GCE 'O' Level Science (Chemistry) Comprehensive Key-Points Study Guide

Experimental Techniques

• Experimental Design

  • Measurement of Time, Temperature, and Mass:

  • Time: Measured using a digital stopwatch (LC Quartz) or an analogue stopwatch in minutes ($min$) and seconds ($s$).

  • Temperature: Measured using a mercury-in-glass thermometer in degrees Celsius ($^\circ\text{C}$).

  • Mass: Measured using an electronic balance or a beam balance in grams ($g$) or kilograms ($kg$).

  • Measurement of Volume of Liquids:

  • Beaker: Used only for estimating volumes; least accurate.

  • Measuring Cylinder: Increasing accuracy, measures to the nearest $0.5\,cm^3$ or $1\,cm^3$.

  • Pipette: Measures accurately fixed volumes of liquids (e.g., $10.0\,cm^3$, $25.0\,cm^3$); accurate to 1 decimal place.

  • Burette: Most accurate piece of apparatus for variable volumes, measuring up to $0.1\,cm^3$ (recorded to 2 decimal places, e.g., $22.50\,cm^3$).

  • Measurement of Volume of Gases:

  • Measured using a graduated gas syringe.

  • Data Recording Rule: Always include units when recording experimental data or carrying out calculations to prevent loss of marks.

  • Collection of Gases:

  • Two key properties determine the collection method: gas density relative to air and solubility in water.

  • Displacement of water: Suitable for gases insoluble or sparingly soluble in water (e.g., $H_2$, $O_2$).

    • Note: Carbon dioxide ($CO_2$) is sparingly soluble in water to form carbonic acid ($H_2CO_3$), so collecting it over water yields a slightly smaller gas volume than theoretical predictions.

  • Displacement of air - Upward delivery: Suitable for gases less dense than air and soluble in water (e.g., $H_2$, $NH_3$).

  • Displacement of air - Downward delivery: Suitable for gases denser than air and soluble in water (e.g., $HCl$, $CO_2$, $Cl_2$, $SO_2$).

  • Worked Questions and Explanations:

  • Question: Which apparatus measures exactly $22.5\,cm^3$ of a liquid?

    • Answer: Burette. Pipettes measure fixed volumes (e.g., $25.0\,cm^3$), whereas burettes accurately deliver variable volumes up to $0.1\,cm^3$.

  • Question: Which gas ($NH_3$, $Cl_2$, $HCl$, $O_2$, $SO_2$) is most suitably collected using a beehive shelf over water?

    • Answer: Oxygen ($O_2$). Ammonia dissolves in water to form an alkaline solution; chlorine, hydrogen chloride, and sulphur dioxide dissolve to form acidic solutions. Oxygen is insoluble.

• Methods of Purification and Analysis

  • Filtration:

  • Used to separate an insoluble solid from a liquid (e.g., sand from a sand-water mixture).

  • Fine pores in filter paper retain large solid particles (residue) while liquid particles pass through (filtrate).

  • Crystallisation:

  • Used to separate a pure soluble solid from an impure solution (e.g., copper(II) sulphate crystals from impure copper(II) sulphate solution).

  • Procedure: 1. Dissolve impure solid in a solvent. 2. Heat solution in an evaporating basin to evaporate most solvent and obtain a saturated solution. 3. Allow hot solution to cool slowly so pure crystals form. 4. Pour off cold solution and dry crystals by pressing between sheets of filter paper.

  • Crystallisation vs. Evaporation to Dryness: Crystallisation partially evaporates solvent so impurities remain in liquid upon filtering. Evaporation to dryness removes all solvent, which can decompose thermally sensitive solids or leave behind solid impurities.

  • Sublimation:

  • Used to separate mixtures containing a component that sublimes upon heating (e.g., separating ammonium chloride from sodium chloride). Ammonium chloride turns directly to vapour and deposits on a cold surface, leaving sodium chloride residue.

  • Simple Distillation:

  • Used to separate a pure liquid (distillate) from a solution containing non-volatile dissolved solids (e.g., pure water from seawater).

  • Setup: Thermometer bulb is placed level with the condenser entrance to measure vapour temperature (distillate boiling point). Anti-bumping granules ensure smooth boiling. Condenser cooling water enters from the bottom and exits at the top to keep the entire tube full and maximum cooling efficiency.

  • Temperature remains flat at $100^\circ\text{C}$ while pure water distils over.

  • Fractional Distillation:

  • Used to separate miscible liquids with different boiling points (e.g., ethanol b.p. $78^\circ\text{C}$ and water b.p. $100^\circ\text{C}$).

  • Components separate according to boiling point; the liquid with the lowest boiling point distils first. Boiling points should be at least $20^\circ\text{C}$ apart.

  • A fractionating column provides a large surface area for repeated condensation and vapourisation cycles.

  • Applications: Separating crude oil components; concentrating fermented liquor.

  • Worked Question: Identifying gases condensing in a $25^\circ\text{C}$ water condenser given boiling points: ammonia ($-33^\circ\text{C}$), butane ($-0.5^\circ\text{C}$), sulphur dioxide ($-10^\circ\text{C}$), hydrogen chloride ($-85^\circ\text{C}$), pentane ($+36^\circ\text{C}$).

