Electronic Configuration Of Elements in Periodic Table-[Explained]

LEARNING OBJECTIVES

In this article, the author has explained electronic configuration of elements in periodic table, Aufbau principle, Pauli exclusion rule, Hund’s rule with electronic configuration of all elements of periodic table.

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The Electronic configuration of elements in the periodic table is a fundamental property of an element. Electronic configurations can vary based on the number of electrons and protons within an atom, which determines how many shells there are.

We will discuss Electronic configuration in detail below as relate it to periodic trends such as atomic radius, ionization energy, electron affinity, electronegativity and boiling point.

These different properties come into play when considering electronic configurations because they determine what type of chemical bonds form between atoms with different Electronic configurations – for example whether or not covalent bonds can form at all!

What are the electronic configurations of elements in the periodic table?

Electronic Configurations are determined by the number of electrons within an atom’s orbitals and their relative energy levels. The shells are occupied one level at a time; first shell (K), second shell (L), third shell (M) etcetera until you reach nine where there is no further Shells. As each new shell is filled, electrons will drop to the next lower energy level.

Why is it important to know how electrons are distributed in an atom for chemistry purposes?

Electrons are the backbone of all chemistry. If you don’t understand how they are distributed in atom, then understanding any other concept in this subject is difficult. It’s important to know where an electron might be found within an atom so as not to confuse it with another element when analyzing a compound or writing chemical formulas

Electrons in an atom are distributed according to the electrons’ stability. Usually, atoms will have one or two unpaired electrons. The distribution of these unpaired electrons is called electron configuration and each type of arrangement has a name:

The most stable arrangement for atomic orbitals is “a bonding orbital” where there’s not any net negative charge; this is when the number of electrons in bonding orbitals equals the number of protons.

The next stable configuration for electron distribution would be a “nonbonding orbital” where there’s not any net positive charge, then an antibonding orbital that has more electrons than protons and finally a lone pair with one unpaired electron on either end.

The electrons in an atom are considered more stable when they’re “bonding orbitals” because this is where there’s no net charge and the electron distribution is symmetrical; if you had to choose between bonding or antibonding, it would always be better for the electrons to be in a bonding orbital.

This doesn’t mean that the electrons will be in a “bonding orbital” the entire time. Electrons can jump between orbitals by absorbing energy from outside sources that’s either transmitted (photon) or absorbed, then released into an orbital of lower potential.

This jumping process is called quantum tunneling and it occurs when there are no other places for the electron to jump to and it’s not stable in the current orbital.

The number of electrons that are distributed between orbitals will determine how many different types of atoms there are, which will also affect chemists’ ability to predict what type of chemical reactions might occur when two or more elements come into contact with one another.

The importance of understanding electron configurations and their effect on chemical properties

Electron configurations are a very important factor in determining the chemical properties of an atom. The outermost electrons around an atom determine where it can and cannot form bonds with other atoms, which determines its stability as well as what kind of reactions it will undergo. For example:

A noble gas like helium has a full shell so it does not have any bonding ability or affinity for forming new compounds. It is highly unreactive and widely used in balloons because it doesn’t cause them to explode when heated due to this lack of reactivity.

The element lithium has a full shell as well but with only two electrons, so it can form bonds and tends to be more reactive than helium. It is commonly used in rechargeable batteries due to its reactivity.

The element boron also has a full orbital of eight electrons and so does not have any bonding ability or affinity for forming new compounds like the noble gas helium; however, it is quite reactive because one electron on each side of the atom are free from being involved in chemical reactions. Boron reacts poorly when exposed to water vapor at high temperatures which will create an oxide coating that prevents further reaction between elements- this type of reaction is called catalysis.

The element oxygen has a full shell with eight electrons, but it can form bonds because of the two free electrons on each side. Oxygen is one of the most reactive elements- that’s why we put people to sleep before surgery using an oxygen mask. It reacts poorly when exposed to water vapor at high temperatures which will create an oxide coating that prevents further reaction between elements and this type of reaction is called catalysis.

