IIT JEE Ch.3. ATOMIC STRUCTURE Core Points for Revision

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.


JEE Syllabus

Atomic structure
Bohr model, spectrum of hydrogen atom, quantum numbers;
Wave-particle duality, de Broglie hypothesis;
Uncertainty principle;
Quantum mechanical picture of hydrogen atom (qualitative treatment), shapes of s, p and d orbitals;
Electronic configurations of elements (up to atomic number 36);
Aufbau principle;
Pauli's exclusion principle and Hund's rule;
----------------

3.1 Fundamental Particles

Electron was discovered by J.J. Thomson at the end of 19th century.

An electron may be defined as a sub-atomic particle which carries one unit negative charge (1.602*10^-19c) and has a mass (9.1*10^-31 kg) equal to 1/1837th of that of hydrogen atom.

Proton is a sub-atomic particle which carries one unit positive charge (1.602*10^19 coulombs) and has mass (1.67*10^-27 kg) which is equal to that of an atom of hydrogen.

3.2 arranging electrons and protons in an atom
3.3 Rutherford's Scattering Experiment
3.4 Concept of Atomic Number and Discovery of Neutron

3.5 Developments Leading to the Bohr Model of Atom

1. Dual behaviour of the electromagnetic radiation.
2. Atomic spectra

Thomson Model

J.J. Thomson studied the properties of cathode rays. Cathode rays were observed in tubes with gas at low pressures when electric charge was applied. The gas starts conducting electricity at low pressure and rays appear. During these studies, Thomson discovered electrons in 1897.

The experiments led to the conclusion that the particles comprising cathode rays were the same irrespective of the material of the cathode and the gas used in discarge tubes, The particles had the same mass and charge. Hence it was concluded that electrons are universal constituents of all matter.


Subsequently proton was also discovered. Rutherford's name can be mentioned in the case of proton as an important researcher.

Thomson proposed that the positive charge is spread over a sphere in which the electrons are embedded. This make the atom neutral. The model was also called Thomson's plum pudding model.

Rutherford model


In 1913, Niels Bohr proposed a model of the atom. He proposed that the electrons in an atom could only be in certain orbits, or energy levels, around the nucleus. Refinement of Bohr theory led to the modern theory of atomic structure based on quantum mechanics.

3.6 Nature of Light and Electromagnetic Radiation

Characteristics of wave motion

1. Wave length
2. Frequency
3. Velocity
4. Amplitude

Electromagnetic spectrum

3.7 Particle Nature of Electromagnetic Radiation and Planck's Quantum Theory

3.8 Atomic Spectra

Emission spectra
Absorption spectra

Emission spectrum from Hydrogen atom


3.9 Failure of Rutherford Model

3.10 Concepts of Energy Levels or Orbits

Bohr's model of the atom

In 1913, Niels Bohr proposed a model of the atom. He proposed that the electrons in an atom could only be in certain orbits, or energy levels, around the nucleus. Refinement of Bohr theory led to the modern theory of atomic structure based on quantum mechanics.

Bohr's model is based on particle theory.

3.11 Modern Concept of Structure of Atom

As wave-particle duality which says that all micromatter particle exhibit duality a new model is proposed.

Debroglie p = h/λ

3.12 Wave Mechanical Model of Atom

Uncertainty principle Δp* Δx ≥h/4π

3.13 Quantum Numbers

The quantum mechanics model

Principal quantum number-Shell,Azimuthal quantum number-sublevel,Magnetic quantum number-orbital, spin quantum number

The orbits are called as shells. The energy level of orbits or shells increases as they increase in distance from the nucleus of the atom. The orbits or shells are represented by numbers as 1,2,3,4,5,6 or 7. They are represented by letters as K,L,M,N,O,P,Q.

Sublevel of an Orbit

The energy levels, or orbits or shells are further divided into sublevels, or subshells. These subshells are designated by letters: s for the first possible sublevel, p for the second possible sublevel, d for the third, f for the fourth, g for the fifth, and from here on they simply go in alphabets.

The number of sublevels of each energy level is equal to the number of the energy levels. This means energy level 1, the K shell will have only one sub levels – s sublevel. The energy level 2, the L shell will have 2 sub levels – s and p.


Orbitals

Sublevels have further divisions called orbitals. Electrons are found in these orbitals. Each orbital contains two electrons.

“s” sublevel has only one orbital. “p” sublevel has 3 orbitals. “d” sublevel has 5 orbitals. “f” sublevel has 7 orbitals.

