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Atomic Structure and Subatomic Particles

• Subatomic Particles:

  • Atoms are composed of three fundamental subatomic particles: protons, neutrons, and electrons.

  • Proton:

  • Symbol: $p^+$ or $p$

  • Relative Charge: $+1$

  • Relative Mass: $1$

  • Neutron:

  • Symbol: $n^0$ or $n$

  • Relative Charge: $0$

  • Relative Mass: $1$

  • Electron:

  • Symbol: $e^-$ or $e$

  • Relative Charge: $-1$

  • Relative Mass: $\frac{1}{1836}$

• Atomic Terminology and Notation:

  • Atomic Number ($Z$): Represents the number of protons present in the nucleus of an atom. Defines the identity of the chemical element.

  • Mass Number / Nucleon Number ($A$): Represents the total number of protons and neutrons in the nucleus ($A = Z + N$).

  • Nuclear Charge: Determined by the number of protons in the nucleus ($+Z$).

  • Isotopes: Atoms of the same element having the same atomic number ($Z$) but different mass numbers ($A$) due to a differing number of neutrons.

  • Ions: Charged species formed when neutral atoms lose or gain electrons.

  • Cations: Positively charged ions formed by electron loss.

  • Anions: Negatively charged ions formed by electron gain.

• Behavior of Charged Particles in Electric Fields:

  • When beams of subatomic particles pass through an electric field between charged plates:

  • Protons are deflected towards the negative plate due to their positive charge.

  • Electrons are deflected towards the positive plate due to their negative charge. Electrons undergo a much greater angle of deflection than protons because of their significantly smaller mass.

  • Neutrons pass straight through un-deflected due to their neutral charge ($0$).

• Relative Atomic, Isotopic, and Molecular Masses:

  • Relative Isotopic Mass: The mass of an atom of an isotope relative to $\frac{1}{12}\text{th}$ of the mass of an atom of carbon-12 ($^{12}\text{C}$).

  • Relative Atomic Mass ($A_r$): The weighted average mass of an atom of an element relative to $\frac{1}{12}\text{th}$ of the mass of an atom of carbon-12 ($^{12}\text{C}$).

  • Relative Molecular Mass ($M_r$): The sum of the relative atomic masses of all atoms present in a molecular formula relative to $\frac{1}{12}\text{th}$ of the mass of carbon-12 ($^{12}\text{C}$).

• Electron Energy Levels and Subshells:

  • Principal Quantum Number ($n$): Defines the main electron shell or principal energy level ($n = 1, 2, 3, 4, \dots$).

  • Principal energy levels are subdivided into subshells: $s$, $p$, $d$, and $f$.

  • Orbital Capacities and Shapes:

  • An atomic orbital is a region of space around the nucleus where there is a high probability of finding up to two electrons with opposite spins.

  • $s$ Subshell: Contains $1$ spherical orbital; holds a maximum of $2$ electrons.

  • $p$ Subshell: Contains $3$ dumbbell-shaped orbitals ($p_x, p_y, p_z$); holds a maximum of $6$ electrons.

  • $d$ Subshell: Contains $5$ orbitals; holds a maximum of $10$ electrons.

  • $f$ Subshell: Contains $7$ orbitals; holds a maximum of $14$ electrons.

  • Maximum electron capacities per principal shell ($n$):

  • $n = 1$: $2$ electrons ($1s^2$)

  • $n = 2$: $8$ electrons ($2s^2 2p^6$)

  • $n = 3$: $18$ electrons ($3s^2 3p^6 3d^{10}$)

  • $n = 4$: $32$ electrons ($4s^2 4p^6 4d^{10} 4f^{14}$)

• Rules for Electronic Configuration:

  • Aufbau Principle: Electrons occupy the orbitals of lowest energy level first ($1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d \rightarrow 4p$). Note that the $4s$ subshell fills before the $3d$ subshell because it is at a lower energy level in an un-ionized atom.

  • Hund's Rule: Orbitals of equal energy (degenerate orbitals) are each occupied by one electron before any orbital is doubly occupied. All single electrons in singly occupied orbitals must have the same spin.

  • Pauli Exclusion Principle: An orbital can hold a maximum of two electrons, and they must have opposite spins (spin-pair repulsion).

  • Ground state electron configurations of selected species:

  • Nitrogen ($\text{N}$, $Z=7$): $1s^2 2s^2 2p^3$

  • Chlorine ($\text{Cl}$, $Z=17$): $1s^2 2s^2 2p^6 3s^2 3p^5$

  • Iron ($\text{Fe}$, $Z=26$): $1s^2 2s^2 2p^6 3s^2 3p^6 3d^6 4s^2$

• Ionisation Energy ($IE$):

  • First Ionisation Energy ($1\text{st } IE$): The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous $1+$ ions.

  • Equation: $\text{X}(g) \rightarrow \text{X}^+(g) + e^-$

  • Second Ionisation Energy ($2\text{nd } IE$): The energy required to remove one mole of electrons from one mole of gaseous $1+$ ions to form one mole of gaseous $2+$ ions.

  • Equation: $\text{X}^+(g) \rightarrow \text{X}^{2+}(g) + e^-$

  • Third Ionisation Energy ($3\text{rd } IE$):

  • Equation: $\text{X}^{2+}(g) \rightarrow \text{X}^{3+}(g) + e^-$

  • Factors Influencing Ionisation Energy:

  • Nuclear Charge: Greater number of protons increases the attraction between the nucleus and the outer electrons, increasing $IE$.

  • Atomic Radius: Greater distance between the outer electron and the nucleus weakens electrostatic attraction, decreasing $IE$.

  • Shielding Effect: Inner shell electrons shield outer electrons from the attraction of the positive nuclear charge, decreasing $IE$.

  • Spin-Pair Repulsion: Paired electrons in the same orbital repel each other, making them slightly easier to remove than unpaired electrons in degenerate orbitals.

  • Successive ionisation energies increase progressively as electrons are removed from increasingly positive ions. Large jumps/spikes in successive $IE$ values indicate a change in the principal quantum shell ($n$), allowing determination of an element's group position in the Periodic Table.

Chemical Stoichiometry and Reacting Quantities

• The Mole Concept and Avogadro's Constant:

  • A mole is the amount of substance containing the same number of elementary entities as there are atoms in exactly $12\,g$ of carbon-12.

  • Avogadro's Constant ($N_A$): $6.02 \times 10^{23}\,mol^{-1}$

  • Number of particles: $\text{Number of particles} = n \times N_A$

• Calculations Involving Moles, Mass, and Volume:

  • Mass Relationship:

  • $n = \frac{m}{M_r}$ or $n = \frac{m}{A_r}$

  • Where $m$ is mass in grams ($g$), and $M_r$ / $A_r$ is molar mass in $g\,mol^{-1}$.

  • Solution Concentration:

  • $c = \frac{n}{V}$

  • Where $c$ is concentration in $mol\,dm^{-3}$, $n$ is moles ($mol$), and $V$ is volume in $dm^3$.

  • Gaseous Volumes:

  • At Standard Temperature and Pressure (STP / RTP): $1\,mol$ of any gas occupies $24.0\,dm^3$ (or $24000\,cm^3$).

  • $n = \frac{V}{24.0}$ (with $V$ in $dm^3$)

• Empirical and Molecular Formulae:

  • Empirical Formula: The simplest whole-number ratio of atoms of each element present in a compound.

  • Molecular Formula: The actual number of atoms of each element present in a single molecule of a compound.

  • Relation: $\text{Molecular Formula} = (\text{Empirical Formula})_n$, where $n = \frac{M_r}{\text{Empirical Mass}}$.

• Hydrated Compounds and Water of Crystallisation:

  • Hydrated salts contain water molecules incorporated into their crystalline lattice structure (water of crystallisation).

  • Example: Hydrated Copper(II) Sulfate: $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$.

  • Heating hydrated compounds yields the anhydrous salt:

  • $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}(s) \rightarrow \text{CuSO}_4(s) + 5\text{H}_2\text{O}(g)$

• Reacting Masses, Volumes, Limiting and Excess Reagents:

  • Limiting Reagent: The reactant that is completely consumed first in a chemical reaction, limiting the theoretical yield of products.

  • Excess Reagent: The reactant present in a quantity greater than required to react fully with the limiting reagent.

