In atomic physics and quantum chemistry, the electron configuration is the distribution of electrons of an atom or molecule (or other physical structure) in atomic or molecular orbitals.

Electrons in an atom are grouped around the nucleus into shells. Shell (electron): A grouping of electrons in an atom according to energy. The farther a shell is from the nucleus, the larger it is, the more electrons it can hold, and the higher the energies of those electrons.

Electron Configuration

Introduction: What are Electron Configurations?

The electron configuration of an element describes how electrons are distributed in its atomic orbitals. Electron configurations of atoms follow a standard notation in which all electron-containing atomic subshells (with the number of electrons they hold written in superscript) are placed in a sequence. For example, the electron configuration of sodium is 1s²2s²2p⁶3s¹.

However, the standard notation often yields lengthy electron configurations (especially for elements having a relatively large atomic number). In such cases, an abbreviated or condensed notation may be used instead of the standard notation. In the abbreviated notation, the sequence of completely filled subshells that correspond to the electronic configuration of a noble gas is replaced with the symbol of that noble gas in square brackets. Therefore, the abbreviated electron configuration of sodium is [Ne]3s¹ (the electron configuration of neon is 1s²2s²2p⁶, which can be abbreviated to [He]2s²2p⁶). This is called spectroscopic notation of electron configurations.

Electron Configurations are useful for:

• Determining the valency of an element.
• Predicting the properties of a group of elements (elements with similar electron configurations tend to exhibit similar properties).
• Interpreting atomic spectra.

This notation for the distribution of electrons in the atomic orbitals of atoms came into practice shortly after the Bohr model of the atom was presented by Ernest Rutherford and Niels Bohr in the year 1913.

Writing Electron Configurations

Shells

The maximum number of electrons that can be accommodated in a shell is based on the principal quantum number (n). It is represented by the formula 2n², where ‘n’ is the shell number. The shells, values of n, and the total number of electrons that can be accommodated are tabulated below.

Shell and ‘n’ value	Max. Electrons in the Electron Configuration
K shell, n=1	2*12 = 2
L shell, n=2	2*22 = 8
M shell, n=3	2*32 = 18
N shell, n=4	2*42 = 32

Figure 6.34 Multi-electron System

Subshells

• The subshells into which electrons are distributed are based on the azimuthal quantum number (denoted by ‘l’).
• This quantum number is dependent on the value of the principal quantum number, n. Therefore, when n has a value of 4, four different subshells are possible.
• When n=4. the subshells correspond to l=0, l=1, l=2, and l=3 and are named the s, p, d, and f subshells respectively.
• The maximum number of electrons that can be accommodated by a subshell is given by the formula 2*(2l + 1).
• Therefore, the s, p, d, and f subshells can accommodate a maximum of 2, 6, 10, and 14 electrons respectively.

All the possible subshells for values of n up to 4 are tabulated below.

Principle Quantum Number Value	Value of Azimuthal Quantum Number	Resulting Subshell in the Electron Configuration
n=1	l=0	1s
n=2	l=0	2s
l=1	2p
n=3	l=0	3s
l=1	3p
l=2	3d
n=4	l=0	4s
l=1	4p
l=2	4d
l=3	4f

Thus, it can be understood that the 1p, 2d, and 3f orbitals do not exist because the value of the azimuthal quantum number is always less than that of the principal quantum number.

Figure 6.35 Energy level of Orbitals

Notation

• The electron configuration of an atom is written with the help of subshell labels.
• These labels contain the shell number (given by the principal quantum number), the subshell name (given by the azimuthal quantum number), and the total number of electrons in the subshell in superscript.
•  If two electrons are filled in the ‘s’ subshell of the first shell, the resulting notation is ‘1s²’.
• With the help of these subshell labels, the electron configuration of magnesium (atomic number 12) can be written as 1s² 2s² 2p⁶ 3s².

Filling of Atomic Orbitals

The Aufbau principle, from the German Aufbau prinzip, also called the aufbau rule, states that in the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy levels before occupying higher levels. For example, the 1s subshell is filled before the 2s subshell is occupied.

