Comprehensive Cambridge IGCSE & O Level Complete Chemistry Study Guide

Syllabus Architecture and Examination Structure

• The curriculum is structured into two academic pathways:
  • Core syllabus content: targeted for foundational mastery; does not include extended syllabus extensions.
  • Extended syllabus content: comprises the complete core syllabus supplemented by extended material.
• Assessment Objectives (AOs) govern examination standards across all papers:
  • AO1 (Knowledge with understanding): Accounts for approximately $50\%$ of the total examination mark. Assesses recall, scientific terminology, and conceptual comprehension.
  • AO2 (Information handling and problem solving): Accounts for approximately $30\%$ of the total examination mark. Assesses candidate performance in unfamiliar contexts, graphical analysis, data evaluation, and mathematical calculations.
  • AO3 (Experimental skills and investigations): Accounts for approximately $20\%$ of the total examination mark. Assesses practical procedures, observation recording, experimental design, and error assessment.
• Examination components for Core syllabus candidates:
  • Paper 1: Multiple Choice format, duration of $45\,\text{minutes}$, comprising $40$ four-choice items, totaling $40\,\text{marks}$ ($30\%$ of total marks). Assesses AO1 and AO2.
  • Paper 3: Core Theory format, duration of $1\,\text{hour}\,15\,\text{minutes}$, comprising short-answer and structured questions, totaling $80\,\text{marks}$ ($50\%$ of total marks). Assesses AO1 and AO2.
• Examination components for Extended syllabus candidates:
  • Paper 2: Multiple Choice format, duration of $45\,\text{minutes}$, comprising $40$ four-choice items, totaling $40\,\text{marks}$ ($30\%$ of total marks). Assesses AO1 and AO2.
  • Paper 4: Extended Theory format, duration of $1\,\text{hour}\,15\,\text{minutes}$ ($75\,\text{minutes}$), comprising short-answer and structured questions, totaling $80\,\text{marks}$ ($50\%$ of total marks). Assesses AO1 and AO2.
• Practical assessment components (applicable to both Core and Extended pathways):
  • Paper 5: Practical Test, duration of $1\,\text{hour}\,15\,\text{minutes}$, involving hands-on laboratory experimentation, totaling $40\,\text{marks}$ ($20\%$ of total marks). Assesses AO3.
  • Paper 6: Alternative to Practical, duration of $1\,\text{hour}\,15\,\text{minutes}$ ($75\,\text{minutes}$), written examination testing experimental competence without physical lab execution, totaling $40\,\text{marks}$ ($20\%$ of total marks). Assesses AO3.

States of Matter and Kinetic Particle Theory

• The fundamental premise of chemistry posits that all matter is composed of submicroscopic entities called particles.
• Empirical evidence supporting the existence and motion of particles:
  • Robert Brown (1827) observed pollen grains suspended in water exhibiting erratic, continuous jiggly movement (Brownian motion) under a microscope.
  • Albert Einstein (1905) mathematically explained this phenomenon: the pollen grains move because they are continuously bombarded by invisible, rapidly moving water particles.
  • Dust specks dancing in sunlit rooms result from constant collisions with energetic gas particles in air.
  • Dissolution of potassium manganate(VII): Placing a purple crystal of $\text{KMnO}_4$ at the bottom of a water beaker causes the purple colour to slowly detach and disperse evenly throughout the liquid without mechanical stirring.
  • Bromine vapour diffusion: When an open gas jar containing a few drops of dense, red-brown liquid bromine ($\text{Br}_2$) is inverted below an open jar of air, the colour migrates upwards against gravity until both jars are homogeneously mixed.
• Three fundamental particle categories:
  • Atoms: The smallest indivisible particles of an element that cannot be broken down further by chemical reactions (e.g., argon gas exists as isolated monoatomic particles).
  • Molecules: Stable entities formed of two or more atoms chemically bonded together (e.g., water, $\text{H}_2\text{O}$; bromine, $\text{Br}_2$; oxygen, $\text{O}_2$; nitrogen, $\text{N}_2$).
  • Ions: Electrically charged particles resulting from the loss or gain of electrons by atoms or groups of atoms (e.g., potassium ions and manganate(VII) ions in $\text{KMnO}_4$).
  • Visualization: Optical light microscopes cannot resolve individual atoms; electron microscopes, which utilize accelerated beams of electrons instead of light waves, can produce direct structural micrographs.
• Macroscopic distinguishing properties of the three physical states:
  • Solids: Possess a fixed shape, fixed volume, cannot flow, and cannot be compressed into a smaller space.
  • Liquids: Possess no fixed shape (adopt the geometry of the bottom of the container), possess a fixed volume, flow readily, and exhibit negligible compressibility.
  • Gases: Possess no fixed shape and no fixed volume (expand completely to fill any enclosing boundary), have low density relative to solids and liquids, and are readily compressed.
• Thermophysical transition temperatures for selected substances:
  • Oxygen: Melting point of $-219\,^\circ\text{C}$, boiling point of $-183\,^\circ\text{C}$; exists as a gas at room temperature ($20\,^\circ\text{C}$).
  • Mercury: Melting point of $-39\,^\circ\text{C}$, boiling point of $357\,^\circ\text{C}$; exists as a liquid metal at room temperature.
  • Ethanol: Melting point of $-15\,^\circ\text{C}$, boiling point of $78\,^\circ\text{C}$; exists as a volatile liquid at room temperature.
  • Water: Melting point of $0\,^\circ\text{C}$, boiling point of $100\,^\circ\text{C}$; exists as a liquid at room temperature.
  • Sodium: Melting point of $98\,^\circ\text{C}$, boiling point of $890\,^\circ\text{C}$; exists as a solid metal at room temperature.
  • Table salt (Sodium chloride): Melting point of $801\,^\circ\text{C}$, boiling point of $1465\,^\circ\text{C}$; exists as an ionic solid at room temperature.
  • Iron: Melting point of $1540\,^\circ\text{C}$, boiling point of $2900\,^\circ\text{C}$; exists as a strong structural metal at room temperature.
  • Diamond: Melting point of $3550\,^\circ\text{C}$, boiling point of $4832\,^\circ\text{C}$; giant covalent macromolecule.
• Microscopic particle arrangements, forces, and kinetic dynamics:
  • Solid state: Arranged in a highly ordered, regular three-dimensional lattice. Held by strong interparticle forces of attraction. Particles cannot undergo translational motion; they vibrate to and fro about fixed equilibrium positions.
  • Liquid state: Particles are closely packed but not in a rigid lattice. Interparticle forces are weaker than in solids. Particles possess sufficient kinetic energy to translate, slide, and roll over one another, allowing fluidity.
  • Gas state: Particles are separated by vast intermolecular distances. Interparticle forces of attraction are virtually non-existent. Particles move at high velocities in random straight lines, colliding elastically with one another and the container walls.
• Kinetic particle mechanisms of phase changes:
  • Melting: Thermal energy increases particle vibrational amplitude. At the melting point, kinetic energy overcomes the lattice binding forces, collapsing the rigid order into a fluid state.
  • Boiling: Rapid thermal vaporisation occurring throughout the bulk liquid at a fixed temperature where vapor pressure equals external pressure; kinetic energy fully overcomes intermolecular forces.
  • Evaporation: Surface phenomenon occurring across a broad temperature range below the boiling point; only surface particles with kinetic energy in the upper tail of the Maxwell-Boltzmann distribution escape intermolecular attraction into the gaseous phase.
  • Condensation: Cooling diminishes gas kinetic energy. As particles slow down, attractive forces capture them upon collision, forming a liquid.