Therefore electron configurations play a major role in determining a chemical’s reactivity and what kind of reactions it is capable of.

How to write Electronic Configurations?

Shells

The maximum number of electrons that can be accommodated in a shell is based on the principal quantum number (n). This formula 2n², where ‘n’ is the shell level. The shells and values for n have been 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

Subshells

Electrons are distributed into subshells depending on their azimuthal quantum number (denoted by ‘l’). This quantum number is dependent upon the value of the principal quantum, n. When n has a value of 4, four different subshells can exist with electrons occupying l=0, 1 , 2 and 3 respectively. The s-subshell accommodates up to 2 electrons; p-subshell: 6; d-subshell  contains 10 while f does not have any restrictions in this regard accommodating 14 total or two less than its maximum occupancy capacity as given by the formula max =2*(2l+1) .

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

1p, 2d and 3f orbitals are not really considered to exist because the quantum number for these orbital is always less than that of a principal quantum.

Aufbau principle and filling of atomic orbitals

The Aufbau principle, is the idea that electrons occupy orbitals of lower energy before occupying higher-energy orbitals. Electrons are filled in order from 1s to 7p and an illustration can be seen below:

Pauli Exclusion Principle and filling of atomic orbitals

In 1931, Wolfgang Pauli postulated an important principle to limit the number of electrons in a single orbital. He found that only two can occupy this space at any given time and they must be oppositely spin for it not to create instability in atoms with more than one electron.

This is also known as his “Principle of No More Than Two Electrons”. The idea behind this was what if there were too many? It would cause extreme chaos and lack stability which could lead to chemical reactions depending on how strong their magnetic fields are (since all electrons have slight magnets).

Hund’s Rule and filling of atomic orbitals

Hund’s rule is a quantum mechanical rule that determines the order in which electrons are filled. It states that every orbital in a given subshell must be singly occupied by an electron before another one enters, and it explains how to maximize total spin of all these electrons.

Following diagram explains the Hund’s rule.

Electronic configuration of elements in the periodic table

The electronic configuration chart given below contains all the elements with their electronic configuration.

