Chemistry is the science that deals with the composition, structure, properties and transformations of matter. Matter is anything that has mass and occupies space. It can be classified as pure substances (elements and compounds) or mixtures (homogeneous and heterogeneous).
Laws of Chemical Combination
Law of Conservation of Mass (Lavoisier), Law of Definite Proportions (Proust), Law of Multiple Proportions (Dalton), Law of Reciprocal Proportions.
Dalton’s Atomic Theory
Matter consists of indivisible atoms. Atoms of same element are identical. Atoms combine in simple whole number ratios. Atoms are neither created nor destroyed in reactions.
Mole Concept
1 mole = 6.022 × 10²³ particles (Avogadro number). Molar mass = mass of 1 mole in grams. Percentage composition, empirical and molecular formulae.
Stoichiometry
Balanced chemical equations give mole ratios. Limiting reagent, theoretical and actual yield.
Concentration Terms
Molarity (M), Molality (m), Mole fraction, Mass percentage, ppm. Significant figures and dimensional analysis.
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Practice Questions (20)
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Q1. The law of conservation of mass was proposed by:
Q2. One mole of any substance contains:
Q3. The empirical formula of C₆H₁₂O₆ is:
Q4. The number of significant figures in 0.00060 is:
Q5. Molarity is defined as:
Q6. In the reaction 2H₂ + O₂ → 2H₂O, if 4 g of H₂ reacts with 32 g of O₂, the limiting reagent is:
Q7. The percentage of oxygen in CO₂ is:
Q8. Avogadro’s number is approximately:
Q9. The molecular mass of H₂SO₄ is:
Q10. Which law explains that 1 g of hydrogen combines with 8 g of oxygen in H₂O and 16 g in H₂O₂?
Q11. The mole fraction of a solute in a solution is:
Q12. 18 g of water contains:
Q13. The SI unit of molar mass is:
Q14. In stoichiometry, the balanced equation gives:
Q15. The empirical formula mass of a compound is 30 u and molecular mass is 60 u. The value of n is:
Q16. Which of the following is an intensive property?
Q17. 1 mole of CO₂ contains how many oxygen atoms?
Q18. The law of definite proportions applies to:
Q19. Molality of a solution is independent of:
Q20. The number of moles in 22 g of CO₂ is:
2. Structure of Atom
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The atom is the smallest particle of an element that can take part in a chemical reaction and retains all the chemical properties of that element. Modern atomic theory developed through contributions of Dalton, Thomson, Rutherford, Bohr, de Broglie, Heisenberg and Schrödinger.
Discovery of Subatomic Particles
Electron (J.J. Thomson, 1897): Charge = –1.6022 × 10⁻¹⁹ C, mass = 9.1094 × 10⁻³¹ kg.
Proton (E. Goldstein, 1886): Charge = +1.6022 × 10⁻¹⁹ C, mass = 1.6726 × 10⁻²⁷ kg.
Neutron (James Chadwick, 1932): Charge = 0, mass = 1.6749 × 10⁻²⁷ kg.
Atomic Number & Mass Number
Atomic number Z = number of protons = number of electrons (neutral atom).
Mass number A = Z + number of neutrons.
Isotopes: same Z, different A. Isobars: same A, different Z. Isotones: same neutrons, different Z.
Atomic Models
Thomson’s Plum Pudding Model → Rutherford’s Nuclear Model (α-scattering) → Bohr’s Model (quantized orbits, explains H-spectrum) → Quantum Mechanical Model (wave nature, uncertainty principle, orbitals).
Quantum Numbers & Electronic Configuration
n, l, mₗ, mₛ. Aufbau principle, Pauli exclusion, Hund’s rule. Exceptions in Cr and Cu for extra stability.
Practice Questions (20)
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Q1. Who discovered the electron?
Q2. The charge on an electron is:
Q3. The nucleus of an atom contains:
Q4. The mass number of an atom is equal to:
Q5. Isotopes of an element have:
Q6. In Rutherford’s α-particle scattering experiment, most α-particles:
Q7. According to Bohr’s model, the radius of the nth orbit is proportional to:
Q8. The energy of an electron in the first Bohr orbit of hydrogen is:
Q9. The de Broglie wavelength is given by:
Q10. Heisenberg’s uncertainty principle states:
Q11. The azimuthal quantum number l for a d-orbital is:
Q12. The maximum number of electrons in an orbital is:
Q13. Which of the following has the maximum number of unpaired electrons?
Q14. The electronic configuration of Cu (Z = 29) is:
Q15. The number of orbitals in the n = 3 shell is:
Q16. Which quantum number determines the shape of an orbital?
Q17. The correct order of increasing energy is:
Q18. The spin quantum number can have values:
Q19. The number of radial nodes in 3p orbital is:
Q20. Which of the following is isoelectronic with O²⁻?
3. Classification of Elements and Periodicity in Properties
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Study Material
The systematic arrangement of elements in order of increasing atomic number is called the Periodic Table. It helps in understanding the properties of elements and predicting their behaviour. The modern periodic law states: The physical and chemical properties of elements are a periodic function of their atomic numbers.
Historical Development
Dobereiner’s Triads (1817), Newlands’ Law of Octaves (1865), Mendeleev’s Periodic Table (1869) based on atomic mass with gaps for undiscovered elements, and the Modern Periodic Table by Moseley (1913) based on atomic number.
Structure of the Modern Periodic Table
• 7 Periods (horizontal rows)
• 18 Groups (vertical columns, IUPAC numbering)
Period 1: 2 elements, Period 2 & 3: 8 elements, Period 4 & 5: 18 elements, Period 6: 32 elements, Period 7: Incomplete.