  • Freezing (Solidification): Further cooling causes liquid particles to lose sufficient kinetic energy such that attractive forces lock them into fixed lattice coordinates.
• Thermodynamic interpretations of heating and cooling curves:
  • Pure substances produce sharp, horizontal phase transition plateaus at invariant temperatures (melting point and boiling point).
  • During phase transitions, heat energy absorbed is utilized solely to break attractive bonds (latent heat) rather than increasing kinetic energy; hence, temperature remains constant.
  • During cooling curves, latent heat of fusion and vaporisation is liberated back to the surroundings as bonds form, balancing heat loss and arresting temperature decline until phase change completion.
  • Presence of impurities disrupts lattice homogeneity, depressing the melting point, elevating the boiling point, and causing phase transitions to occur across a broad temperature interval (yielding slanted rather than horizontal plateaus).
• Gaseous behavior and pressure relations:
  • Gas pressure is generated by the continuous physical bombardment of gas particles against container boundaries, transferring momentum per unit surface area.
  • Effect of temperature at constant pressure: Increasing temperature increases average kinetic energy and velocity of gas particles; collisions become more frequent and forceful, forcing a movable boundary (e.g., piston) outward to maintain pressure equilibrium ($V \propto T$).
  • Effect of pressure at constant temperature: Compressing a gas into a smaller volume decreases spatial separation, causing particles to collide with container walls more frequently, thereby increasing internal pressure ($P \propto \frac{1}{V}$).
  • Liquefaction via compression: Excessive external pressure forces gas particles into such close proximity that intermolecular forces dominate, converting the gas into a liquid.
• Kinetic explanation of diffusion and Graham's Law principles:
  • Diffusion is the net movement of particles from a region of higher concentration to one of lower concentration via random thermal collisions.
  • Relative molecular mass ($M_r$) inversely dictates rate of diffusion:     $\text{Rate} \propto \frac{1}{\sqrt{M_r}}$
  • Laboratory confirmation via the diffusion tube experiment:
  • Cotton wool soaked in concentrated aqueous ammonia ($\text{NH}_3$) is inserted at tube end A; cotton wool soaked in concentrated hydrochloric acid ($\text{HCl}$) is inserted at tube end B.
  • Molar masses: Ammonia ($\text{NH}_3$) has $M_r = 14 + (3 \times 1) = 17$; hydrogen chloride ($\text{HCl}$) has $M_r = 1 + 35.5 = 36.5$.
  • Gaseous particles volatilize and diffuse toward each other, colliding with atmospheric nitrogen and oxygen molecules along their paths.
  • When the two gases meet, they react to form a white ring of solid ammonium chloride:       $\text{NH}_3\,(g) + \text{HCl}\,(g) \rightarrow \text{NH}_4\text{Cl}\,(s)$
  • The white ring precipitates closer to the $HCl$ plug (end B), demonstrating that lighter $\text{NH}_3$ molecules diffuse faster and travel further per unit time than heavier $\text{HCl}$ molecules.

Atomic Structure, Isotopes, and Historical Development

• The subatomic constituents of the atom:
  • Proton: Mass of $1\,\text{atomic mass unit (amu)}$, positive charge of $+1$, situated in the nucleus.
  • Neutron: Mass of $1\,\text{atomic mass unit (amu)}$, uncharged ($0$), situated in the nucleus.
  • Electron: Mass of $0.0005\,\text{amu}$ (negligible, taken as $0$), negative charge of $-1$, situated in quantized concentric shells orbiting the nucleus.
• Spatial geometry: The nucleus is extremely dense and minuscule relative to the total atomic volume; if an atom were scaled to the size of a football stadium, its nucleus would occupy the dimensions of a single pea.
• Fundamental atomic indices:
  • Proton number (Atomic number, $Z$): The quantity of protons in the nucleus of an atom; defines the chemical identity of the element.
  • Nucleon number (Mass number, $A$): The combined total quantity of protons and neutrons in the nucleus of an atom.
  • Standard nuclear notation:     $^A_Z\text{X}$
  • Calculation of neutron count:     $\text{Neutrons} = A - Z$
  • Electrostatic neutrality: In any neutral atom, the number of negatively charged electrons is strictly identical to the number of positively charged protons, neutralizing the net charge to $0$.
• Subatomic composition and nuclear notation of the first $20$ elements:
  • Hydrogen ($\text{H}$): $Z = 1$, $1\,\text{electron}$, $0\,\text{neutrons}$, $A = 1$; notation $^1_1\text{H}$
  • Helium ($\text{He}$): $Z = 2$, $2\,\text{electrons}$, $2\,\text{neutrons}$, $A = 4$; notation $^4_2\text{He}$
  • Lithium ($\text{Li}$): $Z = 3$, $3\,\text{electrons}$, $4\,\text{neutrons}$, $A = 7$; notation $^7_3\text{Li}$
  • Beryllium ($\text{Be}$): $Z = 4$, $4\,\text{electrons}$, $5\,\text{neutrons}$, $A = 9$; notation $^9_4\text{Be}$
  • Boron ($\text{B}$): $Z = 5$, $5\,\text{electrons}$, $6\,\text{neutrons}$, $A = 11$; notation $^{11}_5\text{B}$
  • Carbon ($\text{C}$): $Z = 6$, $6\,\text{electrons}$, $6\,\text{neutrons}$, $A = 12$; notation $^{12}_6\text{C}$
  • Nitrogen ($\text{N}$): $Z = 7$, $7\,\text{electrons}$, $7\,\text{neutrons}$, $A = 14$; notation $^{14}_7\text{N}$
  • Oxygen ($\text{O}$): $Z = 8$, $8\,\text{electrons}$, $8\,\text{neutrons}$, $A = 16$; notation $^{16}_8\text{O}$
  • Fluorine ($\text{F}$): $Z = 9$, $9\,\text{electrons}$, $10\,\text{neutrons}$, $A = 19$; notation $^{19}_9\text{F}$
  • Neon ($\text{Ne}$): $Z = 10$, $10\,\text{electrons}$, $10\,\text{neutrons}$, $A = 20$; notation $^{20}_{10}\text{Ne}$
  • Sodium ($\text{Na}$): $Z = 11$, $11\,\text{electrons}$, $12\,\text{neutrons}$, $A = 23$; notation $^{23}_{11}\text{Na}$
  • Magnesium ($\text{Mg}$): $Z = 12$, $12\,\text{electrons}$, $12\,\text{neutrons}$, $A = 24$; notation $^{24}_{12}\text{Mg}$
  • Aluminium ($\text{Al}$): $Z = 13$, $13\,\text{electrons}$, $14\,\text{neutrons}$, $A = 27$; notation $^{27}_{13}\text{Al}$
  • Silicon ($\text{Si}$): $Z = 14$, $14\,\text{electrons}$, $14\,\text{neutrons}$, $A = 28$; notation $^{28}_{14}\text{Si}$
  • Phosphorus ($\text{P}$): $Z = 15$, $15\,\text{electrons}$, $16\,\text{neutrons}$, $A = 31$; notation $^{31}_{15}\text{P}$
  • Sulfur ($\text{S}$): $Z = 16$, $16\,\text{electrons}$, $16\,\text{neutrons}$, $A = 32$; notation $^{32}_{16}\text{S}$
  • Chlorine ($\text{Cl}$): $Z = 17$, $17\,\text{electrons}$, $18\,\text{neutrons}$, $A = 35$; notation $^{35}_{17}\text{Cl}$
  • Argon ($\text{Ar}$): $Z = 18$, $18\,\text{electrons}$, $22\,\text{neutrons}$, $A = 40$; notation $^{40}_{18}\text{Ar}$
  • Potassium ($\text{K}$): $Z = 19$, $19\,\text{electrons}$, $20\,\text{neutrons}$, $A = 39$; notation $^{39}_{19}\text{K}$
  • Calcium ($\text{Ca}$): $Z = 20$, $20\,\text{electrons}$, $20\,\text{neutrons}$, $A = 40$; notation $^{40}_{20}\text{Ca}$
• Electronic configurations of the first $20$ elements:
  • Filling rules: Shell $1$ holds a maximum of $2\,\text{electrons}$; Shell $2$ holds a maximum of $8\,\text{electrons}$; Shell $3$ holds a maximum of $8\,\text{electrons}$ when functioning as the outer shell (can hold up to $18$ in transition elements); Shell $4$ receives electrons subsequent to Shell $3$ reaching $8$.