The two electons in each orbital spin in different directions.


3.14 Pauli's Exclusion Principles

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

3.15 Shapes of Orbitals
Shape 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.

3.16 Energy Level Diagram for Electrons in an Atom
3.17 Electron Configurations of Atoms


Electrons occupy the lowest energy sublevels that are available. This is known as ‘aufbau’ order or principles.


Hund’s rule says that, for any set of orbitals of equal energy say p orbitals of orbit 2, there is one electron is each orbital before the second electron enters or occupies an orbital.

The energy level of some sublevels at higher orbits is less than the some sublevels at lower orbitals. This fact is to be kept in mind when electron configuration is determined for any atom. The increasing order of energy levels of sublevels is:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f*, 5d, 6p, 7s, 5f*, 6d, 7p, 8s




















Jauhar Chapter Contents

3.1 Fundamental Particles
3.2 arranging electrons and protons in an atom
3.3 Rutherford's Scattering Experiment
3.4 Concept of Atomic Number
3.5 Developments Leading to the Bohr Model of Atom
3.6 Nature of Light and Electromagnetic Radiation
Particle Nature of Electromagnetic Radiation and Planck's Quantum Theory
3.8 Atomic Spectra
3.9 Failure of Rutherford Model
3.10 Concepts of Energy Levels or Orbits
3.11 Modern Concept of Structure of Atom
3.12 Wave Mechanical Model of Atom
3.13 Quantum Numbers
3.14 Pauli's Exclusion Principles
3.15 Shapes of Orbitals
3.16 Energy Level Diagram for Electrons in an Atom
3.17 Electron Configurations of Atoms


Updated 8 January 2020,  21 May 2015
Published first 19 January 2008

Ch.1 Atomic Structure and Chemical Bonding - JEE Main Core Revision Points

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

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,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 ↑ ↓.

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

JEE - Study Guide - 3. Atomic Structure

Text Book
Modern's abc of Chemistry by Dr. S.P. Jauhar for Class XI CBSE

Sections in the Chapter

3.1 Fundamental particles
3.2 Thomson’s Atomic Model
3.3 Rutherford’s Scattering Experiment
3.4 Concept of Atomic Number and Discovery of Neutron
3.5 Developments Leading to the Bohr Model of Atom
3.6 Nature of Light and Electromagnetic Radiation
3.7 Particle Nature of Electromagnetic Radiation and Planck’s Quantum Theory
3.8 Atomic Spectra
3.9 Failure of Rutherford Model
3.10 Concept of Energy Levels or Orbits
3.11 Modern Concept of Structure of an Atom: Quantum Mechanical Model
3.12 Wave Mechanical Model of Atom and Concept of Atomic Orbital
3.13 Quantum Numbers
3.14 Pauli’s Exclusion Principle
3.15 Shapes of Orbitals or Boundary Surface Diagrams
3.16 Energy Level Diagram for Electrons in an Atom
3.17 Electronic Configuration of Atoms


Conceptual Questions with Answers: 35
Additional Numerical Problems for Practice: 8
Revision Exercises
Very Short Answer questions 16
Short Answer Questions 42
Long Answer Questions 3

Competition File
Numerical Problems 20
Objective Questions: 47
Fill in the blanks: 10
True or False: 10


Study Plan

Day 1

3.1 Fundamental particles
3.2 Thomson’s Atomic Model
3.3 Rutherford’s Scattering Experiment

Day 2

3.4 Concept of Atomic Number and Discovery of Neutron
Ex. 3.1 to 3.3
Practice Problems 3.1 to 3.9

Day 3

3.5 Developments Leading to the Bohr Model of Atom
3.6 Nature of Light and Electromagnetic Radiation

Ex. 3.4, 3.5,3.6
P.P. 3.10 to 3.15



Day 4

3.7 Particle Nature of Electromagnetic Radiation and Planck’s Quantum Theory
Ex. 3.7 to 3.16

Day 5
P.P 3.16 to 3.24

Day 6
3.8 Atomic Spectra
Ex. 3.17, 3.18
P.P. 3.25 t 3.27

3.9 Failure of Rutherford Model

Day 7
3.9 Failure of Rutherford Model
3.10 Concept of Energy Levels or Orbits
Ex. 3.19 to 3.23
P.P 3.28 to 3.32

Day 8
Bohr’s Theory and Concept of Quantisation
3.11 Modern Concept of Structure of an Atom: Quantum Mechanical Model
Ex. 3.24 to 3.29
P.P. 3.33, 3.34