Classification of Elements
s-block (Groups 1-2), p-block (Groups 13-18), d-block (Groups 3-12, Transition elements), f-block (Lanthanoids and Actinoids, Inner transition elements).
Periodic Trends in Properties
Atomic Radius: Decreases across a period, increases down a group.
Ionisation Enthalpy: Increases across a period, decreases down a group.
Electron Gain Enthalpy: Most negative for halogens.
Electronegativity: Increases across a period, decreases down a group (F = 4.0).
Metallic Character: Decreases across a period, increases down a group.
Anomalous behaviour of second period elements and diagonal relationship (Li-Mg, Be-Al, B-Si).
Important Points for NEET
Exceptions in IE and electron gain enthalpy due to half-filled and fully-filled subshells. Variable oxidation states in transition elements. Basic nature of oxides decreases across a period.
Practice Questions (20)
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Q1. Modern periodic law is based on:
Q2. The number of elements in the 6th period of the modern periodic table is:
Q3. Which of the following has the smallest atomic radius?
Q4. Ionisation enthalpy is maximum for:
Q5. Electron gain enthalpy is most negative for:
Q6. The element with highest electronegativity is:
Q7. Which group element shows variable valency?
Q8. Diagonal relationship is shown by:
Q9. The correct order of increasing atomic radius is:
Q10. Which of the following has the highest ionisation enthalpy?
Q11. In the periodic table, metallic character increases:
Q12. The number of valence electrons in Group 17 elements is:
Q13. Which of the following is a metalloid?
Q14. The first element of a group shows anomalous behaviour due to:
Q15. The correct order of electron gain enthalpy (most negative to least) is:
Q16. Transition elements are placed in:
Q17. Which property shows a sudden increase from Group 2 to Group 13?
Q18. The element with electronic configuration [Ar] 3d⁵ 4s¹ is:
Q19. In a period, the basic nature of oxides:
Q20. Lanthanoids and Actinoids are called:
4. Chemical Bonding and Molecular Structure
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Study Material
Chemical bonding is the attractive force that holds atoms together in a molecule or crystal. It is responsible for the stability and properties of compounds. Atoms tend to attain stable noble gas configuration (octet rule) by losing, gaining, or sharing electrons.
Types of Chemical Bonds
1. Ionic (Electrovalent) Bond: Complete transfer of electrons (e.g., NaCl).
2. Covalent Bond: Sharing of electrons (single, double, triple bonds).
3. Coordinate (Dative) Bond: One atom donates both electrons.
4. Metallic Bond: Delocalised electrons in a sea of positive ions.
Valence Bond Theory (VBT)
Overlap of atomic orbitals forms σ and π bonds. Hybridisation: sp, sp², sp³, sp³d, sp³d², sp³d³.
Molecular Orbital Theory (MOT)
Formation of bonding and antibonding orbitals. Bond order = (Nb – Na)/2. Explains paramagnetic nature of O₂.
VSEPR Theory
Electron pairs arrange to minimise repulsion. Lone pair repulsion is stronger than bond pair repulsion. Determines molecular shapes.
Hydrogen Bonding
Strong dipole-dipole attraction (X–H…Y where X, Y = N, O, F). Intermolecular and intramolecular types. Affects boiling point and solubility.
Resonance, Formal Charge & Important Trends
Resonance structures, formal charge calculation, bond length, bond energy, exceptions to octet rule (BF₃, PCl₅, SF₆).
Practice Questions (20)
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Q1. The bond formed by complete transfer of electrons is:
Q2. Which molecule has a coordinate bond?
Q3. The hybridisation of carbon in CH₄ is:
Q4. The shape of SF₄ molecule is:
Q5. Bond order of O₂ molecule is:
Q6. Which has the highest lattice energy?
Q7. The number of σ and π bonds in acetylene (C₂H₂) is:
Q8. Which molecule is paramagnetic?
Q9. The correct order of bond angle is:
Q10. Resonance is not shown by:
Q11. The hybridisation of Xe in XeF₄ is:
Q12. Which bond is strongest?
Q13. Intramolecular hydrogen bonding is present in:
Q14. The formal charge on N in NH₄⁺ is:
Q15. Which has the maximum bond angle?
Q16. The number of lone pairs on central atom in XeF₂ is:
Q17. Which theory explains the paramagnetic nature of oxygen?
Q18. The geometry of IF₇ is:
Q19. The most ionic compound among the following is:
Q20. Bond order of NO⁺ is:
5. States of Matter: Gases and Liquids
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Matter exists in three states — solid, liquid and gas. This chapter deals with gaseous and liquid states, their properties, intermolecular forces and gas laws.
Gaseous State
Gases have no definite shape or volume. They are highly compressible and have high diffusibility.
Gas Laws
Boyle’s Law (P ∝ 1/V at constant T), Charles’ Law (V ∝ T at constant P), Gay-Lussac’s Law (P ∝ T at constant V), Avogadro’s Law, Ideal Gas Equation (PV = nRT).
Dalton’s Law of Partial Pressures
Total pressure of a gas mixture is the sum of partial pressures of individual gases.
Kinetic Molecular Theory of Gases
Gases consist of tiny particles in constant random motion. Collisions are elastic. Average kinetic energy is proportional to absolute temperature.
Real Gases & van der Waals Equation
(P + a/V²)(V – b) = RT. Real gases deviate from ideal behaviour at high pressure and low temperature.
Liquid State
Liquids have definite volume but no definite shape. Important properties: vapour pressure, surface tension, viscosity. Intermolecular forces include London dispersion, dipole-dipole and hydrogen bonding.
Practice Questions (20)
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Q1. Boyle’s law is applicable at:
Q2. The ideal gas equation is:
Q3. Which gas law explains that equal volumes of gases contain equal number of molecules?
Q4. van der Waals constant ‘a’ is a measure of:
Q5. The critical temperature of a gas is the temperature:
Q6. Surface tension of a liquid:
Q7. Viscosity of a liquid:
Q8. Which of the following has the highest viscosity?
Q9. RMS velocity of a gas is:
Q10. At constant temperature, pressure of a gas is inversely proportional to:
Q11. Real gases behave ideally at:
Q12. The intermolecular force in ideal gas is:
Q13. Which gas shows maximum deviation from ideal behaviour?
Q14. The unit of van der Waals constant ‘a’ is:
Q15. Vapour pressure of a liquid increases with:
Q16. The correct order of boiling points is:
Q17. Which of the following is not a characteristic of gases?
Q18. The average kinetic energy of gas molecules is proportional to:
Q19. Liquefaction of gas is possible only below its:
Q20. The SI unit of viscosity is:
6. Thermodynamics
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Thermodynamics deals with energy changes in physical and chemical processes. It helps predict the feasibility and direction of reactions.
Important Terms
System, Surroundings, State functions (U, H, S, G), Path functions (q, w).
Types of systems: Open, Closed, Isolated.
First Law of Thermodynamics
ΔU = q + w (Conservation of energy).
At constant volume: ΔU = qᵥ
At constant pressure: ΔH = qₚ (Enthalpy).
Enthalpy Changes
Standard enthalpy of formation, combustion, solution, hydration, neutralisation. Hess’s Law.
Second Law of Thermodynamics
Entropy of the universe increases in spontaneous processes.
Third Law
Entropy of a perfect crystal at 0 K is zero.
Practice Questions (20)
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Q1. Which of the following is a state function?
Q2. First law of thermodynamics is based on:
Q3. For an isolated system, ΔU =
Q4. Enthalpy change at constant pressure is equal to:
Q5. Hess’s law is an application of:
Q6. The criterion for spontaneity is:
Q7. Entropy of the universe in a spontaneous process:
Q8. For a reaction to be spontaneous at all temperatures:
Q9. The relationship between ΔG° and K is:
Q10. Which has maximum entropy?
Q11. Standard enthalpy of formation of an element in its standard state is:
Q12. Work done in isothermal reversible expansion of an ideal gas is:
Q13. The SI unit of entropy is:
Q14. Which process is always spontaneous?
Q15. Gibbs free energy is zero at:
Q16. For the reaction 2H₂ + O₂ → 2H₂O, ΔH is:
Q17. The correct relation is:
Q18. Entropy change is maximum in:
Q19. Which of the following is an extensive property?
Q20. The efficiency of a heat engine is maximum when:
7. Equilibrium
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Equilibrium is the state in which the rate of forward reaction equals the rate of backward reaction. Concentrations of reactants and products become constant. It can be physical equilibrium or chemical equilibrium.
Law of Mass Action
For aA + bB ⇌ cC + dD, Kc = [C]^c [D]^d / [A]^a [B]^b
Le Chatelier’s Principle
If a system at equilibrium is subjected to a change (concentration, pressure, temperature), the equilibrium shifts to counteract the change.
Ionic Equilibrium
Acids, bases, pH, buffers, solubility product (Ksp), common ion effect.
Factors Affecting Equilibrium
Concentration, pressure (gaseous), temperature, catalyst (no effect on position).
Practice Questions (20)
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Q1. Chemical equilibrium is:
Q2. For the reaction N₂ + 3H₂ ⇌ 2NH₃, Kc has the unit:
Q3. Le Chatelier’s principle is applicable to:
Q4. For N₂ + 3H₂ ⇌ 2NH₃, increase in pressure shifts equilibrium to:
Q5. pH of neutral solution at 25°C is:
Q6. Conjugate base of H₂SO₄ is:
Q7. A buffer solution is:
Q8. Solubility product (Ksp) is for:
Q9. Common ion effect is used in:
Q10. pH of 0.001 M HCl is:
Q11. For the reaction N₂ + 3H₂ ⇌ 2NH₃, Kp and Kc are related by:
Q12. pH + pOH at 25°C is:
Q13. Buffer solution resists change in:
Q14. The solubility product Ksp is for:
Q15. Which is a strong acid?
Q16. Common ion effect:
Q17. pH of 0.01 M NaOH is:
Q18. The equilibrium constant K is independent of:
Q19. For exothermic reaction, increase in temperature shifts equilibrium to:
Q20. Henderson-Hasselbalch equation is used for:
8. Redox Reactions
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Study Material
Redox reactions are chemical reactions involving simultaneous oxidation and reduction. Oxidation is loss of electrons (increase in oxidation number) and reduction is gain of electrons (decrease in oxidation number). These reactions are fundamental to electrochemistry, corrosion, respiration, combustion, and many industrial processes.
Oxidation Number Rules
• Free elements = 0
• Hydrogen = +1 (except in metal hydrides = –1)
• Oxygen = –2 (except in peroxides = –1, superoxides = –½, OF₂ = +2)
• Alkali metals = +1, Alkaline earth metals = +2
• Sum in neutral molecule = 0, in ion = charge of the ion.
Types of Redox Reactions
1. Combination reactions
2. Decomposition reactions
3. Displacement reactions (single and double)
4. Disproportionation reactions (same element is both oxidised and reduced).
Balancing Redox Reactions
• Oxidation number method
• Ion-electron method (half-reaction method) – most important for NEET.
Redox Titrations
• Permanganometry (KMnO₄ in acidic medium)
• Dichrometry (K₂Cr₂O₇)
• Iodometry and Iodimetry.
Important Concepts
Oxidising agents: KMnO₄, K₂Cr₂O₇, HNO₃, Cl₂.