The Aufbau principle | Atomic structure and properties | AP Chemistry | Khan Academy

• It states that in ground state of atom, the orbitals are filled in increasing order of their energies.
• So electrons first occupy the lowest energy orbital available to them and enter into higher energy orbitals only after the lower energy orbitals are filled as per Pauli Exclusion principle & Hund Rule
• Now we know that energy of a given orbital depends upon effective nuclear charge and different type of orbitals are affected to different extent. So there is no single ordering of energies of orbitals which will be universally correct for all atoms
• But Following ordering can be used to predict the configuration of electrons in atom
  1s, 2s ,2p, 3s ,3p ,4s ,3d ,4p, 5s,4d,5p,4f,5d,6p

Figure 6.36 Aufbau Principle

Summary:

Type of orbital

Each shell in an atom has a different type of orbital.
For example:

• The innermost shell (n=1) only has an s orbital.
• The second shell (n=2) has an s orbital and a p orbital.
• The third shell (n=3) has s, p, d orbitals.
• From the fourth shell (n>=4) there are s, p, d, f orbitals.

Each orbital has a fixed number of electrons.

• s Orbital can contain 2 electrons.
• p Orbital can contain 6 electrons.
• d Orbital can contain 10 electrons.
• f Orbital can contain 14 electrons.

Each orbital has a slightly different electrical potential energy.

s < p < d < f

The s orbital has the lowest energy so that the electrons can be the most stable.
If both the s orbital and the p orbital of the same shell are empty, the electrons first try to fill the s orbital first.

Rules for electrons to be filled

In a stable atom, electrons are filled from the innermost shell.
In the same shell, electrons are filled in the order of s, p, d, f orbitals.
In other words, as the energy increases in the order of s → p → d → f, the energy level increases.

When the number of shells exceeds 3, the energy boundary begins to become unclear.
For example, a 4s orbital has a lower energy value compared to a 3d orbital. So, before the 3d orbital is filled with electrons, the 4s orbital is filled with electrons first.
Similarly, before the 4d orbital is filled with electrons, it is first filled in the 5s orbital.

Besides, electrons have the property of being in pairs (even) as can as possible.
Also, there are cases where electrons become stable as they fill the upper orbital first. In this case, an exception occurs.

Reference: https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Quantum_Mechanics/10%3A_Multi-electron_Atoms/Electron_Configuration

The above order is based on Madelung’s Rule which states that

1. The sum of the values of the principal quantum numbers (n) and azimuthal quantum number (l) i.e (n+l) determined the energy level of an orbital. Subshell having lower n+l value will be filled first.
2. The lower value of the sum of (n+l) implies that the energy of the orbital is low. If the value of n+l for two orbitals is equal then the orbital with a lower value of n will have a lower energy level.

The table below shows the electron configuration for each element in the periodic table:

Reference: https://chem.libretexts.org/Bookshelves/General_Chemistry/Book%3A_ChemPRIME_(Moore_et_al.)/05%3A_The_Electronic_Structure_of_Atoms/5.17%3A_Electron_Configurations_and_the_Periodic_Table
Table: Atomic Electron Configurations

Z (atomic number)