    • Answer: Pentane. Only pentane has a boiling point above condenser temperature ($25^\circ\text{C}$), enabling its vapour to condense into liquid.

  • Paper Chromatography:

  • Separates substances in a mixture based on solubility in a solvent. More soluble components move further up the paper.

  • Procedure: 1. Draw baseline in pencil. 2. Place small spot of sample on baseline. 3. Suspend paper in solvent bath with baseline above solvent level. 4. Cover beaker to prevent solvent evaporation. 5. Remove chromatogram before solvent reaches paper top.

  • Key Procedural Precautions: Baseline drawn in pencil because graphite is insoluble and won't develop; baseline kept above solvent level so sample spots don't dissolve directly into bath; beaker covered to prevent solvent evaporation from paper during movement.

  • Pure substances produce a single spot; impure mixtures produce multiple spots.

  • Criteria of Purity:

  • Pure substances have sharp, fixed melting and boiling points.

  • Melting Point Determination: Solid sample in capillary melting tube tied to thermometer in liquid bath heated slowly with stirring. Record $T_1$ (melting starts) and $T_2$ (fully melted). Pure solid: $T_1 = T_2$. Impure solid: melts over a temperature range ($T_1 < T_2$) with a depressed melting point.

  • Boiling Point Determination: Pure liquids boil at a fixed temperature. Impurities elevate the boiling point.

  • Worked Question: Solid X melts at $133^\circ\text{C}$. Pure urea melts at $133^\circ\text{C}$. A mixture of X and pure urea melts at $129^\circ\text{C}-132^\circ\text{C}$.

    • Answer: Solid X is not urea. The melting point depression and melting range indicate that X acts as an impurity in pure urea.

• Identification of Ions and Gases

  • Tests for Aqueous Cations:

  • Add aqueous sodium hydroxide ($NaOH(aq)$) or aqueous ammonia ($NH_3(aq)$) dropwise until in excess.

  • Aluminium ($Al^{3+}$):

    • $NaOH(aq)$: White precipitate formed, soluble in excess to form a colourless solution.

    • $NH_3(aq)$: White precipitate formed, insoluble in excess.

  • Ammonium ($NH_4^+$):

    • $NaOH(aq)$: No precipitate; ammonia gas ($NH_3$) produced on warming.

    • $NH_3(aq)$: No reaction.

  • Calcium ($Ca^{2+}$):

    • $NaOH(aq)$: White precipitate formed, insoluble in excess.

    • $NH_3(aq)$: No precipitate or slight white precipitate.

  • Copper(II) ($Cu^{2+}$):

    • $NaOH(aq)$: Light blue precipitate formed, insoluble in excess.

    • $NH_3(aq)$: Light blue precipitate formed, soluble in excess to form a dark blue solution.

  • Iron(II) ($Fe^{2+}$):

    • $NaOH(aq)$: Green precipitate formed, insoluble in excess.

    • $NH_3(aq)$: Green precipitate formed, insoluble in excess.

  • Iron(III) ($Fe^{3+}$):

    • $NaOH(aq)$: Reddish-brown precipitate formed, insoluble in excess.

    • $NH_3(aq)$: Reddish-brown precipitate formed, insoluble in excess.

  • Lead(II) ($Pb^{2+}$):

    • $NaOH(aq)$: White precipitate formed, soluble in excess to form a colourless solution.

    • $NH_3(aq)$: White precipitate formed, insoluble in excess.

  • Zinc ($Zn^{2+}$):

    • $NaOH(aq)$: White precipitate formed, soluble in excess to form a colourless solution.

    • $NH_3(aq)$: White precipitate formed, soluble in excess to form a colourless solution.

  • Distinguishing $Al^{3+}$ and $Pb^{2+}$: Add dilute $HCl(aq)$ or $KI(aq)$. $Al^{3+}$ forms colourless solutions ($AlCl_3(aq)$ or $AlI_3(aq)$). $Pb^{2+}$ forms a white precipitate ($PbCl_2(s)$) with $HCl$ or a bright yellow precipitate ($PbI_2(s)$) with $KI$.

  • Tests for Anions:

  • Carbonate ($CO_3^{2-}$): Add dilute acid -> Effervescence observed; gas produced forms a white precipitate in limewater ($CO_2$).

  • Chloride ($Cl^-$): Acidify with dilute nitric acid ($HNO_3$), then add aqueous silver nitrate ($AgNO_3$) -> White precipitate formed ($AgCl$).

  • Iodide ($I^-$): Acidify with dilute nitric acid ($HNO_3$), then add aqueous lead(II) nitrate ($Pb(NO_3)_2$) -> Yellow precipitate formed ($PbI_2$).

  • Nitrate ($NO_3^-$): Add aqueous sodium hydroxide ($NaOH$), then aluminium foil/powder; warm carefully -> Pungent gas produced turns moist red litmus paper blue ($NH_3$).

  • Sulphate ($SO_4^{2-}$): Acidify with dilute nitric acid ($HNO_3$), then add aqueous barium nitrate ($Ba(NO_3)_2$) -> White precipitate formed ($BaSO_4$).

  • Tests for Gases:

  • Ammonia ($NH_3$): Colourless, pungent gas -> Turns damp red litmus paper blue.

  • Carbon dioxide ($CO_2$): Colourless, odourless gas -> Forms a white precipitate in limewater.

  • Chlorine ($Cl_2$): Yellowish-green, pungent gas -> Bleaches damp litmus paper (turns damp blue litmus red, then bleaches white).

  • Hydrogen ($H_2$): Colourless, odourless gas -> Extinguishes a lighted splint with a 'pop' sound.

  • Oxygen ($O_2$): Colourless, odourless gas -> Relights a glowing splint.

  • Sulphur dioxide ($SO_2$): Colourless, pungent gas -> Turns orange acidified potassium dichromate(VI) ($K_2Cr_2O_7$) green.

The Particulate Nature of Matter

• The Three States of Matter

  • Characteristics:

  • Solid: Fixed shape, fixed volume, incompressible. Particles are closely packed in an orderly arrangement, held by very strong attraction forces, vibrating and rotating about fixed positions.

  • Liquid: No fixed shape (takes container shape), fixed volume, virtually incompressible. Particles are closely packed in a disorderly arrangement, held by strong attraction forces, moving and exchanging partners throughout the liquid.

  • Gas: No fixed shape, no fixed volume, highly compressible. Particles are far apart in a disorderly arrangement, held by weak attraction forces, moving rapidly and randomly in all directions.

  • The Kinetic Theory of Matter: States that all matter is composed of extremely small particles (atoms, ions, or molecules) that are in continuous motion.

• Changes in State

  • Physical changes caused by absorption or loss of thermal energy.

  • Phase Transitions:

  • Melting: Solid absorbs heat, particles vibrate vigorously until gaining sufficient kinetic energy to overcome fixed lattice positions at the melting point ($T_m$). Temperature remains constant during melting.

  • Freezing: Liquid loses kinetic energy, particles slow down and settle into fixed lattice positions at the freezing point ($T_f = T_m$). Temperature remains constant during freezing.

  • Boiling: Liquid absorbs heat at a fixed temperature (boiling point), overcoming intermolecular forces to form gas rapidly throughout the liquid volume with visible bubbles.

  • Evaporation: Liquid converts to gas at any temperature below the boiling point; occurs slowly only at the liquid surface without bubble formation.

  • Condensation: Gas loses thermal energy, particles slow down and draw close to reform liquid structure.

  • Sublimation: Solid transforms directly into gas without passing through liquid state (e.g., solid carbon dioxide / dry ice, iodine crystals $I_2$, mothballs / naphthalene).

  • Heating and Cooling Curves:

  • During state changes (flat horizontal plateaus on temperature-time graphs), temperature stays constant because absorbed heat energy is used to break structural bonds/forces (or released latent heat balances cooling loss).

Atomic Structure and Chemical Bonding

• Atomic Structure

  • Subatomic Particles:

  • Proton: Relative mass = $1$, relative charge = $+1$, located in nucleus.

  • Neutron: Relative mass = $1$, relative charge = $0$, located in nucleus.

  • Electron: Relative mass = $0.0005$ ($\frac{1}{1840}$), relative charge = $-1$, located in electronic shells.

  • Definitions:

  • Proton Number (Atomic Number, $Z$): Number of protons in the nucleus of an atom. Defines the element identity.

  • Nucleon Number (Mass Number, $A$): Total number of protons and neutrons in the nucleus of an atom.

  • Neutral Atom: Number of protons equals number of electrons.

  • Isotopes:

  • Atoms of the same element with the same proton number but different nucleon numbers (different numbers of neutrons).

  • Examples: Hydrogen isotopes ($^1_1H$ protium, $^2_1H$ deuterium, $^3_1H$ tritium); Carbon isotopes ($^{12}_6C$, $^{13}_6C$, $^{14}_6C$); Chlorine isotopes ($^{35}_{17}Cl$, $^{37}_{17}Cl$).

  • Electronic Configuration:

  • Electrons fill shells starting from the shell closest to the nucleus (Shell 1: max $2\,e^-$; Shell 2: max $8\,e^-$; Shell 3: max $8\,e^-$ for first 20 elements).

  • Outer shell electrons are valence electrons.

  • Period number = number of occupied electron shells.

  • Group number = number of valence electrons.

• Structure and Properties of Materials

  • Element: Pure substance containing only one type of atom.

  • Compound: Pure substance composed of two or more different elements chemically combined in a fixed proportion by mass. Has fixed melting and boiling points and distinct chemical properties.

  • Mixture: Two or more substances physically mixed in variable proportions. Retains individual component properties, melts/boils over a temperature range, and easily separates via physical techniques.

• Ionic Bonding

  • Occurs between metals and non-metals via complete electron transfer.

  • Metals lose valence electrons to form positive cations ($M \rightarrow M^{n+} + n e^-$).

  • Non-metals gain valence electrons to form negative anions ($X + m e^- \rightarrow X^{m-}$).

  • Oppositely charged ions are attracted by strong electrostatic forces, forming ionic bonds.

  • Structure: Giant ionic lattice structure (e.g., $NaCl$, $MgCl_2$).

  • Physical Properties:

  • High melting and boiling points: Large thermal energy required to overcome strong electrostatic attraction between ions in lattice.

  • Electrical Conductivity: Conducts electricity when molten or aqueous (mobile ions free to move); non-conductor in solid state (ions locked in fixed lattice positions).

  • Solubility: Generally soluble in water, insoluble in organic solvents.