  • Percentage Yield: $\text{Percentage Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

  • Percentage Purity: $\text{Percentage Purity} = \frac{\text{Mass of Pure Substance}}{\text{Total Mass of Impure Sample}} \times 100\%$

Chemical Bonding and Structure

• Electronegativity ($EN$):

  • Definition: The ability of an atom in a covalent bond to attract the shared pair of electrons towards itself.

  • Factors Affecting Electronegativity:

  • Nuclear Charge: Higher nuclear charge increases $EN$.

  • Atomic Radius: Smaller atomic radius increases $EN$ due to electrons being closer to the nucleus.

  • Shielding: Less shielding increases $EN$.

  • Periodic Trends:

  • Increases across a Period (left to right) due to increasing nuclear charge and decreasing atomic radius with similar shielding.

  • Decreases down a Group (top to bottom) due to increasing atomic radius and increasing shielding.

  • Fluorine ($\text{F}$) is the most electronegative element, followed by Oxygen ($\text{O}$), Nitrogen ($\text{N}$), and Chlorine ($\text{Cl}$).

• Ionic Bonding:

  • Definition: The strong electrostatic force of attraction between oppositely charged ions in a lattice.

  • Formed by the transfer of electrons from a metal (forming cations) to a non-metal (forming anions).

  • Structure: Giant ionic lattice (e.g., Sodium Chloride, $\text{NaCl}$).

  • Properties of Ionic Compounds:

  • High melting and boiling points due to strong electrostatic attraction requiring high thermal energy to overcome.

  • Hard and brittle; applied mechanical force shifts layers of ions causing like charges to align and repel.

  • Electrical conductivity: Non-conductors in solid state (ions fixed in lattice); good conductors when molten ($l$) or in aqueous solution ($aq$) as ions are free to move.

• Metallic Bonding:

  • Definition: The electrostatic attraction between a lattice of positive metal ions (cations) and a delocalised sea of electrons.

  • Structure: Giant metallic lattice.

  • Properties: High electrical and thermal conductivity (delocalised electrons are mobile charge carriers), malleable, ductile.

• Covalent and Coordinate (Dative Covalent) Bonding:

  • Covalent Bond: Electrostatic attraction between two positive nuclei and a shared pair of electrons.

  • Coordinate / Dative Covalent Bond: A covalent bond in which both shared electrons are donated by a single atom.

  • Formed by an atom possessing a lone pair of electrons donating into an empty orbital of an electron-deficient species.

  • Examples:

    • Ammonium ion ($\text{NH}_4^+$): Nitrogen donates a lone pair to $\text{H}^+$: $\text{NH}_3 + \text{H}^+ \rightarrow \text{NH}_4^+$

    • Hydronium ion ($\text{H}_3\text{O}^+$): Oxygen donates a lone pair to $\text{H}^+$: $\text{H}_2\text{O} + \text{H}^+ \rightarrow \text{H}_3\text{O}^+$

    • Dative dimer $\text{Al}_2\text{Cl}_6$: Formed by coordinate bonding between $\text{Cl}$ lone pairs and electron-deficient $\text{Al}$ atoms.

  • Expanded Octet: Elements in Period 3 and beyond can expand their valence shell to accommodate more than 8 electrons using empty $d$ orbitals (e.g., $\text{PCl}_5$, $\text{SF}_6$).

• Orbital Hybridisation:

  • The mixing of atomic orbitals ($s$ and $p$) to form degenerate hybrid orbitals suitable for bonding.

  • $sp^3$ Hybridisation: Mixing of one $s$ and three $p$ orbitals forming four equivalent $sp^3$ hybrid orbitals directed tetrahedrally (bond angle $109.5^\circ$). Example: Methane ($\text{CH}_4$).

  • $sp^2$ Hybridisation: Mixing of one $s$ and two $p$ orbitals forming three equivalent $sp^2$ hybrid orbitals in a trigonal planar arrangement (bond angle $120^\circ$), leaving one unhybridised $p$ orbital. Example: Ethene ($\text{C}_2\text{H}_4$).

  • $sp$ Hybridisation: Mixing of one $s$ and one $p$ orbital forming two linear $sp$ hybrid orbitals (bond angle $180^\circ$), leaving two unhybridised $p$ orbitals. Example: Ethyne ($\text{C}_2\text{H}_2$).

• Sigma ($\sigma$) and Pi ($\pi$) Bonds:

  • Sigma ($\sigma$) Bond: Direct end-on (head-on) overlap of orbitals along the internuclear axis. Possesses electron density concentrated between the two nuclei. Single covalent bonds consist of one $\sigma$ bond.

  • Pi ($\pi$) Bond: Sideways overlap of unhybridised parallel $p$ orbitals above and below the internuclear axis.

  • A double bond consists of $1\,\sigma$ bond and $1\,\pi$ bond.

  • A triple bond consists of $1\,\sigma$ bond and $2\,\pi$ bonds.

• Valence Shell Electron Pair Repulsion (VSEPR) Theory and Molecular Shapes:

  • Electron pairs in the valence shell of a central atom repel each other and position themselves as far apart as possible to minimize repulsion.

  • Order of repulsion strength: $\text{Lone Pair - Lone Pair} > \text{Lone Pair - Bonding Pair} > \text{Bonding Pair - Bonding Pair}$.

  • Shapes and Angles:

  • 2 Bonding Pairs, 0 Lone Pairs: Linear, bond angle $180^\circ$ (e.g., $\text{BeCl}_2$, $\text{CO}_2$).

  • 3 Bonding Pairs, 0 Lone Pairs: Trigonal Planar, bond angle $120^\circ$ (e.g., $\text{BF}_3$).

  • 4 Bonding Pairs, 0 Lone Pairs: Tetrahedral, bond angle $109.5^\circ$ (e.g., $\text{CH}_4$, $\text{NH}_4^+$).

  • 3 Bonding Pairs, 1 Lone Pair: Trigonal Pyramidal, bond angle $107^\circ$ (e.g., $\text{NH}_3$).

  • 2 Bonding Pairs, 2 Lone Pairs: Non-linear / Bent / V-shaped, bond angle $104.5^\circ$ (e.g., $\text{H}_2\text{O}$).

  • 6 Bonding Pairs, 0 Lone Pairs: Octahedral, bond angle $90^\circ$ (e.g., $\text{SF}_6$).

Intermolecular Forces and States of Matter

• Types of Intermolecular Forces (Van der Waals Forces):

  • Instantaneous Dipole - Induced Dipole (id-id) Forces / London Dispersion Forces:

  • Occur in all atoms and molecules due to continuous movement of electrons creating temporary fluctuating dipoles, which induce dipoles in neighboring species.

  • Strength increases with an increasing number of electrons and greater surface area/contact area of molecules.

  • Permanent Dipole - Permanent Dipole (pd-pd) Forces:

  • Occur between polar molecules possessing permanent dipoles due to differences in electronegativity.

  • Hydrogen Bonding:

  • Exceptionally strong form of pd-pd attraction occurring when hydrogen is bonded directly to a highly electronegative atom ($\text{N}$, $\text{O}$, or $\text{F}$) containing at least one lone pair of electrons.

  • Extends between the hydrogen atom of one molecule and the lone pair of $\text{N}$, $\text{O}$, or $\text{F}$ on an adjacent molecule.

  • Responsible for anomalous properties of water (ice is less dense than liquid water due to an open 3D hydrogen-bonded lattice; high boiling point of water).

• Comparison of Structures and Physical Properties:

  • Simple Molecular Crystals: Low melting/boiling points due to weak intermolecular forces (e.g., Iodine $\text{I}_2$, Ice $\text{H}_2\text{O}$, Buckminsterfullerene $\text{C}_{60}$). Electrical non-conductors.

  • Giant Covalent Structure (Macromolecular): High melting/boiling points due to strong covalent bonds extending throughout lattice.

  • Diamond: Tetrahedral $\text{C}$ atoms ($sp^3$), extremely hard, non-conductor.

  • Graphite: Hexagonal layers of $\text{C}$ atoms ($sp^2$) with delocalised electrons between layers; conducts electricity, soft/lubricating.

  • Silicon Dioxide ($\text{SiO}_2$): Giant covalent lattice, high melting point.

• The Gaseous State and Ideal Gases:

  • Ideal Gas Equation:

  • $PV = nRT$

  • Where $P$ is pressure in Pascals ($\text{Pa}$ or $\text{N\,m}^{-2}$), $V$ is volume in cubic meters ($m^3$), $n$ is moles ($mol$), $R$ is the ideal gas constant ($8.31\,J\,K^{-1}\,mol^{-1}$), and $T$ is absolute temperature in Kelvin ($K$).

  • Conversion: T(K) = T(^\circ\text{C}) + 273$.\n * **Deviations from Ideal Behavior (Real Gases vs Ideal Gases):**\n * Ideal gases assume negligible molecular volume and zero intermolecular forces.\n * Real gases deviate significantly from ideal behavior at **high pressure** (molecular volume becomes significant compared to total volume) and **low temperature** (intermolecular forces become significant as kinetic energy drops).\n\n# Chemical Energetics and Thermochemistry\n\n* **Enthalpy Changes (\Delta H):**\n * **Exothermic Reaction:** Energy is released to the surroundings; \Delta H$is negative ($\Delta H < 0). Bonds formed are stronger than bonds broken.\n * **Endothermic Reaction:** Energy is absorbed from the surroundings; \Delta H$is positive ($\Delta H > 0). Bonds broken are stronger than bonds formed.\n* **Standard Enthalpy Conditions:**\n * Temperature: 298\,K$($25^\circ\text{C})\n * Pressure: 101\,kPa$($1\times 10^5\,\text{Pa}$or$1\,\text{bar})\n * Solutions: Concentration of 1.0\,mol\,dm^{-3}\n* **Standard Enthalpy Definitions:**\n * **Standard Enthalpy Change of Reaction (\Delta H_r^\ominus):** The enthalpy change when amounts of reactants as shown in the stoichiometric equation react under standard conditions.\n * **Standard Enthalpy Change of Formation (\Delta H_f^\ominus$):** The enthalpy change when one mole of a compound is formed from its constituent elements in their standard states under standard conditions. (Note:$\Delta H_f^\ominus of an element in its standard state is zero).\n * **Standard Enthalpy Change of Combustion (\Delta H_c^\ominus):** The enthalpy change when one mole of a substance is completely burned in excess oxygen under standard conditions.\n * **Standard Enthalpy Change of Neutralisation (\Delta H_{neut}^\ominus):** The enthalpy change when one mole of water is formed by the reaction of an acid with an alkali under standard conditions.\n* **Calculation of Enthalpy Changes:**\n * **Heat Energy Transferred (q):**\n * q = m c \Delta T\n * Where q$is heat energy in Joules ($J$),$m$is mass of liquid in grams ($g$),$c$is specific heat capacity ($4.18\,J\,g^{-1}\,K^{-1}$for water), and$\Delta T$is temperature change ($K$or$^\circ\text{C}).\n * \Delta H = -\frac{q}{n}$(expressed in$kJ\,mol^{-1}).\n * **Enthalpy Change from Bond Energies:**\n * \Delta H_r = \sum (\text{Bond Energies of Bonds Broken}) - \sum (\text{Bond Energies of Bonds Formed})\n* **Hess's Law:**\n * The total enthalpy change of a chemical reaction is independent of the route taken, provided the initial and final conditions are identical.\n * **Using Enthalpies of Formation (\Delta H_f^\ominus):**\n * \Delta H_r^\ominus = \sum \Delta H_f^\ominus(\text{Products}) - \sum \Delta H_f^\ominus(\text{Reactants})\n * **Using Enthalpies of Combustion (\Delta H_c^\ominus):**\n * \Delta H_r^\ominus = \sum \Delta H_c^\ominus(\text{Reactants}) - \sum \Delta H_c^\ominus(\text{Products})\n\n# Reaction Kinetics and Collision Theory\n\n* **Rate of Reaction:**\n * Defined as the change in concentration of a reactant or product per unit time (mol\,dm^{-3}\,s^{-1}).\n * \text{Rate} = \frac{\Delta[\text{Products}]}{\Delta t} = -\frac{\Delta[\text{Reactants}]}{\Delta t}\n* **Collision Theory and Factors Affecting Reaction Rate:**\n * For a reaction to occur, reactant particles must collide with each other with correct orientation and with kinetic energy equal to or greater than the Activation Energy (E_a).\n * **Activation Energy (E_a):** The minimum kinetic energy required by colliding particles for a reaction to take place.\n * **Factors Influencing Rate:**\n * **Concentration / Pressure:** Increasing concentration (or pressure of gases) increases the number of particles per unit volume, increasing collision frequency, leading to a faster rate.\n * **Surface Area:** Increasing surface area of solid reactants exposes more surface particles to collisions, increasing collision frequency.\n * **Temperature:** Increasing temperature increases the average kinetic energy of particles. A small increase in temperature results in a large increase in the fraction of particles possessing energy \ge E_a, significantly increasing successful collision frequency and reaction rate.\n* **Maxwell-Boltzmann Distribution:**\n * Shows the distribution of molecular kinetic energies in a gas at a given temperature.\n * Curve starts at the origin (no molecules have zero energy), reaches a peak (most probable energy), and drops asymmetrically towards zero at high energy.\n * Effect of Temperature Increase: Curve flattens and shifts to the right; area under the curve beyond E_a increases substantially.\n* **Catalysts:**\n * A catalyst increases the rate of a chemical reaction by providing an alternative reaction pathway with a lower activation energy (E_a), without being consumed in the overall process.\n * **Homogeneous Catalyst:** Catalyst is in the same physical state as the reactants (e.g., \text{NO}2(g) in acid rain catalysis).\n * **Heterogeneous Catalyst:** Catalyst is in a different physical state from the reactants (e.g., solid Iron \text{Fe}(s) in the Haber process; solid Rhodium/Platinum in catalytic converters).\n\n# Chemical Equilibria and Industrial Applications\n\n* **Dynamic Equilibrium:**\n * Reversible reaction where the rate of the forward reaction equals the rate of the reverse reaction, and concentrations of reactants and products remain constant in a closed system.\n* **Equilibrium Constant (K_c$and$K_p):**\n * For a general reversible reaction: a\text{A} + b\text{B} \rightleftharpoons c\text{C} + d\text{D}\n * Concentration Constant (K_c):\n * K_c = \frac{[\text{C}]^c [\text{D}]^d}{[\text{A}]^a [\text{B}]^b}\n * Units depend on the stoichiometric powers in the expression.\n * Partial Pressure Constant (K_p):\n * K_p = \frac{(P\text{C})^c (P_\text{D})^d}{(P_\text{A})^a (P_\text{B})^b}\n * Partial pressure of gas \text{A}$:$P_\text{A} = \chi_\text{A} \times P_{total}$, where mole fraction$\chi_\text{A} = \frac{\text{moles of A}}{\text{total moles of gas}}.\n * Note: Temperature is the **only** factor that changes the numerical value of K_c$and$K_p.\n* **Le Chatelier's Principle:**\n * If a system at dynamic equilibrium is subjected to a change in conditions (concentration, pressure, temperature), the equilibrium position shifts in a direction that minimizes or opposes that change.\n * **Effect of Concentration:** Increasing reactant concentration shifts equilibrium to the right.\n * **Effect of Pressure:** Increasing pressure shifts equilibrium to the side with fewer moles of gas.\n * **Effect of Temperature:** Increasing temperature shifts equilibrium in the endothermic direction; decreasing temperature shifts it in the exothermic direction.\n * **Effect of Catalyst:** Does not alter position of equilibrium or value of K_c/K_p; increases rates of both forward and reverse reactions equally, reaching equilibrium faster.\n* **Industrial Processes:**\n * **The Haber Process (Synthesis of Ammonia):**\n * Equation: \text{N}2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) \quad \Delta H = -92\,kJ\,mol^{-1}\n * Conditions: Pressure 200\,\text{atm}$, Temperature$400 - 450^\circ\text{C}$, Catalyst: Finely divided Iron ($\text{Fe}).\n * **The Contact Process (Production of Sulfuric Acid):**\n * Key reversible step: 2\text{SO}_2(g) + \text{O}_2(g) \rightleftharpoons 2\text{SO}_3(g) \quad \Delta H = -197\,kJ\,mol^{-1}\n * Conditions: Pressure 1 - 2\,\text{atm}$, Temperature$450^\circ\text{C}$, Catalyst: Vanadium(V) Oxide ($\text{V}_2\text{O}_5).\n\n# Brønsted-Lowry Acids, Bases, and pH Chemistry\n\n* **Brønsted-Lowry Theory:**\n * **Acid:** Proton (\text{H}^+) donor.\n * **Base:** Proton (\text{H}^+) acceptor.\n * **Conjugate Acid-Base Pairs:** Consist of two species related by the gain or loss of a single proton (\text{H}^+).\n* **Strong vs Weak Acids and Bases:**\n * **Strong Acid:** Completely dissociates/ionises in aqueous solution (e.g., \text{HCl}$,$\text{HNO}_3$,$\text{H}_2\text{SO}_4).\n * Equation: \text{HCl}(aq) \rightarrow \text{H}^+(aq) + \text{Cl}^-(aq)\n * **Weak Acid:** Partially dissociates/ionises in aqueous solution establishing an equilibrium (e.g., \text{CH}_3\text{COOH}$,$\text{H}_2\text{CO}_3).\n * Equation: \text{CH}_3\text{COOH}(aq) \rightleftharpoons \text{CH}_3\text{COO}^-(aq) + \text{H}^+(aq)\n * **Strong Base:** Completely dissociates in aqueous solution (e.g., \text{NaOH}$,$\text{KOH}).\n * **Weak Base:** Partially reacts/ionises in aqueous solution (e.g., \text{NH}_3).\n * Equation: \text{NH}_3(aq) + \text{H}_2\text{O}(l) \rightleftharpoons \text{NH}_4^+(aq) + \text{OH}^-(aq)\n* **The pH Scale and Neutralisation:**\n * Definition of pH: \text{pH} = -\log{10}[\text{H}^+]\n * Inverse relation: [\text{H}^+] = 10^{-\text{pH}}\n * Neutralisation Ionic Equation: \text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}2\text{O}(l)\n* **Acid-Base Titrations and Indicators:**\n * **Methyl Orange:** Red in acid, Yellow in alkali (pH range 3.1 - 4.4$; changes color at end point around pH$3.7). Suitable for Strong Acid - Strong Base and Strong Acid - Weak Base titrations.\n * **Phenolphthalein:** Colorless in acid, Pink in alkali (pH range 8.3 - 10.0). Suitable for Strong Acid - Strong Base and Weak Acid - Strong Base titrations.\n * **Bromothymol Blue:** Yellow in acid, Green at neutral, Blue in base.\n\n# Redox Processes and Electrochemistry\n\n* **Redox Definitions:**\n * **Oxidation:** Loss of electrons / Increase in oxidation number / Gain of oxygen / Loss of hydrogen.\n * **Reduction:** Gain of electrons / Decrease in oxidation number / Loss of oxygen / Gain of hydrogen.\n * **Oxidising Agent:** Species that oxidises another species by accepting electrons (itself is reduced).\n * **Reducing Agent:** Species that reduces another species by donating electrons (itself is oxidised).\n* **Oxidation Numbers / Oxidation States:**\n * Rules for assigning oxidation numbers:\n 1. Uncombined elements have an oxidation number of 0$(e.g.,$\text{O}_2$,$\text{Fe}$,$\text{Cl}_2).\n 2. Monatomic ions have an oxidation number equal to their ionic charge (e.g., \text{Na}^+ = +1$,$\text{Cl}^- = -1).\n 3. Hydrogen in compounds is usually +1$(except in metal hydrides like$\text{NaH}$where it is$-1).\n 4. Oxygen in compounds is usually -2$(except in peroxides like$\text{H}_2\text{O}_2$where it is$-1$, and in$\text{F}_2\text{O}$where it is$+2).\n 5. Fluorine is always -1 in compounds.\n 6. Neutral compound oxidation numbers sum to 0; polyatomic ion oxidation numbers sum to the charge of the ion.\n * Roman Numerals in Naming: Used to specify the oxidation number of transition metals or non-metals (e.g., Manganese(IV) Oxide \text{MnO}_2$, Chlorate(I) ion$\text{ClO}^-).\n* **Disproportionation Reactions:**\n * A reaction in which a single species is simultaneously oxidised and reduced.\n * Example: Reaction of chlorine with cold dilute \text{NaOH}:\n * \text{Cl}_2(g) + 2\text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{NaClO}(aq) + \text{H}_2\text{O}(l)\n * Chlorine zero state (0$) changes to$-1$in$\text{NaCl}$(reduction) and$+1$in$\text{NaClO} (oxidation).\n* **Balancing Redox Equations using Half-Equations:**\n 1. Identify species oxidised and reduced.\n 2. Write separate half-equations for oxidation and reduction.\n 3. Balance atoms other than \text{H}$and$\text{O}.\n 4. Balance \text{O}$by adding$\text{H}_2\text{O}.\n 5. Balance \text{H}$by adding$\text{H}^+.\n 6. Balance charges by adding electrons (e^-).\n 7. Multiply half-equations by integers so electron count matches, then combine and cancel duplicate species.\n\n# Periodicity and Period 3 Elements\n\n* **Trends Across Period 3 (\text{Na} \rightarrow \text{Ar}):**\n * **Atomic Radius:** Decreases across Period 3. Nuclear charge increases while shielding remains relatively constant (electrons added to the same outer shell n=3), pulling valence electrons closer to the nucleus.\n * **Ionic Radius:** Cations (\text{Na}^+, \text{Mg}^{2+}, \text{Al}^{3+}$) are significantly smaller than their parent atoms because an entire electron shell is lost. Anions ($\text{P}^{3-}, \text{S}^{2-}, \text{Cl}^-$) are larger than cations because extra electrons increase electron-electron repulsion within shell$n=3.\n * **Electrical Conductivity:** Increases from \text{Na}$to$\text{Al}$due to an increasing number of delocalised valence electrons per cation.$\text{Si}$is a semi-conductor. Non-metals ($\text{P}_4, \text{S}_8, \text{Cl}_2, \text{Ar}) are electrical insulators.\n * **Melting Points:**\n * \text{Na}, \text{Mg}, \text{Al}$: Giant metallic structures; melting point increases from$\text{Na}$to$\text{Al} as metallic bond strength increases with smaller cation radius and higher charge density.\n * \text{Si}: Giant molecular covalent structure; extremely high melting point due to strong covalent bonds extending throughout lattice.\n * \text{P}_4, \text{S}_8, \text{Cl}_2, \text{Ar}$: Simple molecular structures; melting points depend on weak instantaneous dipole-induced dipole (id-id) forces. Melting point trend:$\text{S}_8 > \text{P}_4 > \text{Cl}_2 > \text{Ar} due to relative molecular masses and sizes.