Chemistry Essentials

Cheatsheet Content

1. Introduction to Chemistry 1.1. Branches of Chemistry Analytical Chemistry: Identifies substances and measures their quantity. Biochemistry: Studies chemical nature and energy changes in living systems. Inorganic Chemistry: Studies compounds not containing carbon (except carbonates). Physical Chemistry: Studies physical characteristics of materials and their reactions. Organic Chemistry: Studies substances containing carbon. 1.2. Importance of Chemistry Agriculture: Fertilizers (nitrogen, phosphorus) improve soil quality. Industry: Extraction of metals (e.g., copper) from ores. Plastics: Made from crude oil or coal. Medicine: Manufacture of drugs from plants. Home Use: Detergents and soaps. 1.3. Challenges of Chemical Industrial Activities Production of harmful by-products (e.g., sulfur dioxide, carbon monoxide) affecting the environment. 1.4. Laboratory Safety Rules Always wear shoes. Follow instructions carefully. Report accidents/breakages. Use eye goggles. Use only specified, small amounts of substances. Know the substance being used. Do not use broken glassware. Clean apparatus, bench, and hands after work. Check bottle labels. Hold bottles by the bottom, not the neck. Never remove chemicals/equipment from the laboratory. Never fight or play. Smell gases cautiously by wafting fumes. Turn off gas when not using burner. Keep clothes/head away from Bunsen flame. Work steadily and carefully. Never bring flammable substances near a flame; use water for fire if necessary. Wash off chemicals spilled on skin immediately and report. Do not point heated test tubes at yourself or others. Always add concentrated acid to water, not water to concentrated acid. Never eat, drink, or test anything without permission. Report all accidents, however minor. 2. The Particulate Nature of Matter 2.1. Matter and Kinetic Theory Matter: Anything that occupies space and has mass. Basic Units: Atoms, molecules, and ions. 2.2. States of Matter (Particle Arrangement and Movement) State Characteristics How Particles are Arranged Examples Solid Fixed volume and shape. Requires force to change shape. Particles closely packed, vibrating in fixed positions. Stone, ice block, table salt, wood Liquid Fixed volume but not fixed shape. Changes shape to fit container. Particles farther apart than solids, move randomly. Water, paraffin, cooking oil Gas No fixed shape or volume. Expands to fill container. Easily compressed. Particles very far apart, move randomly. 2.3. Changes of State Melting: Solid to liquid due to increased temperature; particle distance increases. Freezing Point: Liquid to solid; particle distance decreases. Boiling Point: Liquid to gas; temperature remains constant until all liquid changes to gas; energy increases particle distance. Condensation: Gas to liquid. Sublimation: Solid to gas or gas to solid directly. Evaporation/Vaporization: Liquid to gas at any temperature below boiling point. 2.4. Absorption and Release of Heat during Changes of State Changing states: Occurs by absorption or release of heat. Exothermic reaction: Release of heat. Endothermic reaction: Absorption of heat. 2.5. Diffusion Definition: Movement of particles from a region of higher concentration to lower concentration. 2.6. Factors Affecting Rate of Diffusion Molecular Mass: Smaller mass = faster diffusion. Temperature: Higher temperature = faster diffusion (particles move faster). Directly proportional. Concentration: Higher concentration difference = faster diffusion. 3. Experimental Techniques 3.1. Measuring Quantities Time: Stopwatch. Temperature: Thermometer. Mass: Electronic/beam balance. Volume: Measuring cylinder, burettes, pipettes, volumetric flasks, gas syringes. 3.2. Purity of Substances Pure Substance: Sharp melting/boiling point, fixed density. Impure Substance (Mixture): Melts/boils over a range of temperatures. Determining Purity: Heat substance with thermometer; pure substances show constant temperature during phase change. Importance of Purity: Crucial for foodstuffs, medicines (health), and industrial components (e.g., silicon chips). 3.3. Separating Mixtures Physical Change: No new substances, reversible, no energy absorption/release. Chemical Change: New substances formed, not reversible by simple means, energy released/absorbed. 3.4. Methods of Separating Mixtures Decantation: Separating a liquid from settled solid particles by carefully pouring off the liquid. Filtration: Separating insoluble solids from liquids using a filter. Filtrate can be evaporated/crystallized. Crystallization: Recovering a solute from a saturated solution by cooling, allowing solid crystals to form. Simple Distillation: Separating a liquid from a solution by evaporating and then condensing it. Fractional Distillation: Separating liquids with different boiling points. Magnetic Separation: Separating magnetic materials from non-magnetic ones. Evaporation: Recovering a solvent from a solution by evaporating the solvent. Sublimation: Separating substances that sublime (change directly from solid to gas) from non-subliming ones. Separating Funnel: Separating immiscible liquids (liquids that form layers). Centrifugation: Separating immiscible liquids by rapid rotation, pushing denser liquids to the bottom. Chromatography: Separating mixtures based on differential movement across a porous surface by a solvent. Ascending Paper Chromatography: Solvent moves upwards by capillary action. Descending Paper Chromatography: Solvent moves downwards by gravity. Radial Chromatography: Solvent moves outwards from a central point. 