Reducing agents: KI, Na₂S₂O₃, FeSO₄, SnCl₂.
Electrochemical series helps predict feasibility of redox reactions.
Practice Questions (20)
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Q1. In a redox reaction, oxidation is:
Q2. The oxidation number of Cr in K₂Cr₂O₇ is:
Q3. Which of the following is a disproportionation reaction?
Q4. In acidic medium, KMnO₄ acts as:
Q5. The change in oxidation number of Mn in MnO₄⁻ to Mn²⁺ is:
Q6. In the reaction 2KI + Cl₂ → 2KCl + I₂, the oxidising agent is:
Q7. The reducing agent in a redox reaction:
Q8. In acidic medium, the equivalent weight of KMnO₄ is:
Q9. Which of the following is a disproportionation reaction?
Q10. The oxidation number of sulphur in H₂SO₄ is:
Q11. In the reaction Zn + CuSO₄ → ZnSO₄ + Cu, the oxidising agent is:
Q12. The equivalent weight of K₂Cr₂O₇ in acidic medium is:
Q13. Which of the following is a reducing agent?
Q14. The oxidation number of oxygen in OF₂ is:
Q15. In iodometric titration, I₂ is liberated by reaction with:
Q16. The change in oxidation number of sulphur in S + O₂ → SO₂ is:
Q17. Which of the following is not a redox reaction?
Q18. The oxidation number of nitrogen in HNO₃ is:
Q19. In the reaction 2FeCl₃ + SnCl₂ → 2FeCl₂ + SnCl₄, the reducing agent is:
Q20. The equivalent weight of KMnO₄ in alkaline medium is:
9. Hydrogen
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Study Material
Hydrogen is the lightest element and the most abundant in the universe. It has a unique position in the periodic table as it resembles both alkali metals and halogens.
Isotopes of Hydrogen
Protium (¹H), Deuterium (²H), Tritium (³H).
Q2. Tritium is the isotope of hydrogen with mass number:
Q3. Heavy water is chemically:
Q4. Hydrogen is prepared in laboratory by the reaction of:
Q5. Hydrogen peroxide acts as:
Q6. The strongest reducing form of hydrogen is:
Q7. Ionic hydrides are formed by:
Q8. The structure of H₂O₂ is:
Q9. Heavy water is used as:
Q10. Hydrogen bonding is strongest in:
Q11. The temporary hardness of water is due to:
Q12. Nascent hydrogen is:
Q13. The oxidation state of hydrogen in metal hydrides is:
Q14. Hydrogen peroxide is stored in:
Q15. The bond angle in water molecule is:
Q16. Which hydride is electron deficient?
Q17. The temporary hardness of water can be removed by:
Q18. Hydrogen peroxide decomposes in the presence of:
Q19. The number of neutrons in deuterium is:
Q20. Which of the following is used as a moderator in nuclear reactors?
10. The s-Block Elements
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The s-block elements consist of Group 1 (Alkali metals) and Group 2 (Alkaline earth metals). They are highly reactive due to low ionisation enthalpy and large atomic size. They form ionic compounds and show +1 and +2 oxidation states respectively.
Group 1 (Alkali Metals): Li, Na, K, Rb, Cs, Fr
• Highly electropositive, soft metals with low density and low melting points.
• React vigorously with water, oxygen and halogens.
• Lithium shows anomalous behaviour and diagonal relationship with Magnesium.
• Important compounds: Na₂CO₃ (washing soda), NaHCO₃ (baking soda), NaOH (caustic soda), NaCl.
Group 2 (Alkaline Earth Metals): Be, Mg, Ca, Sr, Ba, Ra
• Less reactive than Group 1.
• Be and Mg show some covalent character.
• Be shows anomalous behaviour and diagonal relationship with Aluminium.
• Important compounds: CaO (quick lime), Ca(OH)₂ (slaked lime), CaCO₃, MgSO₄ (Epsom salt), Plaster of Paris (CaSO₄·½H₂O).
Biological Importance
Na⁺ and K⁺ maintain nerve impulses and fluid balance.
Mg²⁺ is central in chlorophyll.
Ca²⁺ is essential for bones, teeth and muscle contraction.
Periodic Trends
Atomic radius, reactivity, basic character of oxides and hydroxides increase down the group. Ionisation enthalpy decreases down the group.
Practice Questions (20)
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Q1. Which element shows diagonal relationship with Mg?
Q2. The most reactive alkali metal is:
Q3. Plaster of Paris is chemically:
Q4. Which has the lowest ionisation enthalpy?
Q5. Washing soda is:
Q6. The element that floats on water is:
Q7. Which is used in Solvay process?
Q8. The most basic oxide among the following is:
Q9. Epsom salt is:
Q10. Which metal is used in photoelectric cells?
Q11. The correct order of hydration enthalpy is:
Q12. Which of the following is used as a drying agent?
Q13. The flame colour of potassium is:
Q14. Which is the least soluble in water?
Q15. The correct order of basic strength of oxides is:
Q16. Which metal is stored in kerosene?
Q17. The formula of baking soda is:
Q18. Which of the following is used in fireworks?
Q19. The least reactive alkaline earth metal is:
Q20. Which compound is used for softening hard water?
11. The p-Block Elements (Group 13 and 14)
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The p-block elements are in groups 13 to 18. This chapter covers Group 13 (Boron family) and Group 14 (Carbon family). These elements show a wide range of properties from non-metallic to metallic.
Group 13 (Boron Family): B, Al, Ga, In, Tl
• General configuration: ns²np¹
• Boron is a non-metal, others are metals.
• Anomalous behaviour of Boron and diagonal relationship with Silicon.
