Comprehensive A-Level Chemistry Encyclopedia

CHAPTER 1: ATOMIC STRUCTURE

1.1 Inside the Atom

• Atomic Model Components:

  • Electrons: Revolve around the nucleus in regions of space called orbitals. They do not move in fixed orbits.
  • Nucleus: Primarily composed of protons and neutrons. It contains almost all the mass of the atom because the relative mass of electrons is negligible.
  • Nuclear Charge: The nucleus is positively-charged due to the presence of protons. Negative electrons surround this positive core.

• Sub-atomic Particle Properties:

  • Protons ($p$):
    • Relative Mass: $1$
    • Relative Charge: $+1$
    • Actual Charge: $+1.6 \times 10^{-19}\,C$
  • Neutrons ($n$):
    • Relative Mass: $1$
    • Relative Charge: $0$
    • Actual Charge: $0\,C$
  • Electrons ($e$):
    • Relative Mass: $\frac{1}{1836}$
    • Relative Charge: $-1$
    • Actual Charge: $-1.6 \times 10^{-19}\,C$

• Behavior in Electric Fields:

  • Protons: Deflected towards the negative plate ($+$ charge attracted to $-$).
  • Electrons: Deflected towards the positive plate ($-$ charge attracted to $+$).
  • Neutrons: Not deflected; they continue in their original direction because they are neutral.
  • Angle of Deflection: The angle of deflection for electrons is greater than that of protons. This is because the angle of deflection is inversely proportional to the mass-to-charge ratio ($m/e$); since electrons have much smaller mass, they deflect more easily.

• Definitions of Numbers:

  • Proton Number ($Z$): The total number of protons in an atom (also called atomic number).
  • Nucleon Number ($A$): The total number of protons and neutrons in an atom (also called mass number).
  • Neutral Atoms: The number of protons equals the number of electrons.
  • Cation: A positively-charged ion formed when an atom loses electron(s). Protons > Electrons.
  • Anion: A negatively-charged ion formed when an atom gains electron(s). Electrons > Protons.

• Comparative Terminology:

  • Isoelectronic: Species having the same number of electrons.
  • Isotonic: Species having the same number of neutrons.
  • Isotopic: Atoms of the same element having the same number of protons.

1.2 Isotopes

• Definition: Atoms of the same element with the same number of protons but a different number of neutrons.
• Similarities:
  • Same number of protons and electrons.
  • Identical electronic configuration.
  • Identical chemical properties, as chemical properties are determined by electron counts.
• Differences:
  • Different number of neutrons and nucleon numbers.
  • Variations in mass, density, and molecular speed.
• Stability: Can be stable or unstable; unstable ones are known as radioactive isotopes or radioisotopes.
• Symbolism: Representation as $_{y}^{x}X$, where $x$ is the nucleon number and $y$ is the proton number.

CHAPTER 2: ATOMS, MOLECULES AND STOICHIOMETRY

2.1 Mass of Atoms and Molecules

• Relative Mass Concept: An indication of an atom's mass compared to a standard model, expressed in atomic mass units ($a.m.u.$).
• Carbon-12 Standard: Chosen because it is the most abundant isotope of carbon and is a solid that is easily handled. Its mass is exactly $12\,a.m.u.$.
• Relative Isotopic Mass: The mass of an isotope on a scale where Carbon-12 is exactly $12$ units.
• Relative Atomic Mass ($A_r$): The weighted average relative masses of all isotopes of an element on the C-12 scale.
  • Formula: $A_r = \frac{\text{Average mass of one atom of element}}{\frac{1}{12} \times \text{mass of one atom of C-12}}$
• Relative Molecular Mass ($M_r$): Weighted average of the masses of molecules on the C-12 scale. Calculated by summing the $A_r$ of all atoms in the molecule (e.g., $M_r(H_2O) = 2(1) + 16 = 18$).
• Relative Formula Mass ($M_r$): Weighted average of formula units (used for both ionic and covalent compounds). Example: $M_r(CuSO_4 \cdot 5H_2O) = 64 + 32 + 4(16) + 5[2(1) + 16] = 249.5$.