• Covalent Bonding

  • Occurs between non-metal atoms via sharing of pairs of valence electrons to achieve stable noble gas configurations.

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

  • Examples: Hydrogen ($H_2$), Chlorine ($Cl_2$), Hydrogen chloride ($HCl$), Oxygen ($O_2$), Water ($H_2O$), Methane ($CH_4$), Carbon dioxide ($CO_2$).

  • Structure: Simple molecular structure.

  • Physical Properties:

  • Low melting and boiling points: Weak intermolecular forces (van der Waals forces) between neutral molecules require little energy to overcome. Strong covalent bonds inside molecules are not broken during melting/boiling.

  • Electrical Conductivity: Non-conductors in all states due to absence of mobile ions or free electrons.

  • Solubility: Generally insoluble in water, soluble in organic solvents (e.g., ethanol, petrol, trichloromethane).

Stoichiometry and the Mole Concept

• Formulae and Equations

  • Chemical Formulae:

  • Metals exist as individual atoms (e.g., $Na$, $Mg$, $Fe$).

  • Non-metals exist as diatomic molecules ($H_2$, $N_2$, $O_2$, $F_2$, $Cl_2$, $Br_2$, $I_2$) or polyatomic structures ($P_4$, $S_8$; represented as $P$ and $S$ in chemical equations).

  • Polyatomic ions: Ammonium ($NH_4^+$), Hydroxide ($OH^-$), Nitrate ($NO_3^-$), Carbonate ($CO_3^{2-}$), Sulphate ($SO_4^{2-}$), Phosphate ($PO_4^{3-}$), Hydrogencarbonate ($HCO_3^-$), Dichromate(VI) ($Cr_2O_7^{2-}$), Manganate(VII) ($MnO_4^-$).

  • State Symbols:

  • Solid ($s$), Liquid ($l$), Gas ($g$), Aqueous ($aq$).

  • Ionic Equations:

  • Written by expressing soluble ionic compounds as separated aqueous ions and cancelling spectator ions that remain unchanged on both sides.

• Relative Masses and Percentage Composition

  • Relative Atomic Mass ($A_r$): Weighted average mass of naturally occurring atoms of an element relative to $\frac{1}{12}$ of the mass of a Carbon-12 atom. Has no units.

  • Relative Molecular Mass ($M_r$): Sum of the relative atomic masses ($A_r$) of all atoms shown in a molecular or ionic formula. Has no units.

  • Percentage Mass of an Element in a Compound:     $\%\,\text{Mass of element} = \frac{\text{Number of atoms of element} \times A_r}{\text{Molar mass } M_r \text{ of compound}} \times 100\%$

• The Mole Concept

  • One mole contains $6 \times 10^{23}$ particles (Avogadro Constant).

  • Moles and Mass:     $\text{Number of moles } (mol) = \frac{\text{Mass } (g)}{\text{Molar mass } (g\,mol^{-1})}$

  • Moles and Gas Volume:

  • One mole of any gas occupies $24\,dm^3$ ($24,000\,cm^3$) at room temperature and pressure (r.t.p., $25^\circ\text{C}$ and $1\,atm$).     $\text{Number of moles of gas} = \frac{\text{Volume of gas in } cm^3}{24,000\,cm^3} = \frac{\text{Volume of gas in } dm^3}{24\,dm^3}$

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

  • Calculated by dividing mass (or percentage mass) by $A_r$ for each element, then reducing mole values to the simplest whole-number ratio.

  • Molecular Formula: The actual number of atoms of each element in one molecule.     $\text{Molecular Formula} = (\text{Empirical Formula})_n$     $n = \frac{M_r}{\text{Empirical Formula Mass}}$

  • Solution Concentration:

  • Concentration in $g\,dm^{-3}$:       $\text{Concentration } (g\,dm^{-3}) = \frac{\text{Mass of solute } (g)}{\text{Volume of solution } (dm^3)}$

  • Concentration in $mol\,dm^{-3}$:       $\text{Concentration } (mol\,dm^{-3}) = \frac{\text{Number of moles of solute } (mol)}{\text{Volume of solution } (dm^3)}$

  • Conversion formula:       $\text{Concentration } (g\,dm^{-3}) = \text{Concentration } (mol\,dm^{-3}) \times M_r$

  • Stoichiometric Calculations and Limiting Reactants:

  • Limiting reactant is completely consumed in a reaction, determining the maximum product mass or volume formed.

Chemical Reactions

• Speed of Reaction

  • Defined as the change in amount of reactant or product per unit time:     $\text{Speed of reaction} = \frac{\text{Change in amount of reactant or product}}{\text{Time}}$

  • Experimental Measurement:

  • Measuring volume of gas evolved over time using a gas syringe.

  • Measuring decrease in mass of reaction mixture over time on an electronic balance (loose cotton wool plug allows gas escape while preventing liquid spray).

  • Rate Curves and Gradient:

  • Gradient of volume-time or mass-time graph equals reaction speed.

  • Gradient is steepest at start ($t=0$) when speed is highest; gradient decreases as reactants are used up; gradient becomes zero when reaction stops.

  • Collision Theory:

  • Reaction requires reacting particles to collide with energy greater than or equal to Activation Energy ($E_a$).

  • Collisions yielding products are effective collisions. Speed of reaction is directly proportional to frequency of effective collisions.

  • Factors Affecting Speed of Reaction:

  • Particle Size: Smaller particle size -> Larger total surface area -> Higher collision frequency -> Increased reaction speed.

  • Concentration: Higher concentration -> More particles per unit volume -> Higher collision frequency -> Increased reaction speed.

  • Pressure (Gaseous reactions): Higher pressure -> Gas molecules compressed closer together -> Higher collision frequency -> Increased reaction speed.

  • Temperature: Higher temperature -> Particles gain kinetic energy, move faster AND a significantly larger fraction of particles possess energy $\ge E_a$ -> Higher effective collision frequency -> Increased reaction speed.

  • Catalyst: Substance that speeds up reaction without being chemically altered; provides an alternative reaction pathway with lower activation energy.

• Redox Reactions

  • Oxidation-Reduction Definitions:

  • Oxidation: Gain of oxygen, loss of hydrogen, loss of electrons, increase in oxidation state.

  • Reduction: Loss of oxygen, gain of hydrogen, gain of electrons, decrease in oxidation state.

  • Rules for Assigning Oxidation Numbers:

  • Uncombined elements = $0$.

  • Monoatomic ions = charge of ion.

  • Hydrogen in compounds = $+1$; Oxygen in compounds = $-2$ (except peroxides = $-1$).

  • Group I = $+1$; Group II = $+2$; Group VII = $-1$.

  • Sum of oxidation states in neutral compound = $0$; in polyatomic ion = ion charge.

  • Tests for Oxidising and Reducing Agents:

  • Test for Oxidising Agents: Add aqueous potassium iodide ($KI(aq)$). Iodide ions ($I^-$) are oxidised to iodine ($I_2$). Observation: Colourless solution turns reddish-brown ($2I^- \rightarrow I_2 + 2e^-$).

  • Test for Reducing Agents: Add acidified potassium dichromate(VI) ($K_2Cr_2O_7(aq)$). Dichromate(VI) ions ($Cr_2O_7^{2-}$) are reduced to chromium(III) ions ($Cr^{3+}$). Observation: Solution changes from orange to green ($Cr_2O_7^{2-} \rightarrow Cr^{3+}$).

  • Alternative Test for Reducing Agents: Add acidified potassium manganate(VII) ($KMnO_4(aq)$). Manganate(VII) ions ($MnO_4^-$) are reduced to manganese(II) ions ($Mn^{2+}$). Observation: Solution changes from purple to colourless.

• Energy Changes

  • Exothermic Reaction: Heat energy is released to surroundings ($\Delta H$ is negative). Surroundings temperature rises. Total enthalpy of products < Total enthalpy of reactants.

  • Endothermic Reaction: Heat energy is absorbed from surroundings ($\Delta H$ is positive). Surroundings temperature drops. Total enthalpy of products > Total enthalpy of reactants.

  • Bond Breaking and Making:

  • Bond breaking requires energy (Endothermic process, $\Delta H > 0$).

  • Bond making releases energy (Exothermic process, $\Delta H < 0$).

  • $\Delta H = \text{Energy absorbed to break bonds} - \text{Energy released making bonds}$.

Acids, Bases and Salts

• Characteristic Properties of Acids and Bases

  • Acids:

  • Substances that ionise in water to produce hydrogen ions ($H^+$).

  • Sour taste, pH $< 7$, turns blue litmus paper red.

  • Chemical Reactions:

    • Acid + Metal (above $H$ in reactivity series) -> Salt + Hydrogen gas ($H_2$).

    • Acid + Base / Alkali -> Salt + Water ($H_2O$).

    • Acid + Carbonate -> Salt + Carbon dioxide ($CO_2$) + Water ($H_2O$).

  • Bases and Alkalis:

  • Base: Metal oxide or hydroxide that reacts with acid to give salt and water only.

  • Alkali: Soluble base that dissolves in water producing hydroxide ions ($OH^-$).

  • Bitter taste, slippery feel, pH $> 7$, turns red litmus paper blue.

  • Chemical Reactions:

    • Alkali + Acid -> Salt + Water (Ionic equation: $H^+(aq) + OH^-(aq) \rightarrow H_2O(l)$).

    • Alkali + Ammonium salt -> Salt + Ammonia gas ($NH_3$) + Water ($H_2O$).

  • pH Scale and Universal Indicator:

  • pH $0-6$: Acidic; pH $7$: Neutral; pH $8-14$: Alkaline.

  • Universal Indicator colours: Red (strongly acidic), Orange/Yellow (weakly acidic), Green (neutral), Blue (weakly alkaline), Violet/Purple (strongly alkaline).

  • Control of Soil pH:

  • Excess soil acidity treated with slaked lime (solid calcium hydroxide $Ca(OH)_2$), lime ($CaO$), or limestone ($CaCO_3$).

  • Slaked lime is preferred because it is cheap and sparingly soluble; excess solid remains in soil without making soil overly alkaline.

• Oxides

  • Acidic Oxides: Non-metal oxides (e.g., $CO_2$, $SO_2$, $NO_2$, $SO_3$); dissolve in water forming acidic solutions, react with alkalis to form salt + water.

  • Basic Oxides: Metal oxides (e.g., $Na_2O$, $MgO$, $CaO$, $CuO$); react with acids to form salt + water.

  • Amphoteric Oxides: Metal oxides that react with BOTH acids and alkalis to form salt + water ($ZnO$, $Al_2O_3$, $PbO$).

  • Neutral Oxides: Non-metal oxides that do not react with acids or alkalis ($CO$, $H_2O$, $NO$).

• Preparation of Salts

  • Solubility Rules:

  • All Sodium ($Na^+$), Potassium ($K^+$), Ammonium ($NH_4^+$), and Nitrate ($NO_3^-$) salts are SOLUBLE.

  • All Chlorides are SOLUBLE except Silver chloride ($AgCl$) and Lead(II) chloride ($PbCl_2$).

  • All Sulphates are SOLUBLE except Barium sulphate ($BaSO_4$), Lead(II) sulphate ($PbSO_4$), and Calcium sulphate ($CaSO_4$, sparingly soluble).

  • All Carbonates are INSOLUBLE except Sodium carbonate ($Na_2CO_3$), Potassium carbonate ($K_2CO_3$), and Ammonium carbonate ($(NH_4)_2CO_3$).

  • All Hydroxides/Oxides are INSOLUBLE except Group I, Calcium hydroxide ($Ca(OH)_2$, sparingly soluble).

  • Preparation Methods:

  • Insoluble Salts (Precipitation): Mix two soluble aqueous solutions containing desired ions. Filter insoluble precipitate, wash residue with cold distilled water, dry between filter papers.

  • Soluble Salts (Not Group I or Ammonium): Action of acid on excess insoluble reactant (metal, metal oxide, hydroxide, or carbonate). Warm acid, add excess solid until no more reacts, filter excess solid, evaporate filtrate to saturation, cool to crystallise, filter crystals, dry between filter paper.

  • Soluble Group I and Ammonium Salts (Titration): Pipette alkali into conical flask with indicator. Titrate acid from burette until indicator endpoint colour change. Note exact acid volume. Repeat without indicator using exact volumes. Evaporate resulting solution to saturation, cool to crystallise, dry.

The Periodic Table

• Periodic Trends

  • Elements arranged in order of increasing proton number.

  • Periods (1-7): Horizontal rows representing number of occupied electron shells. Metallic character decreases across a period from left to right.

  • Groups (I-VIII/0): Vertical columns representing number of valence electrons. Elements in the same group share similar chemical properties.

• Group Properties

  • Group I Elements (Alkali Metals): $Li$, $Na$, $K$, $Rb$, $Cs$.

  • Soft metals, low densities (first three float on water), low melting points.

  • Reactivity increases down the group.

  • React vigorously with cold water: $2M(s) + 2H_2O(l) \rightarrow 2MOH(aq) + H_2(g)$. Produces strongly alkaline solutions.

  • Group VII Elements (Halogens): $F_2$, $Cl_2$, $Br_2$, $I_2$, $At_2$.

  • Diatomic non-metals.

  • Trends down group: Melting/boiling points increase; physical states change from gas ($F_2$, $Cl_2$) to liquid ($Br_2$) to solid ($I_2$); colour darkens ($F_2$ pale yellow, $Cl_2$ yellow-green, $Br_2$ red-brown, $I_2$ black solid / purple vapour).

  • Reactivity and oxidising power decrease down group.

  • Displacement Reaction: A more reactive halogen displaces a less reactive halogen from its aqueous halide solution (e.g., $Cl_2(g) + 2KI(aq) \rightarrow 2KCl(aq) + I_2(aq)$).

  • Group 0 Elements (Noble Gases): $He$, $Ne$, $Ar$, $Kr$, $Xe$, $Rn$.

  • Monoatomic unreactive gases.

  • Possess complete octet/duplet valence shells.

  • Uses: Argon in light bulbs and steel manufacturing; Neon in glowing advertising signs; Helium in weather balloons.

  • Transition Metals:

  • High melting points, high densities, hard and strong.

  • Form coloured compounds ($Cu^{2+}$ blue, $Fe^{2+}$ pale green, $Fe^{3+}$ reddish-brown/yellow).

  • Variable oxidation states ($Cu^+/Cu^{2+}$, $Fe^{2+}/Fe^{3+}$).

  • Act as industrial catalysts ($Ni$ in margarine hydrogenation; $Fe$ in Haber process).

Metals

• Properties of Metals and Alloys

  • Pure metals: Malleable and ductile because identical metal atoms are arranged in regular layers that slide over one another when force is applied.

  • Alloys: Mixture of a metal with another element. Atoms of different sizes disrupt the regular lattice, preventing layers from sliding easily -> Increases hardness, strength, and corrosion resistance.

  • Examples: Brass ($Cu + Zn$); Stainless Steel ($Fe + Cr + Ni + C$); Pewter ($Sn + Sb + Cu$).

• Reactivity Series

  • Reactivity Order: Potassium ($K$) > Sodium ($Na$) > Calcium ($Ca$) > Magnesium ($Mg$) > Aluminium ($Al$) > Zinc ($Zn$) > Iron ($Fe$) > Lead ($Pb$) > [Hydrogen ($H$)] > Copper ($Cu$) > Silver ($Ag$) > Gold ($Au$).

  • Reactions with Cold Water, Steam, and Acids:

  • $K, Na, Ca$: React with cold water to form metal hydroxide + $H_2$. Explosive/vigorous reaction with dilute acid.

