BITSAT Chemistry Syllabus

Chemistry syllabus — full chapter and topic list

Expand any chapter to see its topics and sub-topics. This is the whole examinable outline for Chemistry in BITSAT, not a summary of it.

1. 1

   States of Matter

   3 topics

   • Measurement
     • Physical quantities and SI units
     • Dimensional analysis
     • Precision
     • Significant figures
   • Chemical reactions
     • Laws of chemical combination
     • Dalton’s atomic theory
     • Mole concept
     • Atomic, molecular and molar masses
     • Percentage composition empirical & molecular formula
     • Balanced chemical equations & stoichiometry
   • Three states of matter, intermolecular interactions, types of bonding, melting and boiling points
     • Gaseous state: Kinetic energy and molecular speeds
     • Gaseous state: Gas Laws
     • Gaseous state: Ideal behavior
     • Gaseous state: Ideal gas equation
     • Gaseous state: Empirical derivation of gas equation
     • Gaseous state: Avogadro number
     • Gaseous state: Deviation from ideal behaviour – Critical temperature
     • Gaseous state: Liquefaction of gases
     • Gaseous state: van der Waals equation
     • Liquid state: Vapour pressure
     • Liquid state: Surface tension
     • Liquid state: Viscosity
     • Solid state: Classification
     • Solid state: Space lattices & crystal systems
     • Solid state: Unit cell in two dimensional and three dimensional lattices
     • Solid state: Calculation of density of unit cell – Cubic & hexagonal systems
     • Solid state: Close packing
     • Solid state: Crystal structures
     • Solid state: Voids
     • Solid state: Number of atoms per unit cell in a cubic unit cell
     • Solid state: Imperfections- Point defects, non-stoichiometric crystals
     • Solid state: Electrical, magnetic and dielectric properties
     • Solid state: Amorphous solids – qualitative description
     • Solid state: Band theory of metals, conductors, semiconductors and insulators, and n- and p- type semiconductors
2. 2

   Atomic Structure

   3 topics

   • Introduction
     • Subatomic particles
     • Atomic number, isotopes and isobars
     • Thompson’s model and its limitations
     • Rutherford’s picture of atom and its limitations
     • Hydrogen atom spectrum
     • Bohr model and its limitations
   • Quantum mechanics
     • Wave-particle duality – de Broglie relation
     • Uncertainty principle
     • Hydrogen atom: Quantum numbers and wave functions
     • Atomic orbitals and their shapes (s, p, and d)
     • Spin quantum number
   • Many electron atoms
     • Pauli exclusion principle
     • Aufbau principle and the electronic configuration of atoms
     • Hund’s rule
3. 3

   Periodicity and Bonding

   5 topics

   • Brief history of the development of periodic tables Periodic law and the modern periodic table; Types of elements: s, p, d, and f blocks; Periodic trends: ionization energy, atomic, and ionic radii, inert gas radii, electron affinity, electro negativity and valency.
   • Valence electrons, Ionic Bond: Lattice Energy and Born-Haber cycle; Covalent character of ionic bonds and polar character of covalent bond, bond parameters
   • Molecular Structure: Lewis picture & resonance structures, VSEPR model & molecular shapes
   • Covalent Bond: Valence Bond Theory- Orbital overlap, Directionality of bonds & hybridization (s, p & d orbitals only), Resonance; Molecular orbital theory- Methodology, Orbital energy level diagram, Bond order, Magnetic properties for homonuclear diatomic species (qualitative idea only).
   • Dipole moments; Hydrogen Bond
4. 4

   Thermodynamics

   4 topics

   • Basic Concepts
     • Systems and surroundings
     • State functions
     • Intensive & Extensive Properties
     • Zeroth Law and Temperature
   • First Law of Thermodynamics
     • Work, internal energy, heat, enthalpy, heat capacities and specific heats
     • Measurements of ∆U and ∆H
     • Enthalpies of formation, phase transformation, ionization, electron gain
     • Thermochemistry
     • Hess’s Law
     • Enthalpy of bond dissociation, combustion, atomization, sublimation, solution and dilution
   • Second Law
     • Spontaneous and reversible processes
     • Entropy
     • Gibbs free energy related to spontaneity and non-spontaneity, non-mechanical work
     • Standard free energies of formation, free energy change and chemical equilibrium
   • Third Law
     • Introduction
5. 5