  • Configurations: $\text{H}: 1$; $\text{He}: 2$; $\text{Li}: 2,1$; $\text{Be}: 2,2$; $\text{B}: 2,3$; $\text{C}: 2,4$; $\text{N}: 2,5$; $\text{O}: 2,6$; $\text{F}: 2,7$; $\text{Ne}: 2,8$; $\text{Na}: 2,8,1$; $\text{Mg}: 2,8,2$; $\text{Al}: 2,8,3$; $\text{Si}: 2,8,4$; $\text{P}: 2,8,5$; $\text{S}: 2,8,6$; $\text{Cl}: 2,8,7$; $\text{Ar}: 2,8,8$; $\text{K}: 2,8,8,1$; $\text{Ca}: 2,8,8,2$.
  • Post-calcium shell allocation (e.g., Rubidium, $Z = 37$, Period 5, Group I): $2,8,18,8,1$.
• Periodic structure correlates:
  • Period number indicates the total count of occupied electron shells.
  • Group number (I to VII) equals the quantity of electrons located within the valence (outermost) shell.
  • Group VIII elements possess completely filled valence shells ($\text{He}: 2$; $\text{Ne}, \text{Ar}: 8$), conferring unreactive, inert chemical stability.
• Isotopes and Relative Atomic Mass ($A_r$):
  • Isotopes are atoms of the identical element containing equal quantities of protons, but disparate quantities of neutrons.
  • Chemical identity and reactivity remain identical among isotopes because they possess identical electronic configurations and valence shell occupancy.
  • Physical properties (e.g., mass, density, rate of diffusion) exhibit slight variations.
  • Carbon isotopic family: Carbon-12 ($^{12}_6\text{C}$, $6\,\text{protons}$, $6\,\text{neutrons}$); Carbon-13 ($^{13}_6\text{C}$, $6\,\text{protons}$, $7\,\text{neutrons}$); Carbon-14 ($^{14}_6\text{C}$, $6\,\text{protons}$, $8\,\text{neutrons}$, radioactive).
  • Radioisotopes possess unstable nuclei that spontaneously decompose via nuclear emission, yielding radiation and heat energy (e.g., Uranium-235 ($^{235}_{92}\text{U}$) undergoes nuclear fission to drive steam turbines for electrical generation).
  • Carbon-12 standard: The unified atomic mass scale defines an atom of carbon-12 as having an exact mass of $12\,\text{atomic mass units}$; one unit equals $\frac{1}{12}$ the mass of a single carbon-12 atom.
  • Definition of Relative Atomic Mass ($A_r$): The weighted average mass of naturally occurring isotopes of an element relative to $\frac{1}{12}$ of the mass of an atom of carbon-12:     $A_r = \sum \left( \frac{\text{Isotope Mass} \times \text{Percentage Abundance}}{100} \right)$
  • Worked calculation for elemental chlorine:     $A_r(\text{Cl}) = \left( \frac{35 \times 75}{100} \right) + \left( \frac{37 \times 25}{100} \right) = 26.25 + 9.25 = 35.5$
  • Worked calculation for boron: Given $80.1\%$ of $^{11}\text{B}$ and $19.9\%$ of $^{10}\text{B}$:     $A_r(\text{B}) = \left( \frac{11 \times 80.1}{100} \right) + \left( \frac{10 \times 19.9}{100} \right) = 8.811 + 1.990 = 10.8$
• Historical evolution of the atomic model:
  • Democritus ($\sim 750\text{--}150\,\text{BCE}$): Formulated the concept of indivisible particles termed atomos ("uncuttable"), attributing material properties to particle shapes and textures.
  • Persian alchemist Geber (Abu Musa Jabir ibn Hayyan, $721\text{--}815\,\text{CE}$): Developed foundational laboratory purification techniques (crystallisation, distillation) and quantitative reactant measurement.
  • Antoine Lavoisier ($1743\text{--}1794$): Established the Law of Conservation of Mass; experimentally proved mass remains invariant during chemical reactions.
  • J.J. Thomson ($1897$): Investigated cathode rays in evacuated glass tubes, discovering the electron as a subatomic particle possessing negative charge.
  • Henri Becquerel ($1896$): Discovered radioactivity from uranium salts exposing photographic plates in darkness.
  • Ernest Rutherford ($1911$): Alpha particle scattering experiment directed positively charged alpha particles ($^4_2\text{He}^{2+}$, $7000\,\text{times}$ heavier than electrons) at thin gold foil. Most passed unhindered, but a fraction deflected at acute angles and bounced back, demonstrating that the atom is mostly empty space with a dense, positively charged nucleus containing protons.
  • Niels Bohr ($1913$): Proposed quantized planetary orbits, demonstrating that electrons circle the nucleus at defined energy levels (shells) without collapsing into the positive core.
  • Walther Bothe and Herbert Becker ($1930$): Bombarded beryllium with alpha particles, discovering penetrating, uncharged neutral radiation subsequently characterized as neutrons.

Chemical Bonding, Molecular Geometry, and Material Structures

• Distinctions between physical and chemical processes:
  • Physical changes: Involve changes of state, dissolving, or mechanical subdivision. No novel chemical substances form. Intermolecular bonds may alter, but intramolecular bonds remain intact. Generally facile to reverse.
  • Chemical changes (Chemical reactions): Atomic rearrangements resulting in the formation of one or more chemically distinct substances. Characterized by bond breaking and bond making, significant enthalpy changes (exothermic or endothermic), and difficulty in reversal.
  • Classical synthesis of iron(II) sulfide:
  • Mixing grey iron powder ($\text{Fe}$) and yellow elemental sulfur ($\text{S}$) constitutes a physical change; sulfur can be extracted by dissolving in methylbenzene followed by filtration.
  • Heating the mixture initiates an exothermic reaction with persistent incandescence, producing a non-magnetic, solvent-insoluble black compound:       $\text{Fe}\,(s) + \text{S}\,(s) \rightarrow \text{FeS}\,(s)$
• Underlying cause of bonding: Atoms form chemical bonds to achieve electronic stability characterized by a full outer shell of electrons (isoelectronic to Group VIII noble gases; a stable octet, or duet for hydrogen/helium).
• Ionic bonding:
  • Occurs between metallic elements (low electronegativity/electron donors) and non-metallic elements (high electronegativity/electron acceptors).
  • Ionic bond definition: The strong, non-directional electrostatic attraction between oppositely charged cations and anions.