Day 9

3.12 Wave Mechanical Model of Atom and Concept of Atomic Orbital

3.13 Quantum Numbers
3.14 Pauli’s Exclusion Principle


Day 10
Ex. 3.30 to 3.33
P.P. 3.35 to 3.37

3.15 Shapes of Orbitals or Boundary Surface Diagrams


Day 11

3.16 Energy Level Diagram for Electrons in an Atom
3.17 Electronic Configuration of Atoms

Day 12
P.P 3.38 to 3.46

Revision Period

Day 13
Conceptual Questions with Answers: 1 to 18

Day 14

Conceptual Questions with Answers: 19 to 35

Day 15
Additional Numerical Problems for Practice: 8

Day 16
Very Short Answer questions 16

Day 17
Short Answer Questions 42

Day 18
Study Competition File

Day 19
Numerical Problems 1 to 10

Day 20
Numerical Problems 11 to 20
Day 21

Objective Questions: 1 to 24

Day 22
Objective Questions: 25 to 47

Day 23

Fill in the blanks: 10

Day 24

True or False: 10


Updated 1 May 2016,  11 Mar 2009

Ch.3. ATOMIC STRUCTURE - JEE - CBSE Class XI - Revision Problems

_______________

_______________
Paras Thakur

http://www.questionpapers.net.in/chemistry/structure-of-atom-test-1.html

http://examtimequiz.com/chemistry-mcq-on-atomic-structure/


http://makox.com/chemistry/tag/atomic-structure-mcq-series/

http://schools.aglasem.com/58728

http://gotest.pk/chemistry-online-mcqs/atomic-structure-quiz-test/

http://fachschaften.kst.ch/chemie/chicd/kap2/kap2mce.html

http://pattersonscience.weebly.com/uploads/5/1/5/0/5150508/unit_02_review_for_quiz_1.pdf

https://chemistlibrary.files.wordpress.com/2015/02/5-chapter-atomic-structure-mcqs.pdf


http://www.marsd.org/cms/lib7/NJ01000603/Centricity/Domain/304/Atomic_Structure_Stoich_Study_Guide_Answer_Key.pdf

IIT JEE Chemistry-12th Portion - Study Guide - 1. Atomic Structure and Chemical Bonding

IIT JEE Chemistry 12th class portion starts from here.


Based on Book by Jauhar CBSE 12th class

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


Study Plan

Day 1


1.1 Dual Nature of Radiation
      Video Lectures - 1.1 Dual Nature of Radiation Class XII Chemistry
Pratice Problems 1.1 to 1.12

Day 2

1.2 Dual Nature of Matter - de-Broglie Equation
P.P. 1.13 to 1.22

Day 3

1.3 Heisengberg's Uncertainty Principle
P.P. 1.23 to 1.29

Day 4
1.4 Wave Mechanical Model of Atom and Concept of Atomic Orbital
P.P. 1.30 to 1.31

Day 5
1.5 Quantum Numbers
1.6 Pauli's Exclusion Principles
P.P. 1.32 to 1.35


Day 6
1.7 Orbital Wave Functions and Shapes of Orbitals
1.8 Electronic Configurations of Atoms

Day 7

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

Day 8

1.14 Combination of 2s and 2p Atomic Orbitals to form Molecular Orbitals
1.15 Conditions for the Combination of Atomic Orbitals
P.P. 1.36 to 1.37
1.16 Energy Level diagram for Molecular Orbitals
1.17 Rules for Filling Molecualr Orbitals
1.18 Electronic Configurations and Molecular Behavior

Day 9

1.19 Bonding in Some Diatomic Molecules
P.P. 1.38 to 1.45

Day 10
1.20 Metallic Bond

Day 11
1.21 Hybridisation
1.22 Intermolecular Forces

Day 12

1.23 Hydrogen Bonding
P.P. 1.44 to 1.47

Day 13

Additional numerical problems 1 to 20

day 14

Conceptual questions 1 to 40

Day 15
Key facts to remember
Revision exercises very short answer questions 1 to 30

Revision Period (30 minutes per day)

Day 16

Revision exercises very short answer questions 31 to 50

day 17
Revision exercises very short answer questions 51 to 65

Day 18

Revision exercises short answer questions 1 to 15

Day 19
Revision exercises short answer questions 16 to 30

Day 20
Revision exercises short answer questions 31 to 45

Day 21
Revision exercises short answer questions 46 to 60

Day 22
Revision exercises short answer questions 61 to 78

Day 23
Some useful facts for competitive examinations
Nemerical problems for competitive examinations 1 to 5

