Comprehensive University Study Notes: Chemistry (National Curriculum 2022-23 Curriculum of Pakistan Curriculum Pakistan Core Curriculum)

History of Chemistry and Thought Experiments

A thought experiment is defined as an experiment performed in the "laboratory of the human mind." It involves studying a specific hypothetical situation, observing the mental outcome, and reaching a conclusion. Unlike real experiments done in the physical world, these play a vital role in developing scientific theories. Famous examples include the discovery of the atom, Maxwell's Demon, and Schrodinger's Cat.

The Discovery of the Atom

In 500 BCE, Leucippus, known as the father of atomic philosophy, was the first to think about the atom. In 430 BCE, his student Democritus proposed that matter is made of uniform, indivisible particles called "atomos" (Greek for uncuttable). Democritus's thought experiment suggested that atoms existed even for feelings and the soul, and that their shapes determines properties (e.g., sour taste is due to needle-shaped atoms). In 1808, John Dalton converted this philosophy into a scientific theory in his book A New System of Chemical Philosophy, later supported by Gay-Lussac and Amedeo Avogadro.

Maxwell's Demon Experiment

In 1867, James Clerk Maxwell imagined two containers (A and B) filled with gas at the same temperature, separated by a wall with a small window guarded by a "demon." The demon would allow fast molecules to move from A to B and slow ones from B to A, creating a temperature difference without spending energy. This theoretically violated the second law of thermodynamics, which states that heat flows naturally from hot to cold, increasing entropy. The solution is that the demon requires energy to detect speeds and move the window, thus increasing entropy and following the law.

Schrodinger's Cat Experiment

Erwin Schrodinger (1887–1961) proposed this to address problems in quantum mechanics. He imagined a cat in an opaque box with poison, a Geiger counter, and radioactive material. If the material decays, poison kills the cat. In quantum terms, until observed, the material is in a superposition of states (decayed and undecayed), meaning the cat is simultaneously alive and dead. This relates to his equation laying the foundation for quantum chemistry, suggesting sub-atomic particles exist as both particle and wave until measured.

Imam Al-Ghazali and Causality

Abu Hamid Al-Ghazali (1058–1111 AD) was a theologian who challenged Neoplatonic thinkers in his work Tahafut al-Falsafa. In his "Burning Cotton Experiment," he initially proposed "necessary causation," where a burning flame is the necessary cause of cotton burning. However, he later moved to the idea of "continuation of events," arguing that nature follows specific principles but God's will can suspend these habitual successions (miracles). He argued that people confuse the succession of events with causation, a view later supported by David Hume.

Inductive and Deductive Reasoning

Reasoning is stepwise thinking with a goal. There are two main types:

1. Inductive Reasoning: Deriving generalized conclusions from specific observations. It involves some uncertainty (e.g., observing a flame burns cotton and concluding it always will, potentially missing factors like moisture).
2. Deductive Reasoning: Also called "top-down logic," it uses general statements to reach specific, true conclusions.    - Example (Octet Rule): 1. Elements obeying the octet rule exist freely. 2. Chlorine ($Cl_2$) follows the rule. 3. Therefore, $Cl_2$ has independent existence.    - Example (Chemical Reactions): 1. Balanced equations follow the law of conservation of mass. 2. $2H_2 + O_2 \rightarrow 2H_2O$ is balanced. 3. Therefore, it obeys the law of conservation of mass.

Atomic Structure

Sub-Atomic Particles

Fundamental particles include electrons, protons, and neutrons.

• Electron: Negatively charged ($-1.6022 \times 10^{-19}\,C$), relative charge $-1$. Mass is $9.1095 \times 10^{-31}\,kg$. Discovered by J.J. Thomson.
• Proton: Positively charged ($+1.6022 \times 10^{-19}\,C$), relative charge $+1$. Mass is $1.6727 \times 10^{-27}\,kg$ (1836 times heavier than an electron).
• Neutron: No charge. Mass is $1.6750 \times 10^{-27}\,kg$. In an electric field, electrons curve toward the positive pole, protons toward the negative pole (with less curvature due to larger mass), and neutrons continue un-deflected.

Atomic and Mass Numbers

• Atomic Number ($Z$): Number of protons in the nucleus. In a neutral atom, this equals the number of electrons.
• Mass Number ($A$): Total number of protons ($P$) and neutrons ($N$). $A = P + N$.
• Ion Calculations:   - Electrons in a cation = atomic number $-$ magnitude of charge.   - Electrons in an anion = atomic number $+$ magnitude of charge.

Atomic and Ionic Radius

The radius is the average distance from the nucleus to the outermost electrons.

• Trends: In a period, radius decreases left to right as nuclear charge increases (shielding remains constant). In a group, radius increases top to bottom due to increasing shells and shielding.
• Ionic Sizes: Cations are always smaller than their parent atoms (increased effective nuclear charge, decreased repulsion). Anions are always bigger than their parent atoms (increased repulsion, constant nuclear charge).

Quantum Numbers

Four numerals specify an electron's location and energy:

1. Principal Quantum Number ($n$): Represents the shell ($n = 1, 2, 3...$). Corresponds to Bohr's shells (K, L, M, N) and the period number.
2. Azimuthal Quantum Number ($l$): Describes orbital shapes and subshells. Values range from $0$ to $n-1$. ($l=0$ is $s$, $l=1$ is $p$, $l=2$ is $d$, $l=3$ is $f$).
3. Magnetic Quantum Number ($m$): Explains splitting in a magnetic field (orientation). Values range from $-l$ to $+l$. ($s$ has 1 orientation; $p$ has 3: $p_x, p_y, p_z$; $d$ has 5; $f$ has 7).
4. Spin Quantum Number ($s$): Describes electron self-rotation. Values are $+1/2$ ($\uparrow$) or $-1/2$ ($\downarrow$). Electrons in the same orbital must have opposite spins to cancel magnetic fields.

Rules for Electronic Configuration

1. Aufbau Principle: Electrons fill orbitals in order of increasing energy ($1s < 2s < 2p < 3s < 3p < 4s < 3d...$).
2. $n+l$ Rule: Higher $n+l$ value means higher energy. If $n+l$ is the same, higher $n$ means higher energy.
3. Pauli Exclusion Principle: No two electrons in an atom can have the same four quantum numbers; at least spin must differ.
4. Hund's Rule: In degenerate orbitals, electrons reside separately with the same spin before pairing up. Half-filled and completely filled subshells are more stable (e.g., $Cr$ and $Cu$).

Ionization Energy (I.E.)

The minimum energy to remove the outermost electron from a gaseous isolated atom ($Na(g) \rightarrow Na^+(g) + 1e^-$, $I.E = 496\,kJ\,mol^{-1}$).

• Factors: Increases with nuclear charge and decreasing atomic radii (left to right in periods). Decreases with increasing shielding and size (top to bottom in groups).
• Successive I.E.: Large gaps indicate removal from a lower shell. For $Mg$, the jump from 2nd to 3rd I.E. ($1445$ to $7730\,kJ\,mol^{-1}$) shows the 3rd electron is a core electron.

Mass Spectrometry

Analytical technique measuring the mass-to-charge ratio ($m/z$) of ions. Steps include ionization, ion separation by mass analyzer, and detection. It is used to determine isotopic abundance.

• Average Atomic Mass Formula:$\text{Average Atomic Mass} = \frac{\sum (\text{Isotopic Mass} \times \text{Relative Abundance})}{100}$Example: Chlorine has 75.76% $^{35}Cl$ and 24.24% $^{37}Cl$.   $\text{Avg Mass} = \frac{(35 \times 75.76) + (37 \times 24.24)}{100} = 35.48\,amu$

Modern Electronic Materials (Semiconductors)

Semiconductors like Silicon ($Si$, $Z=14$) have an electronic configuration of $1s^2 2s^2 2p^6 3s^2 3p^2$.

• Band Structure: Electrons in the valence band are bound; in the conduction band, they move freely. The band gap allows excitation.
• Doping: Intentionally adding impurities.   - n-type: Doping with Phosphorus (Group 15) adds extra electrons.   - p-type: Doping with Boron (Group 13) creates positive "holes." This strategically controls current flow in modern technology.

Chemical Bonding

Electronegativity and Dipole Moment

Electronegativity is the power of a covalently bonded atom to attract shared electrons. It depends on atomic size, nuclear charge, and screening. On the Pauling scale, values increase left to right in periods and decrease top to bottom in groups.

• Bond Nature: If $\Delta EN > 1.8$, the bond is ionic. If $\Delta EN$ is between $0.4$ and $1.8$, it is polar covalent. If $\Delta EN < 0.4$, it is non-polar covalent.
• Dipole Moment ($\mu$): $\mu = q \times r$ (charge magnitude $\times$ separation distance). Unit is Debye ($1\,D = 3.335 \times 10^{-30}\,Cm$). Symmetrical molecules like $CO_2$ have polar bonds but a net dipole of zero due to cancellation.

Bond Energy and Bond Length

Bond energy ($kJ\,mol^{-1}$) is the energy required to break one mole of a specific bond. It increases with electronegativity difference and shorter bond lengths. Halogen bond energy decreases from $Cl_2$ to $I_2$, though $F_2$ is lower than $Cl_2$ due to repulsion of lone pairs in its small atoms.

VSEPR Theory (Valence Shell Electron Pair Repulsion)

Predicts molecular shapes based on electron pair repulsion. Order of repulsion: $LP-LP > LP-BP > BP-BP$.

• Shapes Examples:   - $AX_2$ ($BeF_2$): 2 bonding pairs (BP), 0 lone pairs (LP) = Linear ($180^{\circ}$).   - $AX_3$ ($BCl_3$): 3 BP, 0 LP = Trigonal Planar ($120^{\circ}$).   - $AX_2E$ ($SO_2$): 2 BP, 1 LP = Bent ($<120^{\circ}$).   - $AX_4$ ($CH_4$): 4 BP, 0 LP = Tetrahedral ($109.5^{\circ}$).   - $AX_3E$ ($NH_3$): 3 BP, 1 LP = Trigonal Pyramidal ($107^{\circ}$).   - $AX_2E_2$ ($H_2O$): 2 BP, 2 LP = Bent/V-Shaped ($104.5^{\circ}$).   - $AX_5$ ($PCl_5$): 5 BP, 0 LP = Trigonal Bipyramidal.   - $AX_6$ ($SF_6$): 6 BP, 0 LP = Octahedral ($90^{\circ}$, Expanded Octet). This theory is crucial in drug design (e.g., Cisplatin is square planar) to predict how molecules interact with enzyme active sites.

Valence Bond Theory (VBT) and Hybridization

VBT states covalent bonds form by orbital overlap.

• Bond Types:   - Sigma ($\sigma$): Head-on overlap (stronger).   - Pi ($\pi$): Parallel overlap (weaker, more diffuse).
• Hybridization: Mixing atomic orbitals to form new degenerate hybrid orbitals.   - sp: Mix 1 s and 1 p. Linear ($180^{\circ}$), e.g., $BeCl_2$, $C_2H_2$.   - $sp^2$: Mix 1 s and 2 p. Trigonal planar ($120^{\circ}$), e.g., $BCl_3$, $C_2H_4$.   - $sp^3$: Mix 1 s and 3 p. Tetrahedral ($109.5^{\circ}$), e.g., $CH_4$, $NH_3$ ($107^{\circ}$ due to LP), $H_2O$ ($104.5^{\circ}$ due to 2 LP).

Molecular Orbital Theory (MOT)

Proposed by Hund and Mullikan. Atomic orbitals combine linearly to form Bonding MOs (BMO - lower energy) and Antibonding MOs (ABMO - higher energy).

• Bond Order:$\text{Bond Order} = \frac{\text{Electrons in BMO} - \text{Electrons in ABMO}}{2}$
• Paramagnetism: $O_2$ is paramagnetic because it has unpaired electrons in $\pi^* 2p_y$ and $\pi^* 2p_z$ orbitals (a fact not explained by VBT). $N_2$ and $F_2$ are diamagnetic.
• Bond Orders: $H_2 = 1$, $He_2 = 0$ (does not exist), $N_2 = 3$, $O_2 = 2$, $F_2 = 1$.

Intermolecular Forces

Present between molecules (intramolecular forces like ionic/covalent bonds are much stronger).

1. Permanent Dipole-Dipole: Between polar molecules (e.g., $HCl$).
2. Instantaneous Dipole-Induced Dipole (London Forces): In non-polar molecules; increases with size and mass.
3. Hydrogen Bonding: Special case where $H$ is bonded to $F, N,$ or $O$. Explains why water has a high B.P. ($100^{\circ}C$) compared to $H_2S$, and why ice is less dense than water (open rigid hexagonal lattice).

Stoichiometry

The Mole and Molar Volume

A mole is $6.023 \times 10^{23}$ particles (Avogadro's number).

• Molar Volume ($V_m$): At STP, one mole of any gas occupies $22.414\,dm^3$.
• Gas Density: Determinable from molar mass and $V_m$.

Stoichiometric Calculations

Use balanced equations to determine mole ratios.

• Example: $2H_2(g) + O_2(g) \rightarrow 2H_2O(g)$. This means 2 moles of $H_2$ react with 1 mole of $O_2$ to produce 2 moles of $H_2O$.
• Solution Stoichiometry (Molarity):$\text{Molarity (M)} = \frac{\text{moles of solute}}{dm^3 \text{ of solution}}$   Formula for mixing solutions: $\frac{M_1 V_1}{n_1} = \frac{M_2 V_2}{n_2}$

Limiting Reagent and Yields

• Limiting Reactant: The reactant completely consumed first, limiting the product amount.
• Theoretical Yield: Maximum product calculated from the equation.
• Actual Yield: Amount obtained experimentally (usually less due to side reactions/incomplete steps).
• Percent Yield:$\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100$

Importance in Medicine

Stoichiometry ensures active ingredients in drugs (e.g., antibiotics, insulin) are accurate. Dosage depends on precise chemical balance to ensure safety and effectiveness.

States and Phases of Matter

Liquid State

Postulates: Molecules are in contact, in constant limited motion, and have stronger attractions than gases.

• Properties: Liquids are nearly incompressible, show slow diffusion, and expand when heated.
• Evaporation: Spontaneous change at the surface; causes cooling as high-energy molecules leave.
• Vapour Pressure: Pressure of vapours in equilibrium with liquid. Depends on temperature and I.M. forces (Stronger forces = lower V.P.).
• Boiling Point: Temperature where V.P. equals external pressure. At sea level, water boils at $100^{\circ}C$. In a pressure cooker (high pressure), boiling point rises; in vacuum distillation (low pressure), it drops.
• Viscosity: Resistance to flow. Measured in Pascal seconds ($Pa.s$) or Poise ($1\,Pa.s = 10\,Poise$).
• Surface Tension: Unbalanced inward pull of surface molecules. Smaller drops are spherical to minimize area.

Phase Changes and Energy

• Molar Heat of Fusion ($\Delta H_f$): Heat to melt 1 mole of solid (Ice = $+6.02\,kJ\,mol^{-1}$).
• Molar Heat of Vaporization ($\Delta H_v$): Heat to vaporize 1 mole of liquid (Water = $+40.7\,kJ\,mol^{-1}$). $\Delta H_v$ is much higher than $\Delta H_f$ because evaporation requires totally overcoming I.M. forces.

Liquid Crystals

In 1888, Frederick Reimitaz discovered substances like cholesteryl benzoate that have a turbid liquid phase between solid and clear liquid. They possess degrees of order found in crystals but the fluidity of liquids. Uses: Watch displays, laptops, medical skin thermography (detecting tumors), and industrial firing resistant coatings.

Solids and Glaciers

Solids have definite shape/volume, only vibrational motion, and strong cohesive forces.

• Types: Crystalline (ordered 3D pattern, sharp M.P.) and Amorphous (random arrangement, no sharp M.P., e.g., glass, plastic).
• Glaciers: Heat of fusion is critical for studying ice sheets and polar caps. Melting is directly proportional to absorbed heat; ice core analysis informs scientists about past climate conditions.

