Importance of Core Revision Points: Core Revision Points are important because if you remember them strongly, many more points related to them will come out of your memory and help you to answer question and problems. Read them many times and make sure you remember them very strongly.
Sections in the Chapter
1.1 Dual Nature of Radiation
1.2 Dual Nature of Matter - de-Broglie Equation
1.3 Heisengberg's Uncertainty Principle
1.4 Wave Mechanical Model of Atom and Concept of Atomic Orbital
1.5 Quantum Numbers
1.6 Pauli's Exclusion Principles
1.7 Orbital Wave Functions and Shapes of Orbitals
1.8 Electronic Configurations of Atoms
1.9 Chemical Bonding
1.10 Review of Valency Bond Theory
1.11 Molecular Orbital Theory
1.12 Linear Combination of Atomic Orbitals (LCAO) Method
1.13 Relative Energies of Bonding and Antibonding Molecular Orbitals
1.14 Combination of 2s and 2p Atomic Orbitals to form Molecular Orbitals
1.15 Conditions for the Combination of Atomic Orbitals
1.16 Energy Level diagram for Molecular Orbitals
1.17 Rules for Filling Molecualr Orbitals
1.18 Electronic Configurations and Molecular Behavior
1.19 Bonding in Some Diatomic Molecules
1.20 Metallic Bond
1.21 Hybridisation
1.22 Intermolecular Forces
1.23 Hydrogen Bonding
1.1 Dual Nature of Radiation
Einstein in 1905 suggested that light has a dual character - Particle nature as well as wave.
Wave like character of light was proposed by Huygens. In 1856, James Maxwell proposed that light and other forms of radiation propagate though space in the form of waves and these waves have electric and magnetic fields associated with it. Therefore, the light which is travelling through radiation is said to be composed of electromagnetic waves.
Planck's Quantum Theory of Radiation
1.2 Dual Nature of Matter - de-Broglie Equation
In 1924, Louis de Broglie suggested that similar to light, all microscopic material particles in motion have dual character.
Uncertainty principle
In 1927, Heisenberg put forward a principle known as Heisenberg’s uncertainty principle.
According it, “it is not possible to measure simultaneously both the position and momentum (or velocity) of a microscopic particle, with absolute accuracy.”
Mathematically, this principle is expressed as:
∆x * ∆p = h/4 π
Where
∆x = uncertainty in position
∆p = uncertainty in momentum
The constancy of the product of uncertainties means that, if the position of the particle is known with more accuracy, there will be large uncertainty in momentum and vice versa.
This uncertainty arises, as all observations are made by impact of light, the microscopic objects suffer a change in position or velocity as a result of impact of light. So there is a disturbance in them due to the measurement.
The principle does not affect the measurement of large objects as in these cases impact of light does not created any appreciable change in their position or velocity.
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.
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,ml, and ms.
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 ( ml): 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 ml are –l through 0 to +l.
There are (2l+1) values of ml for each value of l.
If l = 0, ml has only one value. ml = 0.
If l = 3, ml has 7 values.
ml = -3,-2,-1,0,1,2,3
4. Spin quantum number (ms) : 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 ↑ ↓.
Pauli's exclusion principle: No two electrons can have all four same quantum numbers
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
1.8 Electronic Configurations of Atoms
1. Aufbau principles
2. Pauli's exclusion principle: No two electrons can have all four same quantum numbers
3. Hund's rule of maximum multiplicity
1.9 Chemical Bonding
1. Valency bond theory 2. Molecular orbital theory
1.10 Review of Valency Bond Theory
Valency bond theory was proposed by Heitler and London in 1927 and it was further developed by Linus Pauling.
The basic idea of the theory are:
1. A covalent bond is formed by the overlap of half-filled atomic orbitals of the different atoms.
2. The overlapping atomic orbitals must have electrons with opposite spins.
1.11 Molecular Orbital Theory
This theory was proposed by Hund and Mulliken in 1932. The basic idea of the theory is that atomic orbitals of individual atoms combine to form molecular orbitals.
1.12 Linear Combination of Atomic Orbitals (LCAO) Method
According to LCAO method, the orbitals are formed by the linear combination (addition or subtraction) of atomic orbitals of the atoms which form the molecule.