Element

Configuration

1	H	1s 1
2	He	1s 2
3	Li	[He] 2s 1
4	Be	[He] 2s 2
5	B	[He] 2s 2 2p1
6	C	[He] 2s 2 2p2

Z

Element

Configuration

7	N	[He] 2s 2 2p3
8	0	[He] 2s 2 2p4
9	F	[He] 2s 2 2p5
10	Ne	[He] 2s 2 2p6
11	Na	[Ne] 3s 1
12	Mg	[Ne] 3s 2
13	Al	[Ne] 3s 2 3p1
14	Si	[Ne]3s 2 3p2
15	P	[Ne] 3s 2 3p3
16	S	[Ne] 3s 2 3p4
17	Cl	[Ne] 3s 2 3p5
18	Ar	[Ne] 3s 2 3p6
19	K	[Ar] 4s 1
20	Ca	[Ar] 4s 2
21	Sc	[Ar] 3d 1 4s 2
22	Ti	[Ar] 3d 2 4s 2
23	V	[Ar] 3d 3 4s 2
24	Cr	[Ar] 3d 5 4s 1
25	Mn	[Ar] 3d 5 4s 2
26	Fe	[Ar] 3d 6 4s 2
27	Co	[Ar] 3d 7 4s 2
28	Ni	[Ar] 3d 8 4s 2
29	Cu	[Ar] 3d 10 4s 1
30	Zn	[Ar] 3d 10 4s 2
31	Ga	[Ar] 3d 10 4s 2 4p 1
32	Ge	[Ar] 3d 10 4s 2 4p 2
33	As	[Ar] 3d 10 4s 2 4p 3
34	Se	[Ar] 3d 10 4s 2 4p 4
35	Br	[Ar] 3d 10 4s 2 4p 5
36	Kr	[Ar] 3d 10 4s 2 4p 6
37	Rb	[Kr] 5s 1
38	Sr	[Kr] 5s 2
39	Y	[Kr] 4d 1 5s 2
40	Zr	[Kr] 4d 2 5s 2
41	Nb	[Kr] 4d 4 5s 1
42	Mo	[Kr] 4d 5 5s 1
43	Tc	[Kr] 4d 5 5s 2
44	Ru	[Kr] 4d 7 5s 1
45	Rh	[Kr] 4d 8 5s 1
46	Pd	[Kr] 4d 10
47	Ag	[Kr] 4d 10 5s 1

Z

Element

Configuration

48	Cd	[Kr] 4d 10 5s 2
49	In	[Kr] 4d 10 5s 2 5p 1
50	Sn	[Kr] 4d 10 5s 2 5p 2
51	Sb	[Kr] 4d 10 5s 2 5p 3
52	Te	[Kr] 4d 10 5s 2 5p 4
53	I	[Kr] 4d 10 5s 2 5p 5
54	Xe	[Kr] 4d 10 5s 2 5p 6
55	Cs	[Xe] 6s 1
56	Ba	[Xe] 6s 2
57	La	[Xe] 5d 1 6s 2
58	Ce	[Xe] 4f 1 5d 1 6s 2
59	Pr	[Xe] 4f 3 6s 2
60	Nd	[Xe] 4f 4 6s 2
61	Pm	[Xe] 4f 5 6s 2
62	Sm	[Xe] 4f 6 6s 2
63	Eu	[Xe] 4f 7 6s 2
64	Gd	[Xe] 4f 7 5d 1 6s 2
65	Tb	[Xe] 4f 9 6s 2
66	Dy	[Xe] 4f 10 6s 2
67	Ho	[Xe] 4f 11 6s 2
68	Er	[Xe] 4f 12 6s 2
69	Tm	[Xe] 4f 13 6s 2
70	Yb	[Xe] 4f 14 6s 2
71	Lu	[Xe] 4f 14 5d 1 6s 2
72	Hf	[Xe] 4f 14 5d 2 6s 2
73	Ta	[Xe] 4f 14 5d 3 6s 2
74	W	[Xe] 4f 14 5d 4 6s 2
75	Re	[Xe] 4f 14 5d 5 6s 2
76	0s	[Xe] 4f 14 5d 6 6s 2
77	Ir	[Xe] 4f 14 5d 7 6s 2
78	Pt	[Xe] 4f 14 5d 9 6s 1
79	Au	[Xe] 4f 14 5d 10 6s 1
80	Hg	[Xe] 4f 14 5d 10 6s 2
81	Tl	[Xe] 4f 14 5d 10 6s 2 6p1
82	Pb	[Xe] 4f 14 5d 10 6s 2 6p2
83	Bi	[Xe] 4f 14 5d 10 6s 2 6p 3
84	Po	[Xe] 4f 14 5d 10 6s 2 6p 4
85	At	[Xe] 4f 14 5d 10 6s 2 6p 5
86	Rn	[Xe] 4f 14 5d 10 6s 2 6p 6
87	Fr	[Rn] 7s 1
88	Ra	[Rn] 7s 2