   Physical and Chemical Equilibria

   7 topics

   • Concentration Units: Mole Fraction, Molarity, and Molality
   • Solutions: Solubility of solids and gases in liquids, Vapour Pressure, Raoult’s law, Relative lowering of vapor pressure, depression in freezing point; elevation in boiling point; osmotic pressure, determination of molecular mass; solid solutions, abnormal molecular mass, van’t Hoff factor
   • Equilibrium: Dynamic nature of equilibrium, law of mass action
   • Physical Equilibrium: Equilibria involving physical changes (solid-liquid, liquid-gas, solid-gas)
   • Chemical Equilibria: Equilibrium constants (KP, KC), Factors affecting equilibrium, Le- Chatelier’s principle
   • Ionic Equilibria: Strong and Weak electrolytes, Acids and Bases (Arrhenius, Lewis, Lowry and Bronsted) and their dissociation; degree of ionization, Ionization of Water; ionization of polybasic acids, pH; Buffer solutions; Henderson equation, Acid-base titrations; Hydrolysis; Solubility Product of Sparingly Soluble Salts; Common Ion Effect
   • Factors Affecting Equilibria: Concentration, Temperature, Pressure, Catalysts, Significance of ΔG and ΔG0 in Chemical Equilibria
6. 6

   Electrochemistry

   2 topics

   • Redox Reactions
     • Oxidation-reduction reactions (electron transfer concept)
     • Oxidation number
     • Balancing of redox reactions
     • Electrochemical cells and cell reactions
     • Standard electrode potentials
     • EMF of Galvanic cells
     • Nernst equation
     • Factors affecting the electrode potential
     • Gibbs energy change and cell potential
     • Secondary cells
     • Dry cells
     • Fuel cells
     • Corrosion and its prevention
   • Electrolytic Conduction
     • Electrolytic Conductance
     • Specific and molar conductivities
     • Variations of conductivity with concentration
     • Kolhrausch’s Law and its application
     • Electrolysis
     • Faraday’s laws of electrolysis
     • Electrode potential and electrolysis
7. 7

   Chemical Kinetics

   3 topics

   • Aspects of Kinetics
     • Rate and Rate expression of a reaction
     • Rate constant
     • Order and molecularity of the reaction
     • Integrated rate expressions and half-life for zero and first order reactions
   • Factor Affecting the Rate of the Reactions
     • Concentration of the reactants
     • Catalyst
     • Size of particles
     • Temperature dependence of rate constant concept of collision theory (elementary idea, no mathematical treatment)
     • Activation energy
     • Arrhenius Equation
   • Surface Chemistry
     • Adsorption – physisorption and chemisorption
     • Factors affecting adsorption of gases on solids
     • Catalysis: homogeneous and heterogeneous, activity and selectivity: enzyme catalysis
     • Colloidal state: distinction between true solutions, colloids and suspensions
     • Lyophillic, lyophobic multi molecular and macromolecular colloids
     • Properties of colloids
     • Tyndall effect
     • Brownian movement
     • Electrophoresis
     • Coagulations
     • Emulsions–types of emulsions
8. 8