  • $Mg, Zn, Fe$: React with steam to form metal oxide + $H_2$. React with dilute acid to form salt + H_2$.\n - Pb, Cu, Ag, Au: No reaction with cold water, steam, or dilute acids.\n - **Metal Displacement**: A more reactive metal displaces a less reactive metal ion from its aqueous salt solution.\n\n- **Extraction of Metals**\n - Reactive metals (K, Na, Ca, Mg, Al): Compounds very stable; extracted by electrolysis of molten ores.\n - Moderately reactive metals (Zn, Fe, Pb, Cu): Extracted by reduction of metal oxides using carbon/coke or carbon monoxide.\n - Unreactive metals (Ag, Au): Found native (uncombined element).\n - Metal Recycling: Conserves finite ores, reduces air/water pollution from mining, saves energy, reduces landfill waste.\n\n- **Iron and Steel**\n - Blast Furnace Extraction of Iron from Haematite (Fe_2O_3):\n - Raw materials: Haematite (Fe_2O_3$), Coke ($C$), Limestone ($CaCO_3), Hot air.\n - Reactions:\n 1. C(s) + O_2(g) \rightarrow CO_2(g)\n 2. CO_2(g) + C(s) \rightarrow 2CO(g)\n 3. Fe_2O_3(s) + 3CO(g) \rightarrow 2Fe(l) + 3CO_2(g) (Carbon monoxide acts as reducing agent).\n 4. CaCO_3(s) \rightarrow CaO(s) + CO_2(g) (Limestone decomposes).\n 5. CaO(s) + SiO_2(s) \rightarrow CaSiO_3(l)$(Basic$CaO neutralises acidic silica impurity to form liquid slag floating on molten iron).\n - Steels: Low carbon / mild steel (soft, malleable; car bodies, machinery); High carbon steel (strong, brittle; cutting tools); Stainless steel (corrosion resistant; cutlery, chemical plants, surgical tools).\n\n- **Prevention of Rusting**\n - Rusting conditions: Oxygen AND Water required. Rust formula: Hydrated iron(III) oxide (Fe_2O_3 \cdot xH_2O).\n - Prevention Methods:\n - Barrier methods: Painting, greasing, plastic coating, tin plating.\n - Sacrificial Protection: Attaching or coating iron with a more reactive metal (e.g. Zinc in galvanising; Magnesium/Zinc blocks attached to underground pipes or ship hulls). The more reactive metal corrodes preferentially by losing electrons in place of iron.\n\n- **Aluminium**\n - Protective Layer: Unreactive appearance caused by a thin, tough, non-porous layer of aluminium oxide (Al_2O_3) that seals the metal surface.\n - Properties: Low density, high electrical and thermal conductivity, corrosion resistant, strong.\n - Uses: Aircraft bodies (duralumin alloy), food containers, overhead power cables.\n\n\n# Air and Water\n\n- **Air**\n - Composition of Clean, Dry Air: Nitrogen 79\%$, Oxygen$20\%$, Argon$0.93\%$, Carbon dioxide$0.04\%, variable water vapour.\n - Percentage Oxygen Determination: Air passed repeatedly over heated copper inside gas syringes; volume decreases by approx 20\%$as copper oxidises ($2Cu(s) + O_2(g) \rightarrow 2CuO(s)).\n - Uses of Oxygen: Steel manufacturing, hospital oxygen tents, oxy-acetylene welding flames.\n - Air Pollutants:\n - **Carbon monoxide (CO)**: Incomplete combustion of carbon fuels. Colourless, odourless, toxic gas; binds with haemoglobin forming carboxyhaemoglobin, blocking oxygen transport.\n - **Nitrogen oxides (NO, NO_2$)**: Internal combustion engines and lightning ($N_2 + O_2 \rightarrow 2NO$;$2NO + O_2 \rightarrow 2NO_2$). Causes lung damage and forms acid rain ($HNO_3).\n - **Sulphur dioxide (SO_2$)**: Fossil fuel combustion containing sulphur and volcanic eruptions. Causes breathing irritation and forms acid rain ($H_2SO_4).\n - **Methane (CH_4) & Unburnt Hydrocarbons**: Landfill decay, engine exhaust. Greenhouse gas causing global warming and photochemical smog.\n - **Catalytic Converters**: Exhaust attachments using platinum catalysts to convert CO$,$NO_x$, and hydrocarbons into harmless$CO_2$,$N_2$, and$H_2O$.

• Water

  • Purification Process:

  1. Screening removes large debris.

  2. Addition of Alum (causes fine particles to coagulate) and Lime (reduces acidity) in settling tanks.

  3. Addition of Activated Carbon to eliminate foul odors and taste.

  4. Sand/Gravel Bed Filtration removes remaining solids.

  5. Chlorination kills harmful bacteria.

  • Desalination: Removing salt from seawater via Distillation or Reverse Osmosis (NEWater).

Organic Chemistry

• Crude Oil / Petroleum

  • Non-renewable fossil fuel mixture of hydrocarbons.

