Electronic Structure of Atoms

If we have a look at the periodic table it can be seen that every element has an atomic number. The atomic number tells about the electronic configuration of that element. An element’s atomic number is symbolised by the letter Z. The atomic number tells about the number of protons present in the nucleus of an atom of an element. The first electronic configuration ever obtained was Bohr’s Model of Atom. Electrons revolve around different shells. These shells have different names which will be discussed in the latter half. Isotopes are those elements that have the same atomic number but the number of neutrons is different.

What is a shell? What are the features of a shell?

A shell is basically a group of electrons that have the same energy. Shells can be named as k,l,m,n. Therefore we can say that an element has four shells. The shells also consist of subshells. The subshells can be named s,p,d,f. The subshells are always arranged in increasing order of energy. 

RULES:

• The first shell of an atom of an element can only have one subshell which is p.
• The second shell can have an s and p subshell.
• The third shell can have an s,p, and d subshell.
• The fourth shell can have s,p,d, and f subshells.
• The increasing order of energy is s < p < d < f.

SYMBOLS FOR DENOTING THE ELECTRONIC CONFIGURATION OF AN ATOM

• “n”: denotes the principal quantum number
• “l”: denotes the orbital type.
• “s”: denotes the spin quantum number.

In order to understand the above symbols let’s go through an example:

1s2 : it is read as one s two. From here we can say that the s- subshell has two electrons and the value of l is 0.

HOW TO DETERMINE THE VALUE OF l?

There is a rule which says that there are specific assigned values for l.

”l” is known as the azimuthal quantum number.

• For s the value of l is 0.
• For p the value of l is 1.
• For d the value of l is 2.
• For f the value of l is 3.

NUMBER OF ELECTRONS A CERTAIN SHELL CAN HOLD

• The first shell can hold two electrons.
• The second shell can hold eight electrons.
• The third shell can hold eighteen electrons.
• The fourth shell can hold thirty-two electrons.

AUFBAU PRINCIPLE

The Aufbau principle can be summarised as the lowest energy orbitals need to be filled first before moving onto the higher energies. The maximum number of electrons that can be put in any orbital is two. 

The order in which the orbitals are filled in:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, (8s, 5 g, 6f, 7d, 8p, and 9s)

EXCEPTIONS OF THE AUFBAU PRINCIPLE

Atomic numbers and differences in energy levels are inversely proportional to each other. If the atomic number decreases the differences in the energy levels increase. This happens due to the orbitals which may be fully filled or they may be half-filled. The elements in the exception are Ag (silver) and Pt (Platinum). 

HUND’S RULE

Hund’s rule states that all the orbitals should be filled with at least one electron. After all the orbitals are filled with one electron then another electron can doubly fill it.

Copper (Cu) and Chromium (Cr) are exceptions to Hund’s rule.

PAULI’S EXCLUSION PRINCIPLE

Pauli’s exclusion principle states that the same set of four quantum numbers cannot be obtained by any two electrons.

VIOLATION OF PAULI’S EXCLUSION PRINCIPLE

Earlier it had been discussed that an orbital can have two electrons. So the violation here is that those two electrons should have a different spin.

ION’S ELECTRONIC CONFIGURATION

Ions are formed when an atom gains or loses electron/electrons. They lose electrons/electrons in order to obtain their nearest stable inert gas configuration. There are two types of ions:

• Cations: they are formed by losing electron/electrons.
• Anions: they are formed by gaining electron/electrons.

Now coming to their electronic configuration. Let’s take an example:

For sodium (Na) the electronic configuration is 1s22s22p63s1

Now in order to obtain the nearest noble gas configuration sodium has to lose one electron.

Hence the configuration now becomes 1s22s22p6

The above electronic configuration example was for a cation.

Now the electronic configuration of an anion. We take phosphorus (P) as an example.

The electronic configuration for P is: 1s22s22p63s23p3

Now in order to gain the nearest stable inert gas configuration P needs to gain three electrons 1s22s22p63s23p6

APPLICATIONS OF ELECTRONIC CONFIGURATION

• Electronic configuration helps in easy understanding of concepts related to the periodic tables. The concepts include ionization energy, metallic character, and many more.
• Electronic configuration helps in understanding the bonds between atoms. There are chemical bonds, ionic bonds, van der Waal bonds, metallic bonds, hydrogen bonds, and many more.
• Electronic configuration helps to understand the different properties of elements.

CONCLUSION

Now summarising the points which were discussed throughout the entire section.

• Bohr’s model of the atom states that the nucleus is a positive charge and the negatively charged electrons revolve around it.
• The electronic configuration follows different rules along with exceptions.
• There are symbols for denoting several electronic configurations.
• There are two types of ions that can be formed: cations and anions.
• There are shells and subshells in every electronic configuration.
• There are two types of electrons namely valence electrons and core electrons. The valence electrons are those that are present in the outermost shell. The core electrons are present in the innermost shells.
• There are various applications of electronic configuration.