GCE O-Level Chemistry Matters Comprehensive Study Guide

Measurement of Physical Quantities and Laboratory Apparatus

• Physical Quantities and SI Units:

  • Time: SI unit is the second ($s$). Other common units include minute ($min$) ($1\,min = 60\,s$) and hour ($h$) ($1\,h = 60\,min = 3600\,s$).

    • Digital stopwatch accuracy: $\pm 0.01\,s$ (one hundredth of a second).

    • Analogue stopwatch accuracy: $\pm 0.1\,s$ (one tenth of a second). The smaller dial measures minutes, while the larger dial measures seconds.

  • Temperature: SI unit is the kelvin ($K$). Commonly measured in degrees Celsius (${^\circ}\text{C}$).

    • Conversion formula: $\text{Temperature in } K = \text{Temperature in } {^\circ}\text{C} + 273$.

    • Temperatures measured in kelvin are never negative, and no degree sign (${^\circ}$) is used before $K$.

    • Laboratory apparatus: Alcohol thermometers (range approx. $-10\,{^\circ}\text{C}$ to $110\,{^\circ}\text{C}$), mercury thermometers, digital thermometers, and temperature sensors/probes connected to data loggers.

  • Length: SI unit is the metre ($m$).

    • Unit conversions: $1\,m = 10\,dm = 100\,cm = 1000\,mm$; $1\,\mu m = 10^{-6}\,m$.

    • Laboratory apparatus: Metre rule (accuracy $\pm 0.1\,cm$) and measuring tapes (accuracy ranges from $\pm 0.1\,cm$ to $\pm 0.5\,cm$).

  • Mass: SI unit is the kilogram ($kg$).

    • Unit conversions: $1\,kg = 1000\,g = 1\,000\,000\,mg$; $1\,t = 1000\,kg = 1\,000\,000\,g$; $1\,kt = 1000\,t = 1\,000\,000\,kg$.

    • Laboratory apparatus: Electronic balance (accuracy $\pm 0.01\,g$) and beam balance (accuracy depends on calibration marks and counterweights).

  • Volume: SI unit is the cubic metre ($m^3$).

    • Unit conversions: $1\,m^3 = 1000\,dm^3 = 1\,000\,000\,cm^3$; $1\,dm^3 = 1000\,cm^3 = 1000\,mL = 1\,L$

    • Laboratory apparatus for liquids:

      • Pipette: Measures accurate fixed volumes (e.g., $10.0\,cm^3$ or $25.0\,cm^3$).

      • Volumetric Flask: Measures accurate fixed volumes that are larger (e.g., $100\,cm^3$ or $250\,cm^3$).

      • Measuring Cylinder: Measures a range of volumes to the nearest $0.5\,cm^3$ (e.g., $31.5\,cm^3$).

      • Burette: Measures a range of volumes to the nearest $0.05\,cm^3$ (e.g., $31.55\,cm^3$).

    • Laboratory apparatus for gases: Gas Syringe (typical laboratory capacity up to $100\,cm^3$).

  • Avoiding Parallax Error:

    • Position the eye level aligned with the meniscus curve.

    • Concave meniscus (upward curving, e.g., water): Read from the bottom of the meniscus curve.

    • Convex meniscus (downward curving, e.g., mercury): Read from the top of the meniscus curve.

Collection, Purification, and Separation of Substances

• Methods for Collecting Gases:

  • The collection method depends on gas solubility in water and density relative to air (taking relative molecular mass of air as $\approx 30$).

  • Displacement of Water: Suitable for gases that are insoluble to slightly soluble in water ($H_2$, $O_2$, $CO_2$). Density does not affect collection.

  • Downward Delivery: Suitable for gases that are denser than air ($M_r > 30$) and can be soluble or insoluble in water ($Cl_2$, $HCl$, $SO_2$). Gas displaces air at the bottom of the jar.

  • Upward Delivery: Suitable for gases that are less dense than air ($M_r < 30$) and can be soluble or insoluble in water ($NH_3$). Gas displaces air at the top of the jar.

• Methods for Drying Gases:

  • Concentrated Sulfuric Acid ($H_2SO_4$): Dries most gases including $Cl_2$ and $HCl$. Unsuitable for basic gases like $NH_3$ which react with sulfuric acid.

  • Quicklime / Calcium Oxide ($CaO$): Dries basic gases like $NH_3$. Must be freshly heated before use because it absorbs moisture and $CO_2$ from the air. Unsuitable for acidic gases like $CO_2$ which react with calcium oxide.

  • Fused Calcium Chloride ($CaCl_2$): Dries neutral gases like $H_2$, $N_2$, and $CO_2$. Must be freshly heated. Unsuitable for gases like $NH_3$ which react with calcium chloride.

• Separation Techniques for Solid-Solid Mixtures:

  • Magnetic Attraction: Separates magnetic solids (e.g., iron, cobalt, nickel, steel) from non-magnetic solids (e.g., sulfur, sand, plastic).

  • Sieving: Separates solids with different particle sizes using a mesh/sieve with a specific pore size.

  • Using Suitable Solvents: Separates mixtures where one solid dissolves in a chosen solvent (solute) and the other remains insoluble (e.g., dissolving salt-sand mixture in water to dissolve salt).

  • Sublimation: Separates a solid that changes directly from solid to gas upon heating (e.g., iodine, naphthalene, dry ice $CO_2$) from heat-stable solids. Gas deposits onto a cool surface as a pure sublimate.

• Separation Techniques for Solid-Liquid Mixtures:

  • Filtration: Separates insoluble solids (residue) from liquids (filtrate) using filter funnel and filter paper.

  • Evaporation to Dryness: Separates a heat-stable dissolved solute from its solvent by heating until all liquid vaporises.

  • Crystallisation: Obtains a pure, heat-sensitive solid from its saturated solution (a solution where no more solute can dissolve at a given temperature). The solution is gently heated to saturation, cooled gradually to allow crystals to form, filtered, washed with cold distilled water, and dried between filter papers.

  • Simple Distillation: Separates and collects a pure solvent (distillate) from a solution. Boiling chips ensure smooth boiling. The thermometer bulb is placed at the exit sidearm to measure escaping vapour temperature. Water enters the condenser at the bottom and exits at the top to ensure maximum cooling.

• Separation Techniques for Liquid-Liquid Mixtures:

  • Separating Funnel: Separates heterogeneous mixtures of immiscible liquids (e.g., oil and water). The denser liquid forms the bottom layer/phase and is drained first via the stopcock tap.

  • Paper Chromatography: Separates miscible components with different solubilities in a given solvent.

    • Sample spot is placed on a pencil start line (pencil graphite is insoluble in solvent; ink would dissolve and contaminate).

    • Spot must be kept above the initial solvent level.

    • More soluble components travel faster towards the solvent front.

    • $R_f$ (Retention Factor) value formula: $\text{R}_f = \frac{\text{distance travelled by the substance}}{\text{distance travelled by the solvent}}$ . $\text{R}_f$ value is constant for a specific substance under identical solvent and temperature conditions.

    • Locating Agents: Chemicals sprayed onto chromatograms of colourless compounds (e.g., amino acids, sugars) or viewed under ultraviolet (UV) light to form visible coloured spots.

  • Fractional Distillation: Separates miscible liquids with different boiling points.

    • A fractionating column containing glass beads provides a large surface area for repeated condensation and evaporation.

    • The liquid with the lower boiling point distils over first at its constant boiling point.

    • Industrial applications: Oil refineries (crude oil processing), liquid air separation ($N_2$, $O_2$, $Ar$), ethanol extraction from glucose fermentation.

Determination of Substance Purity

• Purity Concepts:

  • A pure substance has a sharp, fixed melting point and boiling point under standard conditions.

  • Impurities affect physical properties:

    • Lower and broaden the melting point range.

    • Raise and broaden the boiling point range.

  • The greater the amount of impurities, the larger the deviation and range in melting or boiling points.

  • Importance of purity: Silicon for computer chips requires high purity for electrical function; impurities in pharmaceutical drugs cause toxic or harmful side effects.

Kinetic Particle Theory and States of Matter

• Core Postulates of Kinetic Particle Theory:

  • All matter is made up of tiny, discrete particles (atoms, molecules, or ions) that are in constant, random motion.

• Comparison of States of Matter:

  • Solid:

    • Particle Arrangement: Very closely packed in an orderly, regular lattice.

    • Attractive Forces: Very strong.

    • Kinetic Energy: Very low.

    • Motion: Vibrate and rotate only about fixed positions.

    • Shape and Volume: Definite shape and definite volume. Incompressible.

  • Liquid:

    • Particle Arrangement: Closely packed in a disorderly, random manner.

    • Attractive Forces: Strong, but less strong than in solids.

    • Kinetic Energy: Low.

    • Motion: Slide past one another freely throughout the liquid.

    • Shape and Volume: Indefinite shape (takes container shape), definite volume. Incompressible.

  • Gas:

    • Particle Arrangement: Very far apart in a disorderly, random manner.

    • Attractive Forces: Very weak / negligible.

    • Kinetic Energy: High.

    • Motion: Move rapidly and randomly in all directions.

    • Shape and Volume: Indefinite shape, indefinite volume. Compressible.

• Phase Transitions and Energy Changes:

  • Melting (Solid to Liquid): Thermal energy absorbed is converted to kinetic energy until melting point; at melting point, thermal energy is converted to potential energy to overcome attractive forces. Temperature remains constant during phase change.

  • Freezing (Liquid to Solid): Energy is released to surroundings; attractive forces pull particles into an orderly arrangement at constant freezing point.

  • Boiling (Liquid to Gas): Occurs throughout the liquid at a fixed boiling point ($100\,{^\circ}\text{C}$ for pure water; $343\,{^\circ}\text{C}$ for icosane). Thermal energy overcomes intermolecular attractive forces completely.

  • Evaporation: Conversion from liquid to gas at any temperature occurring only at the liquid surface.

  • Condensation (Gas to Liquid): Energy released; particles slow down and move closer together.

  • Sublimation (Solid directly to Gas): Thermal energy absorbed directly overcomes lattice forces (e.g., dry ice $CO_2$ sublimes above $-78\,{^\circ}\text{C}$).

  • Vapour Deposition (Gas directly to Solid): Gas particles lose energy and deposit directly into a solid lattice (e.g., iodine vapour forming crystals on a cold surface).

Diffusion and Movement of Particles

• Diffusion: The net movement of particles from a region of higher concentration to a region of lower concentration down a concentration gradient until uniform concentration is reached.

• Factors Affecting Rate of Diffusion:

  • Temperature: Higher temperature increases the average kinetic energy of particles, causing them to move faster and increasing the rate of diffusion.

  • Relative Molecular Mass ($M_r$): Particles with lower relative molecular mass ($M_r$) move and diffuse faster than heavier particles at the same temperature.

    • Demonstration: Cotton wool soaked in concentrated $HCl(aq)$ (releases $HCl$ gas, $M_r = 36.5$) and concentrated $NH_3(aq)$ (releases $NH_3$ gas, $M_r = 17$) placed at opposite ends of a glass tube. A white ring of ammonium chloride ($NH_4Cl$) forms closer to the $HCl$ end because lighter $NH_3$ molecules diffuse faster than heavier $HCl$ molecules.

Sub-atomic Structure of Atoms

• Sub-atomic Particles:

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

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

  • Electron: Relative mass = $\frac{1}{1840}$, relative charge = $-1$, location = electron shells.

• Atom Characteristics:

  • An atom is the smallest particle that retains the chemical characteristics of an element.

  • Atoms are electrically neutral because the number of protons equals the number of electrons.

  • Nucleus radius is at least $100\,000$ times smaller than the radius of the atom.

• Definitions and Notation:

  • Proton Number ($Z$): The number of protons in the nucleus of an atom (also called atomic number).

  • Nucleon Number ($A$): The total number of protons and neutrons in the nucleus of an atom (also called mass number).

  • $\text{Number of neutrons} = A - Z$.

  • Nuclide Notation: $\text{}^{A}_{Z}\text{X}$.

  • Ions: Formed when atoms gain or lose electrons. Cations ($+$, lost electrons); Anions ($-$, gained electrons).

  • Isotopes: Atoms of the same element with the same proton number ($Z$) but different nucleon numbers ($A$) (i.e., different numbers of neutrons).

    • Chemical Properties: Identical because they have the same number and arrangement of valence electrons.

    • Physical Properties: Differ slightly (e.g., density, melting point, boiling point) due to mass differences.

    • Examples: Hydrogen-1 ($\text{}^{1}_{1}\text{H}$), Hydrogen-2 ($\text{}^{2}_{1}\text{H}$), Hydrogen-3 ($\text{}^{3}_{1}\text{H}$); Chlorine-35 ($\text{}^{35}_{17}\text{Cl}$) and Chlorine-37 ($\text{}^{37}_{17}\text{Cl}$).

Electronic Configuration and the Periodic Table

• Electron Distribution Rules:

  • First electron shell: Maximum $2$ electrons.

  • Second electron shell: Maximum $8$ electrons.

  • Third electron shell: Maximum $8$ electrons (for first 18 elements).

• Periodic Table Correspondence:

  • Period Number: Equals the number of occupied electron shells.

  • Group Number: Equals the number of valence electrons (Groups 1–2 have 1–2; Groups 13–18 have 3–8, except Helium which has 2 valence electrons in Group 18).

  • Elements in the same group possess identical numbers of valence electrons and exhibit similar chemical properties.

Chemical Bonding: Ionic, Covalent, and Metallic

• Driving Force of Bonding:

  • Atoms gain, lose, or share electrons to attain a stable noble gas electronic configuration (duplet or octet configuration).

• Ionic Bonding:

  • Occurs between metal atoms (which lose electrons to form positive cations) and non-metal atoms (which gain electrons to form negative anions).

  • Ionic Bond: The mutual electrostatic attraction between oppositely charged ions.

  • Formed structures exist as a Giant Ionic Crystal Lattice.

  • Polyatomic Ions: Ammonium ($NH_4^+$), Hydroxide ($OH^-$), Nitrate ($NO_3^-$), Sulfate ($SO_4^{2-}$), Carbonate ($CO_3^{2-}$), Phosphate ($PO_4^{3-}$), Manganate(VII) ($MnO_4^-$).

• Covalent Bonding:

  • Occurs between non-metal atoms via the sharing of electron pairs (bonding pairs).

  • Covalent Bond: Strong electrostatic attraction between the shared pair of electrons and the positively charged nuclei of the bonding atoms.

  • Valency: The number of electron pairs an atom shares to achieve a stable octet/duplet.

  • Single bond ($1$ shared pair, e.g., $H_2$, $Cl_2$, $H_2O$), Double bond ($2$ shared pairs, e.g., $O_2$, $CO_2$), Triple bond ($3$ shared pairs, e.g., $N_2$).

• Metallic Bonding:

  • Occurs in solid metals.

  • Metallic Bond: The mutual electrostatic attraction between positively charged metal ions and the surrounding "sea of mobile (delocalised) electrons".

Structure and Properties of Materials

• Classification of Matter:

  • Element: A pure substance that cannot be broken down into simpler substances by chemical methods.

  • Compound: A pure substance containing two or more elements chemically combined in a fixed ratio.

  • Mixture: Two or more substances physically combined in any ratio, retaining their individual properties.

• Properties of Ionic Substances:

  • Structure: Giant ionic crystal lattice.

  • Melting and Boiling Points: Very high, because large amounts of thermal energy are required to overcome the strong electrostatic forces of attraction between oppositely charged ions.

  • Hardness: Hard but brittle. When sufficient force is applied, layers of ions slide relative to each other, bringing ions of like charges together; repulsive forces cause the crystal lattice to shatter.

  • Solubility: Usually soluble in water (aqueous), insoluble in organic solvents.

  • Electrical Conductivity: Conducts electricity in molten ($l$) and aqueous ($aq$) states due to the presence of mobile ions. Does NOT conduct electricity in the solid ($s$) state because ions are fixed in lattice positions.

• Properties of Covalent Substances:

  • Simple Covalent Molecules (e.g., $O_2$, $H_2O$, $CO_2$, $CH_4$, $I_2$):

    • Structure: Simple molecular structure with strong covalent bonds within molecules and weak intermolecular forces between molecules.

    • Melting and Boiling Points: Low, because only a small amount of energy is needed to overcome the weak intermolecular forces of attraction. Volatile.

    • Electrical Conductivity: Do not conduct electricity in any state because they consist of neutral molecules without mobile ions or delocalised electrons (Exception: gases like $HCl$ that ionise when dissolved in water).

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

  • Giant Covalent Structures:

    • Diamond: Each carbon atom is covalently bonded to four other carbon atoms in a rigid three-dimensional tetrahedral network. Extremely hard, high melting point, electrical insulator (no free electrons).

    • Graphite: Allotropes of carbon. Each carbon atom is covalently bonded to three other carbon atoms in hexagonal layers. High melting point. Soft and slippery because weak intermolecular forces between layers allow them to slide past each other easily. Conducts electricity along layers due to one delocalised electron per carbon atom.

    • Silicon Dioxide ($SiO_2$ / Silica): Each silicon atom is covalently bonded to four oxygen atoms, and each oxygen atom is bonded to two silicon atoms ($1:2$ ratio). Hard, high melting point, electrical insulator.

  • Macromolecules (Polymers) (e.g., poly(ethene), nylon, Terylene):

    • Long chains of covalent molecules. Soften over a range of temperatures; non-conductors; insoluble in water, soluble in organic solvents.

• Properties of Metals and Alloys:

  • Pure Metals:

    • Regular lattice arrangement of identical positive ions.

    • Layers of atoms/ions slide past each other easily when force is applied $\rightarrow$ Malleable and Ductile.

    • Good conductors of heat and electricity due to the presence of a sea of mobile (delocalised) electrons.

  • Alloys:

    • A mixture of a metal with one or more other elements (e.g., steel, brass, bronze, stainless steel).

    • Irregular lattice arrangement: Atoms of different sizes disrupt the regular arrangement of metal ions, preventing layers from sliding past each other easily $\rightarrow$ Harder, stronger, less malleable and less ductile than pure metals. Melts over a temperature range.

Chemical Formulae, Valency, and Equations

• Valency and Crossover Method:

  • Valency corresponds to the ionic charge magnitude or the number of shared electron pairs.

  • Group 1 ($1$), Group 2 ($2$), Group 13 ($3$), Group 14 ($4$), Group 15 ($3$), Group 16 ($2$), Group 17 ($1$).

  • Transition metal valencies indicated by Roman numerals: Iron(II) ($Fe^{2+}$), Iron(III) ($Fe^{3+}$), Copper(I) ($Cu^+$), Copper(II) ($Cu^{2+}$), Titanium(IV) ($Ti^{4+}$). Zinc ($Zn^{2+}$) and Silver ($Ag^+$) are fixed.

• Constructing Equations:

  • Balanced Chemical Equations: Atoms of each element are conserved ($\text{Left Side} = \text{Right Side}$).

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

  • Ionic Equations: Dissociate soluble aqueous ionic compounds into their constituent ions. Cancel spectator ions (ions appearing identical in charge and state on both sides).

The Mole Concept, Stoichiometry, and Solutions

• Relative Mass Definitions:

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

  • Relative Molecular Mass ($M_r$): The average mass of one molecule of a substance relative to $\frac{1}{12}$ the mass of an atom of carbon-12.

  • Relative Formula Mass ($M_r$): Used for ionic compounds (sum of $A_r$ of all atoms in the formula unit).

• Molar Relationships and Formulas:

  • Avogadro Constant: $1\text{ mole} = 6.02 \times 10^{23}\text{ particles}$.

    • $\text{Number of moles} = \frac{\text{number of particles}}{6.02 \times 10^{23}}$

  • Molar Mass: Mass of one mole of a substance ($g/mol$).

    • $\text{Number of moles} = \frac{\text{mass in } g}{\text{molar mass in } g/mol}$

  • Molar Volume of Gas: $1\text{ mole}$ of any gas occupies $24\,dm^3$ at room temperature and pressure (r.t.p.: $25\,{^\circ}\text{C}$ and $1\text{ atm}$).

    • $\text{Number of moles of gas} = \frac{\text{volume of gas in } dm^3}{24\,dm^3}$

  • Concentration of Solutions:

    • $\text{Concentration in } g/dm^3 = \frac{\text{mass of solute in } g}{\text{volume of solution in } dm^3}$

    • $\text{Concentration in } mol/dm^3 = \frac{\text{number of moles of solute in } mol}{\text{volume of solution in } dm^3}$

    • $\text{Concentration in } mol/dm^3 = \frac{\text{concentration in } g/dm^3}{\text{molar mass in } g/mol}$

• Stoichiometry and Volumetric Analysis:

  • Mole ratio obtained directly from balanced chemical equations.

  • Avogadro's Law for gases: Equal volumes of gases at the same temperature and pressure contain equal numbers of moles ($\frac{V_1}{V_2} = \frac{n_1}{n_2}$).

  • Limiting Reactant: The reactant that is completely consumed first in a chemical reaction; limits the amount of products formed. The reactant remaining is in excess.

Empirical and Molecular Formula Calculations

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

  • Calculation Steps: Mass or % mass $\rightarrow$ divide by $A_r$ to find moles $\rightarrow$ divide by smallest mole value to find mole ratio $\rightarrow$ simplify to whole numbers.

• Molecular Formula: The exact number of atoms of each element in one molecule of a covalent compound.

  • $\text{Molecular Formula} = (\text{Empirical Formula})_n$

  • $n = \frac{\text{Relative Molecular Mass } (M_r)}{\text{Mass of Empirical Formula}}$

Acid-Base Chemistry: Properties, Strength, and pH

• Acids:

  • Definition: A substance that produces hydrogen ions, $H^+$, in aqueous solution.

  • Physical Properties: Sour taste, conducts electricity in aqueous solution, turns blue litmus paper red.

  • Chemical Reactions:

    1. $\text{Acid} + \text{Reactive Metal} \rightarrow \text{Salt} + \text{Hydrogen Gas } (H_2)$.

       • Test for $H_2$: Extinguishes a burning splint with a 'pop' sound.

       • Unreactive metals ($Cu$, $Ag$) do not react. Lead ($Pb$) forms an insoluble layer ($PbCl_2$ or $PbSO_4$) halting reaction.

    2. $\text{Acid} + \text{Base (metal oxide/hydroxide)} \rightarrow \text{Salt} + \text{Water } (H_2O)$ (Neutralisation).

    3. $\text{Acid} + \text{Carbonate} \rightarrow \text{Salt} + \text{Water} + \text{Carbon Dioxide } (CO_2)$.

       • Test for $CO_2$: Forms a white precipitate ($CaCO_3$) when bubbled into limewater ($Ca(OH)_2(aq)$).

• Acid Strength vs Concentration:

  • Strong Acid: An acid that completely ionises/dissociates in aqueous solution (e.g., $HCl \rightarrow H^+ + Cl^-$, $HNO_3$, $H_2SO_4$).

  • Weak Acid: An acid that only partially ionises in aqueous solution (e.g., $CH_3COOH \rightleftharpoons CH_3COO^- + H^+$, citric acid).

  • Strength is the extent of ionisation; Concentration is the amount of solute dissolved in a unit volume.

• Bases and Alkalis:

  • Base: Any metal oxide or hydroxide ($O^{2-}$ or $OH^-$).

  • Alkali: A base that is soluble in water, producing hydroxide ions ($OH^-$) in aqueous solution.

  • Strong alkali ($NaOH$, $KOH$ fully ionise); Weak alkali ($NH_3(aq)$ partially ionises: $NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$).

  • Reactions of Alkalis:

    1. $\text{Alkali} + \text{Acid} \rightarrow \text{Salt} + \text{Water}$. Ionic equation: $H^+(aq) + OH^-(aq) \rightarrow H_2O(l)$.

    2. $\text{Alkali} + \text{Ammonium Salt} \xrightarrow{\text{heat}} \text{Salt} + \text{Water} + \text{Ammonia Gas } (NH_3)$.

       • Test for $NH_3$: Pungent gas, turns damp red litmus paper blue.

• Indicators and pH Scale:

  • $pH < 7$: Acidic ($[H^+] > [OH^-]$).

  • $pH = 7$: Neutral ($[H^+] = [OH^-]$).

  • $pH > 7$: Alkaline ($[OH^-] > [H^+]$).

  • Indicators:

    • Methyl orange: Red in acid ($pH < 3$), yellow in alkali ($pH > 5$).

    • Screened methyl orange: Violet in acid, green in alkali.

    • Litmus: Red in acid, blue in alkali.

    • Thymolphthalein: Colourless in acid/neutral, blue in alkali ($pH > 9.3$).

Classification of Oxides and Control of Soil Acidity

• Oxide Classifications:

  • Basic Oxides: Metal oxides that react with acids to form a salt and water (e.g., $CuO$, $MgO$, $CaO$). Soluble ones ($Na_2O$, $K_2O$) dissolve in water to form alkalis.

  • Amphoteric Oxides: Metallic oxides that react with BOTH acids and bases to form salts and water. Examples: Zinc oxide ($ZnO$), Aluminium oxide ($Al_2O_3$), Lead(II) oxide ($PbO$).

    • $ZnO + 2HCl \rightarrow ZnCl_2 + H_2O$

    • $ZnO + 2NaOH \rightarrow Na_2ZnO_2 + H_2O$ (sodium zincate)

    • $Al_2O_3 + 6HCl \rightarrow 2AlCl_3 + 3H_2O$

    • $Al_2O_3 + 2NaOH \rightarrow 2NaAlO_2 + H_2O$ (sodium aluminate)

  • Acidic Oxides: Non-metal oxides that react with alkalis to form a salt and water. Most dissolve in water to form acids (e.g., $CO_2$, $SO_2$, $SO_3$, $P_4O_{10}$, $SiO_2$ [solid, reacts with hot conc. $NaOH$]).

  • Neutral Oxides: Non-metal oxides that show neither acidic nor basic properties (e.g., $H_2O$, $CO$, $NO$).

• Soil pH Control:

  • Optimum plant growth occurs around neutral or slightly acidic soil ($pH\,5\text{--}8$).

  • Excessive soil acidity (caused by acid rain, fertilisers) is treated by adding slaked lime (calcium hydroxide, $Ca(OH)_2$) or quicklime (calcium oxide, $CaO$) in a process called liming.

Preparation and Solubility Rules of Salts

• Solubility Rules for Common Salts in Water:

  • All Sodium ($Na^+$), Potassium ($K^+$), and Ammonium ($NH_4^+$) salts: SOLUBLE.

  • All Nitrates ($NO_3^-$): SOLUBLE.

  • All Chlorides ($Cl^-$): SOLUBLE, except Silver chloride ($AgCl$) and Lead(II) chloride ($PbCl_2$).

  • All Sulfates ($SO_4^{2-}$): SOLUBLE, except Barium sulfate ($BaSO_4$), Lead(II) sulfate ($PbSO_4$), and Calcium sulfate ($CaSO_4$ - sparingly soluble).

  • All Carbonates ($CO_3^{2-}$): INSOLUBLE, except $Na_2CO_3$, $K_2CO_3$, and $(NH_4)_2CO_3$.

• Methods of Salt Preparation:

  • Method 1: Reaction of Acid with Excess Insoluble Substance (Excess Metal, Insoluble Base, or Insoluble Carbonate).

    • Suitable for: Soluble salts (EXCEPT Group 1 $Na^+/K^+$ and $NH_4^+$ salts).

    • Procedure: Add excess insoluble reactant to acid, stir and warm, filter to remove excess solid, heat filtrate to saturation, cool to crystallise, filter crystals, wash with cold distilled water, dry between filter paper.

  • Method 2: Titration (Acid + Soluble Alkali or Soluble Carbonate).

    • Suitable for: Soluble $Na^+$, $K^+$, and $NH_4^+$ salts.

    • Procedure: Titrate acid into alkali with indicator to determine exact neutralisation volume, repeat without indicator using exact volumes, evaporate filtrate to saturation, cool to crystallise, filter, wash, dry.

  • Method 3: Precipitation (Mixing two soluble salt solutions).

    • Suitable for: Insoluble salts.

    • Procedure: Mix two soluble solutions containing required cation and anion, filter precipitate, wash residue with distilled water, dry between filter papers.

Industrial Production of Ammonia and Reversible Reactions

• Haber Process:

  • Synthesis Equation: $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$

  • Raw Materials: Nitrogen from fractional distillation of liquid air; Hydrogen from cracking of crude oil fractions.

  • Optimal Conditions: Temperature $450\,{^\circ}\text{C}$, Pressure $250\text{ atm}$ (or $200\text{--}250\text{ atm}$), Catalyst finely divided iron.

  • Yield Characteristics: Higher pressure increases yield and rate; lower temperature increases yield but decreases rate. Compromise conditions selected for optimal yield and cost.

  • Process: Gases compressed to $250\text{ atm}$, passed over iron catalyst at $450\,{^\circ}\text{C}$. Unreacted $N_2$ and $H_2$ are recycled back into the catalyst chamber.

• Reversible Reactions:

  • Proceed in both forward and backward directions.

  • Examples: Hydrated vs anhydrous copper(II) sulfate ($CuSO_4 \cdot 5H_2O(s) \rightleftharpoons CuSO_4(s) + 5H_2O(l)$), thermal decomposition of ammonium chloride ($NH_4Cl(s) \rightleftharpoons NH_3(g) + HCl(g)$).

Qualitative Analysis: Identification of Cations, Anions, and Gases

• Testing for Cations:

  • Aluminium ($Al^{3+}$):

    • With $NaOH(aq)$: White PPT of $Al(OH)_3$, soluble in excess to form a colourless solution.

    • With $NH_3(aq)$: White PPT of $Al(OH)_3$, insoluble in excess.

  • Zinc ($Zn^{2+}$):

    • With $NaOH(aq)$: White PPT of $Zn(OH)_2$, soluble in excess to form a colourless solution.

    • With $NH_3(aq)$: White PPT of $Zn(OH)_2$, soluble in excess to form a colourless solution.

  • Calcium ($Ca^{2+}$):

    • With $NaOH(aq)$: White PPT of $Ca(OH)_2$, insoluble in excess.

    • With $NH_3(aq)$: No PPT.

  • Ammonium ($NH_4^+$):

    • With $NaOH(aq)$: No PPT. On warming, ammonia gas ($NH_3$) is evolved (turns damp red litmus paper blue).

    • With $NH_3(aq)$: No PPT.

  • Copper(II) ($Cu^{2+}$):

    • With $NaOH(aq)$: Light blue PPT of $Cu(OH)_2$, insoluble in excess.

    • With $NH_3(aq)$: Light blue PPT of $Cu(OH)_2$, soluble in excess to form a dark blue solution.

  • Iron(II) ($Fe^{2+}$):

    • With $NaOH(aq)$: Green PPT of $Fe(OH)_2$, insoluble in excess, turns brown on standing.

    • With $NH_3(aq)$: Green PPT of $Fe(OH)_2$, insoluble in excess, turns brown on standing.

  • Iron(III) ($Fe^{3+}$):

    • With $NaOH(aq)$: Red-brown PPT of $Fe(OH)_3$, insoluble in excess.

    • With $NH_3(aq)$: Red-brown PPT of $Fe(OH)_3$, insoluble in excess.

• Testing for Anions:

  • Carbonate ($CO_3^{2-}$): Add dilute acid $\rightarrow$ Effervescence, gas evolved forms white PPT ($CaCO_3$) in limewater.

  • Chloride ($Cl^-$): Add dilute nitric acid ($HNO_3$), then aqueous silver nitrate ($AgNO_3$) $\rightarrow$ White PPT of silver chloride ($AgCl$).

  • Iodide ($I^-$): Add dilute nitric acid ($HNO_3$), then aqueous silver nitrate ($AgNO_3$) $\rightarrow$ Yellow PPT of silver iodide ($AgI$).

  • Sulfate ($SO_4^{2-}$): Add dilute nitric acid ($HNO_3$), then aqueous barium nitrate ($Ba(NO_3)_2$) $\rightarrow$ White PPT of barium sulfate ($BaSO_4$).

  • Nitrate ($NO_3^-$): Add aqueous sodium hydroxide ($NaOH$), then aluminium foil, warm carefully $\rightarrow$ Effervescence, ammonia gas evolved (turns damp red litmus blue).

• Testing for Gases:

  • Hydrogen ($H_2$): Colourless, odourless; extinguishes a burning splint with a 'pop' sound.

  • Oxygen ($O_2$): Colourless, odourless; relights/rekindles a glowing splint.

  • Carbon Dioxide ($CO_2$): Colourless, odourless; turns limewater milky (white PPT of $CaCO_3$).

  • Chlorine ($Cl_2$): Yellow-green gas, pungent; turns damp blue litmus paper red, then bleaches it white.

  • Sulfur Dioxide ($SO_2$): Colourless gas, pungent; decolourises filter paper soaked in acidified potassium manganate(VII) ($KMnO_4$) from purple to colourless.

  • Ammonia ($NH_3$): Colourless gas, pungent; turns damp red litmus paper blue.

Redox Reactions and Oxidation States

• Definitions of Oxidation and Reduction:

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

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

• Rules for Oxidation States:

  • Uncombined element atoms = $0$

  • Group 1 metals in compounds = $+1$

  • Group 2 metals in compounds = $+2$

  • Fluorine in compounds = $-1$

  • Hydrogen in compounds = $+1$ (except in metal hydrides where $-1$)

  • Oxygen in compounds = $-2$ (except in peroxides where $-1$, or $OF_2$ where $+2$)

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

• Oxidising and Reducing Agents:

  • Oxidising Agent: Oxidises another substance and is reduced itself (oxidation state decreases, gains electrons).

    • Test: Add colourless aqueous potassium iodide ($KI$) $\rightarrow$ turns yellow-brown solution/black PPT ($I^-$ oxidised to $I_2$).

  • Reducing Agent: Reduces another substance and is oxidised itself (oxidation state increases, loses electrons).

    • Test: Add purple acidified potassium manganate(VII) ($KMnO_4$) $\rightarrow$ turns from purple to colourless ($Mn^{7+}$ reduced to $Mn^{2+}$).

Electrochemistry: Electrolysis, Industrial Uses, and Electrochemical Cells

• Electrolysis Principles:

  • Conduction of electricity through a molten or aqueous electrolyte leading to chemical decomposition.

  • Anode ($+$, connected to positive terminal): Oxidation occurs (anions lose electrons).

  • Cathode ($-$, connected to negative terminal): Reduction occurs (cations gain electrons).

• Products of Electrolysis:

  • Molten Binary Ionic Compounds (e.g., molten $NaCl$):

    • Anode: $2Cl^-(l) \rightarrow Cl_2(g) + 2e^-$

    • Cathode: $Na^+(l) + e^- \rightarrow Na(l)$

  • Selective Discharge in Aqueous Solutions:

    • Cations at Cathode: Ease of discharge increases down the metal reactivity series ($K^+ < Na^+ < Ca^{2+} < Mg^{2+} < Al^{3+} < Zn^{2+} < Fe^{2+} < Pb^{2+} < H^+ < Cu^{2+} < Ag^+$).

    • Anions at Anode: Ease of discharge in dilute solutions ($SO_4^{2-} < NO_3^- < Cl^- < Br^- < I^- < OH^-$). $4OH^-(aq) \rightarrow O_2(g) + 2H_2O(l) + 4e^-$.

    • Concentration Effect: In concentrated halide solutions ($NaCl$, $HCl$), halide ions ($Cl^-$) are selectively discharged over $OH^-$ to give halogen gas ($Cl_2$).

• Industrial Applications:

  • Purification of Copper:

    • Anode: Impure copper ($Cu(s) \rightarrow Cu^{2+}(aq) + 2e^-$).

    • Cathode: Pure copper ($Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)$).

    • Electrolyte: Aqueous copper(II) sulfate ($CuSO_4(aq)$). Impurities form anode slime.

  • Electroplating:

    • Object to be plated = Cathode.

    • Plating metal = Anode.

    • Electrolyte = Soluble salt solution of plating metal.

• Simple Cells and Fuel Cells:

  • Simple Cell: Converts chemical energy to electrical energy. Consists of two different metals immersed in an electrolyte.

    • More reactive metal = Anode ($-$, oxidised, loses electrons).

    • Less reactive metal = Cathode ($+$, reduction occurs).

    • The greater the difference in reactivity between the two metals, the larger the voltage produced.

  • Hydrogen Fuel Cell: Converts chemical energy from $\text{hydrogen} + \text{oxygen}$ directly into electricity.

    • Overall reaction: $2H_2(g) + O_2(g) \rightarrow 2H_2O(l)$.

    • Anode ($-$, fuel $H_2$): $H_2(g) + 2OH^-(aq) \rightarrow 2H_2O(l) + 2e^-$.

    • Cathode ($+$, oxidiser $O_2$): $O_2(g) + 2H_2O(l) + 4e^- \rightarrow 4OH^-(aq)$.

    • Advantages: Renewable, non-polluting (water product), high efficiency. Disadvantages: Expensive, difficult storage/transport of hydrogen.

Periodic Trends and Group Properties

• Periodic Table Trends:

  • Across Period 3: Metallic character decreases (Group 1, 2, 13 metals $\rightarrow$ Group 14 metalloid $\rightarrow$ Groups 15, 16, 17, 18 non-metals). Oxide character changes from basic $\rightarrow$ amphoteric $\rightarrow$ acidic.

• Group Properties:

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

    • Soft, low density ($Li$, $Na$, $K$ float on water), low melting points.

    • Down the group: Melting point decreases, density increases, reactivity increases (outermost electron is further from nucleus and lost more easily).

    • React with cold water to form alkaline hydroxides and hydrogen gas ($2M(s) + 2H_2O(l) \rightarrow 2MOH(aq) + H_2(g)$).

  • Group 17 (Halogens: $F_2$, $Cl_2$, $Br_2$, $I_2$, $At_2$):

    • Diatomic non-metals. $Cl_2$ (yellow-green gas), $Br_2$ (red-brown liquid), $I_2$ (purple-black solid).

    • Down the group: Melting/boiling points increase, colour intensity increases (darker), reactivity decreases (harder to gain electron).

    • Displacement: More reactive halogen displaces less reactive halogen from its halide solution (e.g., $Cl_2(aq) + 2KBr(aq) \rightarrow 2KCl(aq) + Br_2(aq)$).

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

    • Monoatomic, unreactive (full valence duplet/octet configuration). Used to provide inert atmospheres.

  • Transition Metals (Groups 3–11):

    • High melting points, high densities, variable oxidation states (e.g., $Fe^{2+}/Fe^{3+}$, $Cu^+/Cu^{2+}$), form coloured compounds, act as catalysts ($Fe$ in Haber process, $Ni$ in hydrogenation).

Reactivity Series of Metals and Extraction

• Reactivity Series Order:

  • $K > Na > Ca > Mg > Al > (C) > Zn > Fe > Pb > (H) > Cu > Ag > Au$

• Reactivity Reactions:

  • Cold Water: $K, Na, Ca$ react violently/readily $\rightarrow$ hydroxide + $H_2$. $Mg$ reacts very slowly.

  • Steam: $Mg, Zn, Fe$ react with steam $\rightarrow$ metal oxide + $H_2$ ($Mg(s) + H_2O(g) \rightarrow MgO(s) + H_2(g)$). $Cu, Ag$ do not react.

  • Dilute $HCl$: Metals above hydrogen react $\rightarrow$ metal chloride + $H_2$.

• Displacement and Reduction:

  • More reactive metal displaces less reactive metal from its salt solution or oxide.

  • Reduction of Metal Oxides:

    • Oxides of metals below Carbon ($Zn, Fe, Pb, Cu$) are reduced by heating with carbon ($2ZnO + C \rightarrow 2Zn + CO_2$).

    • Oxides of metals below Hydrogen ($Fe, Pb, Cu$) are reduced by heating with hydrogen ($PbO + H_2 \rightarrow Pb + H_2O$).

• Thermal Decomposition of Carbonates:

  • $K_2CO_3$, $Na_2CO_3$: Unaffected by heat.

  • $CaCO_3$, $MgCO_3$, $ZnCO_3$, $FeCO_3$, $PbCO_3$, $CuCO_3$: Decompose to metal oxide + $CO_2$ ($CaCO_3 \xrightarrow{\text{heat}} CaO + CO_2$).

  • $Ag_2CO_3$: Decomposes to silver metal + $CO_2$ + $O_2$.

• Extraction Methods:

  • High reactivity ($K, Na, Ca, Mg, Al$): Electrolysis of molten ores.

  • Medium reactivity ($Zn, Fe, Pb, Cu$): Reduction of oxide with carbon in furnace.

  • Unreactive ($Ag, Au$): Found native/uncombined.

Corrosion and Prevention of Rusting

• Rusting of Iron:

  • Rust is hydrated iron(III) oxide ($Fe_2O_3 \cdot xH_2O$).

  • Essential conditions: BOTH Oxygen ($O_2$) and Water ($H_2O$) are required. Accelerated by $NaCl$ and acidic pollutants.

• Prevention Methods:

  • Barrier Methods: Exclude air and water (painting, oiling/greasing, plastic coating, tin-plating, electroplating).

  • Sacrificial Protection: Iron is connected to a more reactive metal (e.g., zinc or magnesium). The more reactive metal corrodes preferentially by losing electrons ($Zn \rightarrow Zn^{2+} + 2e^-$) in place of iron.

  • Galvanising: Coating iron with a thin layer of zinc. Combines barrier method and sacrificial protection (zinc protects even if scratched).

Chemical Energetics: Enthalpy Changes and Bond Energies

• Enthalpy Change ($\Delta H$):

  • $\Delta H = H_{\text{products}} - H_{\text{reactants}}$

  • Endothermic Reaction ($\Delta H > 0$, positive): Net absorption of thermal energy from surroundings. Temperature of surroundings decreases. Bond breaking is endothermic.

  • Exothermic Reaction ($\Delta H < 0$, negative): Net release of thermal energy to surroundings. Temperature of surroundings increases. Bond making is exothermic.

• Activation Energy ($E_a$): Minimum energy colliding particles must possess to react.

• Calculating Enthalpy Change from Bond Energies:

  • $\Delta H = \text{Total energy absorbed during bond breaking} - \text{Total energy released during bond making}$

Rate of Reactions and Collision Theory

• Measuring Reaction Rates:

  • Rate is the change in concentration, mass, or volume per unit time.

  • Methods: Volume of gas produced over time (gas syringe); Decrease in mass over time (balance with cotton wool plug).

  • Gradient of curve = rate of reaction (steepest at start, gentler over time, zero when reaction completes).

• Collision Theory and Rate Factors:

  • Effective collision requires energy $\ge E_a$ and correct orientation.

  • Concentration: Higher concentration $\rightarrow$ more particles per unit volume $\rightarrow$ higher collision frequency $\rightarrow$ higher frequency of effective collisions.

  • Pressure (Gases): Higher pressure $\rightarrow$ particles closer together per unit volume $\rightarrow$ higher collision frequency.

  • Particle Size / Surface Area (Solids): Smaller particle size $\rightarrow$ larger exposed surface area $\rightarrow$ higher collision frequency.

  • Temperature: Higher temperature $\rightarrow$ particles have higher kinetic energy (move faster) AND more particles possess energy $\ge E_a \rightarrow$ higher frequency of effective collisions.

• Catalysts and Enzymes:

  • A catalyst increases the reaction rate without being chemically changed at the end of the reaction.

  • Mechanism: Provides an alternative reaction pathway with a lower activation energy ($E_a$).

  • Enzymes are biological catalysts (proteins) that are temperature-sensitive and pH-sensitive with specific active sites.

Fuels, Crude Oil, and Fractional Distillation

• Fossil Fuels:

  • Natural gas (mainly methane, $CH_4$) and crude oil are non-renewable fossil fuels.

• Fractional Distillation of Crude Oil:

  • Separates hydrocarbons based on differences in boiling points. Column is hottest at the bottom and coolest at the top.

  • Fractions from Top to Bottom (increasing boiling point, increasing carbon chain length, increasing viscosity, decreasing flammability):

    1. Petroleum Gas ($C_1\text{--}C_4$): Fuel for cooking and heating.

    2. Petrol / Gasoline ($C_5\text{--}C_{10}$): Fuel for motorcars.

    3. Naphtha ($C_8\text{--}C_{12}$): Petrochemical feedstock for plastics/chemicals.

    4. Kerosene / Paraffin ($C_{10}\text{--}C_{16}$): Fuel for aircraft engines, heating.

    5. Diesel Oil ($C_{15}\text{--}C_{20}$): Fuel for diesel engines in buses/trains.

    6. Lubricating Oil ($C_{20}\text{--}C_{50}$): Lubricating machines, waxes, polishes.

    7. Bitumen ($>C_{50}$): Road surfacing and roofing.

• Biofuels:

  • Renewable alternatives (e.g., bioethanol from sugarcane fermentation). Biofuels are more environmentally sustainable regarding net carbon emissions.

Hydrocarbons: Alkanes, Alkenes, Isomerism, and Cracking

• Homologous Series: Family of organic compounds with the same general formula, same functional group, similar chemical properties, and a gradual trend in physical properties. Each member differs by a $-CH_2-$ unit ($14\,g/mol$).

• Alkanes: General formula $C_n H_{2n+2}$. Saturated hydrocarbons (only single $C-C$ and $C-H$ bonds).

  • Methane ($CH_4$), Ethane ($C_2H_6$), Propane ($C_3H_8$), Butane ($C_4H_{10}$).

  • Reactions: Complete combustion ($\rightarrow CO_2 + H_2O$); Substitution reaction with halogens ($Cl_2, Br_2$) under UV light (e.g., $CH_4 + Cl_2 \xrightarrow{UV} CH_3Cl + HCl$).

• Alkenes: General formula $C_n H_{2n}$ ($n \ge 2$). Unsaturated hydrocarbons containing a $C=C$ double bond.

  • Ethene ($C_2H_4$), Propene ($C_3H_6$), Butene ($C_4H_8$).

  • Reactions:

    1. Combustion: Sootier flame than alkanes.

    2. Addition of Hydrogen (Hydrogenation): $\text{Alkene} + H_2 \xrightarrow{150\,{^\circ}\text{C},\,Ni} \text{Alkane}$.

    3. Addition of Bromine (Bromination): $\text{Alkene} + Br_2(aq) \rightarrow \text{dibromoalkane}$. Brown aqueous bromine decolourises rapidly (test for unsaturation).

    4. Addition of Steam (Hydration): $\text{Alkene} + H_2O(g) \xrightarrow{300\,{^\circ}\text{C},\,60\text{ atm},\,H_3PO_4} \text{Alcohol}$.

    5. Addition Polymerisation.

• Cracking: Thermal/catalytic breakdown of long-chain alkanes into smaller alkanes, alkenes, and/or $H_2$ gas ($500\text{--}700\,{^\circ}\text{C}$, $Al_2O_3 / SiO_2$ catalyst, $1\text{ atm}$).

• Isomerism: Compounds with the same molecular formula but different structural formulae.

• Fats and Margarine: Saturated fats (solids at r.t.p.); Unsaturated oils (liquids at r.t.p., contain $C=C$ double bonds). Margarine manufactured by catalytic hydrogenation of vegetable oils ($150\,{^\circ}\text{C}$, $Ni$ catalyst).

Alcohols, Carboxylic Acids, and Esters

• Alcohols: General formula $C_n H_{2n+1}OH$ ($n \ge 1$). Hydroxyl ($-OH$) group. Methanol ($CH_3OH$), Ethanol ($C_2H_5OH$), Propanol ($C_3H_7OH$), Butanol ($C_4H_9OH$).

  • Reactions: Combustion; Oxidation to carboxylic acid ($C_2H_5OH + 2[O] \xrightarrow{\text{heat, acidified } KMnO_4} CH_3COOH + H_2O$, purple $KMnO_4$ decolourises).

  • Manufacture of Ethanol:

    1. Hydration of ethene: $C_2H_4 + H_2O \xrightarrow{300\,{^\circ}\text{C},\,60\text{ atm},\,H_3PO_4} C_2H_5OH$.

    2. Fermentation of glucose: $C_6H_{12}O_6 \xrightarrow{\text{yeast},\,37\,{^\circ}\text{C},\text{ absence of } O_2} 2C_2H_5OH + 2CO_2$.

• Carboxylic Acids: General formula $C_n H_{2n+1}COOH$ ($n \ge 0$). Carboxyl ($-COOH$) group. Methanoic acid ($HCOOH$), Ethanoic acid ($CH_3COOH$), Propanoic acid ($C_2H_5COOH$), Butanoic acid ($C_3H_7COOH$).

  • Weak acids ($RCOOH \rightleftharpoons RCOO^- + H^+$). Exhibit typical acid reactions with metals, carbonates, and bases.

• Esters: General formula $C_n H_{2n+1}COOC_m H_{2m+1}$. Ester linkage $-COO-$.

  • Formation: Carboxylic acid + Alcohol $\xrightarrow{\text{conc. } H_2SO_4,\text{ heat}} \text{Ester} + \text{Water}$ (Condensation/Esterification).

  • Naming: Alcohol part $\rightarrow$ '-yl'; Carboxylic acid part $\rightarrow$ '-oate' (e.g., Ethanol + Ethanoic acid $\rightarrow$ Ethyl ethanoate).

  • Properties: Sweet, fruity smell, volatile, insoluble in water.