Electronic Configuration and General Trends in Chemical Properties of Elements

For understanding the chemistry of elements, it is important to understand their electronic configuration at the beginning. After all, it is these negatively charged particles that determine whether an element will behave like a cation or an anion or will form covalent bonds. Other than that, the electronic arrangements can also declare what kind of compounds they can form like monoatomic, diatomic, triatomic, and so on. Studying the electronic configuration of elements will help one to understand almost 70% of inorganic chemistry including reactions, types of compounds they can form, ionic forms, reactions with acids and alkalis, and so on.

What is electronic configuration of an element?

Every atom has a certain number of electrons that is equal to the atomic number.

For example, calcium has an atomic number of 20 and therefore, the total number of electrons in a calcium atom will be 20. According to Bohr’s model, these charges are distributed in energy levels like K, L, M, N, and so on.

However, the actual electron configuration can be expressed only in the form of orbitals. There is a completely different rule for arranging the electrons concerning these energy sublevels with which one can determine the general trends in the chemical properties.

Types of orbitals and their capacities

The concept of orbitals can be determined by the quantum principles which states that for a given value of n (quantum number determining the shell), l (Azimuthal number) can have any integral value ranging from 0 to (n-1). For example, if K=1, then n=1 and hence, l will have 0 value. This means that at the K energy level, there will be one subshell or orbital.

Similarly, if we consider the L shell, its quantum number is 2 or n=2. And therefore, l = (2-1) = 1. This means the energy level will have two sub-shells as the values will be 0 and 1. For M shells, the subshell values of L will be 3 or 0, 1, and 2.

Based on the number of l, a particular designation is used for writing the subshells like:

Here, ‘s’ determines Sharp, ‘p’ defines Principal, ‘d’ is for Diffuse, and ‘f’ is for Fundamental. Each orbital value has a different electron capacity using which the electronic configuration of elements can be known with ease.

The number of electrons that can be present in each subshell is determined by the formula of 2(2l+1). Therefore:

1. For n=1, l=0 and it means the s subshell will be there. According to the formula, the number of electrons in ‘s’ can be 2(2.0 + 1) or 2.

2. Similarly, when l=1, it becomes the ‘p’ subshell which can accommodate 2(2.1 + 1) or 6 electrons.

3. If l=2, we consider the ‘d’ subshell and it will have 2(2.2+1) or 10 electrons.

4. If l=3, the ‘f’ subshell is considered and in it, 2(2.3 +1) or 14 electrons can sit here.

Rule of arranging the total number of electrons in the orbitals

The electronic configuration  of the elements is defined by Aufbau’s principle which states that:

The electrons are filled in various orbitals in order of their increasing energies

Therefore, 1s subshell will be filled first followed by 2s, 2p, 3s, 3p, 4s, and so on. There are some exceptions in cases when the d and f subshells are considered because of their diffused electronic states.

Electronic arrangement in the first 20 elements

To understand the electron configuration, let’s have a look at the first twenty elements and learn how their electrons are arranged in the sub-shells.

How to know the chemical behaviors of atoms based on the electronic configurations?

The chemical properties of metals and non-metals vary on the electronic configurations greatly. In this section, we will clarify the trends in the chemical behaviors of elements.

Chemical behavior of metals

1. Metals are highly electropositive and therefore, they can form cations by losing the electrons present in the outermost shell.

2. As they can form positive ions, they tend to combine with the negative anions for forming ionic compounds. For example, Na+ can combine with Cl- to form the ionic compound of NaCl.

3. The ionic bonds are very strong and cannot be broken with ease.

4. All the compounds belonging to the metal carry or have 1, 2, and 3 electrons in the outermost subshell that reacts with the OH- ions to form the bases.

5. They can easily form bonds with the halogen elements like Cl, F, Br, and so on to form metallic halides.

6. The metals also react with oxygen to form ionic compounds like Na2O, CaO, and more.

Chemical behavior of non metals

1. One of the most considered topics is the chemical properties of carbon because this non-metal can form a long chain by bonding with the same atoms.

2. Non-metals usually form covalent bonds because of their high electronegativity with each other.

3. Most of the non-metallic elements form compounds with a gaseous state like CO2, N2O, SiO2, and so on.

4. Non-metals mainly form anions or negative ions by gaining more electrons.

5. The halogen compounds can react with each other to form interhalogen compounds.

6. When the non-metallic anions react with water, they form acids like HCl, HF, H2SO4, and so on.

Conclusion

The chemical properties of metals and non-metals can be known only by their electronic configuration. For example, we know that calcium, magnesium, potassium, and others have one or two electrons in the last orbital. Therefore, they can easily lose those electrons to form positively charged ions called cations. That’s why the electronic configuration of metals and non-metals is important to understand and also remember by following Aufbau’s Principle. There are other guiding principles also like Hund’s principle and Pauli’s exclusion policy. Elements having d and f subshells, however, behave differently, and that’s why their compounds are so different from the nonmetals and metals in groups 1, 2, 13, 14, 15, 16, and 17.