IB MYP Chemistry 4 & 5 Vocabulary Flashcards

Chapter 1: What Is Matter?

• Definition and Foundations of Matter:

  • Matter is anything that has mass and takes up space.

  • The form of matter depends on its underlying physical and chemical properties.

  • Scientific knowledge develops through observation, experimental testing, iterative inquiry, and historical refinement of conceptual models.

• Historical Evolution of Atomic Thought:

  • Early Cultural Concepts: Ancient Babylonian, Arab, Chinese, Hindu, Japanese, and Tibetan traditions developed early framework theories based on four, five, or eight elements.

  • Empedocles (492–432 BCE, Sicily): Proposed that all substances consist of combinations of four basic elements: fire, air, water, and earth. Varying ratios produced properties such as softness or warmth.

  • Democritus (460–370 BCE): Argued there is a limit to how finely matter can be divided. Proposed indestructible, indivisible fundamental units called atomos that retain all properties of the original matter.

  • Aristotle (384–322 BCE, Athens): Sided with Empedocles, asserting that the four elements could transform into one another. Lacking experimental verification, Aristotle's rhetoric dominated philosophical thought for nearly 2000 years, establishing the basis for alchemy (the pursuit of transmuting base metals like lead into gold and discovering the elixir of life).

  • European Enlightenment (17th Century): Social and intellectual conditions shifted toward experimental observation and direct evidence.

  • Evangelista Torricelli (1608–1647): Invented the mercury air barometer. Changes in mercury column height due to atmospheric pressure proved that invisible gas is physical matter.

  • Daniel Bernoulli (1700–1782): Formulated the kinetic theory of gases, modeling gas as tiny, constantly moving particles separated by large relative distances. Gas pressure results from millions of microscopic particle collisions.

  • Joseph Priestley (1733–1804): Isolated pure oxygen by heating mercury oxide mineral and demonstrated its unique chemical properties compared to atmospheric air.

  • Antoine Lavoisier (1743–1794) and Marie-Anne Paulze (1758–1836): Reacted acids with metals to produce hydrogen. Formulated the Law of Conservation of Mass, stating that the total mass of matter remains constant during a chemical reaction.

  • John Dalton (1766–1844): Published A New System of Chemical Philosophy (1808) synthesizing gas observations into an Atomic Theory:

  1. All matter consists of small, ball-like particles called atoms; states of matter depend on particle separation and motion.

  2. All atoms of a given element are identical; different elements have distinct atomic properties.

  3. Compounds form when atoms of different elements combine in fixed, simple whole-number ratios.

  4. Chemical reactions involve the combination, separation, or rearrangement of atoms.

  • Dmitri Mendeleev (1834–1907): At the 1860 Karlsruhe conference, scientists sought a taxonomy of elements. Mendeleev organized known element properties on cards, discovered recurring periodic trends, predicted missing elements, and constructed the foundation of the modern Periodic Table.

• States of Matter and Kinetic Molecular Theory:

  • Solid State: Fixed volume and fixed shape. Particles are closely packed in an ordered arrangement held by strong attractive forces; motion is limited to vibration around fixed lattice positions.

  • Liquid State: Fixed volume, but no fixed shape (takes the shape of its container). Particles are in contact but disordered, free to slide past one another (translational and rotational motion).

  • Gas State: No fixed volume and no fixed shape; fills the entire container. Particles are separated by large distances with negligible attractive forces and possess high kinetic energy, moving rapidly in straight lines.

  • Plasma State: A fourth high-energy state of matter present in extreme thermal environments (e.g., the Sun), consisting of ionized gas atoms and free electrons.

• State Changes and Thermal Energy:

  • State changes (freezing, melting, condensing, evaporating, subliming) are physical, completely reversible transitions involving the absorption or release of thermal energy.

  • During a phase transition (e.g., melting at melting point or boiling at boiling point), measured temperature remains constant (plateau on a heating curve).

  • Constant temperature during state change indicates that average kinetic energy ($E_k$) does not increase; added thermal energy alters particle separation distance and potential energy by strengthening or breaking intermolecular attractive forces.

  • Vapour Pressure and Boiling: A liquid boils when its internal vapour pressure equals the surrounding atmospheric pressure. At higher altitudes, atmospheric pressure is lower, reducing the boiling point of liquids.

• Purity and Classification of Matter:

  • Pure Substance: Consists of only one type of particle (identical atoms or identical molecules). Elements and compounds are pure substances.

  • Element: Pure substance containing only one type of atom (e.g., pure iron $\text{Fe}$, oxygen gas $\text{O}_2$).

  • Compound: Pure substance composed of atoms of two or more elements chemically bonded in defined whole-number ratios (e.g., pure water $\text{H}_2\text{O}$).

  • Mixture: Combination of two or more pure substances that are physically mixed rather than chemically bonded; components retain individual properties and can be present in variable ratios.

• Standardization, Measurement, and SI Units:

  • SI Units (Le Système International d'Unités) and IUPAC (International Union of Pure and Applied Chemistry) establish global reference standards.

  • Standard Temperature and Pressure (STP): $0^\text{o}\text{C}$ ($273.15\text{ K}$) and $100\text{ kPa}$ ($1\text{ bar}$ or ideal atmospheric pressure at sea level).

  • Density Formula:     $\rho = \frac{m}{V}$     where $\rho$ is density (SI unit: $\text{kg}\,\text{m}^{-3}$ or $\text{g}\,\text{cm}^{-3}$), $m$ is mass ($\text{kg}$ or $\text{g}$), and $V$ is volume ($\text{m}^3$ or $\text{cm}^3$).

  • Density Calculations:     $m_2 = \left(\frac{m_1}{V_1}\right) \times V_2$     $V_2 = m_2 \times \left(\frac{V_1}{m_1}\right)$

  • Lord Rayleigh’s Monolayer Experiment (1890): Repeated Benjamin Franklin's 1762 oil spill observation on water. Assuming an oil droplet spreads until it forms a monomolecular layer (1 molecule thick):     $V = A \times h$     where $V$ is droplet volume ($\text{m}^3$), $A$ is surface area of film ($\text{m}^2$), and $h$ is molecular thickness ($\text{m}$). Rayleigh calculated oil molecule length to be approximately $2 \times 10^{-9}\,\text{m}$ ($2\text{ nm}$).

• Experimental Uncertainty and Precision:

  • Replicate trials improve confidence in experimental reliability by establishing mean values and quantifying variance.

  • Meniscus Measurement: Observed at eye level to prevent parallax error. Water forms a concave meniscus because adhesive forces (attraction between water and glass) exceed cohesive forces (attraction of water to itself). Mercury forms a convex meniscus because cohesive forces exceed adhesive forces.

Chapter 2: How Do We Use Matter?

• Classification of Mixtures:

  • Homogeneous Mixture: A mixture with dynamic uniform composition throughout a single phase (e.g., solutions, alloys, silica glass).

  • Heterogeneous Mixture: A mixture with non-uniform composition containing two or more distinct phases (e.g., suspensions, oil-water mixtures, chocolate chip cookies).

  • Structural Forms of Solids:

  • Crystalline Solid: Highly ordered, repeating 3D lattice packing of particles.

  • Polycrystalline Solid: Aggregation of many microscopic crystalline domains (e.g., metals).