• Important compounds: Borax (Na₂B₄O₇·10H₂O), Boric acid (H₃BO₃), Aluminium chloride (AlCl₃), Alum.
Group 14 (Carbon Family): C, Si, Ge, Sn, Pb
• General configuration: ns²np²
• Carbon and Silicon are non-metals, others are metals.
• Allotropy in Carbon (Diamond, Graphite, Fullerenes) and Silicon.
• Important compounds: CO₂, SiO₂, Silicates, Silicones, Zeolites, PbO, PbO₂.
Important Trends
Atomic radius increases down the group.
Ionisation enthalpy decreases down the group.
Electronegativity decreases down the group.
Oxidation states: +3 and +1 (Group 13), +4 and +2 (Group 14) – inert pair effect increases down the group.
Practice Questions (20)
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Q1. The element which shows maximum catenation is:
Q2. Borax is chemically:
Q3. The most stable oxidation state of thallium is:
Q4. Which is the hardest substance?
Q5. The formula of orthoboric acid is:
Q6. Which element has the highest ionisation enthalpy in Group 13?
Q7. The structure of diborane (B₂H₆) contains:
Q8. The most acidic oxide among the following is:
Q9. Graphite is a good conductor due to:
Q10. The inert pair effect is most prominent in:
Q11. Silicones are polymers of:
Q12. The correct order of acidic character of oxides is:
Q13. Diamond is hard because:
Q14. Which of the following is electron deficient?
Q15. The formula of zeolite is:
Q16. The inert pair effect is maximum in:
Q17. Which of the following is used as a semiconductor?
Q18. The correct order of acidic strength of oxides is:
Q19. Aluminium chloride in vapour state exists as:
Q20. Which of the following is used as a water softener?
12. Organic Chemistry - Some Basic Principles and Techniques
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Study Material
Organic chemistry is the study of carbon compounds. Carbon has unique properties like catenation, tetravalency, and isomerism, which lead to a vast number of organic compounds.
Classification of Organic Compounds
• Acyclic (open chain) and Cyclic (closed chain)
• Homocyclic and Heterocyclic
• Aliphatic and Aromatic
IUPAC Nomenclature
Rules for naming alkanes, alkenes, alkynes, alcohols, aldehydes, ketones, carboxylic acids, amines, etc. Priority order of functional groups.
Alkynes (CₙH₂ₙ₋₂)
• sp hybridisation, C≡C bond
• Acidic nature of terminal alkynes (reaction with Na, AgNO₃, Cu₂Cl₂)
• Reactions: Addition, Oxidation
Aromatic Hydrocarbons (Benzene and derivatives)
• Structure of benzene (Kekulé, resonance)
• Electrophilic substitution reactions: Nitration, Halogenation, Friedel-Crafts alkylation/acylation
• Directive influence of substituents
Important Tests
• Baeyer’s reagent for unsaturation
• Bromine water test
• Tollens’ reagent for terminal alkynes
Practice Questions (20)
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Q1. General formula of alkanes is:
Q2. Markovnikov’s rule is applicable to:
Q3. The reagent used to distinguish terminal alkynes is:
Q4. Benzene undergoes electrophilic substitution because of:
Q5. Ozonolysis of ethene gives:
Q6. Anti-Markovnikov addition is observed in presence of:
Q7. The number of sigma bonds in benzene is:
Q8. Which of the following is an aromatic compound?
Q9. Friedel-Crafts alkylation is an example of:
Q10. The product of ozonolysis of 2-butene is:
Q11. Which hydrocarbon has acidic hydrogen?
Q12. The catalyst used in Friedel-Crafts reaction is:
Q13. The structure of benzene was proposed by:
Q14. Pyrolysis of alkanes is also called:
Q15. Which of the following decolourises Baeyer’s reagent?
Q16. The number of isomers of C₄H₈ is:
Q17. Nitration of benzene is carried out using:
Q18. The product of Wurtz reaction of methyl bromide is:
Q19. Which of the following has highest boiling point?
Q20. Benzene reacts with Cl₂ in presence of UV light to give:
14. Environmental Chemistry
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Environmental Chemistry deals with the study of chemical and biochemical phenomena occurring in the environment. It includes the sources, reactions, transport, effects and fate of chemical species in air, water and soil.
Atmospheric Pollution
• Troposphere and Stratosphere
• Major pollutants: CO, CO₂, SO₂, NOₓ, hydrocarbons, particulate matter
• Acid rain, greenhouse effect, global warming, ozone depletion
Water Pollution
• Causes: domestic, industrial and agricultural waste
• Parameters: BOD, COD, DO
• Eutrophication, heavy metals, pesticides
Soil Pollution
• Causes: pesticides, fertilisers, industrial waste
• Control measures: proper use of chemicals, bioremediation
Green Chemistry
• Sustainable chemistry that minimises hazardous substances
• 12 principles of green chemistry
Important Environmental Segments
Atmosphere, Hydrosphere, Lithosphere, Biosphere
Practice Questions (20)
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Q1. The lowest layer of atmosphere is:
Q2. The gas responsible for greenhouse effect is:
Q3. Acid rain is due to:
Q4. Ozone layer is present in:
Q5. BOD stands for:
Q6. Which gas is responsible for ozone depletion?
Q7. Eutrophication is caused by:
Q8. The pH of acid rain is generally:
Q9. Which of the following is a primary pollutant?
Q10. Green chemistry aims at:
Q11. The ozone depleting substance is:
Q12. COD stands for:
Q13. Which gas causes Bhopal gas tragedy?
Q14. The major component of photochemical smog is:
Q15. Minamata disease is caused by:
Q16. Which of the following is a secondary pollutant?
