Introduction to fundamental concepts of chemistry
Empirical formula and Molecular Formula
Combustion Analysis
Concept of mole and Avogadro’s Number
Molar Volume
Stoichiometry
Limiting reactant
Yield
Atomic Structure
Discovery of Positive Rays
Plank’s Quanyum Theory
Bohr’s Atomic Model and Hydrogen Spectrum
Concept of Orbital
Quantum Numbers
Electronic Configuration
Gases
Properties of States of Matter
Boyle’s Law
Charle’s Law
Avogadro’s Law
General Gas Equation
Kinetic Molecular Theory of Gases
Non Ideal Behaviour of Gases
Vander Waal’s Equation
Liquids
Intermolecular Forces
London Dispersions Forces
Hydrogen Bonding
Evaporation
Vapour Pressure
Boiling Point
Solids
Solids and Their Types
Properties of Crystalline Solids
Crystal Lattice and Unit Cell
Crystals and their classification
Ionic Solids
Structure of Sodium Chloride
Factors Affecting the Shape of Ionic Crystals
Lattice Energy
Covalent Solids
Molecular Solids
Chemical Equilibrium
States of Chemical Equilibrium
Equilibrium Constant Expression for Important Reactions
Relationships between Equilibrium Constants
Application of Equilibrium Constants
Le-Chatelier’s Principle
Industrial Applications of Le-Chatelier’s Principle
Ionic Product of Water
pH and pOH
Ionization Constant of Acid and Base
Common Ion Effect
Buffer Solutions
Equilibria of Slightly Soluble Ionic Compounds
Reaction Kinetics
Rate of Reaction and Specific Rate Constant
Order of Reaction
Half Life Period
Methods to Determine Order of Reaction
Rate Determining Step
Determination of Rate of Reaction
Factors Affecting Rate Of Reaction
Energy of Activation
Thermo-chemistry and Energetics of chemical reactions
Exothermic and Endothermic Reactions
Spontaneous and Non Spontaneous Reaction
System, Surrounding and State Function
Internal Energy and Ways of Transferring Energy
First Law of Thermodynamics
Types of Enthalpies
Measurement of Enthalpy of A Reaction
Hess’s Law of Constant Heat Summation
Born Haber Cycle
Electrochemistry
Oxidation Number
Conduction and Electrolytic Cell
Electrolysis of Fused Electrolytes
Electrolysis of Aqueous Solutions of Salts
Galvanic or Voltaic Cell
Standard Electrode Potential and Standard Hydrogen Electrode
Electrochemical Series
Balancing of Redox Equations by Oxidation Number Method
Balancing of Redox Equations by Ion Electron Method
Chemical bonding
Chemical Bond
Atomic Size
Ionization Energy
Electron Affinity
Electronegativity
Ionic bond or electrovalent bond
Covalent Bond
Coordinate Covalent Bond
VSEPR Theory
Valence Bond Theory
Atomic Orbital Hybridization
Bond Energy
Bond Length
Dipole moment
The Effect of Bonding on the Properties of Compound
S and p block elements
Blocks in Periodic Table
Atomic Size
Ionization Energy
Metallic and Non Metallic Character
Melting and Boiling Point
Electrical Conductance
s – Block Elements
Peculiar Behaviour of Lithium
Peculiar Behaviour of Beryllium
General Behaviour of Alkali Metals
Trends in Chemical Properties of Alkaline Earth Metals
General Trends in Properties of Compounds of Alkali and Alkaline earth Metals
Reaction of Group IIIA with Water
Transition Elements
Introduction and General Properties of Transition Elements
General Characteristics
Complex Compounds
Nomenclature of Transition Metal Complexes
Geometry of Complexes
Chelates
Fundamental principles of organic chemistry
Introduction and General Features of Organic Compounds
Source of Organic Compounds
Classification of Organic Compounds
Functional group
Structural Isomerism
Geometric Isomerism
Chemistry of Hydrocarbons
Classification of Hydrocarbons
Nomenclature of Hydrocarbons
Reactivity of Alkanes
Preparation and Reactivity of Alkenes
Preparation and Reactivity of Alkynes
Structure of Benzene
Reactions of Benzene
Orientation in Electrophilic Substitution
Comparison of Reactivity of Alkanes, Alkenes and Benzene
Alkyl halides
Alcohols & phenols
Aldehydes and Ketones
Carboxylic acid
Macromolecules
Relative Atomic Mass
The mass of an atom of an element as compared to 1/12th mass of an atom of carbon taken as 12 is called relative atomic mass (atomic mass unit).
The mass of an atom is very small. It cannot be weighed by any known physical balance. That is why we use the relative atomic mass unit scale.
Examples:
H = 1.008 a.m.u, Cl = 35.5 a.m.u. These are the average of atomic masses of isotopes proportional to their relative abundances.
Unit:
The unit used to express atomic mass is called atomic mass unit (a.m.u).
- 1 a.m.u = 1.661 × 10–27 kg
- 1 a.m.u = 1.661 × 10–24 g
Relative Isotopic Mass
The mass of an isotope of an element as compared to 1/12th mass of an atom of carbon-12 is called relative isotopic mass.
Examples:
- Relative atomic mass of Chlorine-35 = 35 a.m.u.
- Relative atomic mass of Chlorine-37 = 37 a.m.u.
Relative Molecular Mass
The mass of a molecule as compared to 1/12th mass of an atom of carbon-12 is called relative molecular mass.
Examples:
- Relative molecular mass of water (H2O) = 18 a.m.u
- Relative molecular mass of carbon dioxide (CO2) = 44 a.m.u
Relative Formula Mass
The mass of a formula unit as compared to 1/12th mass of an atom of carbon-12 is called relative formula mass.
Examples:
- Relative formula mass of sodium chloride (NaCl) = 58.5 a.m.u
- Relative formula mass of sodium sulphate (Na2SO4) = 142 a.m.u
Note: Carbon-12 is taken as standard on the C-12 scale because:
- It is a highly stable isotope.
- Its mass is exactly a whole number (12.000).
- It can be handled easily because it is a solid.
Average Atomic Masses
The average atomic mass depends on the number of possible isotopes and their natural abundance.
Formula:
Average atomic mass = Σ (isotopic mass × fractional abundance)
Example (Neon isotopes):
- Neon-20: Abundance = 90.92%, Mass = 20 a.m.u
- Neon-21: Abundance = 0.26%, Mass = 21 a.m.u
- Neon-22: Abundance = 8.82%, Mass = 22 a.m.u
Relative atomic mass of Neon = (20 × 0.9092) + (21 × 0.0026) + (22 × 0.0882) = 20.18 a.m.u