Energetics (Thermodynamics)

Basic Concepts

• Exothermic: Heat released ($\Delta H$ is negative).
• Endothermic: Heat absorbed ($\Delta H$ is positive).
• Standard Conditions: $1\,atm$ pressure, $298\,K$ (or $25^{\circ}C$), and $1\,M$ concentrations.
• Specific Heat Capacity ($c$): Energy to raise $1\,g$ of substance by $1\,K$. (Water = $4.18\,J\,g^{-1}\,K^{-1}$).
• Equation: $q = m \times c \times \Delta T$.

Enthalpy Definitions

• Enthalpy of Formation ($\Delta H_f^{\circ}$): Heat change when 1 mole of compound forms from elements.
• Enthalpy of Combustion ($\Delta H_c^{\circ}$): Heat when 1 mole is burnt in excess oxygen.
• Enthalpy of Atomization ($\Delta H_a^{\circ}$): Energy to form 1 mole of gaseous atoms ($\frac{1}{2}Cl_{2(g)} \rightarrow Cl_{(g)}$).
• Enthalpy of Neutralization ($\Delta H_n^{\circ}$): Heat evolved when 1 mole of $H^+$ reacts with 1 mole of $OH^-$. For strong acids/bases, this is roughly $-57.4\,kJ\,mol^{-1}$.

Hess's Law and Born-Haber Cycle

• Hess's Law: Enthalpy change is independent of the path taken. $\sum \Delta H (Cycle) = 0$.
• Born-Haber Cycle: Application of Hess's Law to ionic compounds. Steps for $NaCl$: 1. Sublimation of $Na$ ($+108$), 2. Ionization of $Na$ ($+496$), 3. Atomization of $Cl_2$ ($+121$), 4. Electron Affinity of $Cl$ ($-349$), 5. Lattice Energy ($\Delta H_L$).   $Lattice Energy = -787\,kJ\,mol^{-1}$.
• Lattice Energy Factors: Increases with ionic charge; decreases with ionic radius.

Entropy and Gibbs Free Energy

• Entropy ($S$): Measure of disorder. Gases have higher entropy than solids. $\Delta S = \sum S_{products} - \sum S_{reactants}$.
• Gibbs Free Energy ($G$): $\Delta G = \Delta H - T\Delta S$.   - If $\Delta G < 0$, reaction is spontaneous.   - If $\Delta G > 0$, reaction is non-spontaneous.   - If $\Delta G = 0$, the system is in equilibrium.

Chemical Kinetics

Reaction Rates and Rate Law

Kinetics is the study of reaction rates and mechanisms.

• Rate: Change in concentration per unit time ($mol\,dm^{-3}\,s^{-1}$).
• Rate Law: Energy expression relating rate to concentrations: $\text{Rate} = k[A]^x[B]^y$.
• Order of Reaction: Sum of exponents in the rate law. Determined experimentally, not from balanced equations.   - 0-order: Rate independent of concentration.   - 1-order: Rate proportional to 1st power of reactant.   - Pseudo-first order: Solvent in huge excess (e.g., hydrolysis).

Temperature and Activation Energy ($E_a$)

Rates generally double/triple for every $10\,K$ rise.

• Collision Theory: Particles must collide with sufficient energy ($E_a$) and correct orientation.
• Maxwell-Boltzmann Distribution: Shows the fraction of molecules with enough energy to react. Higher temperature shifts the curve, increasing the fraction with $E \geq E_a$.
• Catalyst: Increases rate by providing an alternative mechanism with lower $E_a$. Does not affect $\Delta H$.

Mechanism and Rate Determining Step

A mechanism is a sequence of elementary steps. The slowest step is the Rate-Determining Step (RDS). For $NO_2 + CO \rightarrow NO + CO_2$, the experimental rate is $k[NO_2]^2$, meaning the RDS involves two $NO_2$ molecules, and $CO$ is not involved in the RDS.

Chemical Equilibrium

Dynamic Equilibrium

Reversible reactions reach a state where forward and reverse rates are equal. Macroscopically, concentrations are constant; microscopically, activity continues.