Q17. The unit of BOD is:
Q18. Montreal Protocol is related to:
Q19. Which of the following is a biodegradable pollutant?
Q20. The 12 principles of green chemistry were given by:
15. The Solid State
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Study Material
The Solid State deals with the study of rigid matter where particles are closely packed with definite shape and volume. Solids are classified as Crystalline and Amorphous.
Classification of Solids
• Crystalline solids: Definite melting point, anisotropic, sharp melting point (Ionic, Molecular, Covalent, Metallic)
• Amorphous solids: No definite melting point, isotropic, considered as supercooled liquids
Crystal Lattice and Unit Cell
• Crystal lattice: Regular 3D arrangement of points
• Unit cell: Smallest repeating unit
• Seven crystal systems and 14 Bravais lattices
Packing in Solids
• Close packing: hcp and ccp (fcc)
• Packing efficiency: Simple cubic (52.4%), BCC (68%), FCC (74%)
Defects in Solids
• Point defects: Vacancy, Interstitial, Schottky, Frenkel
• Impurity defects and non-stoichiometric defects
Electrical and Magnetic Properties
• Conductors, Insulators, Semiconductors (n-type and p-type)
• Diamagnetic, Paramagnetic and Ferromagnetic substances
Practice Questions (20)
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Q1. Which of the following is an amorphous solid?
Q2. The packing efficiency of FCC lattice is:
Q3. Schottky defect is observed in:
Q4. The number of atoms per unit cell in BCC lattice is:
Q5. Frenkel defect is shown by crystals having:
Q6. The coordination number of atoms in BCC lattice is:
Q7. Which of the following is a molecular solid?
Q8. The packing efficiency of simple cubic lattice is:
Q9. n-type semiconductor is formed by doping with:
Q10. The crystal system with maximum symmetry is:
Q11. In Frenkel defect, there is:
Q12. The number of atoms in a face-centred cubic unit cell is:
Q13. Which solid is a good conductor of electricity?
Q14. The edge length of unit cell of NaCl is 552 pm. The radius of Na⁺ ion is:
Q15. Ferromagnetism is shown by:
Q16. The coordination number in hcp and ccp is:
Q17. Which defect increases density of the crystal?
Q18. The Bragg’s equation is:
Q19. p-type semiconductor is formed by doping silicon with:
Q20. Which of the following has highest melting point?
16. Chemistry in Everyday Life
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Study Material
Chemistry in Everyday Life deals with the application of chemistry in medicines, food, cleansing agents and other daily use materials.
Drugs and Their Classification
• Drugs: Chemicals of low molecular mass that interact with macromolecular targets to produce biological response.
• Classification: On the basis of pharmacological effect, drug action, chemical structure, molecular targets.
Therapeutic Action of Different Classes of Drugs
• Antacids, Antihistamines, Tranquilizers, Analgesics, Antimicrobials, Antifertility drugs, Antibiotics, Antiseptics & Disinfectants.
Cleansing Agents
• Soaps and Detergents
• Synthetic detergents: Anionic, Cationic and Non-ionic detergents
• Advantages of synthetic detergents over soaps.
Important Terms
• Chemotherapy, Receptors, Enzyme inhibitors, Broad spectrum antibiotics, Narrow spectrum antibiotics, Antipyretics, Antimalarials, etc.
Practice Questions (20)
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Q1. Drugs that bind to the receptor site and inhibit its natural function are called:
Q2. Aspirin is an example of:
Q3. The artificial sweetener used by diabetic patients is:
Q4. Which of the following is an anionic detergent?
Q5. Tranquilizers are used to treat:
Q6. The pH range of our body is:
Q7. Which of the following is a broad spectrum antibiotic?
Q8. Soaps are sodium or potassium salts of:
Q9. The sweetest artificial sweetener among the following is:
Q10. Antiseptics are applied to:
Q11. Which drug is used as an antacid?
Q12. Detergents are better than soaps because:
Q13. The main constituent of Dettol is:
Q14. Which of the following is a narcotic analgesic?
Q15. The preservative used in food is:
Q16. Which of the following is used as an antimalarial drug?
Q17. Non-ionic detergents are:
Q18. The drug used to bring down body temperature is called:
Q19. Which of the following is used as a food preservative?
Q20. The main component of soap is:
17. Aldehydes, Ketones and Carboxylic Acids
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Study Material
Aldehydes and Ketones contain carbonyl group (C=O). Aldehydes have carbonyl at chain end (R-CHO), ketones have it in between (R-CO-R'). Carboxylic acids contain –COOH group.
Preparation
• Aldehydes: Oxidation of primary alcohols, Rosenmund reduction, Stephen reduction, Gattermann-Koch reaction.
• Ketones: Oxidation of secondary alcohols, Friedel-Crafts acylation, from Grignard reagents.
• Carboxylic acids: Oxidation of primary alcohols/aldehydes, hydrolysis of nitriles, Grignard + CO₂.
Distinguishing Tests
• Tollens’ reagent (silver mirror) for aldehydes
• Fehling’s solution and Benedict’s test for aldehydes
• Iodoform test for methyl ketones and acetaldehyde
• Sodium bicarbonate test for carboxylic acids
Practice Questions (20)
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Q1. The general formula of aldehydes is:
Q2. Which reagent gives silver mirror test?
Q3. Cannizzaro reaction is given by aldehydes having:
Q4. The strongest acid among the following is:
Q5. Iodoform test is given by:
Q6. The reagent used in Clemmensen reduction is:
Q7. Which of the following has highest boiling point?
Q8. The product of reaction of Grignard reagent with CO₂ followed by hydrolysis is:
Q9. Aldol condensation requires:
Q10. Which acid is strongest?