Z

Element

Configuration

89	Ac	[Rn] 6d 1 7s 2
90	Th	[Rn] 6d 2 7s 2
91	Pa	[Rn] 5f 2 6d 1 7s 2
92	U	[Rn] 5f 3 6d 1 7s 2
93	Np	[Rn] 5f 4 6d 1 7s 2
94	Pu	[Rn] 5f 6 7s 2
95	Am	[Rn] 5f 7 7s 2
96	Cm	[Rn] 5f 7 6d 1 s 2
97	Bk	[Rn] 5f 9 s 2
98	Cf	[Rn] 5f 10 s 2
99	Es	[Rn] 5f 11 s 2
100	Fm	[Rn] 5f 12 s 2
101	Md	[Rn] 5f 13 s 2
102	No	[Rn] 5f 14 s 2
103	Lr	[Rn] 5f 14 6d 1 s 2
104	Rf	[Rn] 5f 14 6d 2 s 2

The first period of the periodic table has the elements of H (Hydrogen) and He (Helium). Both are distinguished with the electron configuration of 1 s¹ and 1 s² respectively.

The group 1 A and 2 A will have the general valence electrons subshell of 1 s ^x and 2 s ^y where x and y are number of valence electrons which equals number of the group. The most general formula is n s ¹ and n s ² for group the S block of the periodic table of the elements where n is number of the period.

The second period of the periodic table starts with the element Li (Lithium) that has the electron configuration 1 s ² 2 s¹ followed by Be (Beryllium) with electron configuration 1 s ² 2 s². After these two elements, element Boron (B) which starts filling the p orbital. The electrum configuration of Boron (B) is 1 s ² 2 s² 2 p ¹ or [He] 2 s² 2 p ¹ with the 3 valence electrons hosted in 2 s² 2 p ¹. The 2-p subshell will be filled starting from B (Boron) till Ne (Neon).

The groups 2 A through 8 A will have the general valence electrons subshell of 2 s ^x 2 p ^y where x + y equals the number of valence electrons which equals number of the group.

The third period of the periodic table then starts with Na (Sodium) and Mg (Magnesium) where the 3 s orbital starts to be filled. The electron configuration of Na (sodium) is [Ne] 3 s ¹. Mg (Magnesium) will have the electron configuration of [Ne] 3 s ². Starting with the Al (Aluminum) the 3 p orbitals will start to be filled till Ar (Argon).

The groups 3 A through 8 A will have the general valence electrons subshell of 3 s ^x 3 p ^y where x + y equals the number of valence electrons which equals number of the group. More generally the valence electron subshells have the general formula n s ^x n p ^y where n is number of the period and x + y is number of the valence electrons. This formula can be used for All elements of the P block

The transition metals are hosted in group B and called the D block elements. In these elements the d orbitals will be utilized and filled. The general formula for the electron valence subshell is

(n – 1) d ^x n s ^y where n number of the period and x + y is number of the valence electrons which equals number of the group.

The Lanthanides and the Actinides are utilizing the f subshells.  The general formula for the electron subshell is (n-2) f  ^x (n-1) d ^y n s ^z where x + y + z equals the valence electrons and equals number of the main groups. The Lanthanides and Actinides (inner transition elements) occupy periods 6 and 7 respectively.

The Lanthanides and Actinides occupy the groups extending from the D and the P blocks as can be seen below:

Figure 6.37 Periodic Classification of Electron Configuration

Reference: https://chem.libretexts.org/Bookshelves/General_Chemistry/Book%3A_ChemPRIME_(Moore_et_al.)/05%3A_The_Electronic_Structure_of_Atoms/5.17%3A_Electron_Configurations_and_the_Periodic_TablePau

Pauli Exclusion Principle

A basic principle of modern physics states that for particles such as electrons that possess half-integral values of spin, no two of them can be in identical quantum states within the same system. The quantum state of a particle is defined by the values of its quantum numbers, so what this means is that no two electrons in the same atom can have the same set of quantum numbers. This is known as the Pauli exclusion principle, named after the German physicist Wolfgang Pauli (1900-1958, Nobel Prize 1945).

Hund’s Rule

When the two electrons are entering an orbital, they should enter this orbital singly and parallel to each other (with the same spin).

An example of Hund’s Rule is given below:

Figure 6.38 Hund’s Rule

Below is some electronic arrangements of atoms for period I and II elements.

Figure 6.39 location of electrons in different shells and subshells

Reference: https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Quantum_Mechanics/10%3A_Multi-electron_Atoms/Electron_Configuration

Examples

The electron configurations of a few elements are provided with illustrations in this subsection.

