Electron Configuration: The Key to Chemical Reactivity

Electrons are negatively charged particles orbiting around the positively charged nuclei of atoms. The number and position of an atom’s electrons are referred to as its electron configuration. All elements of the periodic table can be classified into groups based on electronic configurations. For instance, the s-p-d-f blocks of the periodic table are defined based on electron configuration.

Electron configuration is the fundamental concept in electron theory. It describes how electrons are arranged in atoms and molecules. Note that this arrangement is of keen importance in the study of the physical and chemical properties of the elements of the periodic table.

How To Write Electron Configurations

Shells or Orbits

Shells are explained by the principal quantum number (n), which describes the energy of an electron and its most probable location in an atom. It refers to the size of the orbital and the energy level of electron places. The number of subshells describes the shape of the orbital.

The shell number is represented by the formula 2n², where n is the shell number. The electron shells are labeled K, L, M, N, O, P, and Q; or 1, 2, 3, 4, 5, 6, and 7. The 1s and 7s in the outermost shell have more energy than those in the inner shells. Moreover, electrons in outer shells can travel farther from the nucleus than electrons in inner shells.

Subshells or Orbitals

Subshells are electron shells formed by subdividing the principal shell of electrons. Each subshell is separated by an electron orbital. Subshells have labels such as s, p, d, and f. The azimuthal quantum number (also known as ‘ℓ’) is the shape of an atom’s orbital. It is denoted by the symbol (ℓ) and its value is equal to the total number of angular nodes in the orbital. A value of ‘ℓ’ can indicate either an s, p, d, or f subshell that varies in shape. The s, p, d, and f subshells therefore can accommodate a maximum of 2, 6, 10, and 14 electrons respectively.

Rules for Electronic Configuration

Aufbau principle

The principle of orbital filling states that electrons fill lower-energy orbitals before filling the higher-energy ones. The Aufbau principle, which was discovered by the German physicist Arnold Sommerfeld, can be used to understand the location of electrons in an atom and their related energy levels.

For example, carbon has 6 electrons and its electronic configuration is 1s² 2s² 2p². Below is a table of the order in which electrons are filled in atomic orbitals according to the Aufbau principle.

Pauli’s Exclusion Principle

Pauli’s Exclusion Principle is a statement about the way, electrons interact with each other. The principle says that no two electrons in an atom can have identical values for all four of their quantum numbers. The Pauli Exclusion Principle follows two rules:

1. Only two electrons can occupy the same orbital at a time.
2. If two electrons are present in the same orbital, they must have opposite spins.

Pauli’s exclusion principle is a quantum theory that helps us understand the arrangement of electrons in atoms and molecules, as well as the classification of elements in the periodic table.

Hund’s Rule

Electrons in a subshell are always singly occupied with one electron before an orbital has two or more electrons, and all electrons in a singly occupied orbital have the same spin.

According to Hund’s rule,

• Before any orbital in a sub-level is double-occupied, every orbital in that sub-level is singly occupied.
• Total spin is maximized when all electrons in a single occupancy orbital have the same spin.

When an atom has one electron left in its lowest-energy orbital, it cannot pair with another electron in the same orbital because it would need to fill all of its orbitals with equivalent energy. Many unpaired electrons are present in atoms at the ground state due to this phenomenon. If two electrons come in contact with each other, they will try to get as far away from each other as possible before they have to pair up.

Also, the octet rule and lewis structures are in order to write electronic configurations.

Examples of Electron Configurations

Sodium (Na)

The atomic number (Z) of Na is 11. It has therefore 11 electrons which are divided as;

• n = 1 named K shell = 2 electrons
• n = 2 named L shell = 8 electron
• n =3 named M shell = 1 electron

The electronic Configuration with the help of all the above rules can be written as;

1s² 2s² 2p⁶ 3s¹

Oxygen (O)

The atomic number (Z) of O is 8. It has therefore 8 electrons which are divided as;

• n = 1 named K shell = 2 electrons
• n = 2 named L shell = 6 electrons

The electronic Configuration with the help of all the above rules can be written as;

1s² 2s² 2p⁴

Calcium (Ca)

The atomic number (Z) of Ca is 20. It has therefore 20 electrons which are divided as;

• n = 1 named K shell = 2 electrons
• n = 2 named L shell = 8 electrons
• n = 3 named M shell = 8 electron
• n = 4 named N shell = 2 electrons

The electronic Configuration with the help of all the above rules can be written as

1s² 2s² 2p⁶ 3s² 3p⁶ 4s²

The last 2 electrons go to the 4^th orbit, instead of being in ‘3d’ subshell of the 3^rd orbit because of Aufbau principle.

Electronic Configuration For Transition Elements

For writing the electronic configuration of transition elements various factors are involved for the stability of transition elements, therefore, they show some expectational behavior in their arrangement of electrons around the nucleus. A term known as “Exchange Energy” is used to calculate the most stable electronic configuration of such elements. 

Exchange Energy

When two or more electrons of the same spin change place in a subshell’s degenerate orbitals, energy is released. As electrons are not localized and based on the exchange of energy, we decide more stable electron configuration.