2.2 Mass Spectrometer

• Functions: Determines relative isotopic mass, relative abundance of isotopes, relative atomic/molecular mass, and structural formulas.
• Process Steps:
  1. Vaporisation: Atoms are turned into gas.
  2. Ionisation: Gaseous atoms are bombarded with high-energy electrons to form positive ions. The chamber is a vacuum to prevent collisions with air molecules.
  3. Acceleration: Ions are accelerated to the same kinetic energy.
  4. Deflection: A magnetic field deflects ions. Deflection depends on the mass/charge ratio ($m/e$ or $m/z$). Larger mass equals smaller deflection; higher charge equals larger deflection.
  5. Detection: Ions are detected electrically, and a mass spectrum is produced.
• Calculation from Spectrum:
  • $A_r = \frac{(m_1 \times a) + (m_2 \times b) + (m_3 \times c)}{a + b + c}$ where $m_n$ is the mass and $a, b, c$ are abundances.

2.3 Amount of Substance

• The Mole: The amount of substance containing the same number of elementary units as there are atoms in $12\,g$ of C-12.
• Avogadro's Constant ($L$): $6.02 \times 10^{23}\,mol^{-1}$.
• Mole and Mass: $\text{No. of moles} = \frac{\text{Mass (g)}}{\text{Molar mass (g/mol)}}$.
• Gases and Volume:
  • Avogadro's Law: Equal volumes of gases under identical conditions contain equal moles.
  • Room Temperature (20 $^\circ$C, 1 atm): $V_{m} = 24\,dm^3\,mol^{-1}$.
  • s.t.p. (0 $^\circ$C, 1 atm): $V_{m} = 22.4\,dm^3\,mol^{-1}$.
  • $\text{No. of moles} = \frac{\text{Volume of gas (dm}^3)}{\text{Molar volume (dm}^3\,\text{mol}^{-1})}$.
• Combustion:
  • Complete: Produces $CO_2$ and $H_2O$.
  • Incomplete: Produces $CO_2, CO, C(\text{soot})$, and $H_2O$.
• Concentration:
  • $g\,dm^{-3}$ is concentration; $mol\,dm^{-3}$ is molarity ($M$).
  • $\text{Concentration (g/dm}^3) = \frac{\text{Mass (g)}}{\text{Volume (dm}^3)}$
  • $\text{Molarity (mol/dm}^3) = \frac{\text{Concentration (g/dm}^3)}{\text{Molar mass (g/mol)}}$
  • $\text{No. of moles} = \frac{\text{Volume (cm}^3) \times \text{Molarity}}{1000}$

2.4 Empirical and Molecular Formula

• Empirical Formula: Simplest ratio of atoms in a molecule. Ionic compounds are always empirical. Some simple inorganic molecules have identical empirical and molecular formulas.
• Molecular Formula: Actual number of atoms. Requires relative formula mass and empirical formula to deduce.
• Percentage Composition: $\% = \frac{A_r \times \text{No. of moles of element}}{\text{Molar mass of compound}} \times 100\%$
• Finding Empirical Formula: Find mass $\rightarrow$ Find moles ($/A_r$) $\rightarrow$ Find simplest ratio ($/\text{smallest value}$) $\rightarrow$ Construct formula.
• Conservation of Mass: Total mass of reactants equals total mass of products in a closed system.

2.5 Stoichiometry and Equations

• Stoichiometry: Proportion of things reacting or combining (e.g., $H_2O$ has a $2:1$ stoichiometry of $H$ to $O$).
• Ionic Equations:
  1. Balanced molecular equation.
  2. Complete ionic equation (split aqueous ionic compounds only).
  3. Cancel spectator ions (present on both sides but not participating).
  4. Net ionic equation (the leftovers).
  • Rule: Never split solids ($s$), liquids ($l$), gases ($g$), metals, or covalent compounds.

CHAPTER 3: ELECTRONS IN ATOMS

3.1 Sub-shells and Atomic Orbitals

• Principal Quantum Shell ($n$): Electrons arranged in energy levels. Higher $n$ means further from the nucleus, more energy, and weaker nuclear hold. Max electrons = $2n^2$.
• Sub-shells: Quanta split into s, p, d, f. Energies: $s < p < d < f$.
• Atomic Orbitals: Region with $>95\%$ probability of finding an electron. Electron has dual nature (particle and wave).
  • s sub-shell: 1 orbital ($s$), spherical, non-directional. Size increase: $1s < 2s < 3s$.
  • p sub-shell: 3 degenerate orbitals ($p_x, p_y, p_z$), dumb-bell shaped, directional. Perpendicular along axes.
  • d sub-shell: 5 orbitals.

3.2 Electronic Configuration

• Heisenberg Uncertainty Principle: Cannot know path or position with certainty.
• Rules for Filling Orbitals:
  • Aufbau Principle: Fill lowest energy first ($4s$ is slightly lower than $3d$, but this reverses once $3d$ is filled).
  • Pauli's Exclusion Principle: Max 2 electrons per orbital with opposite spins.
  • Hund's Rule: Fill degenerate orbitals singly first to avoid repulsion.
• Specific Configurations:
  • Chromium ($Z=24$): $[Ar]3d^5 4s^1$ (half-filled stability).
  • Copper ($Z=29$): $[Ar]3d^{10} 4s^1$ (full $3d$ stability).
• Ion Formation:
  • Cations: Remove electrons in order of decreasing energy. For d-block, $4s$ is removed before $3d$ because $3d$ electrons repel $4s$ to a higher energy level once filled.

3.3 Ionisation Energy ($I.E.$)

• 1st Ionisation Energy ($\Delta H_{i1}$): Energy to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous $1+$ ions: $X(g) \rightarrow X^+(g) + e^-$.
• Factors Affecting $I.E.$:
  1. Nuclear Charge: Higher charge = stronger attraction = higher $I.E.$.
  2. Distance: Larger radius = weaker attraction = lower $I.E.$.
  3. Shielding: Inner electrons repel outer ones, reducing nuclear pull = lower $I.E.$.
• Trends:
  • Down a Group: Decreases (Size and shielding increase, outweighing nuclear charge).
  • Across a Period: Generally increases (Nuclear charge increases, shielding is constant, distance decreases).
• Exceptions/Drops:
  • Be to B / Mg to Al: Drop because the outer electron is in a higher sub-shell ($2p$ vs $2s$) and experience more shielding.
  • N to O / P to S: Drop due to spin-pair repulsion; removing an electron from a paired orbital is easier.
• Successive $I.E.$: Get larger as ions become more positive. Large jumps indicate a change in principal quantum shell.

CHAPTER 4: CHEMICAL BONDING

4.1 Ionic Bonding

• Definition: Electrostatic force of attraction between oppositely-charged ions formed by complete electron transfer. Non-directional.
• Crystal Lattice: Forms a huge lattice structure. Strong bonds result in high melting points.
• Dot-and-cross Diagrams: Show only outer shells; use square brackets and charge symbols.
• Strength Calculation: $E \propto \frac{(Q^+) (Q^-)}{d^2}$. Higher charge and smaller ions result in higher melting points (e.g., $MgO > NaCl$).

4.2 Covalent Bonding

• Definition: Electrostatic attraction between two nuclei for a shared localized pair of electrons.
• Types:
  • Single: 1 pair ($H-H$).
  • Double: 2 pairs ($O=O$).
  • Triple: 3 pairs ($N\equiv N$).
• Co-ordinate (Dative) Bond: One atom providing both electrons. Needs a lone pair and an electron-deficient species. Example: $NH_4^+$ or $Al_2Cl_6$.
• Octet Rule Exceptions:
  • Octet-deficient: Central atom has $<8$ electrons (e.g., $BF_3, NO$).
  • Expanded Octet: Period 3 elements and beyond can accommodate $>8$ electrons using d-orbitals (e.g., $PCl_5$).