  • Electron transfer mechanisms:
  • Formation of Sodium Chloride ($\text{NaCl}$):       $\text{Na}\,(2,8,1) \rightarrow \text{Na}^+\,([2,8]^+) + e^-$$\text{Cl}\,(2,8,7) + e^- \rightarrow \text{Cl}^-\,([2,8,8]^-)$
  • Formation of Magnesium Oxide ($\text{MgO}$):       $\text{Mg}\,(2,8,2) \rightarrow \text{Mg}^{2+}\,([2,8]^{2+}) + 2e^-$$\text{O}\,(2,6) + 2e^- \rightarrow \text{O}^{2-}\,([2,8]^{2-})$
  • Formation of Magnesium Chloride ($\text{MgCl}_2$):       $\text{Mg}\,(2,8,2) \rightarrow \text{Mg}^{2+}\,([2,8]^{2+}) + 2e^-$$2\text{Cl}\,(2,8,7) + 2e^- \rightarrow 2\text{Cl}^-\,([2,8,8]^-)$
  • Giant ionic lattice: A regular, repeating three-dimensional array of alternating cations and anions maximizing electrostatic attraction while minimizing repulsive forces.
  • Valency and polyatomic ions:
  • Cations: Potassium ($\text{K}^+$), Lithium ($\text{Li}^+$), Calcium ($\text{Ca}^{2+}$), Barium ($\text{Ba}^{2+}$), Aluminium ($\text{Al}^{3+}$), Ammonium ($\text{NH}_4^+$).
  • Anions: Fluoride ($\text{F}^-$), Chloride ($\text{Cl}^-$), Bromide ($\text{Br}^-$), Iodide ($\text{I}^-$), Oxide ($\text{O}^{2-}$), Sulfide ($\text{S}^{2-}$), Hydroxide ($\text{OH}^-$), Nitrate ($\text{NO}_3^-$), Carbonate ($\text{CO}_3^{2-}$), Sulfate ($\text{SO}_4^{2-}$), Sulfite ($\text{SO}_3^{2-}$).
  • Transition metals exhibiting variable oxidation states: Copper(I) ($\text{Cu}^+$ in $\text{Cu}_2\text{O}$), Copper(II) ($\text{Cu}^{2+}$ in $\text{CuO}$), Iron(II) ($\text{Fe}^{2+}$ in $\text{FeCl}_2$), Iron(III) ($\text{Fe}^{3+}$ in $\text{FeCl}_3$).
• Covalent bonding:
  • Occurs when non-metallic elements share pairs of electrons to attain noble gas electron configurations.
  • Single covalent bond: One shared pair of electrons ($2\,\text{electrons}$ total).
  • Double covalent bond: Two shared pairs of electrons ($4\,\text{electrons}$ total).
  • Triple covalent bond: Three shared pairs of electrons ($6\,\text{electrons}$ total).
  • Bonding configurations in homonuclear diatomic non-metals:
  • Hydrogen molecule ($\text{H}_2$): Two hydrogen atoms share one pair of electrons, forming a single covalent bond ($\text{H}-\text{H}$).
  • Chlorine molecule ($\text{Cl}_2$): Two chlorine atoms share one pair of electrons, completing their octets ($\text{Cl}-\text{Cl}$).
  • Oxygen molecule ($\text{O}_2$): Two oxygen atoms share two pairs of electrons, forming a double bond ($\text{O}=\text{O}$).
  • Nitrogen molecule ($\text{N}_2$): Two nitrogen atoms share three pairs of electrons, forming a strong triple bond ($\text{N}\equiv\text{N}$).
  • Covalent simple molecular compounds:
  • Hydrogen chloride ($\text{HCl}$): One shared pair between hydrogen and chlorine ($\text{H}-\text{Cl}$).
  • Water ($\text{H}_2\text{O}$): Two single covalent bonds ($\text{H}-\text{O}-\text{H}$); oxygen retains two non-bonding lone pairs, generating a bent molecular geometry.
  • Ammonia ($\text{NH}_3$): Three single covalent bonds; nitrogen retains one non-bonding lone pair, generating a trigonal pyramidal geometry.
  • Methane ($\text{CH}_4$): Four single covalent bonds directed symmetrically in space; mutually repulsive electron pairs produce a tetrahedral geometry with bond angles of $109.5^\circ$.
  • Methanol ($\text{CH}_3\text{OH}$): Carbon bonds tetrahedrally to three hydrogens and one oxygen; oxygen bonds to one hydrogen.
  • Carbon dioxide ($\text{CO}_2$): Carbon forms two double bonds with two oxygen atoms ($\text{O}=\text{C}=\text{O}$); bonding electron pairs repel to opposite poles, generating a linear geometry.
  • Ethene ($\text{C}_2\text{H}_4$): Contains four single carbon-hydrogen bonds and one carbon-carbon double bond ($\text{H}_2\text{C}=\text{CH}_2$).
• Comparative properties of ionic and simple molecular compounds:
  • Melting and boiling points:
  • Ionic compounds possess high melting and boiling points (e.g., $\text{NaCl}$ melts at $801\,^\circ\text{C}$, $\text{MgO}$ melts at $2852\,^\circ\text{C}$) due to the large amount of thermal energy required to overcome strong omnidirectional electrostatic forces in the giant lattice. Doubling ionic charge ($\text{Mg}^{2+}$, $\text{O}^{2-}$ versus $\text{Na}^+$, $\text{Cl}^-$) quadruples electrostatic attraction, elevating transition temperatures.
  • Molecular covalent compounds possess low melting and boiling points (e.g., carbon monoxide melts at $-199\,^\circ\text{C}$, hexane melts at $-95\,^\circ\text{C}$) because only weak intermolecular forces (van der Waals forces) between discrete molecules need to be broken; covalent bonds inside the molecules remain intact.
  • Electrical conductivity:
  • Ionic compounds do not conduct electricity in the solid state because ions are fixed in lattice coordinates. They become good conductors when molten or dissolved in aqueous solution, where the lattice collapses, freeing ions to migrate toward electrodes.
  • Simple molecular compounds do not conduct electricity in solid, liquid, or dissolved states because they consist of uncharged neutral molecules containing no mobile delocalised electrons or ions.
  • Solubility profiles:
  • Ionic compounds generally dissolve in polar solvents (e.g., water) as polar molecules hydrate and stabilise dissociated ions.
  • Covalent compounds are typically insoluble in water but dissolve readily in non-polar organic solvents (e.g., tetrachloromethane, methylbenzene).
• Giant covalent structures (Macromolecules):
  • Diamond:
  • Allotropic form of pure carbon.
  • Structure: Each carbon atom is bonded to four other carbon atoms by strong covalent bonds in an extensive three-dimensional tetrahedral lattice.
  • Properties: Hardest naturally occurring substance; exceptionally high melting point ($3550\,^\circ\text{C}$); non-conductor of electricity due to absence of mobile electrons.
  • Industrial applications: Heavy-duty drill bits, glass cutters, abrasive stone-cutting saw blades.
  • Graphite:
  • Allotropic form of pure carbon.
  • Structure: Each carbon atom is covalently bonded to three other carbon atoms in flat hexagonal planar rings, forming two-dimensional sheets. The sheets are separated by weak interlayer forces. Each carbon atom contributes one valence electron to an extensive delocalised electron system traversing the layers.
  • Properties: Soft and slippery because planar layers slide easily over one another under shear stress; excellent electrical conductor because delocalised electrons drift freely along the layers.
  • Industrial applications: Solid lubricant for mechanical locks and high-temperature engines, inert electrodes in electrolysis cells, pencil cores (mixed with clay).
  • Silicon(IV) oxide (Silica / Quartz, $\text{SiO}_2$):
  • Primary chemical constituent of quartz and sand.
  • Structure: Giant three-dimensional macromolecular network wherein each silicon atom is tetrahedrally bonded to four oxygen atoms, and each oxygen atom bridges two silicon atoms.
  • Properties: High mechanical hardness, high melting point ($1610\text{--}1710\,^\circ\text{C}$), non-conductor of electricity.
  • Industrial applications: High-temperature refractory lining for furnaces, abrasive sandpaper.