   Hydrogen and s-block elements

   4 topics

   • Hydrogen
     • Element: unique position in periodic table, occurrence, isotopes
     • Dihydrogen: preparation, properties, reactions, and uses
     • Molecular, saline, ionic, covalent, interstitial hydrides
     • Water: Properties
     • Structure and aggregation of water molecules
     • Heavy water
     • Hydrogen peroxide: preparation, reaction, structure & use
     • Hydrogen as a fuel
   • s-block elements
     • Abundance and occurrence
     • Anomalous properties of the first elements in each group
     • Diagonal relationships
     • Trends in the variation of properties (ionization energy, atomic & ionic radii)
   • Alkali metals
     • Lithium, sodium and potassium: occurrence, extraction, reactivity, and electrode potentials
     • Biological importance
     • Reactions with oxygen, hydrogen, halogens and water
     • Basic nature of oxides and hydroxides
     • Halides
     • Properties and uses of compounds such as NaCl, Na2CO3, NaHCO3, NaOH, KCl, and KOH
   • Alkaline earth metals
     • Magnesium and calcium: Occurrence, extraction, reactivity and electrode potentials
     • Reactions with O2, H2O, H2 and halogens
     • Solubility and thermal stability of oxo salts
     • Biological importance of Ca and Mg
     • Preparation, properties and uses of important compounds such as CaO, Ca(OH)2, plaster of Paris, MgSO4, MgCl2, CaCO3, and CaSO4
9. 9

   p-d- and f-block elements

   9 topics

   • General
     • Abundance, distribution, physical and chemical properties, isolation and uses of elements; Trends in chemical reactivity of elements of a group; electronic configuration, oxidation states; anomalous properties of first element of each group.
   • Group 13 elements
     • Boron; Properties and uses of borax, boric acid, boron hydrides & halides. Reaction of aluminum with acids and alkalis;
   • Group 14 elements
     • Carbon: carbon catenation, physical & chemical properties, uses, allotropes (graphite, diamond, fullerenes), oxides, halides and sulphides, carbides; Silicon: Silica, silicates, silicone, silicon tetrachloride, Zeolites, and their uses
   • Group 15 elements
     • Dinitrogen; Preparation, reactivity and uses of nitrogen; Industrial and biological nitrogen fixation; Compound of nitrogen; Ammonia: Haber’s process, properties and reactions; Oxides of nitrogen and their structures; Properties and Ostwald’s process of nitric acid production; Fertilizers – NPK type; Production of phosphorus; Allotropes of phosphorus; Preparation, structure and properties of hydrides, oxides, oxoacids (elementary idea only) and halides of phosphorus, phosphine.
   • Group 16 elements
     • Isolation and chemical reactivity of dioxygen; Acidic, basic and amphoteric oxides; Preparation, structure and properties of ozone; Allotropes of sulphur; Preparation/production properties and uses of sulphur dioxide and sulphuric acid; Structure and properties of oxides, oxoacids (structures only).
   • Group 17 and group 18 elements
     • Structure and properties of hydrides, oxides, oxoacids of halogens (structures only); preparation, properties & uses of chlorine & HCl; Inter halogen compounds; Bleaching Powder; Uses of Group 18 elements, Preparation, structure and reactions of xenon fluorides, oxides, and oxoacids.
   • d-Block elements
     • General trends in the chemistry of first row transition elements; Metallic character; Oxidation state; ionization enthalpy; Ionic radii; Color; Catalytic properties; Magnetic properties; Interstitial compounds; Occurrence and extraction of iron, copper, silver, zinc, and mercury; Alloy formation; Steel and some important alloys; preparation and properties of K2Cr2O7, KMnO4.
   • f-Block elements
     • Lanthanoids and actinoids; Oxidation states and chemical reactivity of lanthanoids compounds; Lanthanide contraction and its consequences, Comparison of actinoids and lanthanoids.
   • Coordination Compounds
     • Coordination number; Ligands; Werner’s coordination theory; IUPAC nomenclature; Application and importance of coordination compounds (in qualitative analysis, extraction of metals and biological systems e.g. chlorophyll, vitamin B12, and hemoglobin); Bonding: Valence-bond approach, Crystal field theory (qualitative); Isomerism including stereoisomerisms.
10. 10

    Principles of Organic Chemistry and Hydrocarbons

    6 topics

    • Classification
      • General Introduction
      • Classification based on functional groups
      • Trivial and IUPAC nomenclature
      • Methods of purification: qualitative and quantitative
    • Electronic displacement in a covalent bond
      • Inductive, resonance effects, and hyperconjugation
      • Free radicals
      • Carbocations
      • Carbanions
      • Nucleophiles and electrophiles
      • Types of organic reactions
      • Free radical halogenations
    • Alkanes
      • Structural isomerism
      • General properties and chemical reactions
      • Free radical halogenation
      • Combustion
      • Pyrolysis
    • Alkenes and alkynes
      • Structure of double and triple bonded compounds
      • General methods of preparation and reactions
      • Physical properties
      • Electrophilic and free radical additions
      • Addition of hydrogen, halogen, water, hydrogen halides
      • Markovnikov’s addition and peroxide effect
      • Ozonolysis
      • Oxidation
      • Mechanism of electrophilic addition
      • Acidic character of alkynes
      • 1,2 and 1,4 addition to dienes
    • Aromatic hydrocarbons
      • Sources
      • Properties
      • Isomerism
      • Resonance delocalization
      • Aromaticity
      • Polynuclear hydrocarbons
      • IUPAC nomenclature
      • Mechanism of electrophilic substitution reaction
      • Directive influence
      • Effect of substituents on reactivity
      • Carcinogenicity
      • Toxicity
    • Haloalkanes and haloarenes
      • Physical properties
      • Nomenclature
      • Optical rotation
      • Chemical reactions
      • Mechanism of substitution reaction
      • Uses and environmental effects
      • Di, tri, tetrachloromethanes
      • Iodoform
      • Freon
      • DDT
11. 11

    Stereochemistry

    2 topics

    • Conformations
      • Ethane conformations
      • Newman and Sawhorse projections
    • Geometrical isomerism in alkenes
12. 12

    Organic Compounds with Functional Groups Containing Oxygen and Nitrogen

    2 topics

    • General
      • Nomenclature
      • Electronic structure
      • Important methods of preparation
      • Identification
      • Important reactions
      • Physical and chemical properties
      • Uses of alcohols, phenols, ethers, aldehydes, ketones, carboxylic acids, nitro compounds, amines, diazonium salts, cyanides and isocyanides
    • Specific
      • Reactivity of alpha-hydrogen in carbonyl compounds
      • Effect of substituents on alpha-carbon on acid strength
      • Comparative reactivity of acid derivatives
      • Mechanism of nucleophilic addition and dehydration
      • Basic character of amines
      • Methods of preparation and their separation
      • Importance of diazonium salts in synthetic organic chemistry
13. 13

    Biological, Industrial and Environmental Chemistry

    7 topics

    • Carbohydrates
      • Classification
      • Monosaccharides
      • Structures of Pentoses and Hexoses
      • Simple Chemical Reactions of Glucose
      • Disaccharides: Reducing and Non-reducing Sugars - Sucrose, Maltose, and Lactose
      • Polysaccharides: Elementary Idea of Structures of Starch, Cellulose, and Glycogen
    • Proteins
      • Amino Acids
      • Peptide Bond
      • Polypeptides
      • Primary Structure of Proteins
      • Simple Idea of Secondary, Tertiary, and Quaternary Structures of Proteins
      • Denaturation of Proteins and Enzymes
    • Nucleic Acids
      • Types of Nucleic Acids
      • Primary Building Blocks of Nucleic Acids (Chemical Composition of DNA & RNA)
      • Primary Structure of DNA and its Double Helix
    • Vitamins
      • Classification, Structure, Functions in Biosystems
      • Hormones
    • Polymers
      • Classification of Polymers
      • General Methods of Polymerization
      • Molecular Mass of Polymers
      • Biopolymers and Biodegradable Polymers
      • Methods of Polymerization (Free Radical, Cationic, and Anionic Addition Polymerizations)
      • Copolymerization
      • Natural Rubber
      • Vulcanization of Rubber
      • Synthetic Rubbers
      • Condensation Polymers
      • Some Important Polymers: Natural and Synthetic like Polythene, Nylon, Polyesters, Bakelite, and Rubber
    • Pollution
      • Environmental Pollutants
      • Soil, Water, and Air Pollution
      • Chemical Reactions in Atmosphere
      • Smog
      • Major Atmospheric Pollutants
      • Acid Rain
      • Ozone and its Reactions
      • Depletion of Ozone Layer and its Effects
      • Industrial Air Pollution
      • Greenhouse Effect and Global Warming
      • Green Chemistry, Study for Control of Environmental Pollution
    • Chemicals in Medicine, Health-care, and Food
      • Analgesics
      • Tranquilizers
      • Antiseptics
      • Disinfectants
      • Anti-microbials
      • Anti-fertility Drugs
      • Antihistamines
      • Antibiotics
      • Antacids
      • Preservatives
      • Artificial Sweetening Agents
      • Antioxidants
      • Soaps and Detergents
14. 14

