Sodium is a s block element

• Group number: Generally, the group number is decided by the number of valence electrons or valence electrons and electrons of the penultimate shell.

1. For s block elements- the group number is equal to the number of valence electrons(n).
2. For p block elements- the group number is equal to the number of valence electrons (n)+ 10.
3. For d block elements- the group number is equal to the number of valence electrons(n) + number of electrons in the penultimate shell (n-1) d orbital.

E.g

F₉ [He],2s²2p⁵; p block elements. So, Group no = 7 + 10 = 17

Na₁₁ [Ne]3s¹; s block elements. So, Group no = 1 

Fe₂₆1s²,2s²2p⁶,3s²3p⁶3d⁶,4s²; d block elements. So, Group no = 6 + 2 = 8 

Video link: 

Practice problems:

Q1. Find the spin magnetic moment of manganese (Mn₂₅)

Answer: (A)

Solution: 

Electronic configuration of Mn = Mn₂₅- [Ar] 3d⁵,4s² or [Ar] 4s²,3d⁵

 number of unpaired electron = 5

Q2. Which ion in their aqueous solution exhibit colour

A. Ti³⁺
B. Sc³⁺
C. Cu⁺
D. Zn²⁺

Answer: (A)

Solution: 

Due to d-d transition generally ions having electronic configuration d¹ to d⁹ exhibit colour and d⁰ to d¹⁰ are colourless.

 Ti-[Ar] 3d²4s² and Ti³⁺-[Ar] 3d¹, number of unpaired electron = 1 & it exhibit colour.

Sc-[Ar] 3d¹4s² and Sc³⁺-[Ar] 3d⁰, number of unpaired electrons = 0 & it doesn’t exhibit colour.

Cu-[Ar] 3d¹⁰4s¹ and Cu⁺-[Ar] 3d¹⁰4s⁰, number of unpaired electrons = 0 & it doesn’t exhibit colour.

Zn-[Ar] 3d¹⁰4s² and Zn²⁺-[Ar] 3d¹⁰4s⁰, number of unpaired electrons = 0 & it doesn’t exhibit colour.

Q3. Condensed electronic configuration of F is 

A. 1s²,2s²2p⁵
B. [He] 2s²2p⁵
C. K - 2, L - 7
D. All of these

Answer: (B)

Condensed electronic configuration of F is [He] 2s²2p⁵

Q 4. Carbon in its ground state is 

A. Monovalent
B. Divalent
C. Trivalent
D. tetravalent

Answer: (B)

Solution: electronic configuration of C is 1s², 2s²2p²

Frequently Asked Questions-FAQs

Question 1. Can we write exactly the correct configurations of all elements by strictly following (n + l) rule, Pauli exclusion principle, and Hund's rule?
Answer: No, we can see many configurations which don’t obey these rules (specially Aufbau). Eg. Cr, Cu, Pd, Pt, etc.

Question 2. Is the outermost electronic configuration of elements along with a group of the periodic tables always the same?
Answer: Not always, we can observe in many cases,

Eg- He - 1s² but Ne - 1s²,2s²2p⁶

We can observe the same in many cases in transition metal series.

Question 3. Is it necessary to start filling any orbitals with an upward arrow?
Answer: No, you can start filling orbitals with upward or downward spin but make sure, no electron pairing takes place in the orbitals in a sub-shell until each orbital is occupied by one electron with parallel spin.

Question 4. What are exchange energy and pairing energy?
Answer: Exchange energy is the energy released when two or more electrons with the same spin exchange their positions in the degenerate orbitals of a subshell. The more the options for exchange, the more the electron’s stability. The number of exchange pairs is maximum in half-filled orbitals, hence it is more stable compared to partially filled orbitals.

Pairing energy refers to the energy released with paired electrons sharing one orbital. The more the paired electrons the more the atom’s stability. The number of pairs is maximum in fully filled orbitals, hence it is more stable compared to partially filled orbitals.

Related topics: