B.Sc. 1st Year Chemistry

Basic Chemistry I | CHE-101 | Complete Theory Notes

University: Tribhuvan University

Course: Basic Chemistry I

Course No.: CHE-101

Full Marks: 100

Pass Marks: 35

Nature: Theory

Note: Practical Chemistry I (CHE-102) is not included here.

Contents

Inorganic Chemistry

1. Atomic Structure 2. Nuclear Chemistry 3. Periodic Classification and Periodic Properties 4. Chemical Bonding 5. Acids, Bases and Buffers 6. Qualitative and Quantitative Analysis

Organic Chemistry

7. Structure and Properties of Organic Compounds 8. Alkanes and Free-Radical Reactions 9. Stereochemistry 10. Alkyl Halides: SN1, SN2, E1 and E2 11. Alcohols and Ethers 12. Alkenes 13. Alkynes

Physical Chemistry

14. Gaseous State 15. Liquid and Solid States 16. Chemical Equilibrium 17. Ionic Equilibrium 18. Colligative Properties 19. Chemical Kinetics 20. Thermochemistry and Thermodynamics

Revision

21. Important Reactions, Formulae and Exam Strategy

1. Atomic Structure

1.1 Fundamental Particles

Particle Charge Relative Mass Location
Electron −1 1/1836 Outside nucleus
Proton +1 1 Nucleus
Neutron 0 1 Nucleus

1.2 Rutherford Nuclear Model

Rutherford proposed that an atom contains a small, dense, positively charged nucleus surrounded by electrons and that most of the atomic volume is empty space.

1.3 Bohr's Atomic Model

ΔE = E₂ − E₁ = hν

1.4 de Broglie Equation

λ = h / mv where: λ = wavelength h = Planck constant m = mass v = velocity

1.5 Heisenberg Uncertainty Principle

Δx × Δp ≥ h / 4π

It is impossible to determine simultaneously the exact position and exact momentum of a microscopic particle.

1.6 Quantum Numbers

Quantum number Symbol Meaning
Principal n Shell and approximate energy
Azimuthal l Subshell and orbital shape
Magnetic m Orientation of orbital
Spin s Electron spin

1.7 Aufbau Principle

Electrons occupy orbitals in order of increasing energy.

1s ↓ 2s ↓ 2p ↓ 3s ↓ 3p ↓ 4s ↓ 3d ↓ 4p ↓ 5s ↓ 4d ↓ 5p

1.8 Hund's Rule

Degenerate orbitals are singly occupied with parallel spins before pairing occurs.

1.9 Pauli Exclusion Principle

No two electrons in an atom can possess the same set of four quantum numbers.

Exam focus: Explain Bohr's model, derive the de Broglie equation, explain quantum numbers, and discuss Heisenberg uncertainty principle.
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2. Nuclear Chemistry

2.1 Radioactivity

Radioactivity is the spontaneous disintegration of unstable atomic nuclei accompanied by emission of radiation.

2.2 Alpha Decay

²³⁸₉₂U → ²³⁴₉₀Th + ⁴₂He

2.3 Beta Decay

n → p + e⁻ + ν̄

2.4 Gamma Emission

Gamma emission involves release of high-energy electromagnetic radiation from an excited nucleus without changing atomic or mass number.

2.5 Radioactive Decay Law

N = N₀e⁻ᵏᵗ ln(N₀/N) = kt t½ = 0.693/k

2.6 Nuclear Fission

A heavy nucleus splits into smaller nuclei with release of energy and neutrons.

2.7 Nuclear Fusion

Two light nuclei combine to form a heavier nucleus with release of very large amounts of energy.

Applications of Radioisotopes

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3. Periodic Classification and Periodic Properties

3.1 Modern Periodic Law

The physical and chemical properties of elements are periodic functions of their atomic numbers.

3.2 Atomic Radius

Atomic radius generally decreases from left to right across a period because effective nuclear charge increases.

Atomic radius increases down a group because additional electron shells are added.

3.3 Ionization Energy

Ionization energy is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom.