The formula for exchange energy is given as;

E_exc = KP

where,

• K = constant
• P = possibility to pair electrons with parallel spin

Value of P = ⁿC₂

where,

ⁿC₂ = n!/2!(n-2)

Note: The symbol “!” is termed ‘vectorial’. It means multiplying n by back counting (If n = 3 then 3x2x1)

Given below is a table of exchange energy (p) with the corresponding no of electrons (n). The more the value of exchange energy, the more the stable electronic configuration.

Examples

Copper (Cu²⁹)

The two possible electronic Configurations of Cu²⁹ are;

1. [Ar]₁₈ 3d⁹ 4s² 

2. [Ar]₁₈ 3d¹⁰ 4s¹ 

1. [Ar]₁₈ 3d⁹ 4s²

n = 5 (5 electrons in ‘3d’ with similar spin);

⁵C₂ = 5.4.3.2.1 / 2.1 (5-2)!

= 120 / 12

= 10K

 

n = 4 (4 electrons in ‘3d’ with similar (opposite) spin);

⁴C₂ = 4.3.2.1 / 2.1 (4-2)!

= 24/4

= 6K

Total Exchange energy for [Ar]₁₈ 3d⁹ 4s² configuration is, therefore;

= 10K + 6K

= 16K

2. [Ar]₁₈ 3d¹⁰ 4s¹

n = 5 (5 electrons in ‘3d’ with similar spin);

⁵C₂ = 5x4x3x2x1 / 2×1 (5-2)!

= 120 / 12

= 10K

n = 5 (5 electrons in ‘3d’ with similar (opposite) spin);

⁵C₂ = 5x4x3x2x1 / 2×1 (5-2)!

= 120 / 12

= 10K

Total Exchange energy for [Ar]₁₈ 3d¹⁰ 4s¹ configuration is, therefore;

= 10K + 10K

= 20K

20K is greater than 16K hence [Ar]₁₈ 3d¹⁰ 4s¹  configuration is more stable.

Note that, the pair spin of the s orbital is not considered because it is a single orbital and has no possibility of parallel spin, and if two electrons are there they cancel the effect of each other.

The same steps will be followed while considering Cr²⁴.

Key Takeaway(s)

In summary, the electron configuration is of great importance while studying the physical properties and chemical reactivities of elements of the periodic table. This is so because electron configuration is the unique arrangement of electrons.

It follows the three basic rules of electronic distribution; Aufbau’s rule for filling the lower orbitals first, Pauli-exclusion explains how the spin of two electrons is opposite in an orbital, and Hund’s rule which describes the maximum multiplicity.

Concepts Berg

What are the three rules that must be followed while writing the electronic configuration of elements?

The electronic configuration of an element is a description of the arrangement of electrons in the atom of that element. Three rules must be followed while writing the electronic configuration of elements:

1. Aufbau principle states that electrons fill orbitals in order of increasing energy, starting with the lowest-energy orbital. In other words, electrons fill the orbitals closest to the nucleus before filling higher energy orbitals.

2. Pauli’s Exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers. In other words, if two electrons occupy the same orbital, they must have opposite spins.

3. Hund’s rule states that when electrons occupy orbitals of equal energy, they prefer to occupy separate orbitals with parallel spins rather than pairing up in the same orbital with opposite spins. This is known as the “maximum multiplicity” rule.

By following these three rules, we can write the electronic configuration of any element systematically and correctly.

What are real-life examples of electronic configuration?

Electronic configuration plays an important role in understanding the properties of elements and how they interact with each other. Here are a few real-life examples of electronic configuration:

1. Chemical reactions: The electronic configuration of an atom determines its reactivity and ability to bond with other atoms. For example, the electronic configuration of sodium (1s² 2s² 2p⁶ 3s¹) has a single electron in its outermost shell, which makes it highly reactive and likely to form bonds with other elements.

2. Color of transition metals: The electronic configuration of transition metals determines their color. For example, the blue color of copper sulfate is due to the electronic configuration of copper (3d⁹ 4s²), which absorbs light in the red region of the spectrum and reflects light in the blue region.

3. Electrical conductivity: The electronic configuration of metals makes them good conductors of electricity. Metals have a “sea” of loosely bound electrons, which are free to move and carry an electric current.

4. Magnetic properties: The electronic configuration of atoms can give rise to magnetic properties. For example, the electronic configuration of iron (1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²) gives rise to its strong magnetic properties, which are important in applications such as electromagnets and MRI machines.

Overall, the electronic configuration of atoms plays a fundamental role in a wide range of chemical and physical properties that we encounter in our daily lives.

How should three 2p orbitals for oxygen be occupied?

Oxygen has 8 electrons, and the electronic configuration of an oxygen atom is 1s² 2s² 2p⁴. The 2p subshell of oxygen contains three orbitals, labeled 2p_x, 2p_y, and 2p_z, each of which can hold a maximum of two electrons. The 2p orbital of oxygen will be filled in such a way that the first three electrons go to 2p_x, 2p_y, and 2p_z with same spins according to Hund’s rule of maximum multiplicity, and the fourth electron can reside on any of these orbitals but with opposite spin.