  • Amorphous Solid: Randomly organized internal particle structure lacking long-range order (e.g., glass, which behaves physically as a super-viscous liquid).

• Phases, Solutions, Suspensions, and Colloids:

  • Solution: Homogeneous mixture consisting of a dissolved solute dispersed evenly in a solvent; single phase.

  • Suspension: Heterogeneous mixture where solute particles are larger than $1000\text{ nm}$ ($1 \times 10^{-6}\,\text{m}$). Particles settle out under gravity over time forming separate phases.

  • Colloidal System: Heterogeneous system containing dispersed phase particles between $1\text{ nm}$ and $1000\text{ nm}$ ($1 \times 10^{-9}\,\text{m}$ to $1 \times 10^{-6}\,\text{m}$) distributed within a continuous phase. Colloids do not settle rapidly.

  • Tyndall Effect: Scattering of a light beam as it passes through a colloidal suspension, making the light path visible.

  • Dispersed Particle Size vs. System Appearance:

  • Particle diameter $> 1 \times 10^{-6}\,\text{m}$ ($> 1000\text{ nm}$): Two visible phases.

  • Particle diameter $100\text{ nm} - 1000\text{ nm}$ ($100 \times 10^{-9}\,\text{m} - 1000 \times 10^{-9}\,\text{m}$): Milky white appearance.

  • Particle diameter $50\text{ nm} - 100\text{ nm}$ ($50 \times 10^{-9}\,\text{m} - 100 \times 10^{-9}\,\text{m}$): Blue-white appearance.

  • Particle diameter $< 50\text{ nm}$ ($< 50 \times 10^{-9}\,\text{m}$): Grey to semi-transparent.

  • True molecular solution ($< 1\text{ nm}$): Transparent.

• Classification Types of Colloids:

  • Solid Aerosol: Solid particles dispersed in gas (e.g., smoke).

  • Liquid Aerosol: Liquid droplets dispersed in gas (e.g., fog, mist, spray aerosols).

  • Foam: Gas dispersed in liquid (e.g., whipped cream, shaving foam).

  • Emulsion: Liquid dispersed in an immiscible liquid (e.g., milk, mayonnaise, face cream). Stabilized by emulsifiers.

  • Solid Foam: Gas dispersed in solid (e.g., pumice stone, Styrofoam).

  • Gel: Liquid dispersed in solid network (e.g., jelly, foot abrasive gel).

  • Solid Sol: Solid dispersed in solid (e.g., ruby glass, alloys).

  • Suspension / Sol: Solid dispersed in liquid (e.g., paint, ink, toothpaste).

• Emulsions and Emulsifiers:

  • Water-in-Oil (W/O) vs. Oil-in-Water (O/W) emulsions.

  • Emulsifier Molecules: Amphiphilic molecules possessing a polar/hydrophilic (water-soluble) head and a non-polar/lipophilic (oil-soluble) hydrocarbon tail. They align at the liquid-liquid interface, lowering surface tension and stabilizing droplets.

  • Sudan III stains oil phases red; methylene blue stains water phases blue.

• Physical Separation Techniques:

  • Decantation: Carefully pouring off a liquid phase to separate it from a denser settled solid sediment.

  • Filtration: Separates an insoluble solid (residue) from a liquid (filtrate) using a porous barrier (filter paper).

  • Evaporation to Dryness: Boiling off a volatile liquid solvent to recover a dissolved non-volatile solid solute.

  • Dialysis: Separation technique using a semi-permeable cellulose membrane containing sub-microscopic pores. Small solute molecules/ions (e.g., urea, $\text{Cu}^{2+}$) pass through down concentration gradients, while large colloidal macromolecules (e.g., milk proteins) are retained. Applied in artificial hemodialysis machines with counter-current flow.

  • Separating Funnel: Separates two immiscible liquids based on density differences. The denser lower layer is drained via a stopcock.

  • Simple Distillation: Separates a volatile liquid solvent from a non-volatile solute or liquids with significantly different boiling points. Liquid is boiled into vapour, passed through a water-cooled Liebig condenser, and collected as distillate.

  • Paper and Thin-Layer Chromatography (TLC):

  • Component separation based on relative affinity for a stationary phase (e.g., cellulose paper, silica TLC plate) versus solubility in a mobile phase (solvent).

  • Retardation / Retention Factor ($R_f$):       $R_f = \frac{\text{distance moved by compound}}{\text{distance moved by solvent front}}$

  • $R_f$ values are constant for a specific compound under defined conditions (temperature, solvent, stationary phase) and are always $\le 1.0$.

  • Two-Dimensional Chromatography: Sequential separation using a square stationary medium rotated $90^\circ$ between two different mobile phase solvents.

Chapter 3: How Do We Map Matter?

• Major Groups of Chemical Elements:

  • Metals:

  • Shiny lustre, high thermal and electrical conductivity, malleable (hammered into sheets), ductile (drawn into wires).

  • Low electronegativity; readily lose valence electrons to form positive cations ($\text{M}^{n+}$).

  • Structure: 3D lattice of positive metal ions surrounded by a mobile, delocalized "sea" of electrons.

  • Mercury ($\text{Hg}$) is liquid at room temperature (melts at $-39\text{ }^\text{o}\text{C}$).

  • Cold working (hammering) deforms lattice regularity, hardening the metal.

  • Alloys: Homogeneous mixtures of a metal with other metals or non-metals (e.g., bronze = $87\text{--}88\%$ copper + $10\text{--}11\%$ tin; steel = iron + carbon). Disruption of lattice regularity by differing atomic sizes prevents layers from slipping easily.

  • Alkali Metals (Group 1: $\text{Li}$, $\text{Na}$, $\text{K}$, $\text{Rb}$, $\text{Cs}$, $\text{Fr}$):

  • Soft, silver-white metals, easily cut with a knife. Highly reactive; stored under mineral oil to prevent oxidation with atmospheric oxygen and water vapour.

  • Form $+1$ cations ($\text{M}^+$) and basic oxides ($\text{M}_2\text{O}$). Reaction with water produces metal hydroxide and hydrogen gas:       $2\text{M}(s) + 2\text{H}_2\text{O}(l) \rightarrow 2\text{MOH}(aq) + \text{H}_2(g)$

  • Lithium is the least dense solid element ($\rho = 0.53\,\text{g}\,\text{cm}^{-3}$).

  • Isolated via electrolysis of molten salts by Sir Humphry Davy in 1807 using a high-voltage voltaic pile.

  • Bunsen and Kirchhoff identified caesium (1860, sky-blue lines) and rubidium (1861, deep red lines) via flame emission spectroscopy.

  • Alkaline Earth Metals (Group 2: $\text{Be}$, $\text{Mg}$, $\text{Ca}$, $\text{Sr}$, $\text{Ba}$, $\text{Ra}$):

  • Reactive metals, forming $+2$ cations ($\text{M}^{2+}$). Form refractory basic oxides ("earths").

  • Isolated by Sir Humphry Davy in 1808 using molten salt electrolysis.

  • Electrolysis Principles: Cations ($\text{M}^{n+}$) undergo reduction (gain electrons) at the negative cathode ($\text{M}^{n+} + ne^- \rightarrow \text{M}$); anions ($\text{X}^{n-}$) undergo oxidation (lose electrons) at the positive anode ($\text{X}^{n-} \rightarrow \text{X} + ne^-$).