• Law of Mass Action: Equilibrium constant ($K_c$) is the ratio of product concentration to reactant concentration products.   $aA + bB \rightleftharpoons cC + dD \implies K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$
• Relationships: $K_p = K_c(RT)^{\Delta n}$, where $\Delta n = \text{moles of products} - \text{moles of reactants}$.

Le Chatelier's Principle

If a system is disturbed, it shifts to oppose the change.

• Concentration: Adding reactant shifts equilibrium right.
• Pressure: Increasing pressure shifts to the side with fewer gas moles.
• Temperature: Increasing temperature shifts toward the endothermic direction. Only temperature changes the value of $K_c$.
• Catalyst: No effect on equilibrium position, only achieves it faster.

Industrial Applications

• Haber's Process: $N_{2(g)} + 3H_{2(g)} \rightleftharpoons 2NH_{3(g)}$ ($\Delta H$ negative). Optimum: $400^{\circ}C$, $200-300\,atm$, Iron catalyst, continual removal of $NH_3$.
• Contact Process: $2SO_{2(g)} + O_{2(g)} \rightleftharpoons 2SO_{3(g)}$ ($\Delta H$ negative). Optimum: $450-500^{\circ}C$, $2\,atm$, $V_2O_5$ catalyst.

Acid-Base Chemistry

Concepts and definitions

• Conjugate Acid-Base Pairs: An acid forms a conjugate base after donating a proton; a base forms a conjugate acid after accepting one.
• Lewis Concept: Acid = electron pair acceptor; Base = electron pair donor.
• pH and $K_w$: $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$ at $25^{\circ}C$. $pH = -\log[H^+]$. $pH + pOH = 14$.
• Strong vs. Weak: Strong acids ($HCl, H_2SO_4$) ionize fully ($large\,K_a$); weak acids ($CH_3COOH$) ionize partially ($small\,K_a$, $high\,pK_a$).

Buffer Solutions and Hydrolysis

• Buffer: Resists pH change when small amounts of acid/base are added.   - Acidic buffer: Weak acid + its salt with strong base ($CH_3COOH + CH_3COONa$).   - Basic buffer: Weak base + its salt with strong acid ($NH_4OH + NH_4Cl$).   - Henderson-Hasselbalch: $pH = pK_a + \log\frac{[salt]}{[acid]}$.
• Salt Hydrolysis: Cations from weak bases create acidic solutions ($NH_4^+$); anions from weak acids create basic solutions ($CH_3COO^-$). Salts from strong acids/bases (e.g., $NaCl$) do not hydrolyze.

Solubility Product ($K_{sp}$)

For sparingly soluble salt $A_mB_n$: $K_{sp} = [A^{n+}]^m[B^{m-}]^n$.

• Precipitation: Occurs if Ion Product ($Q$) $> K_{sp}$.
• Common Ion Effect: Solubility of an electrolyte decreases upon addition of a highly soluble salt containing a common ion (e.g., adding $HCl$ to saturated $NaCl$ causes $NaCl$ to precipitate).

Environmental Chemistry

Atmosphere and Air Pollution

Four layers: Troposphere ($12\,km$, weather), Stratosphere (contains Ozone layer), Mesosphere (meteors burn), Thermosphere (ionosphere/exosphere).

• Pollutants:   - $SO_2, NO_x$: from fossil fuels, lead to acid rain ($pH < 5.6$).   - Volatile Organic Compounds (VOCs): contribute to smog.   - CFCs: deplete the ozone layer.   - Particulate Matter (PM): causes respiratory disease.
• Photochemical Smog: Forms from reactions between $NO_x$, VOCs, and sunlight to create ground-level ozone and PAN (Peroxyacetyl nitrate), an eye irritant.
• Catalytic Converter: Reduces vehicle emissions by converting $CO$ and $NO$ to $CO_2$ and $N_2$.
• Nitrification: Biological oxidation of $NH_3 \rightarrow NO_2^- \rightarrow NO_3^-$. Denitrification: Reduction of $NO_3^- \rightarrow N_2$.