• Metallic bonding and physical properties:
  • Structure: Metals crystallize in a giant lattice consisting of positive metal cations arranged in close-packed layers, surrounded by a shared sea of delocalised valence electrons.
  • Metallic bond definition: The electrostatic attraction between positive metal ions and the mobile sea of delocalised electrons.
  • Microstructure: Polycrystalline aggregate consisting of trillions of microscopic crystal grains joined together.
  • Mechanical properties (Malleability and Ductility):
  • Malleability denotes the capacity to be hammered or rolled into sheets without fracture; ductility denotes the capacity to be pulled into wires.
  • Mechanism: Applied mechanical shear stress causes orderly layers of metal ions to slide past one another. The non-directional sea of delocalised electrons dynamically reorganises around the displaced cations, maintaining electrostatic binding and preventing lattice cleavage.
  • Conduction properties:
  • Electrical conductivity: Delocalised electrons drift freely through the lattice toward the positive terminal when a potential difference is applied.
  • Thermal conductivity: Heating excites delocalised electrons, which rapidly transport kinetic energy through the lattice via high-velocity collisions, supplemented by lattice vibrational waves.

Chemical Stoichiometry and Quantitative Mass Relationships

• Nomenclature conventions for binary compounds:
  • Metal + Non-metal: Metal named first without modification; non-metal root appended with the suffix "-ide" (e.g., sodium chloride, calcium sulfide).
  • Non-metal + Non-metal: Hydrogen written first if present; otherwise, the element residing in the lower Periodic Table group precedes the higher group element; second element takes the "-ide" suffix (e.g., hydrogen chloride, carbon dioxide, silicon dioxide).
• Stoichiometric valency rules:
  • Group I: Valency = $1$
  • Group II: Valency = $2$
  • Group III: Valency = $3$
  • Group IV: Valency = $4$
  • Group V: Valency = $3$
  • Group VI: Valency = $2$
  • Group VII: Valency = $1$
  • Hydrogen: Valency = $1$
  • Valency balancing examples: Hydrogen sulfide ($\text{H}_2\text{S}$, valencies $2 \times 1$ vs $2$); Aluminium oxide ($\text{Al}_2\text{O}_3$, valencies $2 \times 3$ vs $3 \times 2$).
• Chemical equations and balancing:
  • State symbols define the thermodynamic phase of reactants and products: $(s)$ for solid, $(l)$ for liquid, $(g)$ for gas, $(aq)$ for aqueous solution.
  • Law of Conservation of Mass: Chemical equations must be stoichiometrically balanced because atoms cannot be created or destroyed during a chemical reaction, only rearranged.
  • Step-by-step balanced synthesis equations:
  • Calcium oxidation in chlorine:       $\text{Ca}\,(s) + \text{Cl}_2\,(g) \rightarrow \text{CaCl}_2\,(s)$
  • Hydrogen combustion in chlorine:       $\text{H}_2\,(g) + \text{Cl}_2\,(g) \rightarrow 2\text{HCl}\,(g)$
  • Magnesium combustion in oxygen:       $2\text{Mg}\,(s) + \text{O}_2\,(g) \rightarrow 2\text{MgO}\,(s)$
  • Hydrogen combustion in oxygen:       $2\text{H}_2\,(g) + \text{O}_2\,(g) \rightarrow 2\text{H}_2\text{O}\,(l)$
• Formula mass indices:
  • Relative Molecular Mass ($M_r$): Sum of the relative atomic masses of all atoms present in a single discrete molecule.
  • Water ($\text{H}_2\text{O}$): $M_r = (2 \times 1) + 16 = 18$
  • Ammonia ($\text{NH}_3$): $M_r = 14 + (3 \times 1) = 17$
  • Carbon dioxide ($\text{CO}_2$): $M_r = 12 + (2 \times 16) = 44$
  • Butane ($\text{C}_4\text{H}_{10}$): $M_r = (4 \times 12) + (10 \times 1) = 58$
  • Relative Formula Mass ($M_r$): Applied to giant ionic networks; equals the sum of the relative atomic masses of the atoms in the empirical formula unit.
  • Sodium chloride ($\text{NaCl}$): $M_r = 23 + 35.5 = 58.5$
  • Magnesium nitrate ($\text{Mg}(\text{NO}_3)_2$): $M_r = 24 + 2 \times [14 + (3 \times 16)] = 24 + 28 + 96 = 148$
  • Copper(II) sulfate pentahydrate ($\text{CuSO}_4\cdot 5\text{H}_2\text{O}$): $M_r = 64 + 32 + (4 \times 16) + 5 \times [(2 \times 1) + 16] = 64 + 32 + 64 + 90 = 250$
• Law of Constant Proportions and Reacting Mass Ratios:
  • Elements always combine in fixed mass ratios to synthesize specific compounds.
  • Reaction: Copper oxidation to black copper(II) oxide:     $2\text{Cu}\,(s) + \text{O}_2\,(g) \rightarrow 2\text{CuO}\,(s)$
  • Reacting proportions: $64\,\text{g}$ of copper reacts precisely with $16\,\text{g}$ of oxygen to produce $80\,\text{g}$ of copper(II) oxide.
  • Proportional scaling: $128\,\text{g}$ of copper reacts with $32\,\text{g}$ of oxygen; $6.4\,\text{g}$ of copper reacts with $1.6\,\text{g}$ of oxygen to yield $8.0\,\text{g}$ of product.
• Percentage composition calculations:
  • Formula for percentage composition by mass of an element within a compound:     $\%\,\text{Element} = \frac{\text{Number of Atoms of Element} \times A_r}{M_r} \times 100\%$
  • Calculation for sulfur dioxide ($\text{SO}_2$, $M_r = 64$):     $\%\,\text{O} = \frac{2 \times 16}{64} \times 100\% = 50\%$$\%\,\text{S} = \frac{32}{64} \times 100\% = 50\%$
  • Calculation for methane ($\text{CH}_4$, $M_r = 16$):     $\%\,\text{C} = \frac{12}{16} \times 100\% = 75\%$$\%\,\text{H} = \frac{4 \times 1}{16} \times 100\% = 25\%$
  • Calculation for calcium carbonate ($\text{CaCO}_3$, $M_r = 100$):     $\%\,\text{Ca} = \frac{40}{100} \times 100\% = 40\%$$\%\,\text{C} = \frac{12}{100} \times 100\% = 12\%$$\%\,\text{O} = \frac{3 \times 16}{100} \times 100\% = 48\%$
• Percentage purity calculations:
  • Formula for purity determination:     $\%\,\text{Purity} = \frac{\text{Mass of Pure Substance}}{\text{Total Mass of Impure Sample}} \times 100\%$
  • Sample processing: Refining $20\,\text{tonnes}$ of impure copper yields $18\,\text{tonnes}$ of pure copper:     $\%\,\text{Purity} = \frac{18}{20} \times 100\% = 90\%$
  • Industrial requirement: Electronic semiconductor applications require silicon wafers derived from silica to possess a chemical purity of at least $99.9999\%$.

The Mole Concept, Gas Stoichiometry, Solutions, and Chemical Analysis

• Fundamental definitions of the mole and the Avogadro constant:
  • A mole ($\text{mol}$) is the SI unit representing the amount of substance that contains as many elementary entities (atoms, molecules, ions, or formula units) as there are atoms in precisely $12\,\text{g}$ of carbon-12.
  • Avogadro constant ($N_A$): $6.02 \times 10^{23}\,\text{particles}\,\text{mol}^{-1}$.