Day 24
Nemerical problems for competitive examinations 6 to 13

Day 25
Multiple choice questions 1 to 20

Day 26

Multiple choice questions 21 to 40

Day 27
Multiple choice questions 41 to 63

Day 28
Fill in the blanks 1 to 15

Day 29
True or false 15

Day 30
Matching
Concept revision



____________________

____________________
Frankly Chemistry



Updated 31 Jan 2016, 2009

JEE Class XII Ch.1 Atomic Structure and Chemical Bonding - Revision Questions

http://www.gcsescience.com/a43-atomic-structure-revision-questions.htm





Ch.1 Atomic Structure and Chemical Bonding - Revision Questions
2. Solid State - JEE - Revision Question - Class XII
3. Solutions - JEE - CBSE Class XII - Revision Questions
4. Chemical Thermodynamics - JEE - CBSE Class XII - Revision Questions
5. Electrochemistry - JEE - CBSE Class XII - Revision Questions
6. Chemical Kinetics - JEE - CBSE Class XII - Revision Questions

JEE Revision - Atomic structure - Basics

Thomson Model

J.J. Thomson studied the properties of cathode rays. Cathode rays were observed in tubes with gas at low pressures when electric charge was applied. The gas starts conducting electricity at low pressure and rays appear. During these studies, Thomson discovered electrons in 1897.

The experiments led to the conclusion that the particles comprising cathode rays were the same irrespective of the material of the cathode and the gas used in discarge tubes, The particles had the same mass and charge. Hence it was concluded that electrons are universal constituents of all matter.


Subsequently proton was also discovered. Rutherford's name can be mentioned in the case of proton as an important researcher.

Thomson proposed that the positive charge is spread over a sphere in which the electrons are embedded. This make the atom neutral. The model was also called Thomson's plum pudding model.

Rutherford model


In 1913, Niels Bohr proposed a model of the atom. He proposed that the electrons in an atom could only be in certain orbits, or energy levels, around the nucleus. Refinement of Bohr theory led to the modern theory of atomic structure based on quantum mechanics.

Wave mechanical model

Atomic structure and chemical bonding- Study Guide - IIT JEE

Bohr model, spectrum of hydrogen atom, quantum numbers; Wave-particle duality, de Broglie hypothesis; Uncertainty principle; Quantum mechanical picture of hydrogen atom (qualitative treatment), shapes of s, p and d orbitals; Electronic configurations of elements (up to atomic number 36); Aufbau principle; Pauli's exclusion principle and Hund's rule;

Orbital overlap and covalent bond; Hybridisation involving s, p and d orbitals only; Orbital energy diagrams for homonuclear diatomic species; Hydrogen bond; Polarity in molecules, dipole moment (qualitative aspects only); VSEPR model and shapes of molecules (linear, angular, triangular, square planar, pyramidal, square pyramidal, trigonal bipyramidal, tetrahedral and octahedral).





________________

________________

IIT JEE Revision - Ch.3. ATOMIC STRUCTURE Core Points

JEE Syllabus

Atomic structure
Bohr model, spectrum of hydrogen atom, quantum numbers;
Wave-particle duality, de Broglie hypothesis;
Uncertainty principle;
Quantum mechanical picture of hydrogen atom (qualitative treatment), shapes of s, p and d orbitals;
Electronic configurations of elements (up to atomic number 36);
Aufbau principle;
Pauli's exclusion principle and Hund's rule;


In 1913, Niels Bohr proposed a model of the atom. He proposed that the electrons in an atom could only be in certain orbits, or energy levels, around the nucleus. Refinement of Bohr theory led to the modern theory of atomic structure based on quantum mechanics.

Bohr's model is based on particle theory.

As wave-particle duality which says that all micromatter particle exhibit dualitya new model is proposed.

Debroglie p = h/λ

Uncertainty principle Δp* Δx ≥h/4π

The quantum mechanics model

Principal quantum number-Shell,Azimuthal quantum number-sublevel,Magnetic quantum number-orbital, spin quantum number

The orbits are called as shells. The energy level of orbits or shells increases as they increase in distance from the nucleus of the atom. The orbits or shells are represented by numbers as 1,2,3,4,5,6 or 7. They are represented by letters as K,L,M,N,O,P,Q.