It generally increases across a period and decreases down a group.

3.4 Electron Affinity

Electron affinity refers to the energy change when an electron is added to a gaseous atom.

3.5 Electronegativity

Electronegativity is the tendency of an atom in a molecule to attract the shared electron pair toward itself.

3.6 Effective Nuclear Charge

Zeff ≈ Z − shielding effect

3.7 Periodic Trends

Property Across period Down group
Atomic radius Decreases Increases
Ionization energy Generally increases Decreases
Electronegativity Generally increases Decreases
Metallic character Decreases Increases
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4. Chemical Bonding

4.1 Ionic Bond

An ionic bond results from electrostatic attraction between oppositely charged ions.

Na → Na⁺ + e⁻ Cl + e⁻ → Cl⁻ Na⁺ + Cl⁻ → NaCl

4.2 Covalent Bond

A covalent bond is formed by sharing of electron pairs between atoms.

4.3 Lewis Concept

Lewis structures represent valence electrons and bonding electron pairs around atoms.

4.4 VSEPR Theory

Electron pairs around a central atom arrange themselves to minimize electron-pair repulsion.

Species Approximate geometry
BeCl₂ Linear
BF₃ Trigonal planar
CH₄ Tetrahedral
NH₃ Trigonal pyramidal
H₂O Bent

4.5 Hybridization

Hybridization Geometry Example
sp Linear BeCl₂
sp² Trigonal planar BF₃
sp³ Tetrahedral CH₄
sp³d Trigonal bipyramidal PCl₅
sp³d² Octahedral SF₆

4.6 Molecular Orbital Theory

Atomic orbitals combine to form bonding and antibonding molecular orbitals.

Bond order = 1/2 (Nb − Na)

Nb = electrons in bonding molecular orbitals
Na = electrons in antibonding molecular orbitals

4.7 Hydrogen Bonding

Hydrogen bonding occurs when hydrogen bonded to a highly electronegative atom interacts with a lone pair on another electronegative atom.

Examples include H₂O, HF and NH₃.

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5. Acids, Bases and Buffers

5.1 Arrhenius Concept

An Arrhenius acid produces H⁺ in aqueous solution, while an Arrhenius base produces OH⁻.

5.2 Brønsted-Lowry Concept

An acid is a proton donor and a base is a proton acceptor.

HCl + H₂O → H₃O⁺ + Cl⁻ NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

5.3 Lewis Concept

A Lewis acid accepts an electron pair and a Lewis base donates an electron pair.

BF₃ + :NH₃ → F₃B←NH₃

5.4 pH

pH = −log[H⁺] pOH = −log[OH⁻] pH + pOH = 14 at 25°C

5.5 Buffer Solution

A buffer resists significant change in pH when small quantities of acid or base are added.

Acidic Buffer

Weak acid + its salt with a strong base.

pH = pKa + log([salt]/[acid])

Basic Buffer

Weak base + its salt with a strong acid.

pOH = pKb + log([salt]/[base])

Common Ion Effect

Suppression of ionization of a weak electrolyte by addition of a strong electrolyte containing a common ion.

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6. Principles of Qualitative and Quantitative Analysis

6.1 Qualitative Analysis

Qualitative analysis identifies the ions or chemical species present in a sample.

6.2 Quantitative Analysis

Quantitative analysis determines the amount or concentration of a substance.

6.3 Gravimetric Analysis

Sample ↓ Precipitation ↓ Filtration ↓ Washing ↓ Drying / Ignition ↓ Weighing ↓ Calculation

6.4 Volumetric Analysis

Volumetric analysis determines the concentration of an analyte by measuring the volume of a standard solution required for reaction.

M₁V₁ / n₁ = M₂V₂ / n₂

6.5 Acid-Base Titration

HCl + NaOH → NaCl + H₂O

6.6 Oxidation-Reduction Titration

Redox titration is based on electron transfer between oxidizing and reducing agents.

6.7 Indicators

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7. Structure and Properties of Organic Compounds

7.1 Classification

7.2 Functional Groups

Functional group Class
−OH Alcohol
−CHO Aldehyde
>C=O Ketone
−COOH Carboxylic acid
−NH₂ Amine
−X Halo compound

7.3 Homologous Series

A homologous series is a family of organic compounds having the same functional group and general formula, with successive members differing by −CH₂−.

7.4 Electrophiles and Nucleophiles

Electrophile: electron-deficient species that accepts an electron pair.
Nucleophile: electron-rich species that donates an electron pair to an electrophile.

Examples

Electrophiles Nucleophiles
H⁺ OH⁻
NO₂⁺ CN⁻
BF₃ NH₃
Carbocations RO⁻

7.5 Inductive Effect

The inductive effect is the permanent displacement of sigma electron density caused by electronegativity differences.

Electron-withdrawing groups show −I effect while alkyl groups generally show +I effect.

7.6 Resonance

Resonance occurs when a molecule cannot be represented adequately by a single Lewis structure and is represented by contributing structures.

O₃ ↔ resonance hybrid In aromatic systems: π electrons are delocalized over the conjugated framework.

7.7 Hyperconjugation

Hyperconjugation is delocalization involving a σ bond, commonly a C−H bond adjacent to a π system or positively charged carbon.

7.8 Carbocation Stability

Benzylic / allylic > 3° > 2° > 1° > methyl

Resonance stabilization makes benzylic and allylic carbocations particularly stable.

7.9 Free Radical Stability

Benzylic / allylic > 3° > 2° > 1° > methyl
Exam focus: Explain inductive effect, resonance, hyperconjugation and their influence on stability of reactive intermediates.
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8. Alkanes and Free-Radical Reactions

8.1 General Formula

CₙH₂ₙ₊₂

8.2 Preparation by Hydrogenation

CH₂=CH₂ + H₂ Ni / Pt / Pd → CH₃−CH₃

8.3 Wurtz Reaction

2R−X + 2Na dry ether → R−R + 2NaX

Wurtz reaction is useful mainly for preparation of symmetrical higher alkanes.

8.4 Decarboxylation

R−COONa + NaOH CaO, heat → R−H + Na₂CO₃

8.5 Free-Radical Chlorination of Methane

CH₄ + Cl₂ hv → CH₃Cl + HCl

Mechanism

Step 1: Initiation

Cl₂ --hv--> 2 Cl•

Step 2: Propagation

Cl• + CH₄ → HCl + CH₃• CH₃• + Cl₂ → CH₃Cl + Cl•

Step 3: Termination

Cl• + Cl• → Cl₂ CH₃• + Cl• → CH₃Cl CH₃• + CH₃• → C₂H₆

8.6 Relative Reactivity of Hydrogen

Hydrogen abstraction is influenced by the stability of the radical formed.

3° H > 2° H > 1° H > methyl H

8.7 Combustion

CH₄ + 2O₂ → CO₂ + 2H₂O + heat
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9. Stereochemistry

9.1 Stereoisomerism

Stereoisomers have the same molecular formula and connectivity but differ in three-dimensional arrangement.

9.2 Geometrical Isomerism

Geometrical isomerism occurs because rotation around a double bond is restricted.

cis-2-butene CH₃ CH₃ \ / C = C / \ H H

9.3 Optical Isomerism

A chiral molecule is not superimposable on its mirror image.

9.4 Enantiomers

Enantiomers are non-superimposable mirror images.

9.5 Diastereomers

Diastereomers are stereoisomers that are not mirror images.

9.6 Optical Activity

9.7 R/S Configuration

  1. Assign priorities according to atomic number.
  2. Place the lowest priority group away from the observer.
  3. Trace 1 → 2 → 3.
  4. Clockwise = R.
  5. Anticlockwise = S.

9.8 Conformational Isomerism

Conformers are produced by rotation about single σ bonds.

Ethane

Butane

The anti conformation is generally more stable than gauche because bulky methyl groups are farther apart.

Exam focus: Explain enantiomers, diastereomers, chirality, optical activity, R/S configuration and conformational analysis.
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10. Alkyl Halides: SN1, SN2, E1 and E2 Mechanisms

10.1 General Structure

R−X X = F, Cl, Br or I

10.2 Preparation from Alcohols

ROH + HX → RX + H₂O

10.3 Reactivity of Leaving Groups

I⁻ > Br⁻ > Cl⁻ >> F⁻

10.4 SN2 Mechanism

SN2 is a bimolecular nucleophilic substitution reaction occurring through a single concerted step.

CH₃Br + OH⁻ → CH₃OH + Br⁻
Nu⁻ attacks the carbon from the backside. Nu⁻ ↓ C — Br ↑ simultaneous bond formation and bond breaking → C — Nu + Br⁻

Characteristics

10.5 SN1 Mechanism

SN1 occurs through formation of a carbocation intermediate.

(CH₃)₃C−Br + H₂O → (CH₃)₃C−OH + HBr

Step 1: Carbocation formation

(CH₃)₃C−Br → (CH₃)₃C⁺ + Br⁻ Slow / rate-determining step

Step 2: Nucleophilic attack

(CH₃)₃C⁺ + H₂O → (CH₃)₃C−OH₂⁺

Step 3: Deprotonation

(CH₃)₃C−OH₂⁺ → (CH₃)₃C−OH + H⁺

Characteristics

10.6 E1 Mechanism

E1 elimination occurs through a carbocation intermediate.

Step 1: R−CH₂−CH(X)−R → R−CH₂−C⁺H−R + X⁻ Step 2: Base removes β-H C−H electrons form C=C → Alkene

E1 commonly competes with SN1 because both begin with carbocation formation.

10.7 E2 Mechanism

E2 is a one-step bimolecular elimination reaction in which proton abstraction and leaving-group departure occur simultaneously.

Base removes β-H ↓ C−H electrons form C=C ↓ C−X bond breaks ↓ Alkene + X⁻

The preferred transition state has the β-C−H and C−X bonds approximately anti-periplanar. :contentReference[oaicite:2]{index=2}

10.8 Zaitsev Rule

In many elimination reactions, the more substituted alkene is formed preferentially because it is generally more stable.

CH₃−CHBr−CH₂−CH₃ alc. KOH, heat → CH₃−CH=CH−CH₃ + CH₂=CH−CH₂−CH₃ 2-butene = major product

10.9 SN1 vs SN2 vs E1 vs E2

Feature SN1 SN2 E1 E2
Steps 2 1 2 1
Intermediate Carbocation None Carbocation None
Rate k[RX] k[RX][Nu] k[RX] k[RX][Base]
Favored substrate 3° Methyl/1° 3° 2°/3° commonly
Main product Substitution Substitution Alkene Alkene
Master rule:

Primary + strong nucleophile → SN2 is commonly favored.
Tertiary + weak nucleophile / polar protic medium → SN1/E1.
Strong base + suitable substrate → E2 becomes important.

Actual product depends on substrate, nucleophile/base, solvent, temperature and stereochemical requirements. :contentReference[oaicite:3]{index=3}
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11. Alcohols and Ethers

11.1 Classification of Alcohols

11.2 Preparation by Hydration of Alkene

CH₂=CH₂ + H₂O H⁺ → CH₃CH₂OH

11.3 Reaction with Sodium

2ROH + 2Na → 2RONa + H₂↑

11.4 Oxidation

Primary alcohol → Aldehyde → Carboxylic acid Secondary alcohol → Ketone Tertiary alcohol → resistant to mild oxidation

11.5 Dehydration of Alcohol

CH₃CH₂OH conc. H₂SO₄, heat → CH₂=CH₂ + H₂O

Mechanistic idea for secondary/tertiary alcohols

1. Protonation of −OH ROH + H⁺ → ROH₂⁺ 2. Water leaves ROH₂⁺ → R⁺ + H₂O 3. β-H removal Base removes β-H → C=C formation

Primary alcohols often undergo dehydration through pathways that avoid a free primary carbocation, whereas secondary and tertiary alcohols can undergo E1-type pathways under suitable acidic conditions.