  • Halogens (Group 17: $\text{F}$, $\text{Cl}$, $\text{Br}$, $\text{I}$, $\text{At}$):

  • Reactive, toxic non-metals existing as diatomic molecules ($\text{F}_2$, $\text{Cl}_2$, $\text{Br}_2$, $\text{I}_2$). Form $-1$ halide anions ($\text{X}^-$).

  • Fluorine ($\text{F}_2$) is the most powerful oxidizing agent known.

  • Chlorofluorocarbons (CFCs, e.g., CFC-12: $\text{CF}_2\text{Cl}_2$) cause stratospheric ozone depletion, leading to the international Montreal Protocol (1987).

  • Rare Earth Elements (REEs / Lanthanides: $\text{Z} = 57\text{--}71$, plus $\text{Sc}$ and $\text{Y}$):

  • 17 soft, malleable, silvery metals with similar chemical reactivity to magnesium. Essential for high-tech applications: wind turbines, smartphone components, camera lenses, permanent magnets, defense systems, and Euro banknote fraud-prevention dyes.

  • Purification requires processing open-pit ores (generating 2000 tonnes of toxic/radioactive waste containing thorium per tonne of REE).

  • Noble Gases (Group 18: $\text{He}$, $\text{Ne}$, $\text{Ar}$, $\text{Kr}$, $\text{Xe}$, $\text{Rn}$):

  • Monatomic, unreactive, colourless gases with complete outer electron shells. Argon makes up $\approx 0.93\%$ of the atmosphere.

• Element Symbols and Nomenclature:

  • Symbols derived from English or historical Latin/Arabic names (e.g., $\text{Ag}$ = argentum, $\text{Au}$ = aurum, $\text{Cu}$ = cuprum, $\text{Fe}$ = ferrum, $\text{Hg}$ = hydrargyrum, $\text{K}$ = kalium, $\text{Na}$ = natrium, $\text{Pb}$ = plumbum, $\text{Sb}$ = stibium, $\text{Sn}$ = stannum, $\text{W}$ = wolfram).

• Periodic Trends and Atomic Structure:

  • Subatomic Particles:

  • Proton: Mass $\approx 1\text{ amu}$, charge $+1$, located in nucleus. Proton number defines atomic number ($\text{Z}$).

  • Electron: Mass $\approx 0.0005\text{ amu}$ ($\frac{1}{1836}$ of proton mass), charge $-1$, located in electron cloud shells.

  • Neutron: Mass $\approx 1\text{ amu}$, charge $0$, located in nucleus.

  • Mass Number ($\text{A}$) = Protons ($\text{Z}$) + Neutrons ($\text{N}$).

  • Döbereiner’s Triads (1829): Groups of 3 elements where the middle element's relative atomic mass ($\text{A}_r$) equals the average of the other two (e.g., $\text{Li}=6.9$, $\text{Na}=23.0$, $\text{K}=39.1$; average $\frac{6.9+39.1}{2} = 23.0$).

  • Modern Periodic Table Organization:

  • Elements arranged by increasing atomic number ($\text{Z}$).

  • Rows = Periods; Columns = Groups.

  • Across a period: Metallic $\rightarrow$ Non-metallic $\rightarrow$ Noble gas.

  • Down a group: Reactivity increases for metals (larger atomic radius, weaker nuclear pull on valence electrons); reactivity decreases for non-metals (smaller atomic radius, stronger attraction for electrons).

Chapter 4: How Do Atoms Bond?

• Four Primary Chemical Bonding Models:

  • Metallic Bonding:

  • Electrostatic attraction between positive metal cations and a surrounding delocalized sea of valence electrons.

  • Physical Properties: Malleable, ductile, high electrical and thermal conductivity in solid and liquid states, insoluble in water.

  • Ionic Bonding:

  • Electrostatic attraction between oppositely charged positive cations (metals) and negative anions (non-metals) organized in a 3D crystalline lattice.

  • Physical Properties: High melting and boiling points, hard, brittle (shatters under mechanical impact due to repulsion between aligned like-charged ions).

  • Conductivity: Non-conductors in solid state (ions fixed in lattice); excellent electrical conductors when molten or dissolved in aqueous solution (mobile ions).

  • Covalent Molecular Structure:

  • Atoms share pairs of electrons via localized covalent bonds to form discrete molecules.

  • Single bond = 1 shared pair ($2e^-$); Double bond = 2 shared pairs ($4e^-$); Triple bond = 3 shared pairs ($6e^-$).

  • Physical Properties: Low melting and boiling points (weak intermolecular forces between molecules are broken during phase changes, while strong intramolecular covalent bonds remain intact). Non-conductors of electricity in all states.

  • Covalent Network (Macromolecular) Structure:

  • Atoms continuously linked by covalent bonds throughout a giant 3D lattice (e.g., diamond, silicon dioxide $\text{SiO}_2$, silicon carbide $\text{SiC}$).

  • Physical Properties: Extremely high melting points ($> 3000\text{ }^\text{o}\text{C}$), extreme hardness, non-conductors of electricity (except graphite/graphene, which have delocalized $\pi$ electrons moving between 2D sheets).

• Electronegativity and Intermolecular Forces:

  • Electronegativity: Measure of the relative attraction that an atom has for a shared pair of electrons in a covalent bond.

  • Increases across a period (left to right) due to increasing nuclear charge ($\text{Z}$) and decreasing atomic radius.

  • Decreases down a group due to increased electron shielding and larger atomic radius.

  • Fluorine is the most electronegative element ($4.0$ Pauling scale); Caesium is the least ($0.7$).

  • Intermolecular Forces (van der Waals' forces):

  • London Dispersion Forces: Weak attractive forces caused by temporary fluctuations in electron cloud density creating instantaneous and induced dipoles. Present in all molecules; strength increases with molecular size and electron number.

  • Dipole-Dipole Forces: Electrostatic attraction between permanent dipoles of asymmetric polar covalent molecules.

  • Hydrogen Bonding: Unusually strong permanent dipole-dipole attraction occurring when a hydrogen atom is directly bonded to small, highly electronegative atoms ($\text{F}$, $\text{O}$, or $\text{N}$).

• Chemical Nomenclature and Formula Writing:

  • Binary Ionic Compounds: Metal named first, non-metal second with suffix "-ide".

  • Transition metals require Roman numerals to specify oxidation state / cation charge (e.g., $\text{FeCl}_2$ = iron(II) chloride; $\text{FeCl}_3$ = iron(III) chloride).

  • Polyatomic Ions:

  • Ammonium: $\text{NH}_4^+$ (valency 1)

  • Hydroxide: $\text{OH}^-$ (valency 1)

  • Nitrate: $\text{NO}_3^-$ (valency 1); Nitrite: $\text{NO}_2^-$ (valency 1)

  • Hydrogen carbonate: $\text{HCO}_3^-$ (valency 1)

  • Carbonate: $\text{CO}_3^{2-}$ (valency 2)

  • Sulfate: $\text{SO}_4^{2-}$ (valency 2); Sulfite: $\text{SO}_3^{2-}$ (valency 2)

  • Phosphate: $\text{PO}_4^{3-}$ (valency 3)

  • Binary Covalent Compounds: Prefixes indicate atom counts: mono- ($1$), di- ($2$), tri- ($3$), tetra- ($4$), penta- ($5$), hexa- ($6$), hepta- ($7$), octa- ($8$), nona- ($9$), deca- ($10$). Less electronegative element named first.