4. Atoms, Elements, Compounds and Molecules 4.1. Atomic Structure and Periodic Table Atom: Smallest particle of an element that takes part in a chemical reaction. Sub-atomic Particles: Protons: Positive charge, found in nucleus. Neutrons: No charge, found in nucleus. Electrons: Negative charge, move in electron shells/energy levels around the nucleus. 4.2. Relative Masses and Charges of Sub-atomic Particles Particle Symbol Relative mass/amu Relative Charge Protons P 1 +1 Neutrons N 1 0 Electrons E $1/1840$ -1 4.3. Proton (Atomic) Number, Nucleon (Mass) Number, and Nuclide Notations Proton Number (Z): Number of protons. Nucleon Number (A) / Relative Atomic Mass: Sum of protons + neutrons. Nuclide Notation: $_Z^A X$, where X is the element symbol. 4.4. Elements Definition: A substance containing only one kind of atoms. Atoms are the smallest particles of matter that cannot be broken down by chemical means. 4.5. Periodic Table Basics Elements arranged by increasing proton number. Periods: Horizontal rows (7 periods), indicates number of occupied electron shells. Groups: Vertical columns (8 main groups), indicates number of valence electrons and similar properties. 4.6. Isotopes Atoms of the same element with different numbers of neutrons (and thus different mass numbers). Example: Oxygen-16 (8p, 8n), Oxygen-17 (8p, 9n), Oxygen-18 (8p, 10n). 4.7. Relative Atomic Mass Calculation Average atomic mass = $\frac{\sum [(\% \text{ abundance of isotope}) \times (\text{mass of isotope})]}{100}$ Atomic mass unit (amu): $1/12$th the mass of a carbon-12 atom. 4.8. Uses of Radioactive Isotopes Checking for Leaks: Tracers detect leaks in pipes. Treating Cancer (Radiotherapy): Gamma rays kill cancer cells. Killing Germs/Bacteria: Sterilization of medical equipment and food preservation. 4.9. Electron Shells Electrons fill shells: First shell (max 2), second (max 8), etc. 4.10. Compounds Made of atoms of different elements bonded together, represented by a formula. 4.11. Ions (Radicals) An atom becomes an ion by losing or gaining electrons. An ion is a charged particle due to unequal protons and electrons. 4.12. Ionic (Electrovalent) Bonds Forms between ions of opposite charge. Metal atoms lose electrons to non-metal atoms, forming a lattice. Example: Sodium (loses e-) and Chlorine (gains e-) form NaCl. Magnesium (loses 2e-) and Oxygen (gains 2e-) form MgO. 4.13. Covalent Bonds Formed by sharing one or more pairs of electrons between two atoms. Examples: Chlorine ($Cl_2$): Each Cl atom shares one electron to achieve a stable outer shell. Oxygen ($O_2$): Each O atom shares two electrons (double bond). Water ($H_2O$): Oxygen shares electrons with two hydrogen atoms. Methane ($CH_4$): Carbon shares electrons with four hydrogen atoms. Carbon Dioxide ($CO_2$): Carbon shares two pairs of electrons with each oxygen atom (two double bonds). 4.14. Uses of Ionic and Covalent Compounds Ionic Compounds: High melting points, good refractory materials (furnace linings). Examples: Magnesium oxide (MgO) in steel mills, aluminum oxide ($Al_2O_3$) in glass/cement making. Covalent Compounds: Alcohol (Ethanol): Preservative, antiseptic, cleaning agent, fuel. Glass (Silicon Dioxide, $SiO_2$): Main ingredient. Used in nylon ropes. Insulators (bad conductors of electricity). Handles for cooking pans. 4.15. Molecules A group of atoms held together by covalent bonds. Molecular elements (e.g., $H_2$, $O_2$, $N_2$, $Cl_2$, $I_2$) are made of molecules. Diatomic: Molecules containing two atoms (e.g., $H_2$, $O_2$). 4.16. Valence and Valence Electrons Valency: Number of electrons an atom loses, gains, or shares to form a compound. Valence Electrons: Electrons in the outermost shell. Group number (1-7) indicates valence electrons. Deducing Valency: Groups 1-3: Lose electrons. Group 4: Share electrons. Groups 5-7: Gain electrons to complete octet. 4.17. Differences in Properties of Ionic and Covalent Compounds Ionic Compounds Covalent Compounds Not volatile Volatile (distinctive smells) Made of ions Molecules Conduct electricity when melted or in solution Do not conduct electricity Usually solids at room temperature Usually gases or liquids at room temperature High boiling/melting points Low boiling/melting points Do not vaporize easily (strong electrostatic forces) Vaporize easily (weak electrostatic forces) Many are soluble in water Many insoluble in water, but soluble in other covalent liquids (e.g., alcohol, tetrachloromethane) Higher densities Lower densities 4.18. Metallic Bonding Lattice of positive ions in a "sea" of delocalized electrons. Outer electrons break free and move around, acting as glue. Example: Copper ($Cu^{2+}$ ions with free electrons). 4.19. Electrical/Thermal Conductivity of Metals Good conductors of electricity (delocalized electrons carry charge). Good conductors of heat (electrons absorb and transfer heat). High melting points (strong lattice). Ductile (drawn into wire) and Malleable (hammered into shape) because positive ions can slide over each other. 5. Macromolecules 5.1. Giant Covalent Structures: Graphite and Diamond Graphite: Carbon atoms in hexagonal layers, held by weak forces between layers. Each carbon bonds to three others. Diamond: Carbon atoms arranged tetrahedrally. Each carbon bonds to four others, forming a giant covalent structure. 5.2. Uses of Graphite and Diamond Substance Properties Uses Diamond Hardest known substance, does not conduct, sparkles. Drilling/cutting tools, jewelry. Graphite Soft, slippery, dark, conducts electricity. Lubricant, pencil 'lead', electrodes, brushes in generators. 5.3. Macromolecular Structure of Silicon (IV) Oxide ($SiO_2$) Oxygen atoms bonded to silicon atoms tetrahedrally. Forms a giant structure. 5.4. Similarities between Diamond and Silicon Dioxide Atoms held together by covalent bonds tetrahedrally. 5.5. Chemical Formulae from Valences Valency determines the ratio of atoms in a compound. Example: Calcium hydroxide ($Ca(OH)_2$), Carbon dioxide ($CO_2$), Aluminium carbonate ($Al_2(CO_3)_3$). 5.6. Common Radicals and Compounds Name of Radical Chemical Symbol Valency Compound formed with cations Hydroxides $OH^-$ 1 Calcium hydroxide - $Ca(OH)_2$ Ammonium $NH_4^+$ 1 Ammonium sulphate - $(NH_4)_2SO_4$ Carbonates $CO_3^{2-}$ 2 Ammonium carbonate - $(NH_4)_2CO_3$ Sulphates $SO_4^{2-}$ 2 Iron(III) sulphate - $Fe_2(SO_4)_3$ Phosphates $PO_4^{3-}$ 3 Ammonium phosphate - $(NH_4)_3PO_4$ 6. Chemical Formulae and Equations 6.1. Constructing Word Equations Write equation in words. Write using symbols, ensure correct formulae. Check if balanced for each atom type. Add state symbols. 6.2. Balanced Chemical Equations Example: $Ca(s) + Cl_2(g) \rightarrow CaCl_2(s)$ Example: $CaCO_3(s) + 2HCl(aq) \rightarrow CaCl_2(aq) + H_2O(l) + CO_2(g)$ 6.3. Net Ionic Equations Ensure balanced equation. Leave solids, covalently bonded substances, liquids, and gases intact. Separate aqueous compounds (aq) into cation and anion parts. Cancel spectator ions (ions appearing on both sides). Example: $CaCO_3(s) + 2H^+(aq) \rightarrow Ca^{2+}(aq) + H_2O(l) + CO_2(g)$ 7. Acids, Bases and Salts 7.1. Properties of Acids and Bases Acid: Compound producing hydrogen ions ($H^+$) as only positive ions in aqueous solution. Base: Oxide or hydroxide of a metal (including ammonium hydroxide). Alkali: Soluble base producing hydroxide ions ($OH^-$) as only negative ions in aqueous solution. 7.2. Weak, Strong, Dilute, and Concentrated Acids/Alkalis Weak Acid: Partially ionizes in water. Strong Acid: Completely ionizes in water to form hydronium ions. Dilute Acid: Acid mixed with significant percentage of water. Concentrated Acid: Large amount of acid dissolved in 1 dm$^3$ of solution. Weak Alkali: Partially dissociates in water. Strong Alkali: Completely dissociates in water. Dilute Alkali: Not very strong, pH 8-10. Concentrated Alkali: Large amount of alkali dissolved in 1 dm$^3$ of solution. 7.3. pH Scale Scale from 0-14 showing degree of acidity/alkalinity. Neutrality: pH 7. Acidity: pH less than 7 (0-6.9). Alkalinity: pH more than 7 (8-14). 7.4. Determining pH Universal Indicator: Gives different colors for different pH values. pH Meter: Gives precise values. 7.5. Characteristic Properties of Acids Sour taste. Corrosive (burn skin). Turn damp blue litmus paper red. React with metals to form salt and hydrogen gas: Acid + Metal $\rightarrow$ Salt + Hydrogen. (e.g., $2HCl(aq) + Mg(s) \rightarrow MgCl_2(aq) + H_2(g)$) React with bases to form salt and water: Acid + Base $\rightarrow$ Salt + Water. (e.g., $HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)$) React with metal carbonates/hydrogen carbonates to form salt, water, and carbon dioxide: Acid + Carbonate $\rightarrow$ Salt + Water + Carbon Dioxide. (e.g., $2HCl(aq) + CaCO_3(s) \rightarrow CaCl_2(aq) + H_2O(l) + CO_2(g)$) Indicators: Blue litmus paper turns red, methyl orange turns red, bromothymol blue turns yellow, phenolphthalein becomes colorless. 7.6. Characteristic Properties of Bases Bitter taste. Soapy/slippery feel. Corrosive. Neutralized by acids. 7.7. Importance of Acid-Base Reactions Controlling Soil Acidity: Affects crop quality. Acidification from air pollution. Treatment of Indigestion: Antacids (bases) neutralize excess stomach acid. Brushing Teeth: Toothpaste (sodium bicarbonate) neutralizes lactic acid from bacteria, preventing tooth decay. 7.8. Uses of Acids and Bases pH control in agriculture. Soap making. Car batteries. 8. Preparation of Salts 8.1. Definition of Salt Compound formed when hydrogen ions of an acid are fully or partially replaced by metal or ammonium ions. Compound of positive metallic/ammonium ions and any negative ion of an acid. 8.2. Classification of Salts Acid salt, basic salt, normal salt. 8.3. Solubility Rules of Salts Soluble Salts Insoluble Salts All nitrates Carbonates of sodium, potassium, ammonium (Group I carbonates) All other carbonates Most sulphates Lead (II) sulphate, barium sulphate, calcium sulphate Most chlorides Silver chloride, lead (II) chloride, mercury (I) chloride Hydroxides and oxides of alkali metals and ammonium All other oxides and hydroxides All salts of alkali metals and ammonium 8.4. Preparation of Insoluble Salts (e.g., Barium Sulphate) Mix solutions of barium chloride and magnesium sulphate (white precipitate forms). Filter the mixture (barium sulphate remains as residue). Rinse with distilled water. Dry in a warm oven or with porous filter paper. 8.5. Preparation of Soluble Salts (Titration Method) Pipette 25.0 cm$^3$ of 1.0 mol/dm$^3$ HCl into a conical flask. Add 2 drops of phenolphthalein indicator and swirl. Slowly add aqueous sodium hydroxide (known concentration) from burette until pink color appears, note volume. Repeat titration without indicator, adding noted volume of alkali. Heat solution to evaporate water and saturate. Crystallize by cooling. Filter pure sodium chloride crystals. 8.6. Hydrated vs. Anhydrous Salts Hydrated Salts: Contain water of crystallization. Anhydrous Salts: Do not contain water of crystallization. 8.7. Behavior of Salts with Atmosphere Hygroscopic: Absorbs moisture but not enough to form a solution. Deliquescent: Absorbs enough moisture to form a solution. Efflorescent: Gives out water of crystallization to the atmosphere. 9. Types of Oxides 9.1. Classification of Oxides Acidic Oxides: Have acidic properties (e.g., $SO_2$, $CO_2$). Basic Oxides: Oxides of Group I and II metals (e.g., $K_2O$, $MgO$). Neutral Oxides: Some non-metal oxides with neither acidic nor basic properties (e.g., $CO$, $H_2O$). Amphoteric Oxides: Have both acidic and basic properties (e.g., $ZnO$, $Al_2O_3$, $PbO$). 9.2. Test for Aqueous Cations and Anions Aqueous Cations (using NaOH and Ammonia solution): Cations Effect of Aqueous Sodium Hydroxide Effect of Aqueous Ammonia Aluminium ions ($Al^{3+}$) White ppt formed, soluble in excess White ppt formed, insoluble in excess Ammonium ions ($NH_4^+$) Ammonia gas produced on warming - Calcium ions ($Ca^{2+}$) White ppt formed, insoluble in excess No change Copper ions ($Cu^{2+}$) Light blue ppt formed, insoluble in excess Light blue ppt, soluble in excess (dark blue solution) Iron(II) ions ($Fe^{2+}$) Green ppt formed, insoluble in excess Green ppt, insoluble in excess, turns reddish-brown on standing Iron (III) ions ($Fe^{3+}$) Red-brown ppt formed, insoluble in excess Red-brown ppt, insoluble in excess Zinc ions ($Zn^{2+}$) White ppt formed, soluble in excess White ppt, soluble in excess (colorless solution) Anions (using various reagents): Anions Test Test Result Carbonate ($CO_3^{2-}$) Add dilute acid Effervescence, carbon dioxide gas produced Chloride ($Cl^-$) Acidify with dilute nitric acid, then add aqueous silver nitrate White ppt formed Iodide ($I^-$) Acidify with dilute nitric acid, then add aqueous lead (II) nitrate Yellow ppt formed Nitrate ($NO_3^-$) Add aqueous sodium hydroxide, then aluminium foil, warm Ammonia gas produced Sulphate ($SO_4^{2-}$) Acidify with dilute nitric acid, then add aqueous barium nitrate White ppt formed 9.3. Identification of Gases Ammonia ($NH_3$): Damp universal indicator paper turns blue; sharp, pungent smell. Carbon Dioxide ($CO_2$): Bubble through limewater (calcium hydroxide) $\rightarrow$ milky (white precipitate of $CaCO_3$). Chlorine Gas ($Cl_2$): Damp universal indicator paper turns red, then bleaches. Hydrogen Gas ($H_2$): Lighted splint in test tube $\rightarrow$ "pop" sound. Oxygen Gas ($O_2$): Glowing/burning splint in test tube $\rightarrow$ re-lights/re-kindles. Sulphur Dioxide ($SO_2$): Bubble through acidified potassium dichromate (VI) solution $\rightarrow$ orange to green. 10. The Mole Concept 10.1. Relative Masses Relative Atomic Mass: Average mass of one atom of an element compared to $1/12$th mass of carbon-12. Relative Molecular Mass: Average mass of one molecule compared to $1/12$th mass of carbon-12. Relative Formula Mass ($M_r$): Sum of relative atomic masses of atoms in a compound. (e.g., $CaCO_3 = 40 + 12 + (3 \times 16) = 100$) 10.2. The Mole Amount of substance containing same number of particles as 12g of carbon-12. Amount of substance containing $6.02 \times 10^{23}$ particles (Avogadro's constant). 10.3. Physical Masses, Molar Mass, and Volume $n = \text{mass}/M_r$ $n = \text{volume}/V_m$ (where $V_m$ is molar volume) 10.4. Avogadro's Law Equal volumes of gases under same T and P contain same number of molecules. Equal moles of gases occupy same volume under same T and P. Molar gas volume ($V_m$): 24 dm$^3$ (r.t.p., 25°C, 101.3 kPa), 22.4 dm$^3$ (s.t.p., 0°C, 1 atm). 10.5. Concentration and Dilution Law Concentration: Amount of substance dissolved in a unit volume. Mass Concentration: Mass of solute (g) / volume of solution (dm$^3$) (units: g/dm$^3$). Mole Concentration (Molarity): Moles of solute (mol) / volume of solution (dm$^3$) (units: mol/dm$^3$ or M). Dilution Law: $M_1V_1 = M_2V_2$ (moles of solute before dilution = moles after dilution). 10.6. Stoichiometry Calculations Calculations involving reacting moles and volumes of gases and solutions using mole concept. 10.7. Percentage Yield and Purity Percentage Yield: ($\text{actual mass} / \text{expected mass}$) $\times 100\%$. Percentage Purity: ($\text{amount of pure substance} / \text{total amount of mixture}$) $\times 100\%$. 10.8. Limiting Reagent Reactant that is completely consumed first, limiting product formation. Steps to determine: Write balanced equation. Convert given amounts to moles. Choose a product and calculate its theoretical yield from each reactant. The reactant producing the least amount of product is the limiting reagent. 10.9. Acid-Base Titration Reactions Titration: Procedure to determine concentration of a solution by reacting it with a solution of known concentration (standard solution). Calculations involve molar ratios from balanced equations. 11. Empirical and Molecular Formulae 11.1. Percentage Composition Percentage composition of an element = ($\text{mass of element} / \text{formula mass of compound}$) $\times 100\%$. 11.2. Empirical vs. Molecular Formula Molecular Formula (M.F.): Actual number of atoms in a compound. Empirical Formula (E.F.): Simplest whole-number ratio of atoms in a compound. M.F. = (E.F.)$_n$, where $n$ is a whole number. 11.3. Determining Empirical and Molecular Formulae Write elements present. Write percentages or masses. Divide by relative atomic masses to get moles. Divide all moles by the smallest number to get ratio. Round to whole numbers (if necessary). Write empirical formula. If relative formula mass is given, divide M.F. mass by E.F. mass to find $n$. 12. Chemical Reactions 12.1. Definition of Chemical Reaction Chemical change where substances combine to form new substances. 12.2. Rates of Chemical Reactions Measure of how long a reaction takes, or change in concentration of reactants/products over time. Rate = $\text{Change in amount of substance} / \text{Time taken}$. 12.3. Factors Affecting Rate of Reaction Temperature: Higher temperature $\rightarrow$ faster rate (increased kinetic energy, more frequent effective collisions). Concentration: Higher concentration $\rightarrow$ faster rate (more particles per volume, more frequent effective collisions). Surface Area: Larger surface area (smaller particle size) $\rightarrow$ faster rate (more contact points). Catalyst: Substance that changes reaction rate without being consumed. Lowers activation energy. Pressure: Higher pressure (for gases) $\rightarrow$ faster rate (particles packed closer, more frequent collisions). Light Intensity: Higher intensity $\rightarrow$ faster rate (provides energy to break bonds, e.g., methane + chlorine). 12.4. Chemical Equilibrium State where forward and backward reaction rates are equal, and substance concentrations remain constant. Occurs in reversible reactions. Effect of Changing Conditions (Le Chatelier's Principle): Changes in temperature, pressure, concentration shift equilibrium. Catalysts have no effect on equilibrium position. 12.5. Redox Reactions Reactions involving both oxidation and reduction. 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. Oxidizing Agent: Oxidizes another substance (gains electrons, reduces itself). Reducing Agent: Reduces another substance (loses electrons, oxidizes itself). Non-Redox Reaction: No oxidation or reduction involved. Oxidation Numbers: Rules for determining oxidation states. 13. Energetics of Reactions 13.1. Endothermic and Exothermic Reactions Endothermic: Absorbs energy (heat) from surroundings, temperature decreases. $\Delta H$ is positive. Exothermic: Releases energy (heat) to surroundings, temperature increases. $\Delta H$ is negative. 13.2. Enthalpy Change ($\Delta H$) $\Delta H = \text{Energy used in bond breaking (positive)} + \text{Energy released in bond making (negative)}$. Endothermic: Bond breaking > bond making. Exothermic: Bond making > bond breaking. 13.3. Activation Energy Energy required to start a reaction. Catalysts lower activation energy. 13.4. Energy Sources (Fuels) Advantages: Safety, cost, availability, renewable sources. Disadvantages: Non-renewable sources. 13.5. Effects of Fuels on Environment Pollution, greenhouse effect (global warming). 13.6. Silver Halide in Photography Reduction of silver ions to metallic silver by light absorption (endothermic). 13.7. Respiration and Photosynthesis Respiration: Exothermic (oxygen + glucose $\rightarrow$ carbon dioxide + water). Photosynthesis: Endothermic (water + carbon dioxide + light $\rightarrow$ glucose + oxygen). 13.8. Radioactive Isotopes in Energy Changes Used as a source of nuclear energy. 13.9. Batteries as Electrical Energy Source Convenient and portable. 14. The Periodic Table 14.1. Groups and Periods Periods: Horizontal rows. Elements in same period have same number of electron shells. Groups: Vertical columns. Elements in same group have same number of valence electrons and similar properties. 14.2. Classification of Elements Metals, non-metals, metalloids. 14.3. Group I Elements (Alkali Metals) Physical Properties: Soft, silvery, low density, low melting/boiling points, good conductors of heat/electricity. Tarnish quickly. Chemical Properties: One valence electron, lose it to form $1^+$ ions. Highly reactive, reactivity increases down the group. React vigorously with water (forms $H_2$), burn in air (forms oxides). React with halogens (forms neutral salts). Trends: Reactivity, density, melting/boiling points increase down the group. Become softer. 14.4. Group II Elements (Alkaline Earth Metals) Physical Properties: Slightly harder than alkali metals, silvery, low density, higher melting/boiling points, good conductors. Tarnish quickly. Chemical Properties: Two valence electrons, lose them to form $2^+$ ions. Reactive, reactivity increases down the group, but less reactive than alkali metals. React with cold water (less vigorously than alkali metals). Trends: Two valence electrons. Reactivity increases down the group. Melting/boiling points decrease down the group. 14.5. Group VII Elements (Halogens) Physical Properties: Non-metals, diatomic molecules (e.g., $F_2$, $Cl_2$, $Br_2$, $I_2$), high melting/boiling points, poisonous, poor conductors. Colored, darken down the group. Chemical Properties: Seven valence electrons, gain one to form $1^-$ ions. Reactivity decreases down the group. Undergo displacement reactions (more reactive halogen displaces less reactive halide). Uses: Fluorine (toothpaste), chlorine (water treatment, antiseptic), bromine (photographic film), iodine (thyroid gland). Chlorine for bleaching. Harmful Effects of Halides: Negative effects in drugs/pesticides. CFCs cause ozone depletion. 14.6. Noble Gases Uses: Providing inert atmosphere (e.g., argon in electrical lamps, helium in balloons). 14.7. Transition Metals Block elements between Group II and Group III. General Properties: Variable valencies, high densities, high melting points, form colored compounds, act as catalysts. Uses: Catalysts: Iron (Haber process), nickel (margarine production). Alloys: Copper-tin (mirrors), copper-zinc (brass), iron-carbon (steel). Conductors: Copper (electrical wiring), aluminum (pots/pans). Paint manufacturing. Heavy metals not used in paint for health reasons. 15. Metals 15.1. Structure of Pure Metals Atoms arranged in layers, allowing bending/shaping. Strong electrical force between free electrons ("sea" of delocalized electrons) and positive metal ions. Delocalized electrons act as a "glue." 15.2. Physical Properties of Metals Hard and strong. Malleable (deformed into different shapes). Ductile (drawn into wires). Sonorous (make ringing sound). Lustrous (shiny), tarnish easily. High melting/boiling points (except mercury). High densities. Good conductors of heat and electricity. 15.3. Chemical Properties of Metals Electropositive (lose electrons to form positive ions). React with oxygen to form metal oxides. React with water to produce metal oxide/hydroxide and hydrogen gas. React with dilute acids to form metal salt and hydrogen gas. React with chlorine to form ionic metal chloride. Reducing agents (electron donors). 15.4. Reactivity Series of Metals Arrangement of metals in order of decreasing reactivity (K > Na > Ca > Mg > Al > Zn > Fe > Pb > [H] > Cu > Hg > Ag > Au > Pt). More reactive metals displace less reactive ones. Non-Reactivity of Aluminum: Due to thin protective oxide layer ($Al_2O_3$). 15.5. Effects of Heat on Metal Compounds More reactive metals: Compounds are difficult to decompose. Less reactive metals: Compounds decompose easily. Nitrates: Decompose to metal oxide, nitrogen dioxide (brown fumes), oxygen. (e.g., $2Mg(NO_3)_2(s) \rightarrow 2MgO(s) + 4NO_2(g) + O_2(g)$) Hydroxides: Decompose to metal oxide and water. (e.g., $Ca(OH)_2(s) \rightarrow CaO(s) + H_2O(l)$) Carbonates: Above sodium: Stable. Below sodium: Decompose to metal oxide and carbon dioxide. (e.g., $CaCO_3(s) \rightarrow CaO(s) + CO_2(g)$) 15.6. Extraction of Metals Copper: Concentration of ore (copper pyrite $CuFeS_2$) by froth flotation. Roasting in air: $2CuFeS_2(s) + 4O_2(g) \rightarrow Cu_2S(g) + 2FeO(g) + 3SO_2(g)$. $FeO$ removed by heating with $SiO_2$ to form slag. $Cu_2S$ further heated in limited air to form copper blister. Refining by electrolysis: Impure copper as anode, pure copper as cathode. Iron: Roasting of ore (haematite $Fe_2O_3$): Removes water, $CO_2$, impurities. Reduction in blast furnace: Coke burns: $C(s) + O_2(g) \rightarrow CO_2(g)$ (heat). $CO_2$ reduced by coke: $CO_2(g) + C(s) \rightarrow 2CO(g)$. Carbon monoxide reduces iron oxide: $Fe_2O_3(s) + 3CO(g) \rightleftharpoons 2Fe(l) + 3CO_2(g)$. Limestone ($CaCO_3$) decomposes: $CaCO_3(s) \rightarrow CaO(s) + CO_2(g)$. $CaO$ reacts with $SiO_2$ (impurity) to form slag ($CaSiO_3$). Zinc: Concentration of zinc sulphide ore. Roasting in air: $2ZnS(s) + 3O_2(g) \rightarrow 2ZnO(s) + 2SO_2(g)$. Reduction: $ZnO(s) + C(s) \rightarrow Zn(g) + CO(g)$ at $1400^\circ C$. Zinc vapor condenses, $CO$ bubbles off. Aluminium: Purify bauxite ($Al_2O_3 \cdot 2H_2O$) to pure $Al_2O_3$: Dissolve bauxite in NaOH: $Al_2O_3(s) + 3H_2O(l) + 2NaOH(aq) \rightarrow 2NaAl(OH)_4(aq)$. Filter, seed with $Al(OH)_3$: $NaAl(OH)_4(aq) \rightarrow NaOH(aq) + Al(OH)_3(s)$. Heat $Al(OH)_3$: $2Al(OH)_3(s) \rightarrow Al_2O_3(s) + 3H_2O(g)$. Electrolysis: $Al_2O_3$ dissolved in molten cryolite ($Na_3AlF_6$). $Al^{3+}$ ions gain electrons at cathode, $O^{2-}$ ions lose electrons at anode. 15.7. Harmful Effects of Metals Lead poisoning (brain damage). Sodium ions (high blood pressure). Aluminum (Alzheimer's). 15.8. Alloys Mixture of metals, or metal and other element. Atoms of different sizes distort regular arrangement, making them harder. Advantages over Pure Metals: Stronger, lighter, high tensile strength, harder, better resistance to corrosion. Common Uses: Silver (coins). Stainless steel (car parts, sinks, cutlery). Duralumin (aircraft parts). Bronze (statues, ornaments, bells). Solders (joining wires). Brass (musical instruments). Composition: Alloy Compositions Properties Uses Bronze 90% copper, 10% tin Hard, strong, doesn't corrode, shiny Statues, monuments, artistic materials Brass 70% copper, 30% zinc Harder than copper Musical instruments, kitchenware Steel 99% iron, 1% carbon Hard, strong Bridges, car bodies, railway tracks Stainless steel 74% iron, 8% carbon, 18% chromium Shiny, strong, doesn't rust Cutlery, surgical instruments Duralumin 93% aluminium, 3% copper, 3% magnesium, 1% manganese Light, strong Aircraft bodies, bullet trains Pewter 96% tin, 3% copper, 1% antimony Luster, strong Souvenirs 16. Corrosion 16.1. Definition Chemical wearing of metal by atmospheric oxygen in presence of moisture. Rusting: Corrosion of iron. 16.2. Relation to Reactivity More reactive metals corrode easily; less reactive metals do not. 16.3. Preventing Corrosion Sacrificial Protection: Protecting iron/steel with more reactive metal. Painting: Coating with paint. Greasing/Oiling: Applying oil or grease. Galvanizing: Coating iron/steel with zinc. Tinning: Coating iron/steel with tin (less reactive). 17. Non-Metals 17.1. General Properties Not strong, malleable, ductile, or sonorous. Lower melting/boiling points than metals. Poor conductors of electricity (except graphite) and heat. Low densities. React with oxygen to form oxides. Form negative ions. 17.2. Hydrogen Laboratory Preparation: Reactive metals + water/steam, or aluminum/zinc + dilute acids/aqueous potassium/sodium hydroxide. Collected by upward delivery. Test: Lighted splint $\rightarrow$ "pop" sound. Physical Properties: Colorless, odorless, tasteless gas. No effect on litmus paper. Less dense than air. Low boiling/melting points. Chemical Properties: Reacts spontaneously with chlorine (hydrogen chloride). Reacts explosively with oxygen (water). Industrial Manufacture: Cracking, electrolysis of brine, natural gas. Uses: Haber process (ammonia), meteorological balloons, margarine (catalyst), rocket fuel. 17.3. Oxygen Laboratory Preparation: Catalytic decomposition of hydrogen peroxide or potassium chlorate. Collected above water. Test: Glowing splint $\rightarrow$ re-lights. Physical Properties: Colorless, tasteless, odorless gas. Insoluble in water. As dense as air. Chemical Properties: Re-lights glowing splint. Reacts with metals (metal oxides) and non-metals. Industrial Manufacture: Fractional distillation of liquid air. Uses: Burning, welding, steel manufacturing, space (astronauts), respiration, fuel cells. Ozone Layer ($O_3$): Traps UV radiation. Depletion by CFCs causes skin cancer, respiratory diseases. Test for Water: Anhydrous copper (II) sulphate turns blue. Importance of Water: Laundry, drinking, cooking, solvent, coolant, chemical reactant. 17.4. Nitrogen Industrial Manufacture: Fractional distillation of liquid air. Characteristics/Importance: Non-reactive, insoluble gas. Used as refrigerant, food packaging, ammonia manufacturing. Preparation/Test for Ammonia: Base + ammonium salt. Damp red litmus paper turns blue. Manufacture of Ammonia (Haber Process): $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$ (iron catalyst, high T/P). Thermal Dissociation of Ammonium Salts: Liberate ammonia gas. Uses of Ammonia: Explosives, refrigerant, fertilizers. Manufacture of Nitric Acid (Ostwald Process): Catalytic oxidation of ammonia: $4NH_3(g) + 5O_2(g) \rightarrow 4NO(g) + 6H_2O(l)$. $NO$ cooled, reacted with excess air: $2NO(g) + O_2(g) \rightarrow 2NO_2(g)$. $NO_2$ absorbed in water: $3NO_2(g) + H_2O(g) \rightarrow 2HNO_3(aq) + NO(g)$. Nitrogenous Fertilizers: Nitrogen (growth), phosphorus (root development), potassium (seed formation). NPK. Effects of Nitrogenous Fertilizers on Environment: Acid rain, ozone-hydrocarbon compounds (poisonous), eutrophication (algae growth in rivers). 17.5. Chlorine Laboratory Preparation: Hot conc. HCl + manganese (IV) oxide. Collected by downward delivery. Test: Turns damp blue litmus paper red, then bleaches. Physical/Chemical Properties: Greenish-yellow gas. Melting point -100.98°C, boiling point -34.6°C. Density 3.214 g/l. Reacts with almost all elements. Uses: Sterilizing water, PVC, HCl, bleaching agents, disinfectants, drugs, solvents, paints. Chemical warfare agent. Industrial Manufacture: Electrolysis of brine ($NaCl(aq)$). Preparation/Test for Hydrogen Chloride Gas: Conc. $H_2SO_4$ + solid metallic chlorides. Forms white smoke with ammonia. Properties of HCl Gas: Color, odor, density, solubility, poisonous. Preparation of Hydrochloric Acid: Dissolving HCl gas in water. Reactions of Dilute HCl: Alkalis, metals, carbonates, ammonia, silver nitrate. 17.6. Sulphur Formation of Sulphur Dioxide ($SO_2$): Combustion of sulphur, fossil fuels. Laboratory Preparation/Test for $SO_2$: Warm dilute acids + sulphites. Turns acidified potassium dichromate green. Physical/Chemical Properties of $SO_2$: Colorless gas, choking/suffocating odor. Boiling point -10°C. Heavier than air. Toxic. Reacts with water, acts as reducing agent. Uses of $SO_2$: Food preservative, bleaching wood pulp, sulphuric acid manufacturing. Industrial Manufacture of Sulphuric Acid (Contact Process): Make $SO_2$: $S(s) + O_2(g) \rightarrow SO_2(g)$. Convert $SO_2$ to $SO_3$: $2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g)$ ($\Delta H = -196 kJ mol^{-1}$, $V_2O_5$ catalyst, $400-450^\circ C$, 1-2 atm). Convert $SO_3$ to $H_2SO_4$: Dissolve $SO_3$ in conc. $H_2SO_4$ to form oleum ($H_2S_2O_7$), then add water ($H_2S_2O_7(l) + H_2O(l) \rightarrow 2H_2SO_4(l)$). Uses of Sulphuric Acid: Explosives, drying agent, soaps, fertilizers. 17.7. Carbon and Carbonates Allotropes: Different physical forms of the same element. Carbon has diamond and graphite. Physical Properties of Carbon Allotropes: Graphite: Dark grey shiny solid, conducts electricity, soft, slippery, density 2.25 g/cm$^3$. Diamond: Colorless transparent crystal, sparkles, does not conduct electricity, very hard, density 3.51 g/cm$^3$. Carbon Monoxide ($CO$): Formation: Incomplete combustion of carbon compounds, reduction of $CO_2$ by carbon. Properties: Colorless, odorless, density. Insoluble in water. Very poisonous. Reducing agent. Carbon Dioxide ($CO_2$): Laboratory Preparation: Dilute HCl + carbonate/bicarbonate. Collected by downward delivery. Test: Bubbled through limewater $\rightarrow$ milky. Physical Properties: Colorless, odorless, tasteless. Heavier than air. Slightly soluble in water. Solidifies to dry ice when cooled/pressurized. Chemical Properties: Burning magnesium decomposes $CO_2$. Dissolves in water to form carbonic acid. Acidic gas, reacts with alkalis. Uses: Fire extinguishers, dry ice (refrigerant), carbonated drinks, photosynthesis. Lime ($CaO$) from Limestone ($CaCO_3$): $CaCO_3(s) \rightarrow CaO(s) + CO_2(g)$ by heating. Uses of Lime and Slaked Lime: Neutralizing acidic soil, drying agent for ammonia. Uses of Limestone: Manufacturing lime, cement, glass, iron. Greenhouse Effect: Atmospheric gases trap heat, preventing loss from Earth. Increased $CO_2$ causes global warming. 17.8. Silicon Properties: Metalloid. Uses: Semiconductors (transistors, diodes, capacitors). Silicones: Macromolecules (oils, waxes, plastics). Fire Resistance of Silicone Plastics: Produce $SiO_2$ (sand) on combustion, carbon-based macromolecules produce $CO_2$. Uses of Silicon Dioxide (Sand): Glass, fire extinguisher, iron extraction. 18. Organic Chemistry 18.1. Organic Compounds Compounds of carbon (other than oxides, carbonates, carbides). 18.2. Hydrocarbons Binary compounds of carbon and hydrogen only. Saturated: Single C-C covalent bonds (alkanes). Unsaturated: At least one multiple C-C covalent bond (alkenes: double bond; alkynes: triple bond). 18.3. Alkanes Saturated hydrocarbons, general formula $C_nH_{2n+2}$. Names end in -ane. Structures (up to 5 carbons): Number of C atoms (n) Name Molecular Formula Structural Formula 1 Methane $CH_4$ $H-C(H_2)-H$ 2 Ethane $C_2H_6$ $CH_3-CH_3$ 3 Propane $C_3H_8$ $CH_3-CH_2-CH_3$ 4 Butane $C_4H_{10}$ Complete 5 Pentane $C_5H_{12}$ Complete Isomers: Compounds with same molecular formula but different structures. Begins with butane. Chemical Properties: Combustion: Excess air (complete): Form $CO_2$ and $H_2O$. (e.g., $CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(l)$) Limited air (incomplete): Form $CO$ and $H_2O$. (e.g., $CH_4(g) + O_2(g) \rightarrow CO(g) + H_2O(g)$) Halogenation: In presence of sunlight/UV light, react with halogens to form substituted products. Cracking: Decomposition of long-chain hydrocarbons into shorter alkanes and alkenes. Unreactivity: Saturated, no active site for chemical attack. 18.4. Alkenes Unsaturated hydrocarbons with double C=C bond. General formula $C_nH_{2n}$. Names end in -ene. Unsaturation Test: Decolorize red-brown bromine solution rapidly. Structures (up to 5 