4.3 Shapes of Molecules (VSEPR Theory)

• Theory: Electron pairs repel. Repulsion strength: $\text{Lone pair-Lone pair} > \text{Lone pair-Bond pair} > \text{Bond pair-Bond pair}$.
• Common Geometries & Angles:
  • Linear: 2 domains, $180^{\circ}$ ($CO_2$).
  • Trigonal Planar: 3 domains, $120^{\circ}$ ($BF_3$).
  • Tetrahedral: 4 domains, $109.5^{\circ}$ ($CH_4$).
  • Trigonal Pyramid: 3 bond, 1 lone, $107^{\circ}$ ($NH_3$).
  • Bent (Non-linear): 2 bond, 2 lone, $104.5^{\circ}$ ($H_2O$).
  • Trigonal Bipyramid: 5 domains, $90^{\circ}$ and $120^{\circ}$ ($PF_5$).
  • Octahedral: 6 domains, $90^{\circ}$ ($SF_6$).

4.4 Orbital Overlap & Hybridisation

• Sigma ($\sigma$) bond: End-to-end overlap; electron density between nuclei.
• Pi ($\pi$) bond: Sideways overlap of p-orbitals; high density regions above/below nuclei. Weaker than $\sigma$.
• Hybridisation: Mixing orbitals into equivalent hybrid orbitals:
  • $sp^3$: Tetrahedral ($109.5^{\circ}$), e.g., Ethane.
  • $sp^2$: Trigonal planar ($120^{\circ}$), e.g., Ethene.
  • $sp$: Linear ($180^{\circ}$), e.g., Ethyne.

4.5 Intermolecular Forces

• Electronegativity: Ability to attract bond pairs. Trends: Increase across Period, decrease down Group. Fluorine is highest.
• Permanent Dipole-Dipole: Between polar molecules (asymmetrical charge distribution).
• Induced Dipole (London/Dispersion forces): All molecules. Strength increases with number of electrons (molecular size) and contact surface area.
• Hydrogen Bonding: Attraction between $H$ (bonded to $F, O, N$) and a lone pair on $F, O, N$. Explains high boiling points of $H_2O, NH_3, HF$.
  • Water Properties: Ice is less dense than liquid (tetrahedral structure with empty spaces). High surface tension/viscosity.

4.6 Metallic Bonding

• Definition: Attraction between delocalized 'sea of electrons' and positive metal ions.
• Strength Factors: Number of delocalized electrons (more = stronger) and ion size (smaller = stronger).

CHAPTER 5: STATES OF MATTER

5.1 Gaseous State

• Ideal Gas Assumptions:
  1. Zero intermolecular forces.
  2. Negligible particle volume.
  3. Constant random motion, elastic collisions.
  4. Pressure due to collisions with walls.
• Deviation: Real gases deviate most at high pressure and low temperature (volume and forces become non-negligible).
• Equation: $pV = nRT$ ($P$ in $Pa$, $V$ in $m^3$, $T$ in $K$, $R = 8.31\,J\,K^{-1}\,mol^{-1}$).

5.2 Liquids and Vapour Pressure

• Melting/Boiling: Absorbing energy to overcome lattice/intermolecular forces.
• Saturated Vapour Pressure: Equilibrium in a closed container where evaporation rate equals condensation rate. Increases with temperature.

5.3 Solid State Lattices

• Ionic: e.g., $NaCl$ ($6:6$ co-ordinated), $CsCl$ ($8:8$). Hard, brittle, conductive only when molten/aqueous.
• Simple Molecular: e.g., Iodine. Low melting point, weak Van der Waals forces.
• Giant Molecular:
  • Silicon(IV) Oxide: Tetrahedral structure, very high melting point.
  • Diamond: $sp^3$ carbon, tetrahedral network, extremely hard.
  • Graphite: $sp^2$ carbon, layers held by weak forces, conducts electricity parallel to layers (delocalized electrons).
• Metallic: e.g., copper ($12$-co-ordinated). Malleable, ductile, heat/electrical conductors.

