Quantum number : Class 9th Notes -n

Quantum number 

A set of four identification number which are needed to have full information about various electrons present in atoms. To know about their location energy orientation or spins are needed.

Types of quantum number :- 

1. Principal quantum number

2. Azimuthal quantum number

3. Magnetic quantum number

4. Spin quantum number

Principal quantum number

it describes the average distance of the electrons from the nucleus.

• It represents the number of shells

• principal quantum number is denoted by "n"

shell ⇒ K , L , M , N

n ⇒       1 ,  2 ,  3 ,  4

r1  < r2  < r3  < r4

E1 < E2 < E3 < E4

Azimuthal quantum number 

• It tells about the shape of orientation, occupied by the electrons.

• the azimuthal quantum number is denoted by "l"

• the azimuthal quantum number can have all integral value from 0 to n-1.

• each value corresponds to an energy level called sub-shells or sub-level.

K-shell	L-shell	M-shell	N-shell
n = 1	n = 2	n = 3	n = 4
l = 0	l = 0,1	l = 0,1,2	l = 0,1,2,3
S = Sub shell	l ⇒ 0 → S	l ⇒ 0 → S	l ⇒ 0 → S
	1 → P	1 → P	1 → P
		2 → D	2 → D
			3 → F

Magnetic quantum number 

• It indicates about the orbital.

• it is indicated by m.

                m = (-l) to (+l) 

Spin quantum number 

• It tells about a spinning of the electron.

Electronic configuration of an atom :- 

the knowledge about the distribution of electrons in various energy shells, sub-shells and orbitals in an atom in the ground state is known as electronic configuration.

• The electronic configuration of an atom is written by terms of  nl^x

n = number of shells

l = sub-shells

x = number of electrons

E.g.- 6P^s → 5e^- are present in P subshell of 6th shell.

Aufbau principle :- the Subshall with minimum energy is filled up first and when this obtained the maximum quantity of electrons then next subshell of higher energy starts filling.

Chemical bonding :-

• only elements with the noble gas configuration have no tendency to react further, since they have completely filled outermost electronic configuration of their orbitals or shells.

• an electronic structure of noble gas type gives stability to an element making it less reactive. 

• most elements don't have noble gas configuration and consequently aim to achieve it.

• they attend this noble gas configuration by combining with each other either by gaining or losing electrons.

Valency :- 

• The valency is tendency of element two combined with one another since the valency depends upon the electronic configuration of atom.

Lewi's symbol for electron dot symbol :- 

• during formation of molecules The outermost electron of the atom are involved. they are referred valence electrons.

• the inner shell orbits are are fully filled and well protected. thus, the inner shell electrons do not participate in bond formation.

 G.N. Lewis represented the atom by simple symbol and outer shell electron as dots. surrounded of the atom. these symbols are called Lewis symbol for electron dot symbol.

Significance of Lewis symbol :- 

• The number of dots represent the number of electrons present in the valence shell of atom.

Chemical bond 

Types of chemical bond

1. Electrovalent bond or ionic bond

2. Covalent bond 

3. Co-ordinate bond or dative bond

nature of bond formed between two atoms depends upon electropositive and electronegative character of the bond atoms.

Electropositive :- acquire positive charge by losing of electrons.

Example :- metals

Electronegative :- acquire a negative charge by gaining of electron.

Example :- non metals

• Electropositive + electronegative = ionic bond

• electropositive + electropositive = metallic bond

• electronegative + electronegative = covalent bond

Ionic bond :- the atomic bond which are formed by the transfer of electrons between the constituent atom of a compound are known as ionic bond.

Cation + anion = ionic bond

Metal + non metal = ionic bond

Factors affecting the formation of ionic compound :- 

There are three main factors :- 

1. Ionization energy

The amount of energy required to remove one electron from the valence shell of an isolated gaseous atom to form a positive gaseous ions.

• lesser the value of ionization energy, greater is the tendency of the atom to form cation by losing valence electron.

2. Electron affinity or electron gain in tholphy 

Electron affinity of an element is the amount of energy released when an electron is added to an isolated gaseous atom to form and anion.

• higher the value of electron affinity greater will be the tendency of the atom to form anion.

3. Lattice energy

it is the amount of energy released when one mole of ionic solid is formed by combination of requisite number of cation and anion.

• higher the value of lattice energy, greater will be the stability of ionic compound and thus, greater will be the ease of it's formation.