Q11. The reaction of aldehyde with Tollens’ reagent gives:
Q12. Wolff-Kishner reduction converts carbonyl group to:
Q13. The functional group in carboxylic acids is:
Q14. Hell-Volhard-Zelinsky reaction is used for:
Q15. Which compound gives iodoform test?
Q16. The acidity of carboxylic acids is due to:
Q17. The reaction of acetaldehyde with Tollens’ reagent gives:
Q18. Which of the following does not undergo aldol condensation?
Q19. The product of reaction of ketone with hydroxylamine is:
Q20. Which acid is used in the preparation of aspirin?
18. Amines
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Study Material
Amines are derivatives of ammonia in which one or more hydrogen atoms are replaced by alkyl or aryl groups. They are classified as primary (1°), secondary (2°), and tertiary (3°) amines.
Classification and Nomenclature
• Aliphatic and Aromatic amines
• IUPAC names: Alkanamine, N-alkylalkanamine, Benzenamine (Aniline)
Preparation
• Reduction of nitro compounds, nitriles, amides
• Gabriel phthalimide synthesis (for 1° amines)
• Hofmann bromamide degradation
• From alkyl halides (with NH₃)
Chemical Properties
• Basic nature (aliphatic > aromatic)
• Alkylation, acylation, carbylamine reaction, Hinsberg test
• Diazotisation (only for 1° aromatic amines)
• Coupling reactions of diazonium salts
Distinguishing Tests
• Hinsberg test (1°, 2°, 3° amines)
• Carbylamine test (1° amines)
• Azo dye test (1° aromatic amines)
Importance
Amines are used in synthesis of dyes, drugs, polymers and as intermediates in organic synthesis.
Practice Questions (20)
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Q1. The general formula of primary amine is:
Q2. Which test is used to distinguish 1°, 2° and 3° amines?
Q3. Gabriel phthalimide synthesis is used for preparation of:
Q4. The basicity order of amines in aqueous solution is:
Q5. Carbylamine test is given by:
Q6. Aniline on diazotisation gives:
Q7. Which of the following is most basic?
Q8. Hinsberg reagent is:
Q9. The product of Hofmann bromamide reaction is:
Q10. Coupling reaction of benzene diazonium chloride with phenol gives:
Q11. Which amine gives foul smell with CHCl₃ + KOH?
Q12. Aniline is less basic than aliphatic amines because:
Q13. The product of reduction of nitrobenzene with Sn/HCl is:
Q14. Tertiary amines do not react with:
Q15. The strongest base among the following is:
Q16. Sandmeyer reaction is used for preparation of:
Q17. Which amine gives brisk effervescence with nitrous acid?
Q18. The number of hydrogen atoms attached to nitrogen in secondary amine is:
Q19. Which of the following is used in the manufacture of dyes?
Q20. The correct order of basicity is:
19. Biomolecules
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Study Material
Biomolecules are organic compounds present in living organisms. The main classes are carbohydrates, proteins, lipids, nucleic acids and vitamins.
Carbohydrates
• Polyhydroxy aldehydes or ketones or substances that yield them on hydrolysis.
• Classification: Monosaccharides (glucose, fructose), Oligosaccharides (sucrose, maltose), Polysaccharides (starch, cellulose, glycogen).
• Reducing and non-reducing sugars. Glucose exists in open chain and cyclic (α & β) forms.
Proteins
• Polymers of α-amino acids linked by peptide bonds.
• Structure: Primary, Secondary (α-helix, β-sheet), Tertiary, Quaternary.
• Denaturation by heat, pH, heavy metals.
Nucleic Acids
• DNA (deoxyribonucleic acid) and RNA (ribonucleic acid).
• DNA has double helix structure (Watson-Crick model), contains deoxyribose and thymine.
• RNA contains ribose and uracil, mainly single stranded.
Vitamins and Hormones
• Vitamins: Fat soluble (A, D, E, K) and Water soluble (B-complex, C).
• Hormones: Chemical messengers (insulin, adrenaline, thyroxine).
Enzymes
• Biocatalysts, highly specific, protein in nature. Follow lock and key or induced fit model.
Practice Questions (20)
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Q1. Which of the following is a monosaccharide?
Q2. The linkage between two monosaccharides is called:
Q3. Which sugar is a reducing sugar?
Q4. The helical structure of proteins is stabilised by:
Q5. The sugar present in DNA is:
Q6. Which vitamin is water soluble?
Q7. The base present in RNA but not in DNA is:
Q8. Denaturation of proteins is caused by:
Q9. The monomer of starch is:
Q10. Which biomolecule stores genetic information?
Q11. The vitamin whose deficiency causes scurvy is:
Q12. α-helix structure is found in:
Q13. The sugar present in RNA is:
Q14. Which of the following is a zwitterion?
Q15. The deficiency of Vitamin D causes:
Q16. The bond linking amino acids in proteins is:
Q17. Cellulose is a polymer of:
Q18. Which of the following is not a vitamin?
Q19. The nitrogen base absent in DNA is:
Q20. Enzymes are:
20. Polymers
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Study Material
Polymers are high molecular mass macromolecules formed by the repeated joining of small repeating units called monomers. The process is called polymerisation.
Classification of Polymers
• On the basis of origin: Natural (rubber, cellulose, proteins), Semi-synthetic, Synthetic (nylon, PVC, Teflon)
• On the basis of structure: Linear, Branched chain, Cross-linked
• On the basis of mode of polymerisation: Addition (chain growth), Condensation (step growth)
• On the basis of molecular forces: Elastomers, Fibres, Thermoplastics, Thermosetting polymers
Molecular Mass
Number average and Weight average molecular mass. Polydispersity Index (PDI).