11.6 Ether Formation: Williamson Synthesis

RONa + R'−X → R−O−R' + NaX

Mechanism

RO⁻ attacks the carbon bearing the leaving group from the backside. RO⁻ + R'−X → R−O−R' + X⁻ SN2 mechanism

Primary alkyl halides are particularly suitable for Williamson ether synthesis because SN2 displacement is favored.

11.7 Cleavage of Ethers by HI

R−O−R' + HI → RI + R'OH With excess HI: R−O−R' + 2HI → RI + R'I + H₂O
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12. Alkenes

12.1 General Formula

CₙH₂ₙ

12.2 Preparation by Dehydrohalogenation

R−CH₂−CH₂−X alc. KOH → R−CH=CH₂ + KX + H₂O

12.3 Electrophilic Addition

The π bond of an alkene is electron-rich and attacks an electrophile. Electrophilic addition is therefore a major reaction type of alkenes.

12.4 Addition of HX

CH₃−CH=CH₂ + HBr → CH₃−CHBr−CH₃

Mechanism

Step 1: π electrons attack H⁺ CH₃−CH=CH₂ + H⁺ → carbocation The more stable carbocation is preferentially formed. Step 2: Br⁻ attacks carbocation carbocation + Br⁻ → 2-bromopropane

12.5 Markovnikov's Rule

In addition of HX to an unsymmetrical alkene under ordinary ionic conditions, hydrogen generally adds to the carbon already bearing more hydrogen, while X becomes attached to the more substituted carbon.

12.6 Peroxide Effect

In the presence of peroxides, HBr can undergo radical addition giving the anti-Markovnikov product.

Initiation: ROOR → 2RO• RO• + HBr → ROH + Br• Propagation: Br• + alkene → carbon radical carbon radical + HBr → product + Br• The Br• radical is regenerated.
The peroxide effect is characteristic of HBr in the standard undergraduate treatment. It should not be generalized to HCl or HI.

12.7 Halogenation

CH₂=CH₂ + Br₂ CCl₄ → BrCH₂−CH₂Br

Mechanism

Alkene π bond attacks Br₂ ↓ Bromonium ion ↓ Br⁻ attacks from opposite side ↓ Vicinal dibromide Overall: anti addition

12.8 Hydration

CH₃−CH=CH₂ + H₂O H⁺ → CH₃−CHOH−CH₃

12.9 Hydroboration-Oxidation

Alkene 1. BH₃·THF 2. H₂O₂ / OH⁻ → Alcohol

Hydroboration-oxidation gives anti-Markovnikov hydration with overall syn addition of H and OH.

12.10 Ozonolysis

RCH=CHR' 1. O₃ 2. Zn/H₂O → RCHO + R'CHO

Ozonolysis cleaves the carbon-carbon double bond and is useful for locating the position of unsaturation.

12.11 Polymerization

n CH₂=CH₂ → (−CH₂−CH₂−)ₙ polyethylene
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13. Alkynes

13.1 General Formula

CₙH₂ₙ₋₂

13.2 Preparation by Double Dehydrohalogenation

Vicinal / geminal dihalide ↓ strong base ↓ Alkyne

13.3 Hydrogenation

RC≡CR' + H₂ Pd / Pt / Ni → RCH=CHR'

Further hydrogenation can produce the alkane.

13.4 Lindlar Reduction

RC≡CR' H₂ / Lindlar catalyst → cis-RCH=CHR'

13.5 Dissolving Metal Reduction

RC≡CR' Na / NH₃(l) → trans-RCH=CHR'

13.6 Addition of HX

RC≡CH + HBr → vinyl bromide → gem-dibromide

13.7 Hydration of Alkyne

RC≡CH HgSO₄ / H₂SO₄ / H₂O → enol → carbonyl compound

13.8 Hydroboration-Oxidation

Terminal alkyne 1. bulky borane 2. H₂O₂ / OH⁻ → aldehyde

13.9 Acidity of Terminal Alkynes

Terminal alkynes are relatively acidic because the conjugate base places negative charge on an sp-hybridized carbon.