Worked Example 1: Electron Configuration of Hydrogen

The atomic number of hydrogen is 1. Therefore, a hydrogen atom contains 1 electron, which will be placed in the s subshell as shown below:

Hydrogen Electron Configuration

 

Worked Example 2: Electron Configuration of Oxygen

Hydrogen Electron Configuration1s² is the electron configuration of Helium [He}

Electron Configuration of Oxygen based on inert gas core:

[ He] 2s² 2p⁴

The valence electron is equal the number of the group. Oxygen is found in the 6^th group and hence the valence electrons are 6.

The orbitals that are hosting the valence electrons are 2s² 2p⁴

 

Worked Example 3: Electron Configuration of Chlorine

Chlorine Electron Configuration

1s² 2s² 2p⁶ is the electron configuration of Neon [Ne]

Electron Configuration of chlorine based on inert gas core:

[ Ne] 3s² 3p⁵

The valence electron is equal the number of the group. Chlorine is found in the 7^th group and hence the valence electrons are 7.

The orbitals that are hosting the valence electrons are 3s² 3p⁵

Worked Example 4: AUFBAU DIAGRAMS AND SPECTROSCOPIC ELECTRON CONFIGURATION

Give the electron configuration for nitrogen (N) and draw an Aufbau diagram.

Give the number of electrons

Nitrogen has seven electrons.

Place two electrons in the 1s orbital

We start by placing two electrons in the 1s orbital: 1s²

Now we have five electrons left to place in orbitals.

Place two electrons in the 2s orbital

We put two electrons in the 2s orbital: 2s²

There are now three electrons to place in orbitals.

Place three electrons in the 2p orbital

We place three electrons in the 2p orbital: 2p³

Write the final answer

The electron configuration is: 1s²2s²2p³1s²2s²2p3. The Aufbau diagram is given in the step above.

Electron Configuration of Ions:

WORKED EXAMPLE 5: AUFBAU DIAGRAM FOR AN ION

Give the electron configuration for (O²⁻) and draw an Aufbau diagram.

Give the number of electrons

Oxygen has eight electrons. The oxygen anion has gained two electrons and so the total number of electrons is ten.

Place two electrons in the 1s orbital

We start by placing two electrons in the 1s orbital: 1s².

Now we have eight electrons left to place in orbitals.

Place two electrons in the 2s orbital

We put two electrons in the 2s orbital: 2s2.

There are now six electrons to place in orbitals.

Place six electrons in the 2p orbital

We place six electrons in the 2p orbital: 2p6.

Write the final answer

The electron configuration is: 1s22s22p6. The Aufbau diagram is:

Reference: https://intl.siyavula.com/read/science/grade-10/the-atom/04-the-atom-06

How to Write Electron Configuration

The following videos illustrate the electron configuration concepts in details:

How to write electron configuration?

Introduction to Electron Configurations

There are two ways to write the electronic configuration of electrons in atoms. The first is more conceptually cohesive and involves using the Periodic Table to write the notation. The second way is use the configuration chart ((Aufbau Principle). )

Writing Electron Configurations Using Only the Periodic Table

Practice Exercise:

Element or Ion	Electron configuration	Core electrons	Valence electrons
Potassium (K)
Helium (He)
Oxygen ion (O2-)
Magnesium ion (Mg2+)
Copper(Cu)
Mercury(Hg)
Arsenic ion ( As3-)
Lead ion ( Pb4+)
Uranium (U)

Contact your instructor for solutions

Orbital shapes 

Each orbital type has a unique shape based on the energy of its electrons. The s orbital is a spherical shape. The p orbital is a dumbbell shape. There are three p orbitals that differ in orientation along a three-dimensional axis.

Figure 6.40 Shapes of s, p, d orbitals

Ref: commons.wikimedia.org/

s Orbitals (l=0)

Three things happen to s orbitals as n increases

1. They become larger, extending farther from the nucleus.
2. They contain more nodes. This is similar to a standing wave that has regions of significant amplitude separated by nodes, points with zero amplitude.
3. For a given atom, the s orbitals also become higher in energy as n increases because of theirr increased distance from the nucleus.