\n* **Oxidation States of Period 3 Elements in Oxides:**\n * Sodium Oxide (\text{Na}_2\text{O}$):$+1\n * Magnesium Oxide (\text{MgO}$):$+2\n * Aluminium Oxide (\text{Al}_2\text{O}_3$):$+3\n * Silicon Dioxide (\text{SiO}_2$):$+4\n * Tetraphosphorus Decoxide (\text{P}_4\text{O}_{10}$):$+5\n * Sulfur Dioxide (\text{SO}_2$):$+4\n * Sulfur Trioxide (\text{SO}_3$):$+6\n* **Reactions of Period 3 Oxides with Water:**\n * **Basic Oxides:**\n * \text{Na}_2\text{O}(s) + \text{H}_2\text{O}(l) \rightarrow 2\text{NaOH}(aq)$(Strongly alkaline solution, pH$\approx 13 - 14)\n * \text{MgO}(s) + \text{H}_2\text{O}(l) \rightarrow \text{Mg(OH)}_2(aq)$(Slightly soluble, weakly alkaline solution, pH$\approx 9 - 10)\n * **Insoluble Oxide:**\n * \text{SiO}_2(s) + \text{H}_2\text{O}(l) \rightarrow \text{No Reaction} (Insoluble giant covalent structure)\n * **Acidic Oxides:**\n * \text{P}_4\text{O}{10}(s) + 6\text{H}2\text{O}(l) \rightarrow 4\text{H}_3\text{PO}_4(aq)$(Phosphoric acid, pH$\approx 1 - 2)\n * \text{SO}_2(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_2\text{SO}_3(aq)$(Sulfurous acid, pH$\approx 2 - 3)\n * \text{SO}_3(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_2\text{SO}_4(aq)$(Sulfuric acid, pH$\approx 0 - 1)\n* **Amphoteric Nature of Aluminium Oxide (\text{Al}_2\text{O}_3):**\n * Reacts as a base with acid:\n * \text{Al}_2\text{O}_3(s) + 3\text{H}_2\text{SO}_4(aq) \rightarrow \text{Al}_2(\text{SO}_4)_3(aq) + 3\text{H}_2\text{O}(l)\n * Reacts as an acid with hot concentrated base:\n * \text{Al}_2\text{O}_3(s) + 2\text{NaOH}(aq) + 3\text{H}_2\text{O}(l) \rightarrow 2\text{NaAl(OH)}_4(aq)\n* **Reactions of Period 3 Chlorides with Water:**\n * \text{NaCl}(s)$: Ionic lattice dissolves to form neutral solution:$\text{NaCl}(s) \xrightarrow{\text{H}_2\text{O}} \text{Na}^+(aq) + \text{Cl}^-(aq)$(pH$= 7).\n * \text{MgCl}_2(s)$: Ionic lattice dissolves, slight hydrolysis: pH$\approx 6.5.\n * \text{AlCl}_3(s)$/$\text{Al}_2\text{Cl}_6(s)$: Dissolves with hydrolysis producing acidic solution and white fumes of$\text{HCl}$(pH$\approx 3).\n * \text{SiCl}_4(l)$: Hydrolyses vigorously with water forming white precipitate of$\text{SiO}_2$and acidic fumes of$\text{HCl}:\n * \text{SiCl}_4(l) + 2\text{H}_2\text{O}(l) \rightarrow \text{SiO}_2(s) + 4\text{HCl}(g)\n * \text{PCl}_5(l)$: Hydrolyses vigorously with water forming phosphoric acid and$\text{HCl} fumes:\n * \text{PCl}_5(l) + 4\text{H}_2\text{O}(l) \rightarrow \text{H}_3\text{PO}_4(aq) + 5\text{HCl}(g)\n\n# Group 2 Elements (Alkaline Earth Metals)\n\n* **Physical and Chemical Trends Down Group 2 (\text{Mg} \rightarrow \text{Ba}):**\n * Atomic Radius: Increases down the group as extra principal quantum shells are added.\n * First Ionisation Energy: Decreases down the group due to increased shielding and larger atomic radius.\n * Reactivity: Increases down the group as outer electrons are lost more easily.\n* **Reactions of Group 2 Metals:**\n * **With Oxygen:**\n * 2\text{M}(s) + \text{O}_2(g) \rightarrow 2\text{MO}(s) (Forms solid white oxide powders).\n * **With Water:**\n * Magnesium reacts very slowly with cold water forming weak alkaline solution:\n * \text{Mg}(s) + 2\text{H}_2\text{O}(l) \rightarrow \text{Mg(OH)}_2(aq) + \text{H}_2(g)\n * Magnesium reacts vigorously with steam to form Magnesium Oxide and Hydrogen:\n * \text{Mg}(s) + \text{H}_2\text{O}(g) \rightarrow \text{MgO}(s) + \text{H}_2(g)\n * Calcium, Strontium, and Barium react with increasing vigor with cold water:\n * \text{M}(s) + 2\text{H}_2\text{O}(l) \rightarrow \text{M(OH)}_2(aq) + \text{H}_2(g)\n * **With Acids:**\n * \text{M}(s) + 2\text{HX}(aq) \rightarrow \text{MX}_2(aq) + \text{H}_2(g) (Vigorous effervescence).\n* **Solubility Trends of Group 2 Compounds:**\n * **Group 2 Hydroxides (\text{M(OH)}_2$):** Solubility **increases** down the group ($\text{Mg(OH)}_2$is insoluble/sparingly soluble;$\text{Ba(OH)}_2 is soluble). Alkalinity of solution increases down the group.\n * **Group 2 Sulfates (\text{MSO}_4$):** Solubility **decreases** down the group ($\text{MgSO}_4$is soluble;$\text{BaSO}_4 is insoluble white precipitate).\n* **Thermal Decomposition of Group 2 Carbonates and Nitrates:**\n * Thermal stability increases down the group because cations increase in ionic radius, lowering charge density and exerting less polarizing power on the carbonate/nitrate anion.\n * **Decomposition of Carbonates (\text{MCO}_3):**\n * \text{MCO}_3(s) \xrightarrow{\Delta} \text{MO}(s) + \text{CO}_2(g)\n * Order of stability: \text{MgCO}_3 < \text{CaCO}_3 < \text{SrCO}_3 < \text{BaCO}_3\n * **Decomposition of Nitrates (\text{M(NO}_3)_2):**\n * 2\text{M(NO}_3)_2(s) \xrightarrow{\Delta} 2\text{MO}(s) + 4\text{NO}_2(g) + \text{O}_2(g)\n * Observations: Brown gas (\text{NO}_2$) evolved and glowing splint relights ($\text{O}_2).\n\n# Group 17 Elements (Halogens)\n\n* **Physical Properties of Halogens (\text{F}_2, \text{Cl}_2, \text{Br}_2, \text{I}_2):**\n * Diatomic non-polar molecules held by weak instantaneous dipole-induced dipole (id-id) intermolecular forces.\n * Physical States and Colors at Room Temperature:\n * Chlorine (\text{Cl}_2): Yellow-green gas.\n * Bromine (\text{Br}_2): Orange-brown liquid.\n * Iodine (\text{I}_2): Dark grey/black solid (sublimes to purple vapor).\n * Volatility decreases down the group (boiling points increase) due to an increasing number of electrons leading to stronger id-id forces.\n * Halogen Bond Energies (\text{X-X}$): Bond strength decreases from$\text{Cl}_2 \rightarrow \text{Br}_2 \rightarrow \text{I}_2$as atomic radius increases leading to greater bond length. (Note:$\text{F}_2 bond strength is anomalously lower due to strong lone-pair repulsion in the compact molecule).\n* **Reactivity and Oxidising Power:**\n * Oxidising ability decreases down Group 17: \text{F}_2 > \text{Cl}_2 > \text{Br}_2 > \text{I}_2.\n * A more reactive halogen displaces a less reactive halide from solution:\n * \text{Cl}_2(aq) + 2\text{Br}^-(aq) \rightarrow 2\text{Cl}^-(aq) + \text{Br}_2(aq) (Solution turns orange-brown).\n * \text{Cl}_2(aq) + 2\text{I}^-(aq) \rightarrow 2\text{Cl}^-(aq) + \text{I}_2(aq) (Solution turns brown/purple).\n* **Reactions of Halogens with Hydrogen:**\n * Fluorine (\text{F}_2): Reacts explosively even in cold and dark conditions:\n * \text{F}_2(g) + \text{H}_2(g) \rightarrow 2\text{HF}(g)\n * Chlorine (\text{Cl}_2): Reacts explosively in presence of sunlight/UV light:\n * \text{Cl}_2(g) + \text{H}_2(g) \rightarrow 2\text{HCl}(g)\n * Bromine (\text{Br}_2): Reacts slowly on heating:\n * \text{Br}_2(g) + \text{H}_2(g) \rightarrow 2\text{HBr}(g)\n * Iodine (\text{I}_2): Reversible reaction, incomplete on heating:\n * \text{I}_2(g) + \text{H}_2(g) \rightleftharpoons 2\text{HI}(g)\n* **Thermal Stability of Hydrogen Halides (\text{HX}):**\n * Thermal stability **decreases** down Group 17: \text{HF} > \text{HCl} > \text{HBr} > \text{HI}.