  • Formula Balancing (Zero Net Charge): Sum of positive charges equals sum of negative charges.

Chapter 5: What Are the Impacts of Chemical Industry?

• Chemical Equations and Phase Symbols:

  • State Symbols: $(s)$ solid, $(l)$ liquid, $(g)$ gas, $(aq)$ aqueous solution.

  • Reactants $\rightarrow$ Products.

• Reactions in Aqueous Solution:

  • Double Displacement Reaction: Precipitation reaction where aqueous cations and anions exchange partners:     $\text{KI}(aq) + \text{Pb(NO}_3)_2(aq) \rightarrow \text{2KNO}_3(aq) + \text{PbI}_2(s)$

  • Net Ionic Equations: Omits spectator ions (ions that remain dissolved and unchanged):     $\text{Pb}^{2+}(aq) + 2\text{I}^-(aq) \rightarrow \text{PbI}_2(s)$

  • General Solubility Rules for Ionic Compounds in Water:

  • Soluble ($> 10\,\text{g}\,\text{dm}^{-3}$):

    • All Group 1 metal salts and $\text{NH}_4^+$ salts.

    • All nitrate ($\text{NO}_3^-$) salts.

    • Most chloride ($\text{Cl}^-$), bromide ($\text{Br}^-$), and iodide ($\text{I}^-$) salts, except those of $\text{Ag}^+$ and $\text{Pb}^{2+}$.

    • Most sulfate ($\text{SO}_4^{2-}$) salts, except those of $\text{Ba}^{2+}$, $\text{Pb}^{2+}$, $\text{Ca}^{2+}$, $\text{Sr}^{2+}$, and $\text{Ag}^+$.

  • Insoluble ($< 1\,\text{g}\,\text{dm}^{-3}$):

    • Most carbonate ($\text{CO}_3^{2-}$), phosphate ($\text{PO}_4^{3-}$), and sulfite ($\text{SO}_3^{2-}$) salts, except Group 1 and $\text{NH}_4^+$ salts.

    • Most hydroxide ($\text{OH}^-$) and oxide ($\text{O}^{2-}$) salts, except Group 1, $\text{Ba}^{2+}$, and $\text{Ca}^{2+}$ (sparingly soluble).

• Balancing Chemical Equations:

  • Law of Conservation of Mass: Matter cannot be created or destroyed in a chemical reaction; equal numbers of each atomic species must exist on both sides of the equation.

  • Coefficients in front of formulas are adjusted; subscripts within chemical formulas must never be altered.

• Acids, Bases, and pH:

  • Arrhenius Model: Acid yields $\text{H}^+$ ions in water; Base yields $\text{OH}^-$ ions in water.

  • Brønsted-Lowry Model: Acid = proton ($\text{H}^+$) donor; Base = proton ($\text{H}^+$) acceptor.

  • Hydronium Ion ($\text{H}_3\text{O}^+$): Formed when a proton combines with a water molecule:     $\text{HCl}(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_3\text{O}^+(aq) + \text{Cl}^-(aq)$

  • Strong Acids (completely dissociated): $\text{HCl}$, $\text{HNO}_3$, $\text{H}_2\text{SO}_4$.

  • Weak Acids (partially dissociated): $\text{CH}_3\text{COOH}$ (ethanoic acid), $\text{H}_2\text{CO}_3$ (carbonic acid), $\text{C}_6\text{H}_8\text{O}_7$ (citric acid).

  • Strong Bases: $\text{NaOH}$, $\text{KOH}$, $\text{Ba(OH)}_2$.

  • Weak Bases: $\text{NH}_3$ (ammonia), $\text{NaHCO}_3$ (sodium hydrogen carbonate).

  • pH Scale: Negative logarithmic measure of hydrogen ion concentration:     $\text{pH} = -\log_{10}[\text{H}^+]$

  • Neutral solution at $25\text{ }^\text{o}\text{C}$: $\text{pH} = 7.0$ ($[\text{H}^+] = 10^{-7}\,\text{mol}\,\text{dm}^{-3}$).

  • Acidic: $\text{pH} < 7.0$; Alkaline: $\text{pH} > 7.0$

  • A change of 1 pH unit represents a 10-fold ($10\times$) change in $\text{H}^+$ ion concentration.

• Neutralization and Oxide Reactions:

  • Acid + Base (Alkalis) $\rightarrow$ Salt + Water     $\text{HCl}(aq) + \text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l)$

  • Acid + Metal Carbonate $\rightarrow$ Salt + Water + Carbon Dioxide     $2\text{HNO}_3(aq) + \text{CaCO}_3(s) \rightarrow \text{Ca(NO}_3)_2(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)$

  • Oxide Trends in the Periodic Table:

  • Basic Oxides: Formed by Group 1 and 2 metals (e.g., $\text{CaO}(s) + \text{H}_2\text{O}(l) \rightarrow \text{Ca(OH)}_2(aq)$).

  • Acidic Oxides: Formed by non-metals (e.g., $\text{CO}_2(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_2\text{CO}_3(aq)$; $\text{SO}_3(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_2\text{SO}_4(aq)$).

  • Amphoteric Oxides: React with both acids and bases (e.g., $\text{Al}_2\text{O}_3$).

Chapter 6: What Determines Chemical Change?

• Kinetic Theory and Particle Motion:

  • Particle Motion Types: Translational (straight-line movement across space), Rotational (spinning around an axis), Vibrational (rapid back-and-forth quivering around a mean position).

  • Brownian Motion: Random, jerky movement of microscopic particles (e.g., pollen grains, smoke particles) suspended in a fluid, caused by continuous, uneven collisions with invisible fast-moving fluid molecules. Demonstrated experimentally by Robert Brown (1827) and mathematically modeled by Albert Einstein (1905).

  • Maxwell-Boltzmann Distribution Curve:

  • Describes the statistical distribution of molecular kinetic energies in a gas at a specific temperature.

  • Average kinetic energy of gas molecules is directly proportional to absolute temperature ($\text{K}$).

  • Evaporation Cooling Effect: Molecules in a liquid with kinetic energy exceeding the activation threshold escape from the surface into the gas phase, reducing the average kinetic energy and temperature of the remaining liquid.

  • Saturated Vapour Pressure: Dynamic equilibrium pressure exerted by a vapour above its liquid in a closed system.

• Quantitative Chemistry and the Mole Concept:

  • Avogadro’s Constant ($N_A$ or $L$): Defined as the number of atoms in exactly $12\text{ g}$ of carbon-12:     $N_A = 6.02 \times 10^{23}\,\text{mol}^{-1}$

  • Mole ($\text{mol}$): SI base unit for amount of substance containing $6.02 \times 10^{23}$ elementary entities.

  • Molar Mass ($M$): Mass of one mole of a substance expressed in $\text{g}\,\text{mol}^{-1}$.

  • Molar Volume ($V_m$): Volume occupied by one mole of any ideal gas at STP ($0\text{ }^\text{o}\text{C}$ / $273.15\text{ K}$ and $100\text{ kPa}$):     $V_m = 22.7\,\text{dm}^3\,\text{mol}^{-1}$

  • Concentration ($c$): Moles of solute per unit volume of solution:     $c = \frac{n}{V}$     where $c$ is concentration ($\text{mol}\,\text{dm}^{-3}$ or $\text{M}$), $n$ is amount ($\text{mol}$), and $V$ is volume ($\text{dm}^3$).

  • Quantitative Conversions:     $n = \frac{m}{M}$     $n = \frac{V}{V_m}$

  • Gas Particle Velocities: Kinetic energy $E_k = \frac{1}{2} m v^2$. At the same temperature, lighter gas particles move with higher average velocity than heavier particles.

• Electrolysis of Acidified Water (Hofmann Voltameter):

  • Pure water is a weak electrolyte. Dilute $\text{H}_2\text{SO}_4$ is added to provide mobile ions ($\text{H}^+$, $\text{SO}_4^{2-}$).

  • Anode Reaction (Positive Electrode - Oxidation):     $2\text{H}_2\text{O}(l) \rightarrow \text{O}_2(g) + 4\text{H}^+(aq) + 4e^-$

  • Cathode Reaction (Negative Electrode - Reduction):     $4\text{H}^+(aq) + 4e^- \rightarrow 2\text{H}_2(g)$

  • Overall Reaction:     $2\text{H}_2\text{O}(l) \rightarrow 2\text{H}_2(g) + \text{O}_2(g)$

  • Experimental Volume Ratio: Produced gas volume of $\text{H}_2 : \text{O}_2 = 2 : 1$.

• Volumetric Analysis (Titration):

  • Quantitative technique using a calibrated burette to measure the exact volume of a standard solution required to react completely with a known volume of analyte solution in the presence of an acid-base indicator.

• Energetics: Enthalpy and Entropy:

  • Enthalpy ($H$): Total internal heat energy stored in chemical bonds and intermolecular attractions.

  • Enthalpy Change ($\Delta H$):     $\Delta H = H_{\text{products}} - H_{\text{reactants}}$

  • Exothermic Reaction: $\Delta H < 0$ (negative value); releases heat to surroundings, temperature rises.

  • Endothermic Reaction: $\Delta H > 0$ (positive value); absorbs heat from surroundings, temperature drops.

  • Entropy ($S$): Measure of disorder, randomness, or energy dispersion in a system. $\Delta S$ is positive when disorder increases (e.g., solid $\rightarrow$ liquid $\rightarrow$ gas, or dissolving a solid solute).

Chapter 7: What's Inside the Nucleus?

• Atomic Structure and Isotopes:

  • Mass Number ($\text{A}$): Total number of protons and neutrons (nucleons) in a nucleus.

  • Atomic Number ($\text{Z}$): Total number of protons in a nucleus.

  • Nuclide Notation: ${}_Z^A\text{X}$

  • Isotopes: Atoms of the same element containing equal numbers of protons ($\text{Z}$) but different numbers of neutrons ($\text{N} = \text{A} - \text{Z}$).

  • Chemical properties of isotopes are identical because they possess identical electron configurations. Physical properties (density, diffusion rate, melting/boiling points) differ due to mass differences.

  • Mass Spectrometry: Analytical instrument separating vaporized positive ions by mass-to-charge ratio ($m/z$) using a magnetic field to determine isotope masses and relative abundances.

  • Weighted Relative Atomic Mass Calculation:     $\text{A}_r = \frac{(\text{Isotope}_1\text{ Mass} \times \%_1) + (\text{Isotope}_2\text{ Mass} \times \%_2) + \dots}{100}$

• Isotope Ratios and Isoforensics:

  • Stable Isotopes (e.g., ${}^1\text{H}/{}^2\text{H}$, ${}^{12}\text{C}/{}^{13}\text{C}$, ${}^{16}\text{O}/{}^{18}\text{O}$): Fractionate during physical processes (evaporation, precipitation) and biological pathways ($\text{C}_3$ vs $\text{C}_4$ plant photosynthesis).

  • Isoforensics / Isotope Signatures: Analyzing stable isotope ratios in human hair ($1\text{ cm}$ growth per month), bone, and teeth collagen tracks geographic origin, water sources, and diet over time.

  • Antarctic Ice Cores: Measuring ${}^{18}\text{O}/{}^{16}\text{O}$ ratios in trapped ice core samples (e.g., Vostok base) provides historical records of global atmospheric temperatures and $\text{CO}_2$ concentrations dating back hundreds of thousands of years.

• Radioactivity and Nuclear Decay:

  • Radioactive Decay: Spontaneous transformation of an unstable nucleus (radionuclide) into a stable nucleus accompanied by emission of particles or radiation.

  • Types of Radioactive Decay:

  • Alpha ($\alpha$) Decay: Emission of a helium-4 nucleus (${}_2^4\text{He}$ or ${}_2^4\alpha$). Mass number decreases by 4, atomic number decreases by 2:       ${}_{92}^{234}\text{U} \rightarrow {}_{90}^{230}\text{Th} + {}_2^4\text{He}$       High ionizing power; stopped by a sheet of paper or a few centimetres of air.

  • Beta-Minus ($\beta^-$) Decay: Emission of a high-speed electron (${}_{-1}^0e$ or ${}_{-1}^0\beta$) and an antineutrino ($\bar{\nu}$) when a neutron converts into a proton. Mass number unchanged, atomic number increases by 1:       ${}_6^{14}\text{C} \rightarrow {}_7^{14}\text{N} + {}_{-1}^0e + \bar{\nu}$       Moderate penetration; stopped by a few millimetres of aluminium.

  • Beta-Plus ($\beta^+$) Decay / Positron Emission: Emission of a positron (${}_{+1}^0e$) and a neutrino ($\nu$) when a proton converts into a neutron. Atomic number decreases by 1.

  • Gamma ($\gamma$) Decay: Emission of high-energy short-wavelength electromagnetic radiation (${}_0^0\gamma$) from an excited nucleus. No change in mass or atomic number. Extremely high penetration; reduced by thick lead or concrete.

  • Transuranium Elements: Synthetic radioactive elements with $\text{Z} > 92$ produced via nuclear bombardment reactions (e.g., ${}_{92}^{238}\text{U} + {}_0^1n \rightarrow {}_{92}^{239}\text{U} \rightarrow {}_{93}^{239}\text{Np} + {}_{-1}^0e$).

  • Radioactive Half-Life ($t_{1/2}$): Time required for half the radioactive parent nuclei in a sample to decay into daughter products. Decay is a random, exponential process independent of temperature and pressure.

• Discovery of the Atomic Nucleus:

  • Ernest Rutherford, Hans Geiger, and Ernest Marsden (1909): Directed alpha particles at a thin gold foil ($400\text{ atoms}$ thick). Most passed straight through, but $\approx 1 \text{ in } 10000$ deflected by $> 90^\circ$. Proved the atom is mostly empty space with a tiny, dense, positively charged nucleus containing virtually all the atomic mass.

• Stellar Nucleosynthesis:

  • Nuclear fusion in stars combines light nuclei under extreme temperature ($> 10^7\text{ }^\text{o}\text{C}$) and pressure (${}^1\text{H} \rightarrow {}^4\text{He} \rightarrow {}^8\text{Be} \rightarrow {}^{12}\text{C} \rightarrow {}^{16}\text{O}$, up to ${}^{56}\text{Fe}$).