1.13 Relative Energies of Bonding and Antibonding Molecular Orbitals
1.14 Combination of 2s and 2p Atomic Orbitals to form Molecular Orbitals
2s-orbitals combine by addition and subtraction to form bonding and antibonding molecular orbitals.
1.15 Conditions for the Combination of Atomic Orbitals
Main Conditions for the Combination of Atomic Orbitals
1. The combining atomic orbitals should not differ much in energies.
2. The extent of overlapping between the atomic orbitals of two atoms should be large.
3. The combining atomic orbitals between the atomic orbitals of two atoms should be large.
1.16 Energy Level diagram for Molecular Orbitals
1.17 Rules for Filling Molecualr Orbitals
1. Aufbau principles
2. Pauli's exclusion principle: No two electrons can have all four same quantum numbers
3. Hund's rule of maximum multiplicity
1.18 Electronic Configurations and Molecular Behavior
The important information conveyed by Electron Configuration of a molecule is:
1. Stability of a molecule
2. Bond Order
1.19 Bonding in Some Diatomic Molecules
1. Hydrogen molecule.
1.20 Metallic Bond
More than 80 elements in the periodic table are metals.
The force which holds together atoms of metals is called metallic bond.
1.21 Hybridisation
Hybridizastion is the phenomenon of intermixing of the orbitals of slightly different energies so as to redistribute their energies and to give new set of orbitals of equivalent energy and shape.
1.22 Intermolecular Forces
In addition to normal covalent bond, ionic bond, and metallic bond, there are weak attractive intermolecular forces which occur in all kinds of molecular solids. These are present in case of non-polar molecules such as H2, O2, CO2, CH4 etc. also.
These are classified as:
i) Dipole-dipole forces
ii) Dipole induced dipole forces
iii) Instantaneous dipole-instantaneous induced dipole forces (called London forces)
iv) Hydrogen bonding
1.23 Hydrogen Bonding
When hydrogen atom is bonded to atoms of highly electronegative elements such as fluorine, oxygen, or nitrogen, the hydrogen atom forms a weak bond with the electronegative atom of the other molecule.
Updated 4 May 2016
First Posted on 23 May 2015
Sections in the Chapter
1.1 Dual Nature of Radiation
1.2 Dual Nature of Matter - de-Broglie Equation
1.3 Heisengberg's Uncertainty Principle
1.4 Wave Mechanical Model of Atom and Concept of Atomic Orbital
1.5 Quantum Numbers
1.6 Pauli's Exclusion Principles
1.7 Orbital Wave Functions and Shapes of Orbitals
1.8 Electronic Configurations of Atoms
1.9 Chemical Bonding
1.10 Review of Valency Bond Theory
1.11 Molecular Orbital Theory
1.12 Linear Combination of Atomic Orbitals (LCAO) Method
1.13 Relative Energies of Bonding and Antibonding Molecular Orbitals
1.14 Combination of 2s and 2p Atomic Orbitals to form Molecular Orbitals
1.15 Conditions for the Combination of Atomic Orbitals
1.16 Energy Level diagram for Molecular Orbitals
1.17 Rules for Filling Molecualr Orbitals
1.18 Electronic Configurations and Molecular Behavior
1.19 Bonding in Some Diatomic Molecules
1.20 Metallic Bond
1.21 Hybridisation
1.22 Intermolecular Forces
1.23 Hydrogen Bonding
Revision Points for Various Sections in the Chapter
1.1 Dual Nature of Radiation
Einstein in 1905 suggested that light has a dual character - Particle nature as well as wave.
Wave like character of light was proposed by Huygens. In 1856, James Maxwell proposed that light and other forms of radiation propagate though space in the form of waves and these waves have electric and magnetic fields associated with it. Therefore, the light which is travelling through radiation is said to be composed of electromagnetic waves.
Planck's Quantum Theory of Radiation
1.2 Dual Nature of Matter - de-Broglie Equation
In 1924, Louis de Broglie suggested that similar to light, all microscopic material particles in motion have dual character.
1.3 Heisengberg's Uncertainty Principle
Uncertainty principle
In 1927, Heisenberg put forward a principle known as Heisenberg’s uncertainty principle.
According it, “it is not possible to measure simultaneously both the position and momentum (or velocity) of a microscopic particle, with absolute accuracy.”