Practice Questions (20)
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Q1. Polymers are:
Q2. Natural rubber is a polymer of:
Q3. Nylon-6,6 is a:
Q4. Vulcanisation of rubber is done by:
Q5. Teflon is a polymer of:
Q6. Bakelite is a:
Q7. Which polymer is used in non-stick cookware?
Q8. The monomer of natural rubber is:
Q9. Nylon-6 is prepared from:
Q10. Which polymer is biodegradable?
Q11. Addition polymerisation is shown by:
Q12. The polymer used in making bullet-proof vests is:
Q13. LDPE and HDPE differ in:
Q14. The monomer of Buna-S is:
Q15. Which polymer is used for making unbreakable crockery?
Q16. The polymer used in making tyres is:
Q17. Which of the following is a biodegradable polymer?
Q18. The repeating unit in cellulose is:
Q19. Which polymer is used as a substitute for wool?
Q20. The polymer used in making optical lenses is:
21: Chemical Kinetics
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Study Material
Chemical kinetics is the branch of chemistry that deals with the study of rates of chemical reactions, the factors affecting the rate and the mechanism by which the reaction proceeds.
Rate of Reaction: Change in concentration of reactant or product per unit time. Average rate and instantaneous rate. Rate law and rate constant.
Order and Molecularity: Order is the sum of powers in rate law (experimental). Molecularity is the number of molecules colliding in elementary step (theoretical). Zero, first, second and pseudo-first order reactions.
Integrated Rate Equations: For zero order: [R] = [R]0 – kt. First order: ln[R] = ln[R]0 – kt and t1/2 = 0.693/k. Second order reactions.
Arrhenius Equation: k = A e–Ea/RT. Activation energy (Ea), collision theory and transition state theory. Effect of temperature and catalyst on rate.
Factors Affecting Rate: Concentration, temperature, surface area, catalyst. Collision frequency, activation energy and orientation.
Practice Questions (20)
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1. The rate of a reaction is expressed as Rate = k[A]²[B]. The overall order of the reaction is
2. For a first order reaction, half-life period is independent of
3. Unit of rate constant for a first order reaction is
4. Arrhenius equation is
5. For a zero order reaction, half-life is proportional to
6. A catalyst increases the rate of reaction by
7. Rate law of a reaction is determined by
8. Molecularity of a reaction is always
9. Hydrolysis of ester in presence of acid is an example of
10. For a first order reaction, the slope of the plot of ln[A] vs time is
11. Temperature coefficient of most reactions lies between
12. According to collision theory, rate of reaction depends on
13. Unit of rate constant for a second order reaction is
14. Activation energy is the minimum energy required for
15. In a zero order reaction, the rate is
16. Half-life of a second order reaction depends on
17. Most of the reactions follow
18. A catalyst does not change the
19. Rate constant increases with increase in
20. Radioactive decay follows
22: Surface Chemistry
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Surface chemistry deals with the phenomena that occur at the surfaces or interfaces of phases (solid-gas, solid-liquid, liquid-gas, etc.).
Adsorption: Physisorption (weak van der Waals forces, low heat of adsorption) and Chemisorption (chemical bond, high heat of adsorption, highly specific).
Adsorption Isotherms: Freundlich isotherm (empirical) and Langmuir isotherm (monolayer adsorption).
Catalysis: Homogeneous and heterogeneous catalysis. Enzyme catalysis and shape-selective catalysis by zeolites.
Colloids: Lyophilic and lyophobic colloids. Preparation, purification, properties (Tyndall effect, Brownian movement, charge, coagulation, electrophoresis). Emulsions and associated colloids (micelles).
Applications: Adsorption in gas masks, water purification, chromatography; colloids in medicine, photography, rubber industry, cleansing action of soap.
Practice Questions (20)
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1. Adsorption is
2. Physisorption is characterised by
3. Chemisorption is
4. Freundlich adsorption isotherm is
5. Langmuir adsorption isotherm assumes
6. In homogeneous catalysis the catalyst and reactants are in
7. Heterogeneous catalysis occurs at the
8. Enzymes are
9. Zeolites are used as
10. Tyndall effect is shown by
11. Brownian movement is observed in
12. Lyophilic colloids are
13. Coagulation is the process of
14. Electrophoresis is
15. Emulsions are
16. Cleansing action of soap is due to
17. Gold number is a measure of
18. Micelles are
19. Colloids are used in
20. Adsorption is used in
23: Coordination Compounds
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Study Material
Coordination compounds are compounds in which a central metal atom or ion is bonded to a number of ligands by coordinate bonds.
Werner’s Theory: Primary valency (ionisable) and secondary valency (non-ionisable, directional).
IUPAC Naming: Ligands named first, metal name with oxidation state in Roman numerals.
Coordination Number: Number of donor atoms attached to central metal.
Isomerism: Structural (linkage, coordination, ionisation) and Stereoisomerism (geometrical and optical).
Theories of Bonding: Valence Bond Theory (hybridisation), Crystal Field Theory (splitting of d-orbitals, colour and magnetic properties).
Applications: Extraction of metals, catalysis, analytical chemistry, medicine (cis-platin), biological systems (haemoglobin, chlorophyll).
Practice Questions (20)
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1. Coordination compounds contain
2. Werner’s coordination theory explained
3. In IUPAC naming, oxidation state of metal is shown by
4. Coordination number is
5. Chelate complexes are
6. Geometrical isomerism is shown by
7. Valence Bond Theory explains geometry through
8. Crystal Field Theory explains
9. Haemoglobin is a coordination compound of
10. EDTA is a
11. Coordination compounds are used in
12. Linkage isomerism is shown by
13. Optical isomerism is shown by complexes that are