  • Fractional Distillation Fractions (Top to Bottom):

  1. Petroleum Gas ($C_1-C_4$, b.p. $</p></li><li><p>Petrol / Gasoline ($C_5-C_10$, b.p.$40-75^\circ\text{C}$): Fuel for motor cars.</p></li><li><p>Naphtha ($C_7-C_14$, b.p.$75-150^\circ\text{C}$): Chemical feedstock for petrochemical industry.</p></li><li><p>Paraffin / Kerosene ($C_{11}-C_{16}$, b.p.$160-250^\circ\text{C}$): Aircraft fuel, heating/cooking fuel.</p></li><li><p>Diesel ($C_{16}-C_{20}$, b.p.$250-300^\circ\text{C}$): Fuel for diesel engines (lorries, trains).</p></li><li><p>Lubricating Oil ($C_{20}-C_{35}$, b.p.$300-350^\circ\text{C}$): Lubricants, polishes, waxes.</p></li><li><p>Bitumen ($>C_{35}$, b.p.$>350^\circ\text{C}$): Road surfacing.</p></li></ol></li><li><p><strong>Homologous Series</strong></p><ul><li><p>Family of organic compounds sharing the same functional group and general formula.</p></li><li><p>Characteristics: Same general formula, similar chemical properties, gradual change in physical properties (m.p., b.p., density increase down series; flammability decreases; viscosity increases), consecutive members differ by$-CH_2-$($14$mass units).</p></li></ul></li><li><p><strong>Alkanes</strong></p><ul><li><p>General formula$C_n H_{2n+2}$($n \ge 1$). Saturated hydrocarbons containing single$C-C$and$C-H$covalent bonds.</p></li><li><p>Unreactive except for:</p></li><li><p><strong>Combustion</strong>: Complete combustion yields$CO_2 + H_2O$; incomplete combustion yields$CO, C, H_2O$</p></li><li><p><strong>Substitution</strong>: Reacts with chlorine gas in sunlight/UV light ($CH_4 + Cl_2 \xrightarrow{UV} CH_3Cl + HCl$).</p></li></ul></li><li><p><strong>Alkenes</strong></p><ul><li><p>General formula$C_n H_{2n}$($n \ge 2$). Unsaturated hydrocarbons containing carbon-carbon double bonds ($-C=C-).

  2. Cracking: Catalytic breakdown ($Al_2O_3/SiO_2 catalyst, $600^\circ\text{C}$) of long alkane molecules into smaller alkanes, alkenes, or hydrogen gas.\n - **Addition Reactions**:\n - **Bromine Water Test**: Alkenes rapidly decolourise reddish-brown aqueous bromine ($C_2H_4 + Br_2 \rightarrow C_2H_4Br_2$). Alkanes do not react.\n - **Addition of Steam**: $C_2H_4 + H_2O(g) \xrightarrow{300^\circ\text{C}, 65\,atm, H_3PO_4} C_2H_5OH$ (Forms alcohol).\n - **Addition of Hydrogen (Hydrogenation)**: $C_2H_4 + H_2 \xrightarrow{Ni, 200^\circ\text{C}} C_2H_6$ (Forms alkane; used to harden vegetable oils into margarine).\n\n- **Alcohols**\n - General formula $C_n H_{2n+1}OH$ ($n \ge 1$). Functional group $-O-H$. Neutral liquids.\n - Ethanol Manufacture:\n - Catalytic addition of steam to ethene.\n - Fermentation of Glucose: Glucose solution with yeast at $37^\circ\text{C}$ in absence of air ($C_6H_{12}O_6 \xrightarrow{yeast} 2C_2H_5OH + 2CO_2$).\n - Chemical Properties: Combustion; Oxidation to carboxylic acids using acidified $K_2Cr_2O_7$ or atmospheric oxygen/bacteria.\n\n- **Carboxylic Acids**\n - General formula $C_n H_{2n+1}COOH$ ($n \ge 0$). Functional group $-COOH$. Weak organic acids.\n - Ethanoic Acid Preparation: Oxidation of ethanol.\n - **Esterification**: Carboxylic acid + Alcohol $\rightleftharpoons$ Ester + Water (in presence of conc. $H_2SO_4$ catalyst).\n - Example: Ethanoic acid + Ethanol $\rightleftharpoons$ Ethyl ethanoate + Water ($CH_3COOH + C_2H_5OH \rightleftharpoons CH_3COOCH_2CH_3 + H_2O$).\n\n- **Macromolecules**\n - **Addition Polymers**: Polymer formed by joining unsaturated monomers containing $-C=C-$ bonds without losing atoms (e.g., Poly(ethene), Poly(vinyl chloride) / PVC, Poly(tetrafluoroethene) / PTFE, Polystyrene).\n - **Condensation Polymers**:\n - **Terylene (Polyester)**: Formed from diol + dicarboxylic acid with removal of water molecules. Contains ester linkages ($-COO-$).\n - **Nylon (Polyamide)**: Formed from diamine + dicarboxylic acid with removal of water molecules. Contains amide linkages ($-CONH-$).\n - Pollution Problems: Non-biodegradable plastics accumulate in landfills and produce toxic gases upon incineration.\n - **Natural Macromolecules**:\n - **Proteins**: Natural polyamides containing amide linkages ($-CONH-$) joining amino acid monomers.\n - **Fats**: Natural polyesters containing ester linkages ($-COO-$) joining glycerol and fatty acids.

  3. Carbohydrates / Starch: Natural polymers composed of glucose/sugar monomers joined by ether/oxygen linkages.