Sublevel of an Orbit

The energy levels, or orbits or shells are further divided into sublevels, or subshells. These subshells are designated by letters: s for the first possible sublevel, p for the second possible sublevel, d for the third, f for the fourth, g for the fifth, and from here on they simply go in alphabets.

The number of sublevels of each energy level is equal to the number of the energy levels. This means energy level 1, the K shell will have only one sub levels – s sublevel. The energy level 2, the L shell will have 2 sub levels – s and p.


Orbitals

Sublevels have further divisions called orbitals. Electrons are found in these orbitals. Each orbital contains two electrons.

“s” sublevel has only one orbital. “p” sublevel has 3 orbitals. “d” sublevel has 5 orbitals. “f” sublevel has 7 orbitals.

The two electons in each orbital spin in different directions.


Shape 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.

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

Electrons occupy the lowest energy sublevels that are available. This is known as ‘aufbau’ order or principles.


Hund’s rule says that, for any set of orbitals of equal energy say p orbitals of orbit 2, there is one electron is each orbital before the second electron enters or occupies an orbital.

The energy level of some sublevels at higher orbits is less than the some sublevels at lower orbitals. This fact is to be kept in mind when electron configuration is determined for any atom. The increasing order of energy levels of sublevels is:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f*, 5d, 6p, 7s, 5f*, 6d, 7p, 8s

JEE Revision - Bohr model

The most important properties of atomic and molecular structure may be exemplified using a simplified picture of an atom that is called the Bohr Model.

This model was proposed by Niels Bohr in 1915;

It was revised later on as it is not completely correct, but it has many features that are approximately correct and they were used in the later models.

The correct theory of the atom is called quantum mechanics;

The Bohr Model is an approximation to quantum mechanics that has the virtue of being much simpler.

A Planetary Model of the Atom

The Bohr Model is probably familar as the "planetary model" of the atom. In the Bohr Model the neutrons and protons occupy a dense central region called the nucleus, and the electrons orbit the nucleus much like planets orbiting the Sun (but the orbits are not confined to a plane).

This similarity between a planetary model and the Bohr Model of the atom ultimately arises because the attractive gravitational force in a solar system and the attractive Coulomb (electrical) force between the positively charged nucleus and the negatively charged electrons in an atom are mathematically of the same form. (The form is the same, but the intrinsic strength of the Coulomb interaction is much larger than that of the gravitational interaction; in addition, there are positive and negative electrical charges so the Coulomb interaction can be either attractive or repulsive, but gravitation is always attractive in our present Universe.)


Orbits Are Quantized

Quantized energy levels in hydrogen

The basic feature of quantum mechanics that is incorporated in the Bohr Model and that is completely different from the analogous planetary model is that the energy of the particles in the Bohr atom is restricted to certain discrete values.

One says that the energy is quantized. This means that only certain orbits with certain radii are allowed; orbits in between simply don't exist.

These levels are labeled by an integer n that is called a quantum number.

The lowest energy state is generally termed the ground state. The states with successively more energy than the ground state are called the first excited state, the second excited state, and so on.

Beyond an energy called the ionization potential the single electron of the hydrogen atom is no longer bound to the atom. Then the energy levels form a continuum.

In the case of hydrogen, this continuum starts at 13.6 eV above the ground state ("eV" stands for "electron-Volt", a common unit of energy in atomic physics).

Spectrum of Hydrogen Atom

All electromagnetic waves have the same speed (3.0*10^8 m s^-1).
But different types of radiations differ from one another in their wavelengths and therefore, in frequency.

v = c/λ
C = speed = constant
λ = wavelength
v = frequency


The arrangement of different types of electromagnetic radiations in the order of increasing wavelengths (or decreasing frequencies) is known as electromagnetic spectrum.

When white light from sun is passed through a prism, it splits into a series of colour bands know as rainbow of colours: violet, indigo, blue, green, yellow, orange and red (VIBGYOR). This series of colour bands obtained is called a spectrum and in this spectrum of visible light there is continuity of colours, one colour merges into the other without any gap.

We can obtain spectrum of atoms. When the gases of vapour of a chemical substance are heated by electric spark, light is emitted. From such light emitted, the spectrum can be obtained using spectroscope.
It is observed that when radiations emitted by different substances are analysed, the spectrum obtained consists of share well defined lines each corresponding to a definited frequency (or wave length). The spectrum obtained from light emitted by chemical substances is called emission spectrum.