RC≡CH + NaNH₂ → RC≡C⁻Na⁺ + NH₃

13.10 Acetylide Ion as Nucleophile

RC≡C⁻ + R'−X → RC≡C−R' + X⁻

This is an SN2 reaction and works particularly well with methyl and primary alkyl halides.

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14. Gaseous State

14.1 Gas Laws

Boyle's law: PV = constant Charles' law: V/T = constant Avogadro's law: V/n = constant

14.2 Ideal Gas Equation

PV = nRT

14.3 Dalton's Law

Ptotal = P₁ + P₂ + P₃ + ...

14.4 Kinetic Molecular Theory

14.5 van der Waals Equation

(P + an²/V²)(V − nb) = nRT

The constants a and b account for intermolecular attraction and finite molecular volume respectively.

14.6 Critical Phenomena

At the critical temperature, a gas cannot be liquefied merely by increasing pressure.

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15. Liquid and Solid States

15.1 Properties of Liquids

15.2 Surface Tension

Surface tension is the tendency of a liquid surface to behave as an elastic membrane due to intermolecular forces.

15.3 Viscosity

Viscosity is the resistance of a liquid to flow.

15.4 Solids

Solids have definite shape and volume because particles are held relatively strongly in an ordered or disordered structure.

Types of Crystalline Solids

Unit Cell

A unit cell is the smallest repeating structural unit of a crystal lattice.

Bragg's Law

nλ = 2d sinθ

Bragg's law is used in X-ray diffraction to study crystal structures.

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16. Chemical Equilibrium

16.1 Dynamic Equilibrium

Chemical equilibrium is a dynamic state in which forward and reverse reactions occur at equal rates and concentrations remain constant.

16.2 Equilibrium Constant

aA + bB ⇌ cC + dD
Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ

16.3 Le Chatelier's Principle

When a system at equilibrium is disturbed, it shifts in a direction that tends to oppose the disturbance.

Effect of Concentration

Increasing the concentration of a reactant generally shifts equilibrium toward products.

Effect of Pressure

For gaseous equilibria, increasing pressure favors the side with fewer moles of gas.

Effect of Temperature

For an exothermic reaction, increasing temperature favors the reverse direction. For an endothermic reaction, increasing temperature favors the forward direction.

Relation between Kp and Kc

Kp = Kc(RT)^Δn
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17. Ionic Equilibrium

17.1 Strong and Weak Electrolytes

Strong electrolytes ionize almost completely, while weak electrolytes ionize only partially in solution.

17.2 Acid Dissociation Constant

HA + H₂O ⇌ H₃O⁺ + A⁻
Ka = [H⁺][A⁻]/[HA]

17.3 Base Dissociation Constant

BOH ⇌ B⁺ + OH⁻
Kb = [B⁺][OH⁻]/[BOH]

17.4 Ionic Product of Water

Kw = [H⁺][OH⁻] At 25°C: Kw = 1.0 × 10⁻¹⁴

17.5 Henderson-Hasselbalch Equation

pH = pKa + log([A⁻]/[HA])

17.6 Solubility Product

For a sparingly soluble salt, the solubility product is the product of ionic concentrations raised to their stoichiometric powers.

AgCl(s) ⇌ Ag⁺ + Cl⁻
Ksp = [Ag⁺][Cl⁻]

17.7 Common Ion Effect

Addition of a common ion decreases the solubility or ionization of a weak electrolyte.

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18. Colligative Properties

Colligative properties depend primarily on the number of dissolved solute particles rather than their chemical identity.

18.1 Relative Lowering of Vapour Pressure

(P° − P)/P° = Xsolute

18.2 Elevation of Boiling Point

ΔTb = Kb m

18.3 Depression of Freezing Point

ΔTf = Kf m

18.4 Osmotic Pressure

π = CRT For electrolytes: π = iCRT

18.5 van't Hoff Factor

i = observed colligative property / calculated colligative property

The van't Hoff factor accounts for association or dissociation of solute particles.