\n * Reason: Bond length increases down the group (\text{H-F} < \text{H-Cl} < \text{H-Br} < \text{H-I}), decreasing bond strength, making thermal decomposition easier:\n * 2\text{HX}(g) \xrightarrow{\Delta} \text{H}_2(g) + \text{X}_2(g)\n* **Halide Ions as Reducing Agents:**\n * Reducing power **increases** down Group 17: \text{F}^- < \text{Cl}^- < \text{Br}^- < \text{I}^-.\n * Larger ionic radius down the group means outer electrons are lost more readily.\n * **Reaction of Solid Halides with Concentrated Sulfuric Acid (\text{H}_2\text{SO}_4):**\n * **Chlorides (\text{NaCl}$):** No redox reaction occurs ($\text{Cl}^-$is a weak reducing agent). Acid-base reaction yields steamy white fumes of$\text{HCl}:\n * \text{NaCl}(s) + \text{H}_2\text{SO}_4(l) \rightarrow \text{NaHSO}_4(s) + \text{HCl}(g)\n * **Bromides (\text{NaBr}$):**$\text{Br}^-$reduces$\text{H}_2\text{SO}_4$(oxidation state$+6$) to sulfur dioxide$\text{SO}_2$(oxidation state$+4$). Steamy fumes of$\text{HBr}$and orange-brown fumes of$\text{Br}_2 gas are observed:\n * \text{NaBr}(s) + \text{H}_2\text{SO}_4(l) \rightarrow \text{NaHSO}_4(s) + \text{HBr}(g)\n * 2\text{HBr}(g) + \text{H}_2\text{SO}_4(l) \rightarrow \text{Br}_2(g) + \text{SO}_2(g) + 2\text{H}_2\text{O}(l)\n * **Iodides (\text{NaI}$):**$\text{I}^-$is a strong reducing agent that reduces$\text{H}_2\text{SO}_4$to$\text{SO}_2$, elemental Sulfur ($\text{S}$), and Hydrogen Sulfide ($\text{H}_2\text{S}, rotten-egg odor). Purple iodine fumes and yellow sulfur precipitate observed:\n * \text{NaI}(s) + \text{H}_2\text{SO}_4(l) \rightarrow \text{NaHSO}_4(s) + \text{HI}(g)\n * 2\text{HI}(g) + \text{H}_2\text{SO}_4(l) \rightarrow \text{I}_2(g) + \text{SO}_2(g) + 2\text{H}_2\text{O}(l)\n * 6\text{HI}(g) + \text{H}_2\text{SO}_4(l) \rightarrow 3\text{I}_2(g) + \text{S}(s) + 4\text{H}_2\text{O}(l)\n * 8\text{HI}(g) + \text{H}_2\text{SO}_4(l) \rightarrow 4\text{I}_2(g) + \text{H}_2\text{S}(g) + 4\text{H}_2\text{O}(l)\n* **Testing for Halide Ions (\text{Cl}^-, \text{Br}^-, \text{I}^-):**\n * Add dilute nitric acid (\text{HNO}_3$) followed by aqueous silver nitrate ($\text{AgNO}_3(aq)):\n * **Chlorine (\text{Cl}^-$):** White precipitate of Silver Chloride ($\text{AgCl}). Dissolves in **dilute** aqueous ammonia.\n * **Bromine (\text{Br}^-$):** Cream precipitate of Silver Bromide ($\text{AgBr}). Insoluble in dilute ammonia, but dissolves in **concentrated** aqueous ammonia.\n * **Iodine (\text{I}^-$):** Yellow precipitate of Silver Iodide ($\text{AgI}). Insoluble in **both** dilute and concentrated aqueous ammonia.\n * General Ionic Equation: \text{Ag}^+(aq) + \text{X}^-(aq) \rightarrow \text{AgX}(s)\n* **Disproportionation Reactions of Chlorine:**\n * **With Cold Dilute \text{NaOH}$($15^\circ\text{C}):**\n * \text{Cl}_2(g) + 2\text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{NaClO}(aq) + \text{H}_2\text{O}(l)\n * Produces sodium chlorate(I) (\text{NaClO}), used in household bleach.\n * **With Hot Concentrated \text{NaOH}$($70^\circ\text{C}):**\n * 3\text{Cl}_2(g) + 6\text{NaOH}(aq) \rightarrow 5\text{NaCl}(aq) + \text{NaClO}_3(aq) + 3\text{H}_2\text{O}(l)\n * Produces sodium chlorate(V) (\text{NaClO}_3).\n * **Water Purification:**\n * Chlorine reacts with water to form chloric(I) acid (\text{HClO}), which acts as a disinfectant killing bacteria:\n * \text{Cl}_2(g) + \text{H}_2\text{O}(l) \rightleftharpoons \text{HCl}(aq) + \text{HClO}(aq)\n * \text{HClO}(aq) \rightleftharpoons \text{H}^+(aq) + \text{ClO}^-(aq)\n\n# Nitrogen and Sulfur Chemistry\n\n* **Unreactivity of Nitrogen Gas (\text{N}_2):**\n * Nitrogen exists as a diatomic molecule with a triple covalent bond (\text{N}\equiv\text{N}).\n * Possesses an exceptionally high bond dissociation energy (945\,kJ\,mol^{-1}) and lack of polarity, making it extremely unreactive under normal conditions.\n* **Ammonia (\text{NH}_3$) and Ammonium Ion ($\text{NH}_4^+):**\n * Ammonia is a weak base due to the lone pair of electrons on the Nitrogen atom, which can accept a proton (\text{H}^+).\n * Shape of Ammonia (\text{NH}_3$): Trigonal Pyramidal, bond angle$107^\circ.\n * Shape of Ammonium ion (\text{NH}_4^+$): Tetrahedral, bond angle$109.5^\circ. Contains three single covalent bonds and one dative covalent bond.\n * **Displacement of Ammonia from Ammonium Salts:**\n * Heating an ammonium salt with a strong base displaces ammonia gas:\n * \text{NH}_4\text{Cl}(s) + \text{Ca(OH)}_2(s) \xrightarrow{\Delta} \text{CaCl}_2(s) + 2\text{H}_2\text{O}(l) + 2\text{NH}_3(g)\n* **Environmental Impact of Nitrogen Oxides (\text{NO}_x):**\n * Formation of Nitrogen Oxides: High temperatures and pressures in internal combustion car engines cause atmospheric nitrogen and oxygen to react:\n * \text{N}_2(g) + \text{O}_2(g) \rightarrow 2\text{NO}(g)\n * 2\text{NO}(g) + \text{O}_2(g) \rightarrow 2\text{NO}_2(g)\n * **Catalytic Removal:** Catalytic converters containing Rhodium/Platinum catalysts reduce toxic emissions:\n * 2\text{CO}(g) + 2\text{NO}(g) \rightarrow 2\text{CO}_2(g) + \text{N}_2(g)\n * **Photochemical Smog and PAN:** Nitrogen oxides react with unburnt hydrocarbons in the presence of sunlight to form photochemical smog containing Peroxyacetyl Nitrate (PAN).\n * **Role in Acid Rain:**\n * Direct atmospheric contribution: 4\text{NO}_2(g) + 2\text{H}_2\text{O}(l) + \text{O}_2(g) \rightarrow 4\text{HNO}_3(aq)\n * Catalytic conversion of Sulfur Dioxide (\text{SO}_2$) to Sulfur Trioxide ($\text{SO}_3):\n * \text{NO}_2(g) + \text{SO}_2(g) \rightarrow \text{NO}(g) + \text{SO}_3(g)\n * \text{NO}(g) + \frac{1}{2}\text{O}_2(g) \rightarrow \text{NO}_2(g) (Catalyst regenerated)\n * \text{SO}_3(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_2\text{SO}_4(aq)\n\n# Introductory Organic Chemistry: Nomenclature and Isomerism\n\n* **Terminology and Representation:**\n * **Hydrocarbon:** Compound containing hydrogen and carbon atoms only.\n * **Empirical Formula:** Simplest whole-number ratio of elements.\n * **Molecular Formula:** Actual number of atoms of each element in a molecule.\n * **Structural Formula:** Shows how atoms are bonded together carbon by carbon (e.g., \text{CH}_3\text{CH}_2\text{OH}).\n * **Displayed Formula:** Shows all atoms and all covalent bonds in 2D.\n * **Skeletal Formula:** Simplest representation showing carbon backbone as lines, omitting carbon and hydrogen symbols attached to carbons.\n* **Homologous Series:**\n * A family of organic compounds with the same functional group, same general formula, similar chemical properties, and a gradual trend in physical properties, where successive members differ by a -\text{CH}_2- unit.