Hunds Rule

There are multiple rules which we need to understand before we actually go on explaining Hund’s rule. First, we will look into Aufbau’s rule, understand it and then go to Hund’s rule.

In Aufbau’s principle section, we discuss how the electrons fill the lowest energy orbital first and then move up to higher energy orbital is only after they have filled the lower energy orbitals when we look at it from a very deep perspective we can figure out that the 1s orbitals should technically get filled before the 2s orbital because of the lower value that 1s orbital has which in turn gives it lower energy so from this theory, where we see a problem with think of a solution and provide an answer to this question which involves Hund’s rule.

Now Hund’s rule states that every orbit in a sub-level is occupied singly before any orbit is doubly occupied.

To maximise the total spin of the electron, all of the electrons in singly occupied orbitals have the same spin.

When the electrons are assigned to the orbital is, similar energy is required by N electron to fill or the orbital is, which is also referred to as degenerate orbitals, this is before pairing with another electron in a half an orbital. Add the ground states atoms tend to have a huge number of unpaired electrons. 

To understand this we can visualise the process of electrons which exhibit the same behaviour as the same poles of a magnet would if they come into contact, as negatively charged electrons filled orbitals, the first try to get as far as possible from each other before having to better.

Electrons are negatively charged because of which they ripple each other, so we see that electrons tend to minimise repulsion by first occupying their own orbitals rather than sharing an orbital with another electron now according to quantum mechanical calculations it is shown that the electrons in singly occupied orbitals are less effectively screened or shielded from the nucleus, which can be called as electron shielding.

Now explaining the second rule, if it speaks technically, the first electron in a sublevel could be either spin up or spin down. The spins of all the electrons in a given sub-level depend on the span of the first electron which is chosen in a sublevel. So we can simply say that when the spin of the first electron in a sub-level is chosen, the spins of all the other electrons in that particular sub-level depend upon the first spin.

To figure these equations out we need electron configurations, it is very difficult to understand when put in a theoretical manner so it is very important to understand the configuration, orbitals, properties, place on the table, etc.

In any whichever case it is the outermost electron of the atoms that come into contact with one another or valence shell that interacts first and an atom which is least stable (and most reactive) can be seen when its valence shell is not full, it is very important to note that the valence electrons are hugely responsible for elements chemical behaviour because elements that have the same number of valence electrons of all tend to have similar chemical properties. 

Configurations also help us in predicting the stability of an atom, the most stable configuration is that we see are the ones that have full energy levels, which occurs in the noble gases. This particular stability of these elements makes noble gases very neutral and that is why they do not react easily with any other elements. 

When we read about chemical compounds or any gas, it is very important to make predictions about how certain elements will react or what kind of chemical compounds or molecules different elements will form, and with the help of electron configurations, we can do these tasks easily.

Electronic Configuration - Pauli's Exclusion Principle, Aufbau Principle and Hund's Rule

Electrons

Electrons are small compared to protons and neutrons, over 1,800 times less than either a proton or a neutron. Electrons have a relative mass of 0.0005439 such that the electron is compared with the mass of a neutron being one or about 9.109x10^-31kg.

The electron was discovered in the year 1897 by British physicist J.J. Thomason. Basically known as "corpuscles," electrons have a negative (-ve) charge and are electrically pulled to the positively charged protons. The atomic nucleus is surrounded by electrons in pathways called orbitals; this idea was put forth by Erwin Schrödinger, a physicist, in the 1920s. Now, this model is known as the quantum model or the electron cloud system. The inner orbitals surrounding the atom are Spherical, but the outer orbitals are much more complex.

An atom's electron configuration is the orbital description of the positions of the electrons in a typical atom. Using the electron configuration and laws of physics, chemists can predict an atom's properties, such as stability, boiling point, and conductivity.

Electronic Configuration

It is the method or distribution of electrons in the orbitals of an atom. An atom comprises subatomic particles like electrons, protons, and neutrons among which only the number of electrons is considered for electronic configuration. Electrons are supplied in such a way that they achieve a high constant configuration.

The atom consists of s, p, d, and f orbitals in which s orbital can hold a maximum of 2 electrons in them,

The p orbital can hold a maximum number of 6 electrons, 'd' orbital can hold a maximum number of 10 electrons and the f orbitals can hold a maximum number of 14 electrons in the orbital shell.

For Example: Chlorine 17

1s²2s²2p⁶3s²3p⁵. 

In the above distribution of electrons, orbitals in which s orbital can hold a maximum of 2 electrons in them, the p orbital can hold a maximum number of 6 electrons

Shell 

The electron shells are named by K, L, M, N, O, P, and Q; or by 1, 2, 3, 4, 5, 6, and 7; going from innermost shell to outermost shell. Every shell is formed by one or more subshells, which are formed by the composed of atomic orbitals it is called as the subshell 

Electron Spin

Electron spin is a quantum feature of electrons. It is a kind of angular momentum. The magnitude value of this angular momentum is permanent. Like charge and rest mass, spin is a basic, unvarying property of the electron.

As a teaching method, we can sometimes liken electron spin to the earth spinning on its own axis every 24 hours. If the electron spins clockwise on its axis, it is called spin-up and if it is counterclockwise then it is called spin-down. This is a suitable explanation, if not fully justifiable mathematically.

The spin angular momentum linked with electron spin is independent of orbital angular momentum, which is associated with the electrons that travel around the nucleus.