Electronic configuration chart

Atomic Number	Name of element	Electronic configuration
1	Hydrogen	1s1
2	Helium	1s2
3	Lithium	[He]2s1
4	Beryllium	[He]2s2
5	Boron	[He]2s22p1
6	Carbon	[He]2s22p2
7	Nitrogen	[He]2s22p3
8	Oxygen	[He]2s22p4
9	Fluorine	[He]2s22p5
10	Neon	[He]2s22p6
11	Sodium	[Ne]3s1
12	Magnesium	[Ne]3s2
13	Aluminum	[Ne]3s23p1
14	Silicon	[Ne]3s23p2
15	Phosphorus	[Ne]3s23p3
16	Sulfur	[Ne]3s23p4
17	Chlorine	[Ne]3s23p5
18	Argon	[Ne]3s23p6
19	Potassium	[Ar]4s1
20	Calcium	[Ar]4s2
21	Scandium	[Ar]3d14s2
22	Titanium	[Ar]3d24s2
23	Vanadium	[Ar]3d34s2
24	Chromium	[Ar]3d54s1
25	Manganese	[Ar]3d54s2
26	Iron	[Ar]3d64s2
27	Cobalt	[Ar]3d74s2
28	Nickel	[Ar]3d84s2
29	Copper	[Ar]3d104s1
30	Zinc	[Ar]3d104s2
31	Gallium	[Ar]3d104s24p1
32	Germanium	[Ar]3d104s24p2
33	Arsenic	[Ar]3d104s24p3
34	Selenium	[Ar]3d104s24p4
35	Bromine	[Ar]3d104s24p5
36	Krypton	[Ar]3d104s24p6
37	Rubidium	[Kr]5s1
38	Strontium	[Kr]5s2
39	Yttrium	[Kr]4d15s2
40	Zirconium	[Kr]4d25s2
41	Niobium	[Kr]4d45s1
42	Molybdenum	[Kr]4d55s1
43	Technetium	[Kr]4d55s2
44	Ruthenium	[Kr]4d75s1
45	Rhodium	[Kr]4d85s1
46	Palladium	[Kr]4d10
47	Silver	[Kr]4d105s1
48	Cadmium	[Kr]4d105s2
49	Indium	[Kr]4d105s25p1
50	Tin	[Kr]4d105s25p2
51	Antimony	[Kr]4d105s25p3
52	Tellurium	[Kr]4d105s25p4
53	Iodine	[Kr]4d105s25p5
54	Xenon	[Kr]4d105s25p6
55	Cesium	[Xe]6s1
56	Barium	[Xe]6s2
57	Lanthanum	[Xe]5d16s2
58	Cerium	[Xe]4f15d16s2
59	Praseodymium	[Xe]4f36s2
60	Neodymium	[Xe]4f46s2
61	Promethium	[Xe]4f56s2
62	Samarium	[Xe]4f66s2
63	Europium	[Xe]4f76s2
64	Gadolinium	[Xe]4f75d16s2
65	Terbium	[Xe]4f96s2
66	Dysprosium	[Xe]4f106s2
67	Holmium	[Xe]4f116s2
68	Erbium	[Xe]4f126s2
69	Thulium	[Xe]4f136s2
70	Ytterbium	[Xe]4f146s2
71	Lutetium	[Xe]4f145d16s2
72	Hafnium	[Xe]4f145d26s2
73	Tantalum	[Xe]4f145d36s2
74	Tungsten	[Xe]4f145d46s2
75	Rhenium	[Xe]4f145d56s2
76	Osmium	[Xe]4f145d66s2
77	Iridium	[Xe]4f145d76s2
78	Platinum	[Xe]4f145d96s1
79	Gold	[Xe]4f145d106s1
80	Mercury	[Xe]4f145d106s2
81	Thallium	[Xe]4f145d106s26p1
82	Lead	[Xe]4f145d106s26p2
83	Bismuth	[Xe]4f145d106s26p3
84	Polonium	[Xe]4f145d106s26p4
85	Astatine	[Xe]4f145d106s26p5
86	Radon	[Xe]4f145d106s26p6
87	Francium	[Rn]7s1
88	Radium	[Rn]7s2
89	Actinium	[Rn]6d17s2
90	Thorium	[Rn]6d27s2
91	Protactinium	[Rn]5f26d17s2
92	Uranium	[Rn]5f36d17s2
93	Neptunium	[Rn]5f46d17s2
94	Plutonium	[Rn]5f67s2
95	Americium	[Rn]5f77s2
96	Curium	[Rn]5f76d17s2
97	Berkelium	[Rn]5f97s2
98	Californium	[Rn]5f107s2
99	Einsteinium	[Rn]5f117s2
100	Fermium	[Rn]5f127s2
101	Mendelevium	[Rn]5f137s2
102	Nobelium	[Rn]5f147s2
103	Lawrencium	[Rn]5f147s27p1
104	Rutherfordium	[Rn]5f146d27s2
105	Dubnium	*[Rn]5f146d37s2
106	Seaborgium	*[Rn]5f146d47s2
107	Bohrium	*[Rn]5f146d57s2
108	Hassium	*[Rn]5f146d67s2
109	Meitnerium	*[Rn]5f146d77s2
110	Darmstadtium	*[Rn]5f146d97s1
111	Roentgenium	*[Rn]5f146d107s1
112	Copernium	*[Rn]5f146d107s2
113	Nihonium	*[Rn]5f146d107s27p1
114	Flerovium	*[Rn]5f146d107s27p2
115	Moscovium	*[Rn]5f146d107s27p3
116	Livermorium	*[Rn]5f146d107s27p4
117	Tennessine	*[Rn]5f146d107s27p5
118	Oganesson	*[Rn]5f146d107s27p6

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