5.4 Ceramics and Conservation

• Ceramics: Inorganic solids ($SiO_2, MgO, Al_2O_3$). Heat resistant, electrical insulators, chemically unreactive.
• Recycling: Saves resources, reduces waste, saves energy vs. extraction (e.g., $Al, Cu, \text{glass}$).

CHAPTER 6: CHEMICAL ENERGETICS

• Enthalpy Change ($\Delta H$): Measured heat change at constant pressure.
  • Exothermic: Releases heat ($-\Delta H$), surroundings warm up. More stable.
  • Endothermic: Absorbs heat ($+\Delta H$), surroundings cool down.
• Standard Conditions ($\Delta H^{\circ}$): $298\,K, 100\,kPa, 1.0\,mol\,dm^{-3}$.
• Specific Enthalpy Definitions:
  • Formation ($\Delta H_f^{\circ}$): 1 mole of compound from elements.
  • Combustion ($\Delta H_c^{\circ}$): 1 mole burnt in excess oxygen.
  • Neutralisation ($\Delta H_n^{\circ}$): 1 mole of water formed from acid/alkali. Strong acid/base reaction is constant ($-57.1\,kJ\,mol^{-1}$).
  • Atomisation ($\Delta H_{at}^{\circ}$): 1 mole of gaseous atoms from element.
  • Hydration ($\Delta H_{hyd}^{\circ}$): 1 mole of gaseous ions dissolved in water.
• Calorimetry: $\text{Heat} = mc\Delta T$. $\Delta H = \frac{-\text{Heat}}{\text{moles}}$.
• Hess' Law: Total enthalpy change is independent of the route. Useful for calculating theoretical $\Delta H$.
• Bond Energy ($E$): Energy to break 1 mole of covalent bonds in gas phase.

CHAPTER 7: REDOX REACTIONS AND ELECTROLYSIS

• Definitions:
  • Oxidation: Gain $O$, loss $H$, loss $e^-$, increase in oxidation number.
  • Reduction: Loss $O$, gain $H$, gain $e^-$, decrease in oxidation number.
  • Disproportionation: Same atom simultaneously oxidised and reduced.
• Oxidation Number Rules: Free elements = $0$. Ions = charge. $F$ is always $-1$. $H$ is $+1$ (except hydrides). $O$ is $-2$ (except peroxides).
• Electrolysis:
  • Anode ($+$): Oxidation occurs.
  • Cathode ($-$): Reduction occurs.
  • Extraction of Al: Alumina in cryolite ($Na_3AlF_6$) to lower MP and increase conductivity.
  • Brine: Produces $Cl_2$ (anode), $H_2$ (cathode), and $NaOH$ (electrolyte).
  • Purification of Cu: Pure Cu cathode, impure Cu anode; sludge forms at bottom.

CHAPTER 8: CHEMICAL EQUILIBRIA

• Dynamic Equilibrium: Rates of forward/reverse reactions are equal; concentrations stay constant in a closed system.
• Le Chatelier's Principle: System opposes external change.
  • Concentration: Adding reactant shifts equilibrium to products.
  • Pressure: Increase shifts to side with fewer gas molecules.
  • Temperature: Increase favours endothermic direction.
  • Catalyst: No change in position; speeds up reaching equilibrium.
• Equilibrium Constants:
  • $K_c = \frac{[C]^c [D]^d}{[A]^a [B]^b}$
  • $K_p = \frac{(pC)^c (pD)^d}{(pA)^a (pB)^b}$
  • Note: Only temperature affects the numerical value of $K_c$ or $K_p$.
• Acids and Bases (Brønsted-Lowry): Acid is a proton donor; Base is a proton acceptor. Strong acids/bases dissociate completely; weak ones partially.