18.6 Applications

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19. Chemical Kinetics

19.1 Rate of Reaction

Rate = change in concentration / change in time

19.2 Rate Law

aA + bB → products
Rate = k[A]ᵐ[B]ⁿ

m and n are determined experimentally and need not equal the stoichiometric coefficients.

19.3 Order of Reaction

Overall order is the sum of powers of concentration terms in the experimentally determined rate law.

19.4 Zero-Order Reaction

[A]t = [A]₀ − kt t½ = [A]₀ / 2k

19.5 First-Order Reaction

ln([A]₀/[A]) = kt k = 2.303/t log([A]₀/[A]) t½ = 0.693/k

19.6 Second-Order Reaction

1/[A] − 1/[A]₀ = kt t½ = 1/(k[A]₀)

19.7 Arrhenius Equation

k = Ae⁻ᴱᵃ/ᴿᵀ ln k = ln A − Ea/RT

A is the frequency factor and Ea is the activation energy.

19.8 Two-Temperature Form

log(k₂/k₁) = Ea / 2.303R × (T₂ − T₁)/(T₁T₂)

19.9 Activation Energy

Activation energy is the minimum energy barrier that reacting molecules must overcome to reach the transition state.

19.10 Collision Theory

According to collision theory, molecules must collide with suitable orientation and sufficient energy to produce reaction.

19.11 Catalyst

A catalyst provides an alternative reaction pathway with lower activation energy and increases reaction rate without being permanently consumed.

Exam focus: Derive integrated rate equations for zero-, first- and second-order reactions and explain Arrhenius equation and activation energy.
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20. Thermochemistry and Thermodynamics

20.1 System and Surroundings

20.2 Types of System

20.3 First Law of Thermodynamics

ΔU = q + w

Energy can neither be created nor destroyed; it can only be transformed from one form to another.

20.4 Enthalpy

H = U + PV

At constant pressure, heat exchanged is related to enthalpy change.

20.5 Hess's Law

The enthalpy change of a reaction is independent of the pathway and depends only on initial and final states.

20.6 Bond Enthalpy

ΔH ≈ Σ bond energies of bonds broken − Σ bond energies of bonds formed

20.7 Entropy

Entropy is a thermodynamic state function related to the dispersal of energy and the number of accessible microscopic arrangements.
ΔS = qrev / T

20.8 Second Law of Thermodynamics

For a spontaneous process, the entropy of the universe increases.

ΔSuniverse = ΔSsystem + ΔSsurroundings For spontaneous process: ΔSuniverse > 0

20.9 Gibbs Free Energy

ΔG = ΔH − TΔS
ΔG Meaning
Negative Spontaneous in the specified conditions
Positive Non-spontaneous in the specified conditions
Zero Equilibrium

20.10 Relation between Free Energy and Equilibrium

ΔG° = −RT ln K

20.11 Carnot Cycle

Isothermal expansion ↓ Adiabatic expansion ↓ Isothermal compression ↓ Adiabatic compression ↓ Initial state

20.12 Carnot Efficiency

η = 1 − T₂/T₁

T₁ is the temperature of the hot reservoir and T₂ is the temperature of the cold reservoir.

20.13 Spontaneity and Temperature

ΔH ΔS Spontaneity
− + Spontaneous at all temperatures
+ − Non-spontaneous at all temperatures
− − Depends on temperature
+ + Depends on temperature
Exam focus: First law, Hess's law, entropy, second law, Gibbs free energy, Carnot cycle, spontaneity and numerical problems.
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21. Important Reactions, Formulae and Exam Strategy

