XII - 1.8 Electronic Configurations of Atoms - Video Lectures

Electronic configuration of atoms using Aufbau, Pauli's principle and Hund's rule - Chemistry
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Elearin




Atomic orbitals are the sub-stationary states or the regions in space where the electrons revolve around the nucleus in an atom.

Electronic configuration of atoms is representation of the occupation of electrons in the orbitals. In other words, electronic configuration of atoms specifies the order in which electrons fill up the orbitals.  The order in which these electrons are filled into the atomic orbitals are controlled by three principles.

1. Aufbau principle
2. Hund's rule and
3. Pauli's exclusion principle


Aufbau Principle.

According to this principle an electron always occupies the lowest energy orbital first before filling the higher level. For example, an electron always occupies 2s, the lower energy orbital, first instead of the higher 3s orbital.

The Aufbau or building-up principle can be explained with the example of Hydrogen Atom. Hydrogen has one electron. This electron enters the 1s orbital which has the lowest energy.
In other words, building-up principle states that the incoming electrons go to an orbital which has the least (n+l) value. However, the orbital having lower 'n' value will be occupied first, in case any two orbitals have the same (n+l) value.

Consider the example of Silicon whose atomic number is 14. Twelve electrons can be accommodated in 1s, 2s, 2p, 3s orbitals. Now, the last two electrons can enter into either 3p or 4s orbital. The (n+ l) values of these orbitals are the same, that is,
3p orbital has a (n+l) value of 3+1=4 and 4s has (n+l) value of 4+0 = 4

This means, both the orbitals have the same (n+l) value. But the 3p orbital has 'n' value, that is 3, which is less than the n value of 4s, which is 4. Therefore thirteenth and fourteenth electrons occupy the 3p orbital first. Thus the electronic configuration of Si is  3s2  3p2.The superscript represents the number of electrons present in the corresponding orbital.

The second important rule to determine the electronic configuration of an atom is the Hund's Rule. It says electron pairing happens only after all the available degenerate orbitals are occupied by one electron each.

Hund's rule  example: . Consider the element Oxygen with Z=8. I has 8 electrons, the first electron goes into the '1s' orbital of the K-Shell. The second electron will be paired up with the first in the same 1s orbital. Similarly the third and fourth electrons will occupy the 2s orbital of L-Shell. The Fifth electron goes into one of the three 2p orbitals of L-Shell. Let that be 2px. Since the three p-orbitals i.e., 2px,2py and 2pz are degenerate , the sixth electron goes into 2py or 2pz but not 2px. Let us say it goes to 2py. Since 2pz is a degenerate orbital, the seventh electron goes to 2pz instead of pairing up with electron in 2px or 2py.

Now, since all the 3 sub-orbitals have one electron each, the eighth electron can pair up with any of the three electrons in 2px, 2py and 2pz orbitals. Thus the electronic configuration of Oxygen can be written as 1s1.2s2. 2px2 .2py1. 2pz1. The arrows indicate electrons with spin +1/2 and -1/2. Let us consider the nitrogen atom. it has 7 electrons. The first six electrons have the same arrangement as that of carbon atom 1s1.2s2. 2px1 .2py1. The seventh electron will enter only in 2pz but can not enter into 2px or 2py orbital. Thus the configuration is 1s1.2s2. 2px1 .2py1 2pz1

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The third important rule for electronic configuration, the Pauli's Exclusion Principle states that no two electrons will have the four quantum numbers same. This means that two electrons can ever have any identical values of n, l, m and s values. Because of this rule a single orbital can have only 2 electrons.

XII - 1.7 Orbital Wave Functions and Shapes of Orbitals - Video Lectures

XII -
1.7 Orbital Wave Functions and Shapes of Orbitals - Video Lectures

1.7 Orbital Wave Functions and Shapes of Orbitals

1. Spherical shape for s.
2. Dumbbell shape for orbitals of p.
3. Four-lobed shape for orbitals of d.
4. Complex shape for all orbitals of higher sublevels

The Wave Behavior of Matter (Part 1 of 2 for Atomic Orbitals)

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dcaulf

Atomic Orbitals Explained (Sequel to Wave Behavior of Matter)

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dcaulf

Shapes of Atomic Orbitals - IIT JEE Main and Advanced Chemistry Video Lecture

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RaoIITAcademy

XI - 1.11 Mole Concept - Video Concepts

XI - 1.11 Mole Concept - Video Concepts

Mole concept
A mole is measuring unit like for example dozen.
A mole is a collection of 6.022*10²³ particles

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Askiitians


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ExamFearVideos

XII - 1.6 Pauli's Exclusion Principle - Video Lectures

XII - 1.6 Pauli's Exclusion Principle - Video Lectures

Pauli's exclusion principle: No two electrons can have all four same quantum numbers

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iProfIndia

XII - 1.5 Quantum Numbers - Video Lectures

XII -

1.5 Quantum Numbers - Video Lectures

According to quantum mechanical model or wave mechanical model of atom, orbitals represent regions in space around the nucleus where the probability of finding electrons is maximum. A large number of orbitals are possible in an atom.

