Wednesday, May 3, 2017

1.9 Atoms and Molecules - Revision Points




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



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Tuesday, May 2, 2017

1.7 Dalton's Atomic Theory Revision points

7. Dalton’s atomic theory

To provide theoretical justification to the laws of chemical combination which are experimentally verified, John Dalton postulated a simple theory of matter. The basic postulates of Dalton’s atomic theory are:

a. Matter is made up of extremely small indivisible and indestructible ultimate particles called atoms.
b. Atoms the same element are identical in all respects ie., in shape, size, mass and chemical properties.
c. Atoms of different elements are different in all respects and have different masses and chemical properties.
d. Atom is the smallest unit that takes part in chemical combinations.
d. Atoms of two or more elements combine in a simpler whole number ratio to form compound atoms (molecules).
e. Atoms can neither be created nor destroyed during any physical or chemical change.
f. Chemical reactions involve only combinations, separation or rearrangement of atoms.

Modern atomic theory
As a result of new discoveries made after Dalton developed his postulates, some modifications were done to atomic theory. They are:

1. Atom is no longer considered to be indivisible: It is found that atom is made up of subatomic particles such as electrons, protons and neutrons. We now state how many electrons are there, protons are there in an atom.

2. Atoms of same element may not be similar in all respects. Atoms of same elements have different atomic masses. These different atoms are called isotopes.

3. Atoms of different elements may have similar one or more properties. Atomic mass of calcium and argon (40 a.m.u.) are same. So the property of atomic mass is same for atoms of different elements. Isobars or elements or atoms having the same atomic mass.

4. Atom is the smallest unit which takes part in chemical reactions. Though electrons and protons are there, it is atom which takes part in chemical reactions and electrons exchange takes place between atoms.

5. The ratio in which the different atoms combine may be fixed and integral but may not always be simple. For example in sugar molecule the ratio of C,H and O atoms is 12:22:11, which is not simple.

6. Atom of one elements may be changed into atoms of other element. Transmutation is the process by which atoms one element can be changed inot elements of other elements by subjecting it to alpha rays.

7. The mass of atom can changed into energy. Mass and energy are inconvertible. The equation give for such conversion is E mc². Hence we cannot say that mass is not destructible. But in chemical reactions, atom remains unchanged and its mass is not destroyed to liberate energy.


Updated 4 May 2017, 4 feb 2008

1.18 Stochiometry of Reactions in Solutions



Many reactions are carried out in aqueous solutions. In this case their concentration is important measure. The amounts of the products of a reaction can be calculated from the volumes of the solutions of the reactants and their concentrations. In the book, the calculations are illustrated through examples.



Stochiometry of Reactions in Solutions
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More to read

http://www.science.uwaterloo.ca/~cchieh/cact/c120/sltnstoich.html


http://jsmith.cis.byuh.edu/books/principles-of-general-chemistry-v1.0/s08-03-stoichiometry-of-reactions-in-.html  - Open book project.  Refer to the site for more topics.

https://www2.chemistry.msu.edu/courses/cem151/chap4lect_2011.pdf

Tuesday, May 3, 2016

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,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

Sunday, May 1, 2016

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

IIT JEE Chemistry - Class XII - Study Guide - 3. Solutions

Sections in the Chapter

3.1 Types of solutions
3.2 Methods for expressing the concentration of a solution: units of solution
P.P. 3.1 to 3.20
3.3 Solubilioty of gases and solids in liquids
3.4 Vapour pressure of solutions
3.5 Ideal and Nonideal solutions
P.P. 3.21 to 3.27
3.6 Colligative properties
3.7 Relative lowering of vapour pressure
P.P. 3.28 to 3.34
3.8 Elevation in boiling point
P.P. 3.35 to 3.40
3.9 Depression in freezing point
3.41 to 3.48
3.10 Osmosis and osmotic pressure
P.P. 3.49 to 3.57
3.11 Electrolytic solutions – Abnormal molar masses
P.P. 3.58 to 3.68


Additional Numerical Problems for Practice: 21

Conceptual Questions with Answers: 29
Key facts to remember
Additional Numerical Problems for Practice: 8
Revision Exercises
Very Short Answer questions 21
Short Answer Questions 66
Long Answer Questions 21

Competition File
Numerical Problems 29
Multiple choice questions: 40
Fill in the blanks: 10
True or False: 10


Study Plan

May 26 to 30

Day 1

3.1 Types of solutions
3.2 Methods for expressing the concentration of a solution: units of solution

Day 2

P.P. 3.1 to 3.20

Day 3
3.3 Solubility of gases and solids in liquids
3.4 Vapour pressure of solutions
3.5 Ideal and Nonideal solutions