S orbitals

Figure 6.41 s-orbitals

Ref: commons.wikimedia.org/

p Orbitals (l=1)

Only s orbitals are spherically symmetrical. As the value of l increases, the numberr of orbitals in a given subshell increases, and the shapes of the orbitals become more complex. Because the 2p subshell has l = 1, with three values of m_l (−1, 0, and +1), there are three 2p orbitals.

Figure 6.42 p-orbitals

Ref: commons.wikimedia.org/

d Orbitals (l=2)

Subshells with l = 2 have five d orbitals; the first principal shell to have a d subshell corresponds to n = 3. The five d orbitals have m_l values of −2, −1, 0, +1, and +2.

Figure 6.43 d-orbitals

Ref: commons.wikimedia.org/

Reference: https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_A_Molecular_Approach_(Tro)/07%3A_The_Quantum-Mechanical_Model_of_the_Atom/7.06%3A_The_Shape_of_Atomic_Orbitals

Core and valence electrons

Electrons in the outermost energy level of an atom are called valence electrons. The electrons that are in the energy shells closer to the nucleus are called core electrons. Core electrons are all the electrons in an atom, excluding the valence electrons. An element that has its valence energy level full is more stable and less likely to react than other elements with a valence energy level that is not full

Using the Electron Configuration Chart

The importance of understanding electron configuration 

By this stage, you may well be wondering why it is important for you to understand how electrons are arranged around the nucleus of an atom. Remember that during chemical reactions, when atoms come into contact with one another, it is the electrons of these atoms that will interact first. More specifically, it is the valence electrons of the atoms that will determine how they react with one another.

To take this a step further, an atom is at its most stable (and therefore unreactive) when all its orbitals are full. On the other hand, an atom is least stable (and therefore most reactive) when its valence electron orbitals are not full. This will make more sense when we go on to look at chemical bonding in a later chapter. To put it simply, the valence electrons are largely responsible for an element’s chemical behaviour and elements that have the same number of valence electrons often have similar chemical properties.

The most stable configurations are the ones that have full energy levels. These configurations occur in the noble gases. The noble gases are very stable elements that do not react easily (if at all) with any other elements. This is due to the full energy levels. All elements would like to reach the most stable electron configurations, i.e. all elements want to be noble gases.

Electron Configurations and the Periodic Table

 Main group elements (also known as representative elements:

These elements occupy s and/or p orbitals in their outermost shells. These elements are called S and P blocks elements.

Transition elements or transition metals.

These elements are metals that occupy the d orbitals in their outermost shells. These elements are called transition metals.

Inner transition elements are metals called Lanthanides and Actinides and occupy the f and d orbitals in their outermost shells. The lanthanide series: Lanthanide (La) through lutetium (Lu). The actinide series: Actinide (Ac) through lawrencium (Lr).

Lanthanides and actinides exhibit similar chemical and physical properties and therefore they are grouped within the same series. Also Lanthanides and Actinides exhibit similar chemical and physical properties similar to the transition elements with the similar corresponding groups although these transition elements have no f orbitals in their outermost shells.

Electron Configurations of Ions

The electron configurations of ions follow the same electron configurations of the neutral elements except adding extra valence electrons in case the ion is negative (anion) or subtracting extra valence electrons in case of the ion is positive (cation).

Key Concepts and Summary

The relative energies of the subshells dictate the order of the atomic orbitals to be filled (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p etc). Electron configurations and orbital diagrams can be determined by applying the Pauli exclusion principle which states that no two electrons can have the same set of four quantum numbers and Hund’s rule what states that electrons keep the unpaired spins status in degenerate orbitals.

Electrons in the outermost orbitals, called valence electrons, are responsible for most of the chemical and physical properties and reactions. The number of the valence electrons equals the number of the group that element is found in. It means that all elements within the same groups will exhibit similar physical and chemical properties.

In the periodic table, elements with similar valence electron configurations are grouped in the same group. There are some exceptions to the predicted filling order, particularly when half-filled or completely filled orbitals can be formed. The periodic table can be divided into four blocks based on the orbital in which the last electron to be added is placed: main group elements (s and porbitals), transition elements (d orbitals), and inner transition elements (forbitals).