\n* **Functional Groups and Suffixes:**\n * Alkane (-\text{C-C}-): suffix **-ane**\n * Alkene (-\text{C=C}-): suffix **-ene**\n * Alcohol (-\text{OH}): suffix **-ol**\n * Halogenoalkane (-\text{F}, -\text{Cl}, -\text{Br}, -\text{I}): prefix **fluoro-**, **chloro-**, **bromo-**, **iodo-**\n * Aldehyde (-\text{CHO}): suffix **-al**\n * Ketone (-\text{C=O}): suffix **-one**\n * Carboxylic Acid (-\text{COOH}): suffix **-oic acid**\n * Ester (-\text{COOR}): suffix **-oate**\n * Nitrile (-\text{C}\equiv\text{N}): suffix **-nitrile**\n * Amine (-\text{NH}_2): suffix **-amine**\n* **Types of Organic Reactions:**\n * **Free Radical:** Species with an unpaired electron.\n * **Electrophile:** An electron-pair acceptor (positively charged or electron-deficient species).\n * **Nucleophile:** An electron-pair donor (negatively charged or species with a lone pair of electrons).\n * **Addition:** Two molecules combine to form a single product.\n * **Substitution:** An atom or group of atoms is replaced by another atom or group.\n * **Elimination:** Removal of a small molecule from a reactant to form a double bond.\n * **Hydrolysis:** Breakdown of a molecule by reaction with water.\n * **Condensation:** Combination of two molecules with the elimination of a small molecule like water.\n* **Isomerism:**\n * **Structural Isomerism:** Compounds with the same molecular formula but different structural formulae.\n * **Chain Isomers:** Differences in carbon skeleton branching (e.g., pentane vs 2-methylbutane).\n * **Position Isomers:** Same functional group located at different positions on carbon chain (e.g., pentan-1-ol vs pentan-2-ol).\n * **Functional Group Isomers:** Different functional groups (e.g., Alkenes and Cycloalkanes; Alcohols and Ethers; Aldehydes and Ketones; Carboxylic Acids and Esters).\n * **Stereoisomerism:** Compounds with the same molecular and structural formula but different spatial arrangement of atoms.\n * **Geometrical Isomerism (cis/trans$):** Occurs in alkenes due to restricted rotation about the$\text{C=C} double bond and two different groups attached to each carbon of the double bond.\n * **cis Isomer:** High-priority / identical groups on the same side of the double bond.\n * **trans Isomer:** High-priority / identical groups on opposite sides of the double bond.\n * **Optical Isomerism:** Occurs in molecules possessing a chiral center (a carbon atom bonded to four different groups).\n * Forms non-superimposable mirror images called **enantiomers**.\n * Enantiomers rotate plane-polarized light in opposite directions.\n\n# Alkanes: Hydrocarbon Chemistry and Reactions\n\n* **Structure and General Formula:**\n * General Formula: \text{C}_n\text{H}{2n+2}\n * Saturated hydrocarbons containing only single \text{C-C}$and$\text{C-H}$$\sigma bonds.\n * Tetrahedral carbon geometry (sp^3$hybridised, bond angle$109.5^\circ).\n* **Unreactivity of Alkanes:**\n * Generally unreactive due to high strength of non-polar \text{C-C}$and$\text{C-H} bonds.\n* **Combustion of Alkanes:**\n * **Complete Combustion:** Excess oxygen produces \text{CO}2$and$\text{H}_2\text{O}:\n * \text{C}_x\text{H}_y + \left(x + \frac{y}{4}\right)\text{O}_2 \rightarrow x\text{CO}_2 + \frac{y}{2}\text{H}_2\text{O}\n * **Incomplete Combustion:** Limited oxygen produces poisonous Carbon Monoxide (\text{CO}$) or soot ($\text{C}).\n* **Free-Radical Substitution Mechanism (Halogenation of Alkanes):**\n * Reaction of methane with chlorine in presence of Ultraviolet (UV) light:\n * Overall Equation: \text{CH}_4 + \text{Cl}_2 \xrightarrow{UV} \text{CH}_3\text{Cl} + \text{HCl}\n * **Step 1: Initiation:** Homolytic fission of halogen bond by UV light producing free radicals:\n * \text{Cl}_2 \xrightarrow{UV} 2\text{Cl}^\bullet\n * **Step 2: Propagation:** Radical reacts with neutral molecule generating new radicals:\n * \text{Cl}^\bullet + \text{CH}_4 \rightarrow \text{CH}_3^\bullet + \text{HCl}\n * \text{CH}_3^\bullet + \text{Cl}_2 \rightarrow \text{CH}_3\text{Cl} + \text{Cl}^\bullet\n * **Step 3: Termination:** Two free radicals collide and combine to form a stable molecule:\n * \text{Cl}^\bullet + \text{Cl}^\bullet \rightarrow \text{Cl}_2\n * \text{CH}_3^\bullet + \text{Cl}^\bullet \rightarrow \text{CH}_3\text{Cl}\n * \text{CH}_3^\bullet + \text{CH}_3^\bullet \rightarrow \text{C}_2\text{H}_6\n* **Cracking of Long-Chain Alkanes:**\n * Thermal or catalytic breakdown of long-chain alkanes into smaller, more useful short-chain alkanes and alkenes:\n * Example: \text{C}{10}\text{H}{22} \rightarrow \text{C}_8\text{H}{18} + \text{C}2\text{H}_4\n\n# Alkenes: Structure, Addition Reactions, and Oxidation\n\n* **Structure and General Formula:**\n * General Formula: \text{C}_n\text{H}{2n}\n * Unsaturated hydrocarbons containing at least one \text{C=C}$double bond ($1\,\sigma$bond and$1\,\pi bond).\n * Trigonal planar geometry around double bond carbons (sp^2$hybridised, bond angle$120^\circ).\n* **Electrophilic Addition Reactions:**\n * **1. Hydrogenation (Addition of \text{H}_2):**\n * Conditions: Nickel (\text{Ni}$) catalyst,$150^\circ\text{C}.\n * Equation: \text{CH}_2\text{=CH}_2 + \text{H}_2 \xrightarrow{\text{Ni}, \Delta} \text{CH}_3\text{CH}_3 (Produces alkane).\n * **2. Halogenation (Addition of \text{X}_2):**\n * Bromine water test: Orange-brown bromine water is decolorised (turns colorless) in the presence of alkenes.\n * Equation: \text{CH}_2\text{=CH}_2 + \text{Br}_2 \rightarrow \text{CH}_2\text{BrCH}_2\text{Br} (1,2-dibromoethane).\n * **3. Addition of Hydrogen Halides (\text{HX}):**\n * Equation: \text{CH}_2\text{=CH}_2 + \text{HBr} \rightarrow \text{CH}_3\text{CH}_2\text{Br}\n * **Markovnikov's Rule:** In electrophilic addition to unsymmetrical alkenes, the hydrogen atom attaches to the carbon atom of the double bond that already has more hydrogen atoms attached, forming the more stable carbocation intermediate.\n * **Carbocation Stability Order:** Tertiary (3^\circ$) > Secondary ($2^\circ$) > Primary ($1^\circ) due to electron-donating inductive effect of alkyl groups.\n * **4. Hydration (Addition of Steam \text{H}_2\text{O}):**\n * Conditions: Concentrated Phosphoric Acid (\text{H}_3\text{PO}_4$) catalyst,$300^\circ\text{C}$,$60\,\text{atm}.\n * Equation: \text{CH}_2\text{=CH}_2 + \text{H}_2\text{O}(g) \xrightarrow{\text{H}_3\text{PO}_4} \text{CH}_3\text{CH}_2\text{OH} (Produces alcohol).\n* **Oxidation Reactions of Alkenes:**\n * **Cold Dilute Acidified Potassium Manganate(VII) (\text{KMnO}_4):**\n * Mild oxidation; purple solution turns colorless, forming a 1,2-diol (dihydric alcohol).\n * Equation: \text{R-CH=CH}_2 + [\text{O}] + \text{H}_2\text{O} \xrightarrow{\text{cold KMnO}_4} \text{R-CH(OH)-CH}_2\text{OH}\n * **Hot Concentrated Acidified Potassium Manganate(VII) (\text{KMnO}_4):**\n * Cleaves the \text{C=C} double bond completely:\n * =\text{CH}_2$group oxidises to Carbon Dioxide ($\text{CO}_2$) and Water ($\text{H}_2\text{O}).\n * =\text{CHR}$group oxidises to an Aldehyde which further oxidises to a Carboxylic Acid ($\text{R-COOH}).