Electronic Configuration: Complete guide to understanding the distribution of electrons in atoms

The electronic configuration is a fundamental aspect to understand the distribution of electrons in atoms.

It is based on principles such as Aufbau and Pauli Exclusion, which determine how electrons are organized in orbitals. Furthermore, sublevels and quantum numbers play a crucial role in this configuration, relating energy to the distribution of electrons.

In this article, we will explore the relationship between electron configuration and the periodic table, as well as practical examples and exercises to practice. We will also address special electronic configurations and their influence on reactivity.

If you want to know more about electrons I recommend reading the article about the electron. There we tell you in detail what they are, their importance in physics and chemistry and their technological applications.

Electronic configuration is a fundamental concept in the study of the chemistry and physics of atoms. It allows us to understand how electrons are distributed around the nucleus of an atom and how this affects the properties and behavior of the element.

The distribution of electrons in atoms follows a specific pattern based on the Aufbau principle. This principle states that electrons fill the lower energy orbitals first before moving on to the higher energy ones. The orbitals are divided into sublevels (s, p, d and f) and each one has a maximum number of electrons it can contain.

The Aufbau principles dictates that electrons are distributed in the sublevels respecting an ascending order of energy. The sublevels are sequentially filled with electrons until reaching their maximum capacity.

For example, the s sublevel can contain a maximum of 2 electrons, the p sublevel can contain a maximum of 6 electrons, and so on. This organization of orbitals and sublevels gives us information about the electronic configuration of an atom.

The Pauli Exclusion Principle establishes that a maximum of two electrons can exist in an orbital, and they must have opposite spins. This ensures that the electrons occupy different quantum states and prevents electronic repulsion. Therefore, in each sublevel, electrons are distributed into different orbitals before adding another electron into an already occupied orbital.

In electronic configuration, sublevels and quantum numbers play a fundamental role in the organization of electrons in atoms. These sublevels represent different regions within an energy level and are denoted by letters such as s, p, d, and f.

Each sublevel has a maximum number of electrons that it can hold according to the formula 2n², where n represents the main number. The s sublevel can contain up to 2 electrons, the p up to 6 electrons, the d up to 10 electrons and the f up to 14 electrons.

The assignment of electrons in the sublevels follows a specific rule called the Aufbau Principle. According to this principle, electrons first fill the lower energy sublevels before moving to higher energy sublevels.

The electronic configuration of an atom is related to the energy of the electrons. As the sublevels are filled with electrons, the energy of the atom increases. This is due to electrostatic repulsion between electrons in the same energy shell.

To better understand the mentioned concepts, let's look at some examples of electronic configuration. For example, the oxygen atom has an electron configuration of 1s²2s²2p⁴, meaning it has 2 electrons in the 1s shell, 2 electrons in the 2s shell, and 4 electrons in the 2p shell.

Another example is the iron atom, which has an electronic configuration of 1s²2s²2p⁶3s²3p⁶4s²3d⁶. This indicates that it has 2 electrons in the 1s shell, 2 electrons in the 2s shell, 6 electrons in the 2p shell, 2 electrons in the 3s shell, 6 electrons in the 3p shell, and 6 electrons in the 3d shell.

The Periodic Table of Elements is a fundamental tool in chemistry, which provides us with information about the electronic configuration of the different elements. Through it, we can understand how electrons are distributed in atoms and how this distribution is related to periodic properties.

Each chemical element has a unique electronic configuration, which is determined by its atomic number. Following an ordered sequence, we can observe patterns in the distribution of electrons as we move down the periodic table.

In general, elements are arranged in increasing order of their atomic numbers, starting at hydrogen (H) and ending at oganeson (Og). To determine the electronic configuration of an element, we must consider the number of electrons it has and distribute them in the different sublevels and orbitals.

The sublevels and periods in the periodic table are closely related to the electronic configuration of the elements. Each period represents a new shell of electrons around the nucleus of the atom, while the subshells represent the different types of orbitals within each shell.

For example, the first period (Period 1) contains only the s sublevels and can hold a maximum of 2 electrons. The second period (Period 2) includes the s and p sublevels, and can accommodate up to 8 electrons.

As we move down the periodic table, each new period introduces a new sublevel and increases the total capacity of electrons it can hold.

The electronic configuration of elements plays a fundamental role in determining periodic properties. By understanding how electrons are distributed in atoms, we can predict characteristics such as reactivity, electrical conductivity, atomic size, and ionization energy.

For example, elements in the same group or family of the periodic table share similar electronic configurations in their valence shell, which gives them analogous chemical properties and behaviors.

Likewise, the stability of atoms is related to the occupation of the energy sublevels and the number of electrons in the valence shell. These factors determine the tendency of atoms to gain, lose, or share electrons to reach a more stable configuration.

Special electronic configurations allow us to understand particular phenomena related to electrons in atoms. Next, we will explore three key aspects of this topic: valence electrons and their influence on reactivity, electronic configuration in transition elements, and anomalous electronic configurations and exceptions.

Valence electrons are those found in the outermost shell of an atom. These electrons are essential to understanding how elements participate in chemical reactions and form bonds with other elements.

The number of valence electrons determines the reactivity and chemical properties of an element. By understanding its distribution and behavior, we can predict an element's ability to gain, lose, or share electrons during a chemical reaction.

The structure of valence electrons is especially important in the study of chemical compounds and bond formation. These electrons determine the ability of an element to combine with others and form stable molecules.

The reactivity of an element can vary depending on the number of valence electrons and their arrangement in the different sublevels. This allows us to predict the chemical properties and tendency of an element to react with other elements.

Transition elements are those found in the d blocks of the periodic table. These elements have special electronic configurations due to the presence of electrons in the d orbitals.

The distribution of these electrons in the d sublevels and their interaction with the valence electrons influence the unique chemical and physical properties of these elements.

The electronic configuration of the transition elements can vary due to the ability of the d electrons to move between orbitals. These elements usually exhibit multiple oxidation states and form compounds with different structures and properties. Understanding their electronic configuration helps us understand the special characteristics of transition elements and their behavior in chemical reactions.

Although there are general rules for determining the electronic configuration of atoms, sometimes anomalies and exceptions occur. These situations are due to the stability of certain configurations and the energy of the subatomic orbitals.

Some examples of anomalous electronic configurations include internal transition elements, such as lanthanides and actinides, which have special configurations due to the structure of the f orbitals.

Additionally, certain elements can have anomalous electronic configurations due to electronic interactions within the atom. These anomalies can influence the chemical and physical properties of the elements, as well as their reactivity in different situations.

Understanding these special configurations and exceptions allows us to have a more complete view of the electronic distribution and its impact on the chemical behavior of the elements.

To ensure that you write electron configurations accurately, it is essential to follow a few key steps. Here are some practical tips to help you master this process:

Below, we present a series of exercises that will allow you to put your knowledge of electronic configurations into practice. Try to solve them yourself and then check your answers to become even more familiar with this process: