UNIT 1: Some Basic Concepts of Chemistry

1. Define mole concept and calculate the number of molecules in 2 moles of CO₂.

 The mole concept states that one mole of any substance contains Avogadro's number of
entities (atoms, molecules, or ions), which is 6.022 × 10²³.

 For CO₂:

o 1 mole contains 6.022 × 10²³ molecules.

o 2 moles contain 2 × (6.022 × 10²³) = 1.2044 × 10²⁴ molecules.

2. State and explain the Law of Conservation of Mass with an example.

 The Law of Conservation of Mass states that mass can neither be created nor destroyed in a
chemical reaction.

 Example:

o In the reaction H₂ + O₂ → H₂O, the total mass of reactants equals the total mass of
products.

3. Differentiate between empirical formula and molecular formula with examples.

 Empirical formula: The simplest ratio of atoms in a compound (e.g., CH for benzene).

 Molecular formula: The actual number of atoms in a molecule (e.g., C₆H₆ for benzene).

4. Calculate the molar mass of the following:

 Na₂CO₃: (2 × 23) + (12) + (3 × 16) = 106 g/mol

 CH₃COOH: (2 × 12) + (4 × 1) + (2 × 16) = 60 g/mol

5. A compound contains 40% Carbon, 6.67% Hydrogen, and 53.33% Oxygen. Find its empirical
formula.

 Divide each % by atomic mass:

o C: 40/12 = 3.33

o H: 6.67/1 = 6.67

o O: 53.33/16 = 3.33

 Simplest ratio: C₁H₂O₁ (Empirical formula: CH₂O)

UNIT 2: Structure of Atom

6. Explain Bohr’s Model of the atom with postulates.

 Electrons revolve around the nucleus in fixed orbits.

 Energy levels are quantized.

 Electrons do not lose energy in stable orbits.

7. Write the electronic configuration of the following:

 Fe (Z = 26): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶

 Cl (Z = 17): 1s² 2s² 2p⁶ 3s² 3p⁵

8. Define quantum numbers and explain their significance.

 Describe the position and energy of an electron in an atom:

o Principal (n) - Energy level

o Azimuthal (l) - Shape of orbital

o Magnetic (m₁) - Orientation of orbital

o Spin (s) - Spin direction

9. What is Hund’s Rule? Explain with an example.

 Electrons fill degenerate orbitals singly before pairing.

 Example: Carbon (1s² 2s² 2p²) fills two separate p orbitals first.

10. Calculate the wavelength of a photon having energy 3 × 10⁻¹⁹ J.

 Formula: λ = hc/E

 λ = (6.626 × 10⁻³⁴ × 3 × 10⁸) / (3 × 10⁻¹⁹)

 λ = 6.626 × 10⁻⁷ m

UNIT 3: Classification of Elements & Periodicity

11. Write the modern periodic law and explain its significance.

 Properties of elements are periodic functions of atomic number.

12. Define electronegativity and explain its trends.

 Tendency of an atom to attract electrons.

 Increases across a period, decreases down a group.

13. Why does atomic size decrease across a period and increase down a group?

 Across a period: Effective nuclear charge increases, pulling electrons closer.

 Down a group: Addition of shells increases atomic size.

14. Differentiate between ionization energy and electron affinity.

 Ionization energy: Energy to remove an electron (Endothermic).

 Electron affinity: Energy change when an electron is added (Exothermic).

15. Arrange in increasing order of atomic size: Na, Mg, Al, Si.

 Si < Al < Mg < Na

UNIT 4: Chemical Bonding & Molecular Structure

16. Explain ionic bonding with the example of NaCl.

 Na loses an electron, Cl gains an electron, forming Na⁺ and Cl⁻ ions.

 Strong electrostatic attraction forms an ionic bond.

17. Define VSEPR theory and explain the shape of H₂O and NH₃.

 VSEPR: Electron pairs repel, determining molecular shape.

 H₂O: Bent (104.5°), NH₃: Trigonal pyramidal (107°).

18. What is hydrogen bonding? Give two examples.

 Weak bond between H and electronegative atoms.

 Examples: H₂O, NH₃.

19. Compare covalent bond and ionic bond with examples.

 Covalent bond: Sharing of electrons (H₂, CH₄).

 Ionic bond: Transfer of electrons (NaCl, MgO).

20. Explain dipole moment and its significance.

 Measure of bond polarity.

 H₂O has a dipole moment, making it polar.

Here are the answers to the questions across the specified units:

UNIT 5: States of Matter

Boyle’s Law:

 Definition: Boyle’s Law states that the pressure of a given amount of gas is inversely
proportional to its volume at a constant temperature.

 Mathematical expression: P∝1VorPV=constantP \propto \frac{1}{V} \quad \text{or} \quad

PV = \text{constant}

Charles’ Law:

 Definition: Charles’ Law states that the volume of a given amount of gas is directly
proportional to its absolute temperature at constant pressure.

 Mathematical expression: V∝TorVT=constantV \propto T \quad \text{or} \quad \frac{V}{T}

= \text{constant}

Derivation of Ideal Gas Equation (PV = nRT):

  From Boyle’s Law: PV=constantPV = \text{constant}

 From Charles’ Law: VT=constant\frac{V}{T} = \text{constant}

 Combining these, we get PV=nRTPV = nRT, where:

o PP = pressure,

o VV = volume,

o nn = number of moles,

o RR = ideal gas constant,

o TT = temperature in Kelvin.

Density of a Gas at STP (Molar mass = 44 g/mol):

 At STP, P=1atm,T=273.15KP = 1 \text{atm}, T = 273.15 \text{K}.

 Ideal gas law: PV=nRTPV = nRT

 Density ρ=mV=PMRT\rho = \frac{m}{V} = \frac{PM}{RT}, where MM is the molar mass. For

M=44g/molM = 44 \text{g/mol}, ρ=1×440.0821×273.15≈1.77 g/L\rho = \frac{1 \times 44}
{0.0821 \times 273.15} \approx 1.77 \, \text{g/L}

Critical Temperature and Critical Pressure:

 Critical temperature: The temperature above which a substance cannot exist as a liquid,
regardless of pressure.

 Critical pressure: The pressure required to liquefy a substance at its critical temperature.

Kinetic Molecular Theory of Gases:

 Gas molecules are in constant random motion.

 The volume of gas molecules is negligible compared to the volume of the container.

 The average kinetic energy is proportional to the temperature of the gas.

 Collisions between gas molecules are perfectly elastic.

UNIT 6: Thermodynamics

First Law of Thermodynamics:

 Statement: Energy cannot be created or destroyed, only transferred or converted from one
form to another.

 Mathematical expression: ΔU=Q−W\Delta U = Q - W where ΔU\Delta U is the change in

internal energy, QQ is heat absorbed, and WW is work done by the system.

Enthalpy, Entropy, and Gibbs Free Energy:

 Enthalpy (H): The total heat content of a system.

 Entropy (S): A measure of the disorder or randomness of a system.

  Gibbs Free Energy (G): The energy available to do work. G=H−TSG = H - TS

Conditions for Spontaneous Reactions:

 A reaction is spontaneous if ΔG<0\Delta G < 0.

Calculate ΔG for a reaction (ΔH = -150 kJ, ΔS = 200 J/K, T = 300 K):

ΔG=ΔH−TΔS=−150 kJ−(300 K×0.2 kJ/K)=−150 kJ−60 kJ=−210 kJ\Delta G = \Delta H - T \Delta S = -150 \,
\text{kJ} - (300 \, \text{K} \times 0.2 \, \text{kJ/K}) = -150 \, \text{kJ} - 60 \, \text{kJ} = -210 \, \text{kJ}

Exothermic vs Endothermic Reactions:

 Exothermic: Releases heat, e.g., combustion of methane.

 Endothermic: Absorbs heat, e.g., melting of ice.

UNIT 7: Equilibrium

Le Chatelier’s Principle:

 Definition: When a system at equilibrium is disturbed, it will shift to counteract the

disturbance.

 Example: N2(g)+3H2(g)⇌2NH3(g)N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) If pressure

is increased, the system shifts to the side with fewer moles of gas.

Equilibrium Constant (Kc):

 The equilibrium constant for a reaction aA+bB⇌cC+dDaA + bB \rightleftharpoons cC + dD is:

Kc=[C]c[D]d[A]a[B]bK_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}

Interpretation of Kc=0.04K_c = 0.04:

 A KcK_c value of 0.04 suggests that the concentration of reactants is greater than that of
products at equilibrium.

pH and pOH:

 pH: A measure of the acidity of a solution. pH=−log⁡[H+]\text{pH} = -\log[H^+]

 pOH: A measure of the alkalinity of a solution. pOH=−log⁡[OH−]\text{pOH} = -\log[OH^-]

 pH of 0.001 M HCl: pH=−log⁡[0.001]=3\text{pH} = -\log[0.001] = 3

Acidic, Basic, and Neutral Solutions:

 Acidic: pH < 7

 Neutral: pH = 7

 Basic: pH > 7

UNIT 8: Redox Reactions

Oxidation and Reduction:

  Oxidation: Loss of electrons (e.g., Fe→Fe2++2e−Fe \to Fe^{2+} + 2e^-).

 Reduction: Gain of electrons (e.g., Cu2++2e−→CuCu^{2+} + 2e^- \to Cu).

Oxidizing and Reducing Agents in CuO+H2→Cu+H2OCuO + H_2 \to Cu + H_2O:

 Oxidizing agent: CuOCuO (it gains electrons).

 Reducing agent: H2H_2 (it loses electrons).

Oxidation Number:

 Definition: The charge an atom would have if the compound was composed of ions.

 Example: In H2OH_2O, the oxidation number of hydrogen is +1 and oxygen is -2.

Balance Redox Reaction (MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺):

 This is done using the ion-electron method by balancing atoms and charges.

Disproportionation vs Comproportionation:

 Disproportionation: A substance is simultaneously oxidized and reduced.

 Comproportionation: Two substances combine, and one is oxidized and the other reduced.

UNIT 9: Hydrogen

Physical and Chemical Properties of Hydrogen:

 Physical: Colorless, odorless, tasteless, highly flammable.

 Chemical: Reacts with oxygen to form water, forms hydrides with metals.

Heavy Water (D₂O):

 Water where both hydrogen atoms are replaced by deuterium. Used as a moderator in
nuclear reactors.

Electrolysis of Water:

 Water is split into hydrogen and oxygen by passing an electric current through it:
2H2O→2H2+O22H_2O \to 2H_2 + O_2

Hydrogen Economy:

 A concept where hydrogen is used as a primary energy source, replacing fossil fuels.

Hydrides:

 Definition: Compounds of hydrogen with metals or non-metals.

 Types: Ionic (e.g., NaH), covalent (e.g., CH₄), and metallic (e.g., LaH₃).

UNIT 10: The s-Block Elements

Reactivity of Alkali Metals:

  Alkali metals are highly reactive due to their low ionization energy.

Diagonal Relationship between Li and Mg:

 Both have similar properties, such as forming basic oxides and reacting with water (though
Mg is less reactive).

Chemical Properties of Sodium:

 Sodium reacts vigorously with water to form sodium hydroxide and hydrogen gas.

BeCl₂ vs MgCl₂:

 BeCl₂ is covalent due to small size and high ionization energy of Be, while MgCl₂ is ionic due
to the larger size of Mg.

Uses of CaCO₃ and Ca(OH)₂:

 CaCO₃: Used in cement, limestone, and as an antacid.

 Ca(OH)₂: Used in the preparation of lime water and in the treatment of acidic soils.

UNIT 11: Some p-Block Elements

Structure and Uses of Borax:

 Borax (Na₂B₄O₇.10H₂O) is a crystalline solid used in glassmaking and as a cleaning agent.

Allotropes of Carbon:

 Diamond: Hard, used in cutting tools.

 Graphite: Soft, used in pencils and lubricants.

Fullerenes:

 Molecules of carbon in spherical, ellipsoidal, or cylindrical shapes, used in nanotechnology

and medicine.

Acidic Nature of CO₂:

 CO₂ dissolves in water to form carbonic acid (H₂CO₃), making it acidic.

Reactivity of Boron and Aluminum:

 Boron is less reactive than aluminum due to its smaller atomic size and higher ionization
energy.

UNIT 12: Organic Chemistry – Some Basic Principles

Homologous Series:

 A series of compounds with the same functional group and similar chemical properties, e.g.,
alkanes (CₙH₂ₙ₊₂).

Structural Isomers:
  Compounds with the same molecular formula but different structures, e.g., butane and
isobutane.

Inductive and Resonance Effects:

 Inductive effect: Electron withdrawal or donation through sigma bonds.

 Resonance effect: Delocalization of electrons in conjugated systems.

IUPAC Names:

 CH₃CH₂CH₂OH: Propanol

 CH₃COCH₃: Acetone

Difference between Alkanes, Alkenes, and Alkynes:

 Alkanes: Saturated hydrocarbons (single bonds).

 Alkenes: Unsaturated hydrocarbons with one or more double bonds.

 Alkynes: Unsaturated hydrocarbons with one or more triple bonds.

UNIT 13: Hydrocarbons

General Formula for Alkanes, Alkenes, and Alkynes:

 Alkanes: CnH2n+2CₙH₂ₙ₊₂

 Alkenes: CnH2nCₙH₂ₙ

 Alkynes: CnH2n−2CₙH₂ₙ₋₂

Markovnikov’s Rule:

 When HX is added to an alkene, the hydrogen atom bonds to the carbon with the greatest
number of hydrogen atoms.

Preparation and Properties of Ethene (C₂H₄):

 Ethene is produced by cracking hydrocarbons.

 It is used in the production of plastics and as a plant hormone.

Aromatic Hydrocarbons:

 Hydrocarbons with conjugated ring systems, e.g., benzene.

Benzene vs Cyclohexane:

 Benzene is an aromatic compound with alternating double bonds, while cyclohexane is a

saturated hydrocarbon.