CHAPTER 9: RATE OF REACTION

• Rate: Change in concentration per unit time. Gradient of conc-time graph.
• Collision Theory: Particles must collide with correct orientation and energy $\ge E_a$ (Activation Energy).
• Factors:
  • Surface Area: More area = higher frequency of collisions.
  • Concentration/Pressure: More particles per volume = higher frequency.
  • Temperature: Increase in temperature increases average $K.E.$. Maxwell-Boltzmann curve shifts right; more particles exceed $E_a$.
• Catalysts: Provide alternative route with lower $E_a$. Not used up. enzymes are biological protein catalysts using 'lock-and-key' mechanism.

CHAPTER 10-13: PERIODICITY & GROUPS

• Period 3 Trends:
  • Radius: Atomic radius decreases across (nuclear charge increases, shielding constant).
  • Oxides: Basic ($Na, Mg$) $\rightarrow$ Amphoteric ($Al$) $\rightarrow$ Acidic ($Si, P, S$).
  • Chlorides: Ionic chlorides dissolve ($NaCl, MgCl_2$); covalent halides hydrolyse ($SiCl_4, PCl_5$) producing $HCl$ fumes.
• Group II: Atomic radius increases down. Reactivity with water increases down. Thermal stability of carbonates/nitrates increases down (lower polarising power of larger cations).
• Group VII (Halogens): Volatility decreases down. Oxidising power decreases down ($F_2$ is strongest). Bond enthalpy decreases down (except $F_2$ due to lone pair repulsion). $\text{Chlorine} + \text{Cold NaOH} \rightarrow Cl^-, ClO^-$ (disproportionation).
• Nitrogen: Very unreactive due to strong triple bond ($944\,kJ\,mol^{-1}$). Haber Process: $N_2 + 3H_2 \rightleftharpoons 2NH_3$ ($450\,^{\circ}C, 200\,atm, \text{Fe catalyst}$).
• Sulfur: Contact Process: $2SO_2 + O_2 \rightleftharpoons 2SO_3$ ($V_2O_5\, \text{catalyst}$). $SO_3$ in $H_2SO_4 \rightarrow \text{oleum}$, then water to form concentrated acid.

CHAPTER 14-19: ORGANIC CHEMISTRY

• Formulae: Structural, Displayed, Skeletal. IUPAC naming rules.
• Isomerism:
  • Structural: Chain, Positional, Functional.
  • Stereoisomerism: Geometrical (cis-trans due to restricted bond rotation); Optical (chiral center with 4 different groups, non-superimposable mirror images).
• Mechanisms:
  • Free-radical substitution (Alkanes): Initiation (UV homolytic fission), Propagation, Termination.
  • Electrophilic Addition (Alkenes): Markovnikov's rule (H attaches to carbon with most H atoms for carbocation stability). Reactions: Hydrogenation ($Ni$), Hydration ($H_3PO_4$), Bromination (test for $\text{C=C}$).
  • Nucleophilic Substitution (Halogenoalkanes): $S_N 1$ (tertiary, via carbocation) and $S_N 2$ (primary, via transition state). Hydrolysis rate: $I > Br > Cl$.
  • Alcohol Reactions: Oxidation ($K_2Cr_2O_7$: Primary $\rightarrow$ Aldehyde/Acid; Secondary $\rightarrow$ Ketone). Esterification with carboxylic acid (acid catalyst).
• Carbonyls: Tested with 2,4-DNPH (orange ppt). Aldehydes only: Tollens' (silver mirror) or Fehling's (red ppt).
• Carboxylic Acids: Weak acids. React with metals, bases, and carbonates (produces $CO_2$). Form acyl chlorides with $PCl_5, PCl_3, \text{or } SOCl_2$.
• Esters: Sweet-smelling. Solvent uses. Hydrolysis (Acidic $\rightarrow \text{reforms acid/alcohol}$; Alkaline $\rightarrow \text{reforms salt/alcohol}$).