To describe each electron in an atom in different orbitals, four quantum numbers are used. They are designated as n,l,m_l, and m_s.



1. Principal quantum number (n) This quantum number determines the main energy shell or level in which the electron is present. It can have whole number values starting from 1 in an atom.

The principle quantum number indicates the average distance of the electron from the nucleus. If n = 1, it is closest to the nucleus and has lowest energy.

Eariest practice was to number shells as K,L,M,N etc.
Shell with principal quantum number n = 1 is called K.
Shell with principal quantum number n = 2 is called etc.

2. Azimuthal quantum number or angular quantum number (l): This number determines the angular momentum of the electron.

It can have positive integer values from zero to (n-1) where n is the principal quantum number. For each value of n, there are n possible values of l.

For n =3, l has three values: l = 0,1,2

The earlier practice is to designate l as subshell and refer it by letters s,p,d,f,….

l=0 = s; l=1=p; l=2=d, l=3=f etc.

The energy of subshell increases with increasing value of l.

3. Magnetic quantum number ( m_l): Magnetic field acts on moving electrical charges. ( from chapters on magnetism in physics syllabus). On revolving electrons external magnetic field of the earth acts. Therefore, the electrons in a given subshell orient themselves in certain preferred regions space around the nucleus. These are called orbitals. This quantum number gives the number of orbitals for given angular quantum number l or in a given subshell.

The allowed values of m_l are –l through 0 to +l.

There are (2l+1) values of m_l for each value of l.

If l = 0, m_l has only one value. m_l = 0.

If l = 3, m_l has 7 values.
m_l = -3,-2,-1,0,1,2,3

4. Spin quantum number (m_s) : It is observed that the electron in an atom is not only revolving around the nucleus but is also spinning around its own axis. This quantum number describes the spin orientation of the electron.

The electron can spin in two ways – clockwise and anticlockwise.
Values of +1/2 and -1/2 are given to this quantum number. Its value is not dependent on other quantum numbers.

The orientations of spin are also designated by up and down arrows ↑ ↓.


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Rao IIT Academy

XII - 1.4 Wave Mechanical Model of Atom and Concept of Atomic Orbital - Video Lectures

XII -

1.4 Wave Mechanical Model of Atom and Concept of Atomic Orbital - Video Lectures

1.4 Wave Mechanical Model of Atom and Concept of Atomic Orbital

Quantum mechanics or wave mechanics is a theoretical science which deals with the study of the motion of the microscopic objects (like electron) which have both observable wave like and particle like properties.

Quantum mechanics was developed indepdendently in 1926 by Werner Heisenberg and Erwin Schrodinger. In 1927, Schrodinger wave equation was published.

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ChemTutor BYU

XI - 1.10 Atomic and Molecular Mass - Video Lectures

XI -

1.10 Atomic and Molecular Mass - Video Lectures

1.9 Atoms and Molecules

Atoms are the smallest particle of an element which may or may not have independent existence, but it takes part in chemical reactions.

Molecule is the smallest particle of a substance (element or compound) capable of independent existence.

There are 112 elements in nature as per the present knowledge. Obviously there are 112 different types of atoms.

Molecules may have two or more atoms. Molecules are divided into two types:

1. Homoatomic molecules

2. Heteroatomic molecules

1.10 Atomic and Molecular Mass

Atomic Mass: An atom is such a small particle that its mass cannot be determined with the help of any available balance.  By an indirect method, the absolute mass of hydrogen atom has been found to be 1.66 X 10 to the power of -24 g.

Chemists have defined the atomic mass of hydrogen as one and expressed the atomic mass of other elements as multiples of hydrogen's atomic mass. The relative atomic masses expressed in terms of hydrogen as termed atomic weights.

Atomic weights of some common elements

Metals

Magnesium  24.3
Aluminium  27.0
Iron              55.8
Zinc             65.4
Tin             118.7

Gram Atomic Mass: Gram atomic mass is the quantity of an element whose mass in grams is numerically equal to its atomic mass.

Molecular Mass: Measuring the mass of a single molecule is not possible. Hence as in the case of atoms,  molecular mass is defined as the average relative mass of its molecule as compared to the mass of an atom of carbon (C12) having mass number 12.

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