Mathematically, this principle is expressed as:
∆x * ∆p = h/4 π
Where
∆x = uncertainty in position
∆p = uncertainty in momentum
The constancy of the product of uncertainties means that, if the position of the particle is known with more accuracy, there will be large uncertainty in momentum and vice versa.
This uncertainty arises, as all observations are made by impact of light, the microscopic objects suffer a change in position or velocity as a result of impact of light. So there is a disturbance in them due to the measurement.
The principle does not affect the measurement of large objects as in these cases impact of light does not created any appreciable change in their position or velocity.
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.
1.5 Quantum Numbers
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,ml, and ms.
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 ( ml): 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 ml are –l through 0 to +l.
There are (2l+1) values of ml for each value of l.
If l = 0, ml has only one value. ml = 0.
If l = 3, ml has 7 values.
ml = -3,-2,-1,0,1,2,3
4. Spin quantum number (ms) : 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 ↑ ↓.
1.6 Pauli's Exclusion Principles
Pauli's exclusion principle: No two electrons can have all four same quantum numbers
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
1.8 Electronic Configurations of Atoms
1. Aufbau principles
2. Pauli's exclusion principle: No two electrons can have all four same quantum numbers
3. Hund's rule of maximum multiplicity
1.9 Chemical Bonding
1. Valency bond theory 2. Molecular orbital theory
1.10 Review of Valency Bond Theory
Valency bond theory was proposed by Heitler and London in 1927 and it was further developed by Linus Pauling.
The basic idea of the theory are:
1. A covalent bond is formed by the overlap of half-filled atomic orbitals of the different atoms.
2. The overlapping atomic orbitals must have electrons with opposite spins.
1.11 Molecular Orbital Theory
This theory was proposed by Hund and Mulliken in 1932. The basic idea of the theory is that atomic orbitals of individual atoms combine to form molecular orbitals.
1.12 Linear Combination of Atomic Orbitals (LCAO) Method
According to LCAO method, the orbitals are formed by the linear combination (addition or subtraction) of atomic orbitals of the atoms which form the molecule.
1.13 Relative Energies of Bonding and Antibonding Molecular Orbitals
1.14 Combination of 2s and 2p Atomic Orbitals to form Molecular Orbitals
2s-orbitals combine by addition and subtraction to form bonding and antibonding molecular orbitals.
1.15 Conditions for the Combination of Atomic Orbitals
Main Conditions for the Combination of Atomic Orbitals
1. The combining atomic orbitals should not differ much in energies.
2. The extent of overlapping between the atomic orbitals of two atoms should be large.
3. The combining atomic orbitals between the atomic orbitals of two atoms should be large.
1.16 Energy Level diagram for Molecular Orbitals
1.17 Rules for Filling Molecualr Orbitals
1. Aufbau principles
2. Pauli's exclusion principle: No two electrons can have all four same quantum numbers
3. Hund's rule of maximum multiplicity
1.18 Electronic Configurations and Molecular Behavior
The important information conveyed by Electron Configuration of a molecule is:
1. Stability of a molecule
2. Bond Order
1.19 Bonding in Some Diatomic Molecules
1. Hydrogen molecule.
1.20 Metallic Bond
More than 80 elements in the periodic table are metals.
The force which holds together atoms of metals is called metallic bond.
1.21 Hybridisation
Hybridizastion is the phenomenon of intermixing of the orbitals of slightly different energies so as to redistribute their energies and to give new set of orbitals of equivalent energy and shape.
1.22 Intermolecular Forces
In addition to normal covalent bond, ionic bond, and metallic bond, there are weak attractive intermolecular forces which occur in all kinds of molecular solids. These are present in case of non-polar molecules such as H2, O2, CO2, CH4 etc. also.
These are classified as:
i) Dipole-dipole forces
ii) Dipole induced dipole forces
iii) Instantaneous dipole-instantaneous induced dipole forces (called London forces)
iv) Hydrogen bonding
1.23 Hydrogen Bonding
When hydrogen atom is bonded to atoms of highly electronegative elements such as fluorine, oxygen, or nitrogen, the hydrogen atom forms a weak bond with the electronegative atom of the other molecule.
Updated 4 May 2016
First Posted on 23 May 2015