\n * =\text{CR}_1\text{R}_2$group oxidises to a Ketone ($\text{R}_1\text{-CO-R}_2).\n\n# Halogenoalkanes: Substitution and Elimination Mechanisms\n\n* **Structure and Classification:**\n * Contains halogen atom attached to sp^3 carbon.\n * Primary (1^\circ): Halogen carbon attached to 1 alkyl group.\n * Secondary (2^\circ): Halogen carbon attached to 2 alkyl groups.\n * Tertiary (3^\circ): Halogen carbon attached to 3 alkyl groups.\n* **Nucleophilic Substitution Reactions:**\n * **1. Hydrolysis with Aqueous Alkali (\text{NaOH}(aq)$or$\text{KOH}(aq)):**\n * Heat under reflux.\n * Equation: \text{R-X} + \text{OH}^-(aq) \rightarrow \text{R-OH} + \text{X}^-(aq) (Forms alcohol).\n * **2. Reaction with Potassium Cyanide (\text{KCN} in ethanol):**\n * Heat under reflux.\n * Equation: \text{R-X} + \text{CN}^- \xrightarrow{\text{ethanol}} \text{R-CN} + \text{X}^- (Forms nitrile, extends carbon chain by 1 carbon).\n * **3. Reaction with Ammonia (\text{NH}_3 in ethanol under pressure):**\n * Heat in sealed tube.\n * Equation: \text{R-X} + 2\text{NH}_3 \xrightarrow{\text{ethanol}, \text{pressure}} \text{R-NH}_2 + \text{NH}_4\text{X} (Forms primary amine).\n* **Nucleophilic Substitution Mechanisms (S_N1$vs$S_N2):**\n * **S_N2 Mechanism:**\n * Second-order nucleophilic substitution; bimolecular rate-determining step.\n * Favored by **primary halogenoalkanes**.\n * Occurs via a single-step concerted process with a 5-coordinate transition state, leading to **inversion of configuration**.\n * **S_N1 Mechanism:**\n * First-order nucleophilic substitution; unimolecular rate-determining step.\n * Favored by **tertiary halogenoalkanes** due to steric hindrance and stability of tertiary carbocations.\n * Occurs in two steps: Step 1 (slow) involves heterolytic cleavage of \text{C-X} bond to form a planar carbocation intermediate; Step 2 (fast) involves nucleophilic attack, producing a **racemic mixture** (loss of optical activity).\n* **Elimination Reactions:**\n * Reaction of halogenoalkane with ethanolic Sodium Hydroxide (\text{NaOH} in ethanol) under heat/reflux:\n * Equation: \text{CH}_3\text{CH}_2\text{Br} + \text{NaOH}(\text{ethanol}) \xrightarrow{\Delta} \text{CH}_2\text{=CH}_2 + \text{NaBr} + \text{H}_2\text{O} (Forms alkene).\n* **Reactivity Trend of Halogenoalkanes:**\n * Reactivity order: \text{R-I} > \text{R-Br} > \text{R-Cl} > \text{R-F}.\n * Controlled by \text{C-X}$bond enthalpy (strength), not bond polarity.$\text{C-I} bond is the weakest and breaks most easily.\n * Verification using aqueous \text{AgNO}_3$in ethanol: Iodoalkanes form yellow precipitate of$\text{AgI} fastest; chloroalkanes form white precipitate slowest.\n\n# Alcohols and Hydroxy Compounds\n\n* **Classification:**\n * Primary (1^\circ$):$\text{R-CH}_2\text{OH}\n * Secondary (2^\circ$):$\text{R}_2\text{CHOH}\n * Tertiary (3^\circ$):$\text{R}_3\text{COH}\n* **Formation of Alcohols:**\n * Electrophilic hydration of alkenes using steam and \text{H}_3\text{PO}_4.\n * Nucleophilic substitution of halogenoalkanes with \text{NaOH}(aq).\n * Reduction of aldehydes, ketones, or carboxylic acids using reducing agents like Lithium Aluminium Hydride (\text{LiAlH}_4$) or Sodium Borohydride ($\text{NaBH}_4).\n* **Chemical Reactions of Alcohols:**\n * **1. Combustion:** Complete burning in \text{O}_2$produces$\text{CO}_2$and$\text{H}_2\text{O}.\n * **2. Reaction with Sodium Metal:**\n * Reacts to form sodium alkoxide and hydrogen gas (effervescence):\n * 2\text{R-OH} + 2\text{Na} \rightarrow 2\text{R-O}^-\text{Na}^+ + \text{H}_2(g)\n * **3. Substitution to Form Halogenoalkanes:**\n * With \text{HX}$:$\text{R-OH} + \text{HX} \rightarrow \text{R-X} + \text{H}_2\text{O}\n * With Phosphorus Pentachloride (\text{PCl}_5$):$\text{R-OH} + \text{PCl}_5 \rightarrow \text{R-Cl} + \text{POCl}_3 + \text{HCl}(g)$(Steamy fumes of$\text{HCl}$test for$-\text{OH} group).\n * With Phosphorus Trichloride (\text{PCl}_3$under heat):$3\text{R-OH} + \text{PCl}_3 \xrightarrow{\Delta} 3\text{R-Cl} + \text{H}_3\text{PO}_3\n * With Thionyl Chloride (\text{SOCl}_2$):$\text{R-OH} + \text{SOCl}_2 \rightarrow \text{R-Cl} + \text{SO}_2(g) + \text{HCl}(g)\n * **4. Dehydration to Alkenes:**\n * Heated with concentrated \text{H}_2\text{SO}_4$at$170^\circ\text{C}$or passed over heated Aluminium Oxide ($\text{Al}_2\text{O}_3):\n * \text{CH}_3\text{CH}_2\text{OH} \xrightarrow{\text{conc. } \text{H}_2\text{SO}_4, 170^\circ\text{C}} \text{CH}_2\text{=CH}_2 + \text{H}_2\text{O}\n * **5. Oxidation of Alcohols:**\n * Reagents: Acidified Potassium Dichromate (\text{K}_2\text{Cr}_2\text{O}_7 / \text{H}^+$), color change from orange ($\text{Cr}_2\text{O}_7^{2-}$) to green ($\text{Cr}^{3+}).\n * **Primary (1^\circ) Alcohols:**\n * Distillation yields an **Aldehyde**: \text{R-CH}_2\text{OH} + [\text{O}] \xrightarrow{\text{distil}} \text{R-CHO} + \text{H}_2\text{O}\n * Reflux with excess oxidant yields a **Carboxylic Acid**: \text{R-CH}_2\text{OH} + 2[\text{O}] \xrightarrow{\text{reflux}} \text{R-COOH} + \text{H}_2\text{O}\n * **Secondary (2^\circ) Alcohols:**\n * Reflux yields a **Ketone**: \text{R-CH(OH)-R'} + [\text{O}] \xrightarrow{\text{reflux}} \text{R-CO-R'} + \text{H}_2\text{O}\n * **Tertiary (3^\circ$) Alcohols:** Resistant to oxidation (no hydrogen attached to the carbon bearing the$-\text{OH} group); solution remains orange.\n * **6. Tri-iodomethane (Iodoform) Test:**\n * Reagent: Aqueous Iodine (\text{I}_2$) in Sodium Hydroxide ($\text{NaOH}).\n * Positive test: Pale yellow precipitate of tri-iodomethane (\text{CHI}_3) with a characteristic antiseptic smell.\n * Identifies alcohols possessing the specific structure \text{CH}_3\text{CH(OH)}- (e.g., Ethanol and all secondary methyl alcohols like propan-2-ol).\n\n# Carbonyl Compounds: Aldehydes and Ketones\n\n* **Structure:**\n * Contain the polar carbonyl group (\text{C=O}).\n * Aldehydes (\text{R-CHO}): Carbonyl carbon attached to at least one hydrogen atom.\n * Ketones (\text{R-CO-R'}): Carbonyl carbon attached to two alkyl groups.\n* **Reduction of Carbonyl Compounds:**\n * Reducing agents: Sodium Borohydride (\text{NaBH}_4$) in aqueous/alcoholic solution or Lithium Aluminium Hydride ($\text{LiAlH}_4) in dry ether.\n * Reduction of Aldehydes produces Primary Alcohols:\n * \text{R-CHO} + 2[\text{H}] \rightarrow \text{R-CH}_2\text{OH}\n * Reduction of Ketones produces Secondary Alcohols:\n * \text{R-CO-R'} + 2[\text{H}] \rightarrow \text{R-CH(OH)-R'}\n* **Nucleophilic Addition with Hydrogen Cyanide (\text{HCN}):**\n * Reagent: \text{HCN}$in presence of$\text{KCN}$(or$\text{NaCN} + \text{H}_2\text{SO}_4) as catalyst at pH 8.\n * Produces Hydroxynitriles (Cyanohydrins):\n * \text{R-CHO} + \text{HCN} \rightarrow \text{R-CH(OH)CN}$$

  • Mechanism: Cyan