    Theoretical Principles of Experimental Chemistry

    7 topics

    • Volumetric Analysis
      • Principles
      • Standard solutions of sodium carbonate and oxalic acid
      • Acid-base titrations
      • Redox reactions involving KI, H2SO4, Na2SO3, Na2S2O3 and H2S
      • Potassium permanganate in acidic, basic and neutral media
      • Titrations of oxalic acid, ferrous ammonium sulphate with KMnO4, K2 Cr2O7/Na2S2O3, Cu(II)/Na2S2O3
    • Qualitative analysis of Inorganic Salts
      • Principles in the determination of the cations Pb2+, Cu2+, As3+, Mn2+, Al3+, Zn2+, Co2+, Ca2+, Sr2+, Ba2+, Mg2+, NH4+, Fe3+, Ni2+ and the anions CO32-, S2-, SO42-, SO32-, NO2-, NO3-, Cl-, Br-, I-, PO43-, CH3COO-, C2O42-
    • Physical Chemistry Experiments
      • Preparation and crystallization of alum, copper sulphate, Benzoic acid, ferrous sulphate, double salt of alum and ferrous sulphate, potassium ferric sulphate
      • Temperature vs. solubility
      • Study of pH changes by common ion effect in case of weak acids and weak bases
      • pH measurements of some solutions obtained from fruit juices, solutions of known and varied concentrations of acids, bases and salts using pH paper or universal indicator
      • Lyophilic and lyophobic sols
      • Dialysis
      • Role of emulsifying agents in emulsification
      • Equilibrium studies involving ferric and thiocyanate ions, [Co(H2O)6]2+ and chloride ions
      • Enthalpy determination for strong acid vs. strong base neutralization reaction, hydrogen bonding interaction between acetone and chloroform
      • Rates of the reaction between sodium thiosulphate and hydrochloric acid, potassium iodate and sodium sulphite, iodide vs. hydrogen peroxide, concentration and temperature effects in these reactions
    • Purification Methods
      • Filtration
      • Crystallization
      • Sublimation
      • Distillation
      • Differential extraction
      • Chromatography
      • Principles of melting point and boiling point determination
      • Principles of paper chromatographic separation – Rf values
    • Qualitative Analysis of Organic Compounds
      • Detection of nitrogen, sulphur, phosphorous and halogens
      • Detection of carbohydrates, fats and proteins in foodstuff
      • Detection of alcoholic, phenolic, aldehydic, ketonic, carboxylic, amino groups and unsaturation
    • Principles of Organic Chemistry Experiments
      • Preparation of acetanilide, p-nitro acetanilide, di-benzayl acetone, aniline yellow, beta-naphthol-aniline dye
    • Basic Laboratory Technique
      • Cutting glass tube and glass rod
      • Bending a glass tube
      • Drawing out a glass jet
      • Boring of cork

Chemistry flashcards for BITSAT

21 of 51 cards from the Chemistry deck — real questions with worked answers.

1. State the mole concept: what is one mole and what is the value of Avogadro's number?

   One mole is the amount of substance containing as many entities as there are atoms in 12 g of carbon-12. Avogadro's number = 6.022 x 10^23 entities per mole.

2. What is the difference between empirical formula and molecular formula?

   Empirical formula gives the simplest whole-number ratio of atoms (e.g. CH for benzene). Molecular formula gives the actual number of atoms (e.g. C6H6). Molecular formula = n x empirical formula, where n = molar mass / empirical formula mass.