  • Operational relationship between mass and moles:     $\text{Mass}\,(g) = \text{Moles}\,(n) \times \text{Molar Mass}\,(M)$$\text{Moles}\,(n) = \frac{\text{Mass}\,(g)}{\text{Molar Mass}\,(M)}$
• Stoichiometry and reacting masses from balanced equations:
  • The coefficients in a balanced equation represent the stoichiometric mole ratios of the reactants and products.
  • Stoichiometric calculation: Hydrogen combustion to yield water:     $2\text{H}_2\,(g) + \text{O}_2\,(g) \rightarrow 2\text{H}_2\text{O}\,(l)$$2\,\text{mol}\,\text{H}_2\,(4\,\text{g}) + 1\,\text{mol}\,\text{O}_2\,(32\,\text{g}) \rightarrow 2\,\text{mol}\,\text{H}_2\text{O}\,(36\,\text{g})$$1\,\text{g of}\;\text{H}_2\;\text{reacts with}\;\frac{32}{4} = 8\,\text{g of}\;\text{O}_2\;\text{to yield}\;\frac{36}{4} = 9\,\text{g of}\;\text{H}_2\text{O}$
  • Limiting reactant analysis:
  • The limiting reactant is completely consumed first, dictating the maximum theoretical yield of products.
  • Example: $1.4\,\text{g}$ of iron added to aqueous copper(II) sulfate produces $1.6\,\text{g}$ of copper:       $n(\text{Fe}) = \frac{1.4}{56} = 0.025\,\text{mol}$$n(\text{Cu}) = \frac{1.6}{64} = 0.025\,\text{mol}$$\text{Mole Ratio} = \frac{n(\text{Fe})}{n(\text{Cu})} = \frac{0.025}{0.025} = 1:1$$\text{Equation}: \text{Fe}\,(s) + \text{CuSO}_4\,(aq) \rightarrow \text{Cu}\,(s) + \text{FeSO}_4\,(aq)$
• Gas stoichiometry and Avogadro's Law:
  • Avogadro's Law (1811): Equal volumes of all gases under identical conditions of temperature and pressure contain equal numbers of molecules.
  • Room Temperature and Pressure ($\text{rtp}$): Defined as $20\,^\circ\text{C}$ ($293\,\text{K}$) and $1\,\text{atmosphere}$ ($101.3\,\text{kPa}$).
  • Molar gas volume: $1\,\text{mole}$ of any ideal gas occupies precisely $24\,\text{dm}^3$ ($24\,000\,\text{cm}^3$) at $\text{rtp}$.
  • Operational conversions for gases:     $\text{Volume at rtp}\,(\text{dm}^3) = n \times 24\,\text{dm}^3\,\text{mol}^{-1}$$\text{Volume at rtp}\,(\text{cm}^3) = n \times 24\,000\,\text{cm}^3\,\text{mol}^{-1}$
  • Sample gas calculation: Volume occupied by $22\,\text{g}$ of $\text{CO}_2$ ($M_r = 44$):     $n = \frac{22}{44} = 0.5\,\text{mol}$$\text{Volume} = 0.5 \times 24 = 12\,\text{dm}^3$
  • Combustion of sulfur:     $\text{S}\,(s) + \text{O}_2\,(g) \rightarrow \text{SO}_2\,(g)$$n(\text{S}) = \frac{1\,\text{g}}{32\,\text{g}\,\text{mol}^{-1}} = 0.03125\,\text{mol}$$V(\text{SO}_2) = 0.03125 \times 24 = 0.75\,\text{dm}^3 = 750\,\text{cm}^3$
• Solution stoichiometry and volumetric concentration:
  • Terminology: Solute is the dissolved chemical entity; solvent is the dissolving liquid medium; solution is the homogeneous mixture.
  • Concentration definitions and mathematical expressions:     $\text{Concentration}\,(\text{g}\,\text{dm}^{-3}) = \frac{\text{Mass of Solute}\,(g)}{\text{Volume of Solution}\,(\text{dm}^3)}$$\text{Concentration}\,(\text{mol}\,\text{dm}^{-3}) = \frac{\text{Moles of Solute}\,(n)}{\text{Volume of Solution}\,(\text{dm}^3)}$
  • Interconversion between concentration metrics:     $\text{Concentration}\,(\text{g}\,\text{dm}^{-3}) = \text{Concentration}\,(\text{mol}\,\text{dm}^{-3}) \times M_r$
  • Molar calculation: Mass of $\text{NaOH}$ in $25\,\text{cm}^3$ of $0.5\,\text{mol}\,\text{dm}^{-3}$ solution ($M_r = 40$):     $n = 0.5 \times \frac{25}{1000} = 0.0125\,\text{mol}$$\text{Mass} = 0.0125 \times 40 = 0.5\,\text{g}$
• Empirical and molecular formula determination:
  • Empirical formula: The simplest whole-number ratio of the atoms or ions of each element present in a chemical compound.
  • Molecular formula: The actual number of atoms of each element present in one molecule of a covalent compound.
  • Experimental crucible determination of magnesium oxide:
  1. Measure mass of empty crucible plus lid ($m_1$).
  2. Introduce coiled magnesium ribbon and measure combined mass ($m_2$); calculate mass of $\text{Mg} = m_2 - m_1$.
  3. Heat crucible strongly using a Bunsen burner, intermittently lifting the lid using tongs to admit oxygen while preventing escaping oxide smoke.
  4. Cool crucible with lid secured to ambient temperature, and record final mass ($m_3$); mass of oxygen reacting = $m_3 - m_2$.
  5. Empirical derivation: If $2.4\,\text{g of}\;\text{Mg}$ combines with $1.6\,\text{g of}\;\text{O}$:        $n(\text{Mg}) = \frac{2.4}{24} = 0.1\,\text{mol}$$n(\text{O}) = \frac{1.6}{16} = 0.1\,\text{mol}$$\text{Ratio} = 1:1 \implies \text{Empirical Formula} = \text{MgO}$
  • Molecular formula derivation protocol:     $n = \frac{\text{Relative Molecular Mass}\,(M_r)}{\text{Empirical Formula Mass}}$$\text{Molecular Formula} = (\text{Empirical Formula})_n$
  • Worked example (Hydrogen Peroxide): Empirical formula $\text{HO}$ (empirical mass $1 + 16 = 17$), experimental $M_r = 34$:     $n = \frac{34}{17} = 2 \implies \text{Molecular Formula} = \text{H}_2\text{O}_2$
  • Worked example (Octane): Combustion analysis gives $84.2\%\,\text{C}$ and $15.8\%\,\text{H}$, $M_r = 114$:     $n(\text{C}) = \frac{84.2}{12} = 7.02\,\text{mol}$$n(\text{H}) = \frac{15.8}{1} = 15.8\,\text{mol}$$\text{Ratio} = \frac{15.8}{7.02} = 2.25 = \frac{9}{4} \implies \text{Empirical Formula} = \text{C}_4\text{H}_9$$\text{Empirical Mass} = (4 \times 12) + (9 \times 1) = 57$$n = \frac{114}{57} = 2 \implies \text{Molecular Formula} = \text{C}_8\text{H}_{18}$
• Percentage yield and purity calculations:
  • Formula for percentage yield:     $\%\,\text{Yield} = \frac{\text{Actual Mass of Product Obtained}}{\text{Theoretical Maximum Mass Calculated}} \times 100\%$
  • Pharmaceutical synthesis of aspirin from salicylic acid:     $\text{C}_7\text{H}_6\text{O}_3 + \text{reagents} \rightarrow \text{C}_9\text{H}_8\text{O}_4$$M_r(\text{Salicylic acid}) = 138, \quad M_r(\text{Aspirin}) = 180$$100.0\,\text{g of salicylic acid} = \frac{100}{138} = 0.7246\,\text{mol}$$\text{Theoretical Mass of Aspirin} = 0.7246 \times 180 = 130.43\,\text{g}$$\text{Actual Mass Recovered} = 121.2\,\text{g}$$\%\,\text{Yield} = \frac{121.2}{130.43} \times 100\% = 92.9\%$
  • Quantitative volumetric analysis of synthesized aspirin purity:     $\text{C}_9\text{H}_8\text{O}_4\,(aq) + \text{NaOH}\,(aq) \rightarrow \text{C}_9\text{H}_7\text{O}_4\text{Na}\,(aq) + \text{H}_2\text{O}\,(l)$$4.00\,\text{g sample titrated with}\;17.5\,\text{cm}^3\;\text{of}\;1.0\,\text{mol}\,\text{dm}^{-3}\,\text{NaOH}$$n(\text{NaOH}) = 1.0 \times \frac{17.5}{1000} = 0.0175\,\text{mol}$$n(\text{Pure Aspirin}) = 0.0175\,\text{mol}$$\text{Mass of Pure Aspirin} = 0.0175 \times 180 = 3.15\,\text{g}$$\%\,\text{Purity} = \frac{3.15}{4.00} \times 100\% = 78.75\%$
  • Quantitative purity assessment of chalk via gas collection:     $\text{CaCO}_3\,(s) + 2\text{HCl}\,(aq) \rightarrow \text{CaCl}_2\,(aq) + \text{H}_2\text{O}\,(l) + \text{CO}_2\,(g)$$10.0\,\text{g chalk sample yields}\;2280\,\text{cm}^3\;\text{of}\;\text{CO}_2\;\text{at rtp}$$n(\text{CO}_2) = \frac{2280}{24\,000} = 0.095\,\text{mol}$$n(\text{Pure}\;\text{CaCO}_3) = 0.095\,\text{mol}$$\text{Mass of Pure}\;\text{CaCO}_3 = 0.095 \times 100 = 9.5\,\text{g}$$\%\,\text{Purity} = \frac{9.5}{10.0} \times 100\% = 95.0\%$

The Periodic Table: Structure, Trends, and Element Groups

• Architecture and governing principles:
  • Elements are arranged in sequential order of increasing atomic (proton) number ($Z$).