Day 4

P.P. 3.21 to 3.27
3.6 Colligative properties
3.7 Relative lowering of vapour pressure
P.P. 3.28 to 3.34

Day 5

3.8 Elevation in boiling point
P.P. 3.35 to 3.40

June 1 to 10

Day 6
3.9 Depression in freezing point
3.41 to 3.48

Day 7
3.10 Osmosis and osmotic pressure
P.P. 3.49 to 3.57

Day 8
3.11 Electrolytic solutions – Abnormal molar masses
P.P. 3.58 to 3.68

Day 9

Additional Numerical Problems for Practice: 21


Day 10
Conceptual Questions with Answers: 29

Day 11
Key facts to remember
Additional Numerical Problems for Practice: 8

Day 12
Revision Exercises: Very Short Answer questions 21

Day 13
Revision Exercises: Short Answer Questions 1 to 3366

Day 14
Revision Exercises: Short Answer Questions 34 to 66

Day 15
Competition File: Numerical Problems 1to 20

June 11 onwards

Revision Period
Day 16
Competition File: Numerical Problems 21to 29

Day 17
Multiple choice questions: 1 to 20

Day 18
Multiple choice questions: 21 to 40

Day 19
Fill in the blanks: 10

Day 20
True or False: 10

Day 21 to 30
Concept revision, formula revision, Test paper problem solving



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Updated 1 May 2016,  3 Jan 2016, 11 March 2009

JEE - Study Guide - 2. Solid State



JEE Syllabus (2015)

Solid State: Classification of solids: molecular, ionic, covalent and metallic solids, amorphous and crystalline solids (elementary idea); Bragg’s Law and its applications; Unit cell and lattices, packing in solids (fcc, bcc and hcp lattices), voids, calculations involving unit cell parameters, imperfection in solids; Electrical, magnetic and dielectric properties.





Jauhar, CBSE XII class

Sections in the Chapter

2.1 Space Lattices and Unit Cell
2.2 Close Packing in Crystalline Solids
2.3 Interstitial Sites or Interstitial Voids
2.4 Types of Cubic Crystals and Number of Atoms per Unit Cell
2.5 Experimental Methods of Determining Crystal Structure: X Rays Diffraction
2.6 Coordination Number and Radius Ratio
2.7 Ionic Radii
2.8 Calculation of Density of a Crystal from its Structure
2.9 Strctures of Ionic Compounds
2.10 Imperfections in solids
2.11 Properties of solids
2.12 Amorphous solids



Additional numerical problems for practice 12
Conceptual Questions with Answers: 13
Key facts to remember
Revision Exercises: Very Short Answer questions 40
Short Answer Questions : 50
Long Answer Questions : 6

Competition File
Some useful facts
numerical problems 9

Objective Questions: Multiple choice 25
Fill in the blanks: 10
Matching type question 1

Study Plan

Day 1

2.1 Space Lattices and Unit Cell
2.2 Close Packing in Crystalline Solids
2.3 Interstitial Sites or Interstitial Voids

Day 2
2.4 Types of Cubic Crystals and Number of Atoms per Unit Cell
Practice problems 2.1 to 2.3

Day 3

2.5 Experimental Methods of Determining Crystal Structure: X Rays Diffraction
P.P. 2.4,2.5
2.6 Coordination Number and Radius Ratio
P.P. 2.14 to 2.17

Day 4

2.7 Ionic Radii
2.8 Calculation of Density of a Crystal from its Structure

Day 5
P.P. 2.8 to 2.19

Day 6

2.9 Structures of Ionic Compounds
P.P 2.20 to 2.25

Day 7
2.10 Imperfections in solids
2.11 Properties of solids
2.12 Amorphous solids

Day 8


Concept revision
Additional numerical problems for practice 12

Day 9


Conceptual Questions with Answers: 13
Key facts to remember

Day 10

Revision Exercises: Very Short Answer questions 40

Revision period

Day 11
Revision Exercises:Short Answer Questions : 1 to 15

Day 12
Revision Exercises:Short Answer Questions : 16 to 30

Day 13
Revision Exercises:Short Answer Questions : 31 to 50

Day 14

Competition File
Some useful facts
numerical problems 9

Day 15

Objective Questions: Multiple choice 25

Day 16
Fill in the blanks: 10
Matching type question 1

Day 17 to 20

Concept revision and test paper questions/problems




Updated   1 May 2016,  6 June 2015
Originally published 11 March 2009