  • Elements heavier than iron are synthesized during supernova explosions via rapid ($r$) and slow ($s$) neutron capture processes.

Chapter 8: Why Do Electrons Matter?

• Flame Tests, Photons, and Atomic Energy Levels:

  • Flame Tests: Volatilized metal cations absorb thermal energy, promoting valence electrons to higher energy orbits. Upon returning to lower ground state levels, energy is emitted as photons of visible light of characteristic wavelengths:

  • Lithium ($\text{Li}^+$): Crimson / Red

  • Sodium ($\text{Na}^+$): Intense Yellow

  • Potassium ($\text{K}^+$): Lilac / Violet

  • Calcium ($\text{Ca}^{2+}$): Brick Red / Orange

  • Barium ($\text{Ba}^{2+}$): Apple Green

  • Copper ($\text{Cu}^{2+}$): Blue-Green

  • Quantum Light Equations:     $c = f\lambda$     where $c = 3.00 \times 10^8\,\text{m}\,\text{s}^{-1}$ (speed of light), $f$ is frequency ($\text{Hz}$ or $\text{s}^{-1}$), and $\lambda$ is wavelength ($\text{m}$).     $E = hf$     where $h = 6.63 \times 10^{-34}\,\text{J}\,\text{s}$ (Planck's constant) and $E$ is energy ($\text{J}$).     $E_{\text{mol}} = N_A h f$

• Historical Atomic Models and Shells:

  • J.J. Thomson (1897): Discovered the electron using cathode ray tubes. Proposed the "plum pudding" model: electrons embedded in a sphere of positive charge.

  • Niels Bohr (1913): Proposed that electrons orbit the nucleus in fixed, quantized principal energy levels (shells $n = 1, 2, 3, 4$ holding a maximum of $2n^2$ electrons: $2, 8, 18, 32$

  • Outer valence shell holds up to 8 electrons (Octet Rule).

  • Electron shell configurations for $\text{Z} = 1\text{--}20$:

  • $\text{Na}$ ($\text{Z}=11$): $2, 8, 1$

  • $\text{Ca}$ ($\text{Z}=20$): $2, 8, 8, 2$

  • $\text{Cl}^-$ anion: $2, 8, 8$

• Periodic Trends in Atomic Properties:

  • Atomic Radius: Decreases across a period (left to right) due to increasing nuclear charge ($\text{Z}$) attracting electron shells closer. Increases down a group due to addition of electron shells.

  • First Ionization Energy: Energy required to remove one mole of electrons from one mole of gaseous atoms:     $\text{X}(g) \rightarrow \text{X}^+(g) + e^-$     Increases across a period; decreases down a group.

• Electron Representation Conventions:

  • Dot-and-Cross Diagrams: Shows valence electron transfer in ionic bonds or electron sharing in covalent bonds.

  • Lewis Electron Dot Structures: Atomic symbol surrounded by dots representing valence electrons. Unpaired dots represent bonding capacity/oxidation state; paired dots represent non-bonding lone pairs.

• Reactivity Series and Voltaic Cells:

  • Activity Series of Metals:     $\text{K} > \text{Na} > \text{Li} > \text{Ca} > \text{Mg} > \text{Al} > \text{Zn} > \text{Fe} > \text{Pb} > \text{H} > \text{Cu} > \text{Ag} > \text{Au}$

  • Metal Displacement Reaction: A more reactive metal oxidizes and displaces a less reactive metal cation from solution:     $\text{Zn}(s) + \text{Cu}^{2+}(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{Cu}(s)$

  • Oxidation half-equation: $\text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2e^-$

  • Reduction half-equation: $\text{Cu}^{2+}(aq) + 2e^- \rightarrow \text{Cu}(s)$

  • Electrochemical / Voltaic Cell:

  • Converts chemical energy from a spontaneous redox reaction into electrical energy.

  • Anode (Negative electrode): Oxidation site ($\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-$).

  • Cathode (Positive electrode): Reduction site ($\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$).

  • Salt Bridge: Contains aqueous electrolyte (e.g., $\text{KNO}_3$) allowing ion migration to maintain electrical neutrality.

  • Cell Potential Formula:       $E_{\text{cell}}^\text{o} = E_{\text{cathode}}^\text{o} - E_{\text{anode}}^\text{o}$

  • Corrosion of Iron (Rusting):

  • Requires both oxygen and liquid water. Hydrated iron(III) oxide formed: $\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$.

  • Sacrificial Protection / Galvanizing: Coating iron with zinc ($\text{Zn}$). Zinc is more reactive, so it oxidizes preferentially, donating electrons to the iron.

• Industrial Metal Extraction:

  • Smelting in Blast Furnace (Extraction of Iron from Haematite $\text{Fe}_2\text{O}_3$):

  • Coke ($\text{C}$) burns in air to produce $\text{CO}_2$: $\text{C}(s) + \text{O}_2(g) \rightarrow \text{CO}_2(g)$

  • $\text{CO}_2$ reacts with excess coke to form carbon monoxide ($\text{CO}$): $\text{C}(s) + \text{CO}_2(g) \rightarrow 2\text{CO}(g)$

  • Carbon monoxide reduces iron(III) oxide: $\text{Fe}_2\text{O}_3(s) + 3\text{CO}(g) \rightarrow 2\text{Fe}(l) + 3\text{CO}_2(g)$

  • Limestone ($\text{CaCO}_3$) decomposes to $\text{CaO}$, reacting with silica impurities ($\text{SiO}_2$) to form molten slag ($\text{CaSiO}_3$).

  • Electrolysis of Molten Alumina ($\text{Al}_2\text{O}_3$ in Cryolite $\text{Na}_3\text{AlF}_6$):

  • Cathode: $\text{Al}^{3+}(l) + 3e^- \rightarrow \text{Al}(l)$

  • Anode: $2\text{O}^{2-}(l) \rightarrow \text{O}_2(g) + 4e^-$

Chapter 9: How Are Environmental Systems Sustained by Their Chemistry?

• Atmospheric Origins and Evolution:

  • Primordial atmosphere formed by volcanic outgassing: $\text{CO}_2$, $\text{H}_2\text{O}$, $\text{N}_2$, $\text{NH}_3$.

  • Miller-Urey Experiment (1953): Subjected $\text{CH}_4$, $\text{NH}_3$, $\text{H}_2$, and $\text{H}_2\text{O}$ to electrical sparks, synthesizing amino acids and organic molecules.

  • Atmospheric Composition Today: $78\%\text{ N}_2$, $21\%\text{ O}_2$, $0.93\%\text{ Ar}$, $0.04\%\text{ CO}_2$.

  • Industrial Fractional Distillation of Liquid Air: Air is cooled to $-200\text{ }^\text{o}\text{C}$, pressurized, and distilled by boiling points: Nitrogen (b.p. $-196\text{ }^\text{o}\text{C}$), Argon (b.p. $-186\text{ }^\text{o}\text{C}$), Oxygen (b.p. $-183\text{ }^\text{o}\text{C}$).

• Atmospheric Layers:

  • Troposphere ($0\text{--}12\text{ km}$): Contains $90\%$ of atmospheric gas mass; weather occurs here; temperature decreases with altitude.

  • Stratosphere ($12\text{--}50\text{ km}$): Contains ozone layer ($\text{O}_3$); temperature increases with altitude due to UV absorption.

  • Mesosphere ($50\text{--}80\text{ km}$): Temperature decreases with altitude.

  • Thermosphere ($> 80\text{ km}$): Ionizing solar radiation increases temperature.

• Greenhouse Effect and Climate Change:

  • Greenhouse gases ($\text{CO}_2$, $\text{H}_2\text{O}$, $\text{CH}_4$, $\text{N}_2\text{O}$, CFCs) allow solar visible radiation to reach Earth, but absorb and re-emit outgoing infrared radiation.

  • Global Warming Potential (GWP) relative to $\text{CO}_2 = 1$:

  • Methane ($\text{CH}_4$): $25\times$

  • Nitrous oxide ($\text{N}_2\text{O}$): $300\times$

  • CFCs: $1000\times - 10000\times$

  • Carbon Footprint: Total greenhouse gas emissions caused by an individual, event, or product, expressed as tonnes of $\text{CO}_2$ equivalent ($\text{CO}_2\text{e}$).

• Ozone Depletion and Chlorofluorocarbons (CFCs):

  • Ozone ($\text{O}_3$) in stratosphere filters harmful UV-B radiation ($290\text{--}320\text{ nm}$).

  • CFC Photolysis in Stratosphere:     $\text{CF}_2\text{Cl}_2 + hf \rightarrow \text{CF}_2\text{Cl}^\bullet + \text{Cl}^\bullet$     $\text{Cl}^\bullet + \text{O}_3 \rightarrow \text{ClO}^\bullet + \text{O}_2$     $\text{ClO}^\bullet + \text{O}_3 \rightarrow \text{Cl}^\bullet + 2\text{O}_2$     A single chlorine radical destroys thousands of ozone molecules. International action led to the Montreal Protocol (1987).

• Biogeochemical Nutrient Cycles:

  • Carbon Cycle: Balanced by photosynthesis ($6\text{CO}_2 + 6\text{H}_2\text{O} \rightarrow \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2$) and respiration/combustion ($\text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2 \rightarrow 6\text{CO}_2 + 6\text{H}_2\text{O}$).

  • Nitrogen Cycle: Unreactive $\text{N}_2$ gas fixed into nitrates ($\text{NO}_3^-$) via nitrogen-fixing bacteria, lightning, or the industrial Haber process ($\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3$).

  • Phosphorus Cycle: Weathering of rocks releases inorganic phosphate ($\text{PO}_4^{3-}$). Non-renewable agricultural resource.

Chapter 10: How Can Our Energy Resources Be Accessed Fairly?

• History and Principles of Combustion:

  • Phlogiston Theory (Becher & Stahl): Obsolete theory claiming flammables release phlogiston when burned. Disproved by Lavoisier showing mass increase during metal oxidation.

  • Flash Point: Lowest temperature at which a liquid fuel produces sufficient vapour to ignite in air.

  • Fire Triangle: Fuel, Oxygen, Heat (Activation Energy).

  • Complete Combustion: Fuel + $\text{O}_2 \rightarrow \text{CO}_2 + \text{H}_2\text{O}$ ($\Delta H < 0$).

  • Incomplete Combustion: Fuel + limited $\text{O}_2 \rightarrow \text{CO} + \text{H}_2\text{O}$ or $\text{C}(s) + \text{H}_2\text{O}$. Carbon monoxide ($\text{CO}$) binds irreversibly to blood hemoglobin.

• Calorimetry and Energy Calculations:

  • Heat Energy Transferred ($Q$):     $Q = m c \Delta T$     where $Q$ is energy ($\text{J}$), $m$ is mass of water ($\text{g}$), $c$ is specific heat capacity of water ($4.18\,\text{J}\,\text{g}^{-1}\,^\text{o}\text{C}^{-1}$), and $\Delta T$ is temperature change ($^\text{o}\text{C}$).

  • Molar Enthalpy Change ($\Delta H$):     $\Delta H = -\frac{Q}{n}$

  • Bond Energy Calculations:     $\Delta H = \sum (\text{Bonds Broken}) - \sum (\text{Bonds Formed})$

  • Breaking bonds is endothermic ($\Delta H > 0$).

  • Forming bonds is exothermic ($\Delta H < 0$).

• Gibbs Free Energy and Thermodynamic Spontaneity:   $\Delta G = \Delta H - T\Delta S$   where $\Delta G$ is Gibbs free energy change ($\text{kJ}\,\text{mol}^{-1}$), $\Delta H$ is enthalpy change ($\text{kJ}\,\text{mol}^{-1}$), $T$ is absolute temperature ($\text{K}$), and $\Delta S$ is entropy change ($\text{kJ}\,\text{K}^{-1}\,\text{mol}^{-1}$).

  • A chemical reaction is spontaneous if $\Delta G < 0$

Chapter 11: How Can We Shift the Balance of a Reaction?

• Reaction Kinetics and Collision Theory:

  • Reaction Rate: Change in concentration of reactants or products per unit time ($\text{mol}\,\text{dm}^{-3}\,\text{s}^{-1}$ or $\text{g}\,\text{s}^{-1}$, $\text{cm}^3\,\text{s}^{-1}$).

  • Tangent to concentration-time curve yields instantaneous reaction rate ($\text{slope} = \frac{\Delta c}{\Delta t}$).

  • Collision Theory Requirements: Particles must collide with energy $\ge$ Activation Energy ($E_a$) and correct steric collision geometry.

  • Factors Increasing Reaction Rate:

  1. Temperature: Increases molecular speed/kinetic energy; exponentially increases fraction of particles with $E \ge E_a$.

  2. Concentration / Pressure: Increases particle density per unit volume, increasing collision frequency.

  3. Surface Area: Dividing solid reactants into smaller pieces increases available collision sites.

  4. Catalysts: Provide alternative reaction pathway with lower activation energy ($E_a$). Catalysts increase rate without being consumed or altering $\Delta H$ or final equilibrium yield.

• Reversible Reactions and Dynamic Equilibrium:

  • Reversible Reaction ($\rightleftharpoons$): Forward and reverse reactions occur simultaneously.

  • Dynamic Equilibrium: In a closed system, rate of forward reaction equals rate of backward reaction; concentrations of reactants and products remain constant.

  • Le Chatelier’s Principle: If a system at dynamic equilibrium is subjected to a change in conditions, the position of equilibrium shifts to counteract that change.

  • Temperature Increase: Shifts equilibrium in the endothermic direction.

  • Temperature Decrease: Shifts equilibrium in the exothermic direction.

  • Pressure Increase: Shifts equilibrium toward the side with fewer moles of gas.

  • Pressure Decrease: Shifts equilibrium toward the side with more moles of gas.

  • Concentration Increase: Shifts equilibrium away from the added component.

• Industrial Case Study - The Haber Process:   $\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) \quad (\Delta H = -92.4\,\text{kJ}\,\text{mol}^{-1})$

  • Industrial Compromise Conditions:

  • Temperature: $450\text{ }^\text{o}\text{C}$ (Compromise: low temperature favours high equilibrium yield, but high temperature increases reaction rate).

  • Pressure: $200\text{ atm}$ ($20\text{ MPa}$) (High pressure shifts equilibrium right toward 2 moles gas vs 4 moles gas).