  • Periodicity describes the systematic recurrence of elements with similar chemical and physical properties at fixed numerical intervals.
  • Columns represent Groups (numbered I through VIII): Elements within a group share the same valence shell electron count, conferring analogous chemical reactivities.
  • Rows represent Periods (numbered 1 through 7): Period number denotes the quantity of occupied electron energy levels (shells).
  • The stepped zig-zag boundary line separates metallic elements on the left from non-metallic elements on the right; approximately $80\%$ of all elements are metals.
  • Hydrogen ($Z = 1$) sits isolated: possesses $1$ valence electron and forms the $H^+$ cation like Group I alkali metals, yet exists as a diatomic non-metallic gas ($\text{H}_2$) exhibiting non-metal reaction patterns.
  • Natural versus synthetic elements: $94$ of the $118$ known elements occur naturally on Earth; the remaining elements are laboratory synthesised, highly unstable radioisotopes localized predominantly within Period 7.
• Historical evolution of the Periodic Table:
  • John Newlands (1865): Formulated the Law of Octaves upon arranging the $56$ known elements in order of increasing atomic weight (relative atomic mass), noting that chemical properties repeated every eighth element (e.g., lithium, sodium, potassium). The model failed and was rejected by the Chemical Society because transition metals forced dissimilar elements into identical groups (e.g., placing unreactive copper with reactive sodium; grouping didymium, which was later found to be a mixture).
  • Dmitri Ivanovich Mendeleev (1869): Published the foundational Periodic Table arranging $63$ known elements by atomic weight and chemical properties. Crucially, Mendeleev left deliberate vacancies for undiscovered elements and accurately predicted their atomic weights and properties:
  • Eka-boron (discovered as Scandium, 1879)
  • Eka-aluminium (discovered as Gallium, 1875)
  • Eka-silicon (discovered as Germanium, 1886)
  • Anomalous mass orderings corrected by Mendeleev: Mendeleev inverted the positions of tellurium and iodine, and argon and potassium, placing potassium ($A_r = 39$) after argon ($A_r = 40$) to align with alkali metal reactivity. The discovery of the nuclear proton in 1911 validated these inversions, confirming proton number ($Z$) as the true organizing metric.
• Group I: The Alkali Metals:
  • Elements: Lithium ($\text{Li}$), Sodium ($\text{Na}$), Potassium ($\text{K}$), Rubidium ($\text{Rb}$), Caesium ($\text{Cs}$), Francium ($\text{Fr}$, unstable/radioactive).
  • Physical properties: Exceptionally soft silvery metals (can be cut with a laboratory scalpel), low mechanical densities (lithium, sodium, and potassium float on water), low melting points relative to typical structural metals.
  • Measured physical property trends:
  • Lithium ($\text{Li}$): Density = $0.53\,\text{g}\,\text{cm}^{-3}$, melting point = $181\,^\circ\text{C}$
  • Sodium ($\text{Na}$): Density = $0.97\,\text{g}\,\text{cm}^{-3}$, melting point = $98\,^\circ\text{C}$
  • Potassium ($\text{K}$): Density = $0.86\,\text{g}\,\text{cm}^{-3}$, melting point = $63\,^\circ\text{C}$
  • Rubidium ($\text{Rb}$): Density = $1.53\,\text{g}\,\text{cm}^{-3}$, melting point = $39\,^\circ\text{C}$
  • Caesium ($\text{Cs}$): Density = $1.88\,\text{g}\,\text{cm}^{-3}$, melting point = $29\,^\circ\text{C}$
  • Systematic trends descending Group I:
  • Density generally increases (potassium exhibits a minor anomaly, being less dense than sodium).
  • Melting points systematically decrease.
  • Chemical reactivity dramatically increases.
  • Storage: Maintained submerged under paraffin oil to exclude atmospheric oxygen and moisture.
  • Chemical reactions:
  • Reaction with water: Exothermic reduction yielding an aqueous alkali metal hydroxide and effervescing hydrogen gas:       $2\text{M}\,(s) + 2\text{H}_2\text{O}\,(l) \rightarrow 2\text{MOH}\,(aq) + \text{H}_2\,(g)$
    • Lithium: Floats, fizzes vigorously with effervescence.
    • Sodium: Melts into a silver spherical ball due to reaction exothermicity, skitters across the aqueous surface, effervesces rapidly.
    • Potassium: Melts, moves vigorously, ignites the liberated hydrogen gas, burning with a characteristic lilac flame.
  • Reaction with chlorine: Vigorous combustion upon insertion into chlorine gas jars, producing white crystalline halide salts:       $2\text{Na}\,(s) + \text{Cl}_2\,(g) \rightarrow 2\text{NaCl}\,(s)$
  • Reaction with oxygen: Heated alkali metals combust fiercely to yield white solid basic metal oxides:       $4\text{Na}\,(s) + \text{O}_2\,(g) \rightarrow 2\text{Na}_2\text{O}\,(s)$
  • Electronic rationale for reactivity: Alkali metal atoms possess a single valence electron ($ns^1$). Chemical oxidation requires the loss of this single electron to produce a stable $1+$ cation ($\text{M}^+$). Descending the group, additional quantum shells shield the valence electron from nuclear pull and increase the atomic radius. The weaker electrostatic attraction facilitates easier loss of the outer electron, increasing reactivity.
• Group VII: The Halogens:
  • Elements: Fluorine ($\text{F}_2$), Chlorine ($\text{Cl}_2$), Bromine ($\text{Br}_2$), Iodine ($\text{I}_2$), Astatine ($\text{At}_2$, radioactive).