Laws

They are three important laws that fulfill these electrons namely

1. Pauli's exclusion principle

2. Aufbau principle

3. Hund's rule.

Pauli's Exclusion Principle

According to this law, an orbital cannot have both the electrons in the same spin motion (half-integer spin); electrons will be in either positive half spin (+1/2) or negative half spin (-1/2)

For example, argon's electron configuration:

1s² 2s² 2p⁶ 3s² 3p⁶

The 1s level can accommodate two electrons with the same n, l, and ml quantum numbers. Argon's pair of electrons in the 1s orbital meet the exclusion principle because they have opposite spins, determining they have different spin quantum numbers, ms. One spin is +½, the other is -½. (Instead of saying +½ or -½ often the electrons are said to be spin-up up arrow or spin-down down arrow.)

The 2s level electrons have a separate principal quantum number to these in the 1s orbital. A couple of 2s electrons differ from each other because they have different spins.

The 2p level electrons have a different orbital angular impulse number from those in the s orbitals, hence the letter p rather than s. There are three p orbitals of similar energy, px, py, and pz. These orbitals are different from one another because they have different bearings in place. Each of the px, py, and pz orbitals can contain a pair of electrons with opposite spins.

The 3s level rises to a greater principal quantum number; this orbital accommodates an electron pair with opposite spins.

The 3p level's information is similar to that for 2p, but the principal quantum number is higher: 3p lies at higher energy than 2p.

Aufbau Principle

This principle explains filling up electrons in rising orbital energy.

For example, 1s orbital should be fulfilled before 2s orbital for 1s is lower in energy than 2s orbital. 

By regarding these three rules, the electron configuration of an atom is composed.

For example, the electron configuration of the Carbon atom. 

Carbon is a p block element that includes 6 electrons. It comprises s and p orbitals. Hence by grasping the three rules the electronic configuration of the carbon atom can be written as, 1s²2s²2p²

The electron configuration for the carbon atom is recorded as. The total no of 6 electrons is disposed over 1s, 2s, and 2p orbitals. s orbitals can hold two electrons and p orbital holds 2 electrons by following Hund's rule of highest multiplicity.

Hund's Rule

According to this principle, for a given electronic configuration, the paring of the particle is done after each subshell is filled with a single electron. In other words, the under subshell should have maximum multiplicity.

 Hund's rule states that:

• Each orbital in a subshell is only obtained before any orbital is double involved.

• All of the electrons in singly occupied orbitals have a similar spin (to maximize total spin).

When allowing electrons to orbitals, an electron first seeks to fill all the orbitals with comparable energy (also called degenerate orbitals) before joining with another electron in a half-filled orbital. Atoms at ground states tend to have as many unpaired electrons as likely. 

In reflecting this process, consider how electrons show the same behavior as the same poles on an attraction would if they came into contact with each other; as the negatively charged electrons fill orbitals, they first try to get as far as possible from each other before having to match up.

According to the first principle, electrons always start with an empty orbital before they join up. Electrons are negatively charged and, as a result, they resist each other. Electrons tend to reduce objection by occupying their own orbitals, rather than receiving or accepting an orbital with another electron. 

Furthermore, quantum-mechanical computations have shown that the electrons in only filled orbitals are small, adequately screened, or shielded from the nucleus. Electron shielding is further discussed at the next level.

For the second principle, unpaired electrons in only filled orbitals have similar spins. Technically speaking, the first electron in a subshell could be either "spin-up" or "spin-down." 

Once the spin of the first electron in a subshell is chosen, however, the spins of all of the separate electrons in that sub-shell depend on that first spin. To avoid interference, scientists typically draw the first electron, and any other unpaired electron, in an orbital as "spin-up."

For Example:

Carbon and Oxygen 

Considering the electron configuration for carbon atoms: 1s²2s²2p², the two 2s electrons will fill the similar orbital, whereas the two 2p electrons will be in various orbital (and aligned in the same direction) in accordance with Hund's rule.

Consider the electron configuration of oxygen. Oxygen has 8 electrons. The electron configuration can be written as 1s²2s²2p⁴. To draw the orbital diagram, begin with the subsequent observations: the first two electrons will pair up in the 1s orbital shell; the next two electrons will pair up in the 2s orbital shell. That leaves 4 electrons, which must be placed in the 2p orbital shell. 

According to Hund’s rule, all orbitals will be once filled before an electron is double filled. Therefore, two p orbital get one electron and one will have 2 electrons. Hund's rule also specifies that all of the unpaired electrons must have the same spin. In keeping with practice, the unpaired electrons are drawn as "spin-up".

Answer the following question:

1. What is an electron?

2. Define the term electron spin?

3. State Hund's rule?

4. Explain Pauli’s exclusion principle?

Fill in the blanks:

1. The electron was discovered by __________, a British physicist in the year 1897. (Ans: J.J. Thomason)

 2. Every shell is formed by one or more subshells, which are formed by the composed of atomic orbitals it is called as ________ (Ans: subshell)

3. Unpaired electrons in only filled orbitals have similar spins. Technically speaking, the first electron in a subshell could be either ______________. (Ans: "spin-up" or "spin-down.")

4. The p orbital can hold a maximum number of ___ electrons, 'd' orbital can hold a maximum number of __ electrons and f orbitals can hold a maximum number of _____ electrons in the orbital shell. (Ans: 6, 10, 14)

Conclusion

This is all about Hund’s Rule and the principles related to it. Understand the explanation well and find out how electrons are configured in the shells outside the nucleus of an atom.