  • Catalyst: Finely divided iron ($\text{Fe}$) catalyst.

  • Unreacted $\text{N}_2$ and $\text{H}_2$ are recycled; $\text{NH}_3$ is condensed to liquid and removed.

Chapter 12: Organic Chemistry and Molecular Synthesis

• Carbon Bonding and Catenation:

  • Organic chemistry is the study of carbon compounds (excluding $\text{CO}$, $\text{CO}_2$, carbonates, and cyanides).

  • Carbon ($\text{Z}=6$) has 4 valence electrons, forming 4 covalent bonds. Catenation is the ability of carbon to form stable covalent chains, branched networks, and rings.

• Homologous Series and Functional Groups:

  • Homologous Series: Family of organic compounds sharing the same functional group, same general formula, similar chemical properties, and showing a gradual trend in physical properties (e.g., increasing boiling point with chain length).

  • Alkanes: Saturated hydrocarbons with single bonds. General formula: $\text{C}_n\text{H}_{2n+2}$. Suffix "-ane" (methane $\text{CH}_4$, ethane $\text{C}_2\text{H}_6$, propane $\text{C}_3\text{H}_8$, butane $\text{C}_4\text{H}_{10}$, pentane $\text{C}_5\text{H}_{12}$, hexane $\text{C}_6\text{H}_{14}$).

  • Alkenes: Unsaturated hydrocarbons with one $\text{C}=\text{C}$ double bond. General formula: $\text{C}_n\text{H}_{2n}$. Suffix "-ene" (ethene $\text{C}_2\text{H}_4$, propene $\text{C}_3\text{H}_6$, boot-1-ene $\text{C}_4\text{H}_8$). Undergo addition reactions across the double bond (e.g., addition of $\text{Br}_2$ turns orange bromine water colourless).

  • Alkynes: Unsaturated hydrocarbons with one $\text{C}\equiv\text{C}$ triple bond. General formula: $\text{C}_n\text{H}_{2n-2}$.

  • Alcohols: Contain hydroxyl group ($-\text{OH}$). General formula: $\text{C}_n\text{H}_{2n+1}\text{OH}$. Suffix "-ol" (methanol, ethanol). Hydration of ethene: $\text{C}_2\text{H}_4 + \text{H}_2\text{O} \xrightarrow{\text{H}_3\text{PO}_4} \text{C}_2\text{H}_5\text{OH}$.

  • Carboxylic Acids: Contain carboxyl group ($-\text{COOH}$). General formula: $\text{C}_n\text{H}_{2n+1}\text{COOH}$. Suffix "-oic acid" (methanoic acid $\text{HCOOH}$, ethanoic acid $\text{CH}_3\text{COOH}$). Weak organic acids.

  • Esters: Formed by esterification (condensation) between a carboxylic acid and an alcohol using concentrated $\text{H}_2\text{SO}_4$ catalyst:     $\text{Carboxylic Acid} + \text{Alcohol} \rightleftharpoons \text{Ester} + \text{Water}$     $\text{CH}_3\text{COOH} + \text{C}_2\text{H}_5\text{OH} \rightleftharpoons \text{CH}_3\text{COOC}_2\text{H}_5 + \text{H}_2\text{O}$     Sweet, fruity fragrances.

• Isomerism:

  • Structural Isomers: Molecules with identical molecular formulas but different structural connectivity (e.g., butane and 2-methylpropane, both $\text{C}_4\text{H}_{10}$).

  • Geometric (Cis-Trans) Isomers: Caused by restricted rotation around $\text{C}=\text{C}$ double bonds. Cis-isomers have substituents on the same side; trans-isomers have substituents on opposite sides.

  • Enantiomers (Optical Isomers): Non-superimposable mirror-image molecules containing an asymmetric/chiral carbon atom bonded to 4 different groups. Enantiomers rotate plane-polarized light in opposite directions ($d/+$ and $l/-$).

• Crude Oil Refining and Polymers:

  • Fractional Distillation of Crude Oil: Separates hydrocarbons by boiling point range in a fractionating column (Refinery gas, Gasoline/petrol, Kerosene/paraffin, Diesel oil, Fuel oil, Lubricating wax, Bitumen).

  • Cracking: Thermal or catalytic breakdown of long-chain unreactive alkanes into shorter, useful alkanes and reactive alkenes.

  • Addition Polymers: Monomers containing $\text{C}=\text{C}$ double bonds join via repeated addition reactions (e.g., poly(ethene), PVC, PTFE/Teflon, polystyrene).

  • Condensation Polymers: Monomers with two functional groups react, eliminating a small molecule (e.g., $\text{H}_2\text{O}$) at each linkage:

  • Polyamides / Nylon: Monomers linked by amide bonds ($-\text{CO-NH}-$).

  • Polyesters / Terylene: Monomers linked by ester bonds ($-\text{CO-O}-$).

Glossary of Essential Terms

• Acids: Substances forming hydrogen ions ($\text{H}^+$) in aqueous solution, with $\text{pH} < 7.0$.

• Activation Energy ($E_a$): Minimum kinetic energy required by colliding particles to initiate a chemical reaction.

• Alkali: Water-soluble base forming hydroxide ions ($\text{OH}^-$) in solution, with $\text{pH} > 7.0$

• Allotropes: Different structural forms of the same element in the same physical state (e.g., diamond, graphite, and $\text{C}_{60}$ fullerene).

• Amphiprotic: Molecule or ion capable of acting as both a proton donor and proton acceptor (e.g., $\text{H}_2\text{O}$).

• Amphoteric: Oxide or hydroxide capable of reacting with both acids and bases (e.g., $\text{Al}_2\text{O}_3$).

• Anion: Negatively charged ion attracted to the positive anode during electrolysis.

• Anode: Electrode where oxidation (loss of electrons) occurs.

• Atomic Number ($\text{Z}$): Number of protons in the nucleus of an atom.

• Avogadro's Constant ($N_A$): Number of particles in one mole of substance ($6.02 \times 10^{23}\,\text{mol}^{-1}$).

• Cathode: Electrode where reduction (gain of electrons) occurs.

• Cation: Positively charged ion attracted to the negative cathode during electrolysis.

• Electronegativity: Measure of the relative attraction an atom has for a shared pair of electrons in a covalent bond.

• Enthalpy Change ($\Delta H$): Heat energy absorbed or released by a system during a reaction at constant pressure.

• Entropy ($S$): Quantitative measure of disorder or randomness in a chemical system.

• Isotopes: Atoms of the same element with identical proton number ($\text{Z}$) but different neutron numbers ($\text{N}$).

• Mass Number ($\text{A}$): Sum of protons and neutrons in an atomic nucleus.

• Molar Volume ($V_m$): Volume occupied by one mole of ideal gas at STP ($22.7\,\text{dm}^3\,\text{mol}^{-1}$).

• Oxidation: Loss of electrons, gain of oxygen, loss of hydrogen, or increase in oxidation state.

• Reduction: Gain of electrons, loss of oxygen, gain of hydrogen, or decrease in oxidation state.

• Standard Temperature and Pressure (STP): Defined reference conditions of $0\text{ }^\text{o}\text{C}$ ($273.15\text{ K}$) and $100\text{ kPa}$ ($1\text{ bar}$).