21.1 Must-Know Organic Reactions

Reaction Important reagent / condition
Wurtz reaction Na / dry ether
Free-radical halogenation Cl₂ / hν
SN2 Strong nucleophile, suitable primary substrate
SN1 Stable carbocation, polar protic medium
E2 Strong base
Alcohol dehydration Conc. H₂SO₄ / heat
Williamson ether synthesis RO⁻ + primary R−X
HX addition Markovnikov orientation normally
HBr peroxide effect Anti-Markovnikov radical addition
Halogen addition Br₂ / inert solvent
Hydroboration oxidation BH₃·THF; H₂O₂/OH⁻
Ozonolysis O₃ followed by reductive work-up
Lindlar reduction cis-alkene
Na/NH₃ reduction trans-alkene

21.2 Mechanisms You Must Practice Drawing

  1. Free-radical chlorination of methane
  2. SN1 mechanism
  3. SN2 mechanism
  4. E1 mechanism
  5. E2 mechanism
  6. Acid-catalyzed dehydration of alcohol
  7. Williamson ether synthesis
  8. Electrophilic addition of HBr to alkene
  9. Peroxide effect of HBr
  10. Bromination of alkene through bromonium ion
  11. Hydroboration-oxidation

21.3 Important Physical Chemistry Formulae

PV = nRT Kp = Kc(RT)^Δn ΔTb = Kb m ΔTf = Kf m π = iCRT ln([A]₀/[A]) = kt t½(first order) = 0.693/k k = Ae⁻ᴱᵃ/ᴿᵀ ΔU = q + w H = U + PV ΔG = ΔH − TΔS ΔG° = −RT ln K η = 1 − T₂/T₁ nλ = 2d sinθ

21.4 High-Value 10-Mark Questions

  1. Discuss Bohr's atomic model and explain its limitations.
  2. Explain quantum numbers and electronic configuration.
  3. Discuss periodic trends in atomic radius, ionization energy, electron affinity and electronegativity.
  4. Explain VSEPR theory and hybridization with suitable examples.
  5. Explain molecular orbital theory and calculate bond order.
  6. Explain SN1 and SN2 mechanisms with comparison.
  7. Explain E1 and E2 mechanisms and discuss factors affecting elimination.
  8. Explain free-radical halogenation of alkanes with initiation, propagation and termination steps.
  9. Discuss electrophilic addition reactions of alkenes with mechanisms.
  10. Explain Markovnikov's rule and peroxide effect.
  11. Discuss stereochemistry including optical and geometrical isomerism.
  12. Derive integrated rate equations for different orders of reactions.
  13. Explain Arrhenius equation and activation energy.
  14. Discuss first and second laws of thermodynamics.
  15. Explain Gibbs free energy and its relation to spontaneity and equilibrium.

21.5 Important 5-Mark Questions

21.6 How to Write a 10-Mark Chemistry Answer

For theory:

1. Definition / introduction
2. Principle or theory
3. Explanation
4. Equation / derivation
5. Diagram or structure where appropriate
6. Examples
7. Applications / significance
8. Limitations if relevant
9. Short conclusion

For an organic reaction:

1. Starting compound
2. Reagent and conditions
3. Major product
4. Mechanism
5. Reactive intermediate
6. Electron movement
7. Regioselectivity
8. Stereochemical outcome
9. Factors affecting the reaction

21.7 Organic Chemistry Master Revision

ALKANE ↓ Free-radical substitution ALKYL HALIDE ↓ SN1 / SN2 ↓ Substitution ALKYL HALIDE ↓ E1 / E2 ↓ Alkene ALCOHOL ↓ Oxidation / substitution / dehydration ALKENE ↓ Electrophilic addition ↓ Alcohol / Haloalkane / Dihalide / etc. ALKYNE ↓ Addition / reduction / hydration REACTION MECHANISM ↓ Identify electron-rich site ↓ Identify electron-poor site ↓ Identify leaving group ↓ Draw electron movement ↓ Identify intermediate ↓ Determine major product
Final exam reminder:

In Organic Chemistry, do not memorize only the final product. Learn:

Reagent → attacking species → electron movement → intermediate → product → stereochemistry.

This approach makes it much easier to solve unfamiliar reaction questions rather than only reproducing memorized equations.
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