Water Pollution and Treatment

• Sources: Agriculture (pesticides, high-nutrient runoff causing Eutrophication/algal bloom), Industrial effluents (heavy metals like $Pb, Hg, Cd$), Oil spills.
• Raw Water Treatment: Screening, Coagulation/Flocculation (using alum), Sedimentation, Filtration, Disinfection (Chlorine, Ozone, UV).
• Conservation: Smart irrigation, water harvesting, and legislation like Pakistan's Water Act 1998.

Organic Chemistry

Fundamentals

Carbon's ability for Catenation (self-linking) creates millions of compounds.

• Formulae Types: Molecular (actual atoms), Empirical (simplest ratio), Condensed structural, Displayed (2D showing all bonds), Skeletal (lines representing carbons/H assumed).
• Homologous Series: Group with same functional group, differing by $-CH_2-$, similar chemical properties.
• Reagents: Electrophile ($+$ charge or electron deficient), Nucleophile ($-$ charge or lone pair), Free Radical (unpaired electron).
• Fission: Homolytic (equal splitting, creates radicals), Heterolytic (unequal, creates ions).

Isomerism

• Structural Isomerism: Same molecular formula, different bond arrangement.   - Chain: Different chain length.   - Positional: Functional group at different location.   - Functional Group: Different functional group (alcohol vs ether).   - Metamerism: Different alkyl distribution around group.   - Tautomerism: Proton migration (keto-enol).
• Stereoisomerism: Same bonds, different 3D arrangement.   - Geometric (cis-trans): Restricted rotation around double bond or ring.   - Optical: Non-superimposable mirror images (enantiomers) due to a Chiral/Asymmetric carbon (bonded to 4 different groups).

Reactions and Synthesis

• Alkanes: Undergo Free Radical Substitution.
• Alkenes: Undergo Electrophilic Addition (Hydrogenation, Hydration, Halogenation, Ozonolysis). Follows Markovnikov's Rule (H attaches to carbon with more H's, mediated by carbocation stability: $3^{\circ} > 2^{\circ} > 1^{\circ}$).
• Alcohols: Produced by hydration of alkenes or reduction of carbonyls. Oxidize to aldehydes/ketones/acids. Respond to Iodoform reaction if they have $CH_3CH(OH)-$ group.
• Carbonyls: Aldehydes (reduce to 1-alcohols, oxidize to acids, give positive Tollen's test); Ketones (reduce to 2-alcohols, resist weak oxidation). Detected by 2,4-DNPH (orange ppt).
• Amines: Ammonia derivatives. Secondary amines are more basic than primary due to inductive effect.
• Azo Compounds: Formed from diazonium salts; stable aromatic ones used as dyes.
• Retrosynthesis: Working backward from target molecule to find simple precursors. Add carbon atoms using $CN^-$ nitriles.

Energy Sources

Petroleum Refining

Refining separates crude oil into fractions using Fractional Distillation based on boiling points.

• Fractions (Low to High B.P.): LPG ($C_1-C_4$), Petrol ($C_5-C_{10}$, fuel for cars), Naphtha (chemical feedstock), Kerosene (jet fuel), Diesel, Lubricating oils, Fuel oil, Bitumen (roads).
• Petrochemical Technology:   - Cracking: High M.W. hydrocarbons $\rightarrow$ smaller ones. Thermal ($450-750^{\circ}C$) or Catalytic (using zeolites).   - Steam Reforming: Hydrocarbons + steam $\rightarrow H_2 + CO$.

Nuclear and Solar Energy

• Nuclear Fission: Splitting heavy nucleus ($^{235}U$) by neutron absorption.
• Nuclear Fusion: Combining light nuclei ($H$ isotopes) to form Helium; releases massive energy, no greenhouse gases, no long-lived waste.
• Solar Energy:   - Solar Thermal: Captures heat to fluid.   - Photovoltaic (PV) Cells: Silicon cells convert photons to electricity (DC), then inverted to AC. Sustainable, no emissions, low maintenance.
• Carbon Footprint: Total greenhouse gas weight emitted by activity ($Tons\,of\,CO_2$ equivalent). Reduction involves sustainable energy, efficiency, and reforestation.