3. Define molarity, molality, and mole fraction.

   Molarity (M) = moles of solute per litre of solution. Molality (m) = moles of solute per kg of solvent. Mole fraction = moles of a component divided by total moles of all components.

4. State the de Broglie equation and Heisenberg's uncertainty principle.

   de Broglie: lambda = h/(mv). Heisenberg: delta-x times delta-p is greater than or equal to h/(4 pi), meaning position and momentum cannot both be known exactly.

5. What are the four quantum numbers and what does each describe?

   Principal (n): shell/energy. Azimuthal (l): subshell/shape (0 to n-1). Magnetic (m_l): orbital orientation (-l to +l). Spin (m_s): electron spin (+1/2 or -1/2).

6. State Hund's rule, the Pauli exclusion principle, and the Aufbau principle.

   Hund's rule: orbitals of equal energy are singly filled before pairing. Pauli: no two electrons in an atom have the same four quantum numbers. Aufbau: orbitals fill in order of increasing energy (n+l rule).

7. Give the Rydberg formula for the wavenumber of spectral lines of hydrogen.

   1/lambda = R_H (1/n1^2 - 1/n2^2), where R_H = 1.097 x 10^7 m^-1 and n2 > n1. Lyman (n1=1), Balmer (n1=2), Paschen (n1=3).

8. State the ideal gas equation and the value of R in SI units.

   PV = nRT. R = 8.314 J K^-1 mol^-1 (or 0.0821 L atm K^-1 mol^-1).

9. What are the van der Waals equation corrections for real gases?

   (P + an^2/V^2)(V - nb) = nRT. The 'a' term corrects for intermolecular attraction (pressure), and 'b' corrects for the finite volume of molecules.

10. State the first law of thermodynamics and define enthalpy.

    First law: delta-U = q + w (internal energy change equals heat added plus work done on system). Enthalpy H = U + PV; at constant pressure delta-H = q_p.

11. State Hess's law of constant heat summation.

    The total enthalpy change for a reaction is the same whether it occurs in one step or several steps, since enthalpy is a state function. Allows adding enthalpies of intermediate reactions.

12. Write the Gibbs free energy equation and the criterion for spontaneity.

    delta-G = delta-H - T delta-S. A process is spontaneous when delta-G < 0, at equilibrium when delta-G = 0, and non-spontaneous when delta-G > 0.

13. State Le Chatelier's principle.

    If a system at equilibrium is disturbed by a change in concentration, temperature, or pressure, the equilibrium shifts in the direction that counteracts the disturbance.

14. What is the relationship between Kp and Kc?

    Kp = Kc (RT)^(delta-n), where delta-n = (moles of gaseous products) - (moles of gaseous reactants).

15. Define pH, pOH, and the ionic product of water Kw.

    pH = -log[H+], pOH = -log[OH-]. Kw = [H+][OH-] = 1.0 x 10^-14 at 25 C, and pH + pOH = 14.

16. State the Henderson-Hasselbalch equation for a buffer.

    pH = pKa + log([salt]/[acid]) for an acidic buffer. For a basic buffer, pOH = pKb + log([salt]/[base]).

17. State Faraday's two laws of electrolysis.

    First law: mass deposited is proportional to charge passed (m = Z I t). Second law: masses of different substances deposited by the same charge are proportional to their equivalent weights.

18. Write the Nernst equation for an electrode at 298 K.

    E = E_standard - (0.0591/n) log Q, where n is the number of electrons transferred and Q is the reaction quotient.

19. Relate standard cell potential to Gibbs free energy and equilibrium constant.

    delta-G_standard = -nFE_cell_standard, and E_cell_standard = (0.0591/n) log K at 298 K, where F = 96500 C/mol.

20. Give the integrated rate law for a first-order reaction and its half-life.

    k = (2.303/t) log([A]0/[A]). Half-life t(1/2) = 0.693/k, which is independent of initial concentration.

21. State the Arrhenius equation and what it shows about temperature and rate.

    k = A e^(-Ea/RT). Rate constant increases with temperature and decreases with activation energy Ea; A is the frequency/pre-exponential factor.

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