Electronic Configuration in Periods and Groups

The electronic configuration of a molecule refers to the distribution of electrons in various molecular orbitals. It is critical to understand the molecule. The number of electrons in bonding and antibonding molecular orbitals of a molecule or molecular ion can be calculated from its electronic configuration.

An electronic configuration, also known as an electronic structure, is the arrangement of electrons at different energy levels around an atomic nucleus. 

Long electron configurations are typically produced by the conventional notation (especially for elements having a relatively large atomic number). In such instances, a shortened or condensed notation may be employed instead of the normal notation. In shortened notation, the sequence of entirely filled subshells that correspond to a noble gas’s electronic configuration is replaced by the noble gas’s symbol in square brackets. As a result, sodium’s abbreviated electron configuration is [Ne] 3s¹ (the electron configuration of neon is 1s² 2s² 2p⁶, which can be abbreviated to [He] 2s² 2p⁶). As a result, electron configurations can be used for:

1. Determining an element’s valency
2. Predicting the qualities of a group of elements (elements with similar electron configurations tend to exhibit similar properties).
3. The interpretation of atomic spectra.

Electronic Configuration in Periods

1. The period of the element is the value of n, the primary quantum number, for the valence shell.
2. The number of electrons that can be accommodated by different energy levels varies.
3. 2n², where n is the energy level, is the greatest number of electrons that a given energy level can allow. So the first energy level (K shell) can hold up to 2 electrons, the second (L shell) up to 8 electrons, the third (M shell) up to 18 electrons, and so on.
4. The second period begins with Lithium and Beryllium, both of which have three and four electrons, respectively, and so the final electrons reach the level twos.
5. The third period begins with Sodium and finishes with Argon, filling the 3s and 3p orbitals in that order. There are eight elements in this period as well.
6. The level 4s are filled first in the fourth period with n = 4. It all starts with potassium. However, we know that the 3d orbital must be full before the 4p orbital can be filled. Scandium is the first of the three-dimensional transition elements. The 3d orbital is filled with zinc.
7. The level 5s are filled first in the fifth period with n = 5. The 4d transition series, which begins with the Yttrium, dominates this time. The 5p orbital is completely filled by Xenon at the end of the period.
8. With n = 6, the sixth period has 32 elements, with electrons filling the 6s, 4f, 5d, and 6p orbitals. Cerium signifies the entry of electrons into the 4f orbital, resulting in the lanthanide series of 4f-inner transition elements.
9. The man-made radioactive elements with electrons filling the 7s, 5f, 6d, and 7p orbitals belong to the seventh period with n = 7. Similar to period 6, this period causes electrons to fill the 5f orbital, giving rise to the actinide series of 5f-inner transition elements.

Electronic Configuration in Groups

The outermost shells of elements in the same group have the same number of electrons, resulting in identical valence shell electrical configurations. As a result, the characteristics and chemistry of elements in the same group follow a similar pattern.

Filling of Atomic Orbitals

Aufbau Principle

1. The verb Aufbeen, which means “to build up” in German, is the source of this principle’s name.
2. Electrons will occupy lower energy orbitals before moving on to higher energy orbitals, according to the Aufbau principle.
3. The energy of an orbital is calculated by adding its primary and azimuthal quantum numbers.
4. According to this principle, electrons are filled in the following order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.
5. There are a few exceptions to the Aufbau principle, such as chromium and copper. These exceptions can occasionally be explained by the stability offered by half-filled or fully filled subshells.

Pauli Exclusion Principle

1. According to the Pauli exclusion principle, an orbital can only hold a maximum of two electrons with opposite spins.
2. This principle can be stated another way that no two electrons in the same atom have the same values for all four quantum numbers.
3. As a result, if two electrons have the same principle, azimuthal, and magnetic numbers, they must have opposite spins.

Hund’s Rule

1. This rule specifies the order in which electrons are filled in all subshell orbitals.
2. It asserts that every orbital in a particular subshell is solely occupied by electrons before a second electron is put in it.
3. The electrons in orbitals with only one electron all have the same spin to maximize the total spin (or the same values of the spin quantum number).

Electronic Configurations of first 20 elements of the periodic table:

Element	Symbol	Atomic Number	Electronic Configuration
Hydrogen	H	1	1s1
Helium	He	2	1s2
Lithium	Li	3	[He] 2s1
Beryllium	Be	4	[He] 2s2
Boron	B	5	[He] 2s2 2p1
Carbon	C	6	[He] 2s2 2p2
Nitrogen	N	7	[He] 2s2 2p3
Oxygen	O	8	[He] 2s2 2p4
Fluorine	F	9	[He] 2s2 2p5
Neon	Ne	10	[He] 2s2 2p6
Sodium	Na	11	[Ne] 3s1
Magnesium	Mg	12	[Ne] 3s2
Aluminium	Al	13	[Ne] 3s2 3p1
Silicon	Si	14	[Ne] 3s2 3p2
Phosphorus	P	15	[Ne] 3s2 3p3
Sulphur	S	16	[Ne] 3s2 3p4
Chlorine	Cl	17	[Ne] 3s2 3p5
Argon	Ar	18	[Ne] 3s2 3p6
Potassium	K	19	[Ar] 4s1
Calcium	Ca	20	[Ar] 4s2

Sample Questions

Question 1: What is the electronic configuration of an element?

Answer:

An element’s electronic configuration is a symbolic representation of how its atoms’ electrons are arranged across different atomic orbitals.

Question 2: What are the three rules to be followed while writing the electronic configuration of elements?

Answer: 

1. According to the Aufbau principle, electrons must completely fill the atomic orbitals of the previous energy level before occupying an orbital associated with a higher energy level. Electrons occupy orbitals in ascending order of increasing orbital energy level.
2. No two electrons may have the same values for all four quantum numbers, according to Pauli’s exclusion principle. As a result, each orbital subshell can only hold a maximum of two electrons, both of which must have opposite spins.
3. According to Hund’s maximum multiplicity rule, all subshells in an orbital must be occupied singly before any subshell can be occupied twice. In addition, all electrons in singly occupied subshells must have the same spin (in order to maximise the overall spin).

Question 3: Why are electronic configurations important?

Answer:

Electron configurations provide insight into the chemical behaviour of elements by assisting in the determination of an atom’s valence electrons. It also aids in the separation of elements into distinct blocks (such as the s-block elements, the p-block elements, the d-block elements, and the f-block elements). This makes it simple to look into the properties of the components as a whole.

Question 4: What is the Pauli exclusion principle?

Answer:

An orbital can only hold a maximum of two electrons with opposite spins, according to the Pauli exclusion principle.

Question 5: What is Hund’s rule?

Answer:

Hund’s rule specifies the order in which electrons are filled in all of a subshell’s orbitals.