What is the electronic configuration of the first 20 elements?

1. Hydrogen: 1s¹
2. Helium: 1s²
3. Lithium: 1s² 2s¹
4. Beryllium: 1s² 2s²
5. Boron: 1s² 2s² 2p¹
6. Carbon: 1s² 2s² 2p²
7. Nitrogen: 1s² 2s² 2p³
8. Oxygen: 1s² 2s² 2p⁴
9. Fluorine: 1s² 2s² 2p⁵
10. Neon: 1s² 2s² 2p⁶
11. Sodium: 1s² 2s² 2p⁶ 3s¹
12. Magnesium: 1s² 2s² 2p⁶ 3s²
13. Aluminum: 1s² 2s² 2p⁶ 3s² 3p¹
14. Silicon: 1s² 2s² 2p⁶ 3s² 3p²
15. Phosphorus: 1s² 2s² 2p⁶ 3s² 3p³
16. Sulfur: 1s² 2s² 2p⁶ 3s² 3p⁴
17. Chlorine: 1s² 2s² 2p⁶ 3s² 3p⁵
18. Argon: 1s² 2s² 2p⁶ 3s² 3p⁶
19. Potassium: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹
20. Calcium: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²

Note that each element has a unique electronic configuration, which determines its chemical properties and behavior. The electronic configuration is determined by the number and arrangement of electrons in the atom, and it can be used to predict the element’s reactivity, bonding behavior, and other properties.

What is the full form of s, p, d, and f?

What are exceptional cases in electron configuration?

There are a few exceptional cases in an electron configuration that arise due to the stability of half-filled and fully-filled subshells. These include:

1. Chromium (Cr): The expected electron configuration of chromium, based on its atomic number (24), would be [Ar] 4s² 3d⁴. However, the actual electron configuration is [Ar] 4s¹ 3d^5. This is because having a half-filled 3d subshell is more stable than having a partially-filled 3d subshell.

2. Copper (Cu): The expected electron configuration of copper, based on its atomic number (29), would be [Ar] 4s² 3d^9. However, the actual electron configuration is [Ar] 4s¹ 3d¹⁰. This is because having a fully-filled 3d subshell is more stable than having a partially-filled 3d subshell.

3. Silver (Ag): The expected electron configuration of silver, based on its atomic number (47), would be [Kr] 5s² 4d⁹. However, the actual electron configuration is [Kr] 5s¹ 4d¹⁰. This is also due to the stability of a fully-filled 4d subshell.

4. Molybdenum (Mo): The expected electron configuration of molybdenum, based on its atomic number (42), would be [Kr] 5s² 4d⁴. However, the actual electron configuration is [Kr] 5s¹ 4d⁵. This is because having a half-filled 5p subshell is more stable than having a partially-filled 4d subshell.

5. Tungsten (W): The expected electron configuration of tungsten, based on its atomic number (74), would be [Xe] 6s² 4f¹⁴ 5d⁴. However, the actual electron configuration is [Xe] 6s¹ 4f¹⁴ 5d⁵. This is also due to the stability of a half-filled 5d subshell.

What is the electronic configuration of lead?

The atomic number of lead (Pb) is 82. Therefore, the electronic configuration of lead can be written as:

1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s² 4f¹⁴ 5d¹⁰ 6p²

This electronic configuration indicates that lead has a total of 82 electrons, which are distributed among the various orbitals and subshells in its electron shell. The configuration shows that lead has a filled 6s subshell, a filled 4f subshell, a filled 5d subshell, and two electrons in the 6p subshell. Lead is a heavy element with a relatively complex electron configuration due to the presence of many subshells and orbitals.

What is the electronic configuration of P²⁺?

A P²⁺ ion means that two electrons have been removed from a neutral phosphorus (P) atom. The atomic number of neutral phosphorus is 15, so it has 13 electrons distributed in its electronic configuration. When two electrons are removed to form the P²⁺ ion, the electronic configuration becomes;

1s² 2s² 2p⁶ 3s² 3p¹

What is the electron configuration for I?

The element I stands for iodine, which has an atomic number of 53. Therefore, the electron configuration for I can be written as:

1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁵

This configuration shows how the 53 electrons of the iodine atom are distributed among the various orbitals and subshells in its electron shell. The electron configuration indicates that iodine has a completely filled 5s and 4d subshell, and it has five electrons in the 5p subshell. The outermost valence shell of iodine contains 7 electrons, which makes it a halogen in group 17 of the periodic table.

What is the electronic configuration of the O^2- ion?

O^-2 is the oxide ion, which means that it is formed by the addition of two electrons to a neutral oxygen (O) molecule. The electronic configuration of the neutral O2 molecule is:

1s² 2s² 2p⁴

When two electrons are added to this configuration to form the oxide ion, the two electrons occupy the highest available energy level, which is the 2p subshell. Therefore, the electronic configuration of the O^-2 ion is:

1s² 2s² 2p⁶

This configuration shows that the oxide ion has 10 electrons, which are distributed among the various orbitals and subshells in its electron shell. The electronic configuration of the O^-2 ion is isoelectronic with the neon (Ne) atom, which also has 10 electrons and the same electronic configuration.