  • Physical states and appearances at room temperature ($20\,^\circ\text{C}$):
  • Chlorine ($\text{Cl}_2$): Pale yellow-green toxic gas, boiling point = $-35\,^\circ\text{C}$
  • Bromine ($\text{Br}_2$): Volatile red-brown liquid giving off dense red-brown vapour, boiling point = $59\,^\circ\text{C}$
  • Iodine ($\text{I}_2$): Lustrous grey-black crystalline solid that sublimes upon warming into a purple vapour, boiling point = $184\,^\circ\text{C}$
  • Physical trends descending Group VII:
  • Molecular mass, density, and boiling point systematically increase (intermolecular van der Waals forces strengthen as electron counts rise).
  • Physical state shifts progressively from gas $\rightarrow$ liquid $\rightarrow$ solid.
  • Electronegativity and chemical reactivity systematically decrease.
  • Chemical reactivity with iron wool:
  • Heated iron wool reacts with passing halogen vapour to form iron(III) halides (with iodine yielding iron(II) iodide):       $2\text{Fe}\,(s) + 3\text{Cl}_2\,(g) \rightarrow 2\text{FeCl}_3\,(s)\quad (\text{Vigorous incandescence; yellow-brown solid})$$2\text{Fe}\,(s) + 3\text{Br}_2\,(g) \rightarrow 2\text{FeBr}_3\,(s)\quad (\text{Moderate glow; red-brown solid})$$\text{Fe}\,(s) + \text{I}_2\,(g) \rightarrow \text{FeI}_2\,(s)\quad (\text{Faint red glow; black solid})$
  • Aqueous displacement reactions:
  • A more reactive halogen displaces a less reactive halide ion from its aqueous solution.
  • Addition of pale yellow chlorine water ($\text{Cl}_2\,(aq)$) to colourless aqueous potassium bromide ($\text{KBr}$) displaces bromine, turning the solution orange:       $\text{Cl}_2\,(aq) + 2\text{KBr}\,(aq) \rightarrow 2\text{KCl}\,(aq) + \text{Br}_2\,(aq)$$\text{Ionic equation}: \text{Cl}_2\,(aq) + 2\text{Br}^-\,(aq) \rightarrow 2\text{Cl}^-\,(aq) + \text{Br}_2\,(aq)$
  • Addition of chlorine water or bromine water ($\text{Br}_2\,(aq)$) to colourless potassium iodide ($\text{KI}$) displaces iodine, turning the solution deep red-brown:       $\text{Cl}_2\,(aq) + 2\text{KI}\,(aq) \rightarrow 2\text{KCl}\,(aq) + \text{I}_2\,(aq)$$\text{Br}_2\,(aq) + 2\text{KI}\,(aq) \rightarrow 2\text{KBr}\,(aq) + \text{I}_2\,(aq)$
  • Adding iodine to solutions containing chloride or bromide ions results in no displacement (no reaction).
  • Electronic rationale for reactivity: Halogen atoms possess seven valence electrons ($ns^2 np^5$). Reactivity depends on their ability to capture one external electron to form a $1-$ halide anion ($\text{X}^-$). Descending the group, increasing atomic radii and electron shielding weaken the electrostatic pull of the positive nucleus on incoming electrons, reducing reactivity.
• Periodicity and Property Gradients Across Period 3:
  • Survey of elements across Period 3:
  • Sodium ($\text{Na}$): Group I metal, boiling point = $883\,^\circ\text{C}$, forms basic oxide $\text{Na}_2\text{O}$, chloride $\text{NaCl}$.
  • Magnesium ($\text{Mg}$): Group II metal, boiling point = $1107\,^\circ\text{C}$, forms basic oxide $\text{MgO}$, chloride $\text{MgCl}_2$.
  • Aluminium ($\text{Al}$): Group III metal, boiling point = $2467\,^\circ\text{C}$, forms amphoteric oxide $\text{Al}_2\text{O}_3$, chloride $\text{AlCl}_3$.
  • Silicon ($\text{Si}$): Group IV metalloid, boiling point = $2355\,^\circ\text{C}$, forms acidic giant covalent oxide $\text{SiO}_2$, chloride $\text{SiCl}_4$.
  • Phosphorus ($\text{P}$): Group V non-metal, ignites/boils at $280\,^\circ\text{C}$, forms acidic oxide $\text{P}_4\text{O}_{10}$, hydride $\text{PH}_3$.
  • Sulfur ($\text{S}$): Group VI non-metal, boiling point = $445\,^\circ\text{C}$, forms acidic oxide $\text{SO}_2$, hydride $\text{H}_2\text{S}$.
  • Chlorine ($\text{Cl}$): Group VII non-metal, boiling point = $-35\,^\circ\text{C}$, forms acidic oxide $\text{Cl}_2\text{O}_7$, hydride $\text{HCl}$.
  • Argon ($\text{Ar}$): Group VIII noble gas, boiling point = $-186\,^\circ\text{C}$, chemically inert monoatomic gas.
  • Systematic trends across Period 3:
  • Structure and bonding: Shifts from giant metallic lattice ($\text{Na}, \text{Mg}, \text{Al}$) to giant covalent network ($\text{Si}$) to discrete simple molecular lattices ($\text{P}_4, \text{S}_8, \text{Cl}_2$) and isolated monoatomic atoms ($\text{Ar}$).
  • Electrical conductivity: High in metallic conductors, intermediate in metalloid semiconductor silicon, zero in covalent non-metals.
  • Valency states with chlorine: Increases sequentially from $1$ to $4$ ($\text{NaCl}, \text{MgCl}_2, \text{AlCl}_3, \text{SiCl}_4$).
  • Oxide acid-base character: Basic metal oxides ($\text{Na}_2\text{O}, \text{MgO}$) react with acids to yield salts; amphoteric oxide ($\text{Al}_2\text{O}_3$) reacts with both acids and alkalis; non-metal oxides ($\text{SiO}_2, \text{P}_4\text{O}_{10}, \text{SO}_2, \text{Cl}_2\text{O}_7$) are acidic and neutralize bases.
• The Transition Elements:
  • Physical properties:
  • High mechanical densities: Iron has a density of $7.9\,\text{g}\,\text{cm}^{-3}$; copper and nickel have densities of $8.9\,\text{g}\,\text{cm}^{-3}$ (compared with $0.97\,\text{g}\,\text{cm}^{-3}$ for Group I sodium).
  • High melting points: Iron melts at $1535\,^\circ\text{C}$; copper melts at $1083\,^\circ\text{C}$; nickel melts at $1455\,^\circ\text{C}$. (Mercury is an anomalous liquid transition metal melting at $-39\,^\circ\text{C}$).
  • High mechanical strength and hardness.
  • Superior electrical and thermal conductivities: Silver exhibits the highest electrical conductivity of all metals, followed closely by copper.
  • Chemical properties:
  • Marked chemical inertness relative to Group I; do not react vigorously with ambient cold water.
  • Formation of brightly coloured chemical compounds: Utilized extensively in ceramic glazes, glass pigments, and protective coatings.
  • Catalytic properties: Transition metals and their coordination complexes function as industrial catalysts by providing accessible low-activation-energy pathways (e.g., finely divided iron in the Haber Process for ammonia synthesis).
  • Variable oxidation states: Transition metals readily lose variable numbers of electrons from their outer shells, giving multiple stable ions with distinctive names and colours (e.g., Copper(I) oxide, $\text{Cu}_2\text{O}$ vs Copper(II) oxide, $\text{CuO}$; Iron(II) chloride, $\text{FeCl}_2$ vs Iron(III) chloride, $\text{FeCl}_3$).