Covalent bond :- atomic bond formed by sharing of electrons is called covalent bond

Example :- 

Types of covalent bond :- 

On the basis of number of bond

1. Single bond

Example:- H₂, Cl₂

2. Double bond

Example:- O₂

3. Triple bond

Example:- N₂

Polar and nonpolar covalent bond :- 

When the electronegativity of the atoms differs, a dipole may form and the bond is called polar bond.

Example :- H₄, HCl

• a polar bond depends on the difference between the electronegativity value of the two atoms.

• the greater the electronegativity difference, the more polar the bond.

Non-polar bond :-

when the two atoms in covalent bond have the same electronegativity no dipole is formed and the bond is called non-polar bond.

Example :- Cl₂, O₂, H₂

Co-ordinate bond or Dative bond or semi-polar bond :-

a covalent bond in which both electron of the shared pair are contributed by one of the two atoms.

= A ➝ B

• the atom which contributes electron pair is called donor while the atom which accept it is called acceptor.

• co-ordinate bond is a combination of electrovalent and covalent bond.

Properties of ionic compound :- 

1. Crystalline in nature 

• Electrovalent compound are usually crystalline in nature.

• the constituent unit in an ionic crystal are ions not molecules. These ions are arranged together in a regular way in an ionic lattice.

• force of attraction between the ions is non-directional and extended in all direction.

• electrovalent compounds consists of three dimensional solid aggregates.

• electrovalent compounds have closed packed structured and ions have no freedom of movement.

• co-ordinate number :- each ion is surrounded by a number of oppositely charged ions and this number is called co-ordinate number.

2. Melting and boiling point 

• electrovalent compound process high melting and boiling point due to strong electrostatic force of attraction. ions are held tightly in their positions in the crystal lattice.

3. Hard and brittle 

• electrovalent compounds are hard in nature. this hardness is due to strong force of attraction between oppositely charged ions, which keep them in their ellotod positions.

• the brittleness of crystal is due to movement of a large crystal on the other layer by application of external force. when like ions comes in front of each other the force of repulsion come into play the breaking of crystal in account of repulsion.

H₂SO₄, H₂SO₃, H₃PO₄, N₂O₃, HNO₃

N₂O₃ ⇒ N₂O + O₂

Two essential condition for the formation of a chemical bond :- 

1. there should be maximum overlapping of the involved atomic orbitals of the two atoms.

2. each of the two involved atomic orbitals must have one unpaired electron with an opposite spin.

Sigma bond :- a bond formed between two atoms by the overlap of singlly occupied orbitals along their axis ( end to end overlap) is called sigma bond.

Or

bond orbital which is symmetrical with the line joining the two nuclei is known as sigma bond.

• It is formed by head on or axial overlap.

• in sigma bond maximum overlap is possible between electron cloud so it is a strong bond.

Sigma bond are formed by 3 types of overlapping 

1. S-S overlapping 

E.g.- H₂ → H + H

2. S-P overlapping

E.g.- HF → H = 1S¹  F = 1S² 2S² 2P⁵

3. P-P overlapping

E.g.- F₂

Pi bond 

• Sidewise overlapping or literally overlapping. 

Comparison of Sigma and Pie bond 

Sigma bond 

• the bond is formed by the overlapping of orbitals along their axis.

• it includes . S-S overlapping, S-P overlapping, P-P overlapping

• it is a strong bond

• electron cloud is symmetrical about the line joining the two nuclei

• they can be free rotation of atom around this bond.

• they are less reactive.

• the shape of the molecule is determined by these bond.

• Sigma electron are referred as localised.

• sigma bond can have an independent existence.

Pie bond

• The bond is formed by sidewise overlapping of orbitals.

• it includes P-P overlapping.

• it is a weak bond.

• electron cloud is unsymmetrical.

• free rotation is not possible around this bond.

• these are more reactive.

• these bonds do not affect the shape of the molecule

• Pie electron are referred as mobile electron.

• pie bond always exist along with a sigma bond.

All single bond :- ---____###@@@

Double bond :- ___$$##@@@@@

Bond energy :- increases from a single bond to a double bond.

Single bond < double bond < triple bond

Bond strength :- increases from a single bond to a triple bond.

Single bond < double bond < triple bond

Bond length :- multiple bond (double bond or triple bond) is always shorter than the corresponding single bond.

Single bond > double bond > triple bond 

Reactivity :- a multiple bond is always more reactive than single bond this is due to the fact that Pi electron are mobile in nature.

Triple bond > double bond > single bond