When electromagnetic radiation for example say white light is allowed to pass through a gas or a solution of some salt and the light that comes out is sent through a spectroscope, we obtain a spectrum. In this spectrum some dark lines are observed, in an otherwise continuous spectrum. This indicates that some radiations of specific wavelengths are absorbed by the substance. This spectrum is called absorption spectrum

Spectrum of hydrogen atom

The spectrum of hydrogen atom can be obtained by passing an electric discharge through the hydrogen gas in a discharge tube under low pressure.

It is observed that the spectrum consists of a large number of lines appearing in different regions of wavelengths. Some of the lines are present in visible region while others in ultra-violet and infra-red regions.

In 1885, J.J. Balmer developed a simple relationship among the different wavelengths of the series of visible lines in the hydrogen spectrum.

The relationship developed by Balmer for frequencies observed in hydrogen spe trum is:

1/ λ = v (in cm^-1) = 109677 (1/2² - 1/n²)

Where n is an integer that takes values equal to or greater than 3.
109677 cm^-1 is called Rydberg constant.

The lines in the hydrogen spectrum in various regions of the electromagnetic spectrum are given different names.

Lyman series, Balmer series, paschen series, Brackett series, and Pfund series

Lyman series appears in the ultraviolet region.
Balmer series appear in the visible region.
Other three series appear in the infra red region.

A more general formula for hydrogen spectral lines is

1/ λ = v (in cm^-1) = 109677 (1/nf² - 1/ni²)

ni>nf

ni = the initial level of the electron
nf = the final level of the electron

The above equation is called Rydberg equation.

For lyman series nf = 1
For Balmer series nf = 2
For Paschen series nf = 3
For Brackett series nf = 4
For Pfund series nf = 5

We have to note that above equation is valid for hydrogen spectrum lines only.

Emission of radiation occurs from atoms when electrons in the atom goes into an excited state and then returns to a lower energy state.

For example, if an hydrogen atom goes into an excited state, say the electron goes to fourth energy level and then returns to second energy level, radiation is emitted. The frequency of the emitted radiation can be found out by

1/ λ = v (in cm^-1) = 109677 (1/2² - 1/4²) = 486 nm.

This frequency corresponds to bluish green in visible region.

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,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 ↑ ↓.

Wave-particle duality

The earliest view of light, due to Newton, regarded light as made up of particles (commonly termed as corpuscles of light). The particle nature of light explained some of the experimental facts such as reflection and refraction of light. However, using particle nature of light, phenomena of interference and diffraction could not be explained. (See relevant chapters in Physics text and link that material with this idea).

Huygens proposed wave like character of light. With the help of wave theory of light, Huygens explained the phenomena of interference and diffraction as well as reflection and refraction. Hence the corpuscular theory was discarded and wave theory was adopted.

But once again, phenomena such black body radiation and photoelectric effect needed explanation. Wave theory could not explain these phenomena.

Einstein has suggested that light can behave as a wave as well as a particle i.e., it has dual character.


Characteristics of electromagnetic radiations

(i) The consists of electric and magnetic fields that oscillate in the directions perpendicular to each other and perpendicular to the direction in which the wave is traveling

(ii) These electromagnetic radiations do not require any medium for propagation.

(iii) All waves are characterized in terms of their wavelengths, frequency, velocity and amplitude.

Velocity of a wave = wavelength * frequency

JEE Revision - de Broglie hypothesis

de Broglie hypothesis




Wave nature of light could not explain the photoelectric effect. Einstein has suggested that light can behave as a wave as well as a particle i.e., it has dual character. In 1924, de-Broglie extended the wave nature to matter. He suggested that all microscopic particles such as electrons, protons, atoms and molecules have dual character similar to light.

Thus according to de Broglie hypothesis, “all material particles in motion possess wave characteristics.”

The wavelength(λ), mass of the particle (m) and velocity (v) of the particle are related to each other by the relation

λ = h/mv

where
h = Planck’s constant = 6.63 *10^-34 kg m² s^-1

The equation was experimentally verified by Davisson and Germer.

The wave like chacter of electron helped in making electron microscope. The electron microscope utilizes the wave like behaviour of electrons.

The dual nature of matter is applicable to all material objects but it is significant for microscopic bodies only. For bigger bodies, wave lengths calculated using the above formula are so small that they have no significance.

JEE revision - Heisenberg 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.

JEE Revision -Shapes of s, p and d orbitals

Shape 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.

JEE Revision - Pauli's exclusion principle

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

JEE Revision - Hund's rule

Hund’s rule says that, for any set of orbitals of equal energy say p orbitals of orbit 2, there is one electron is each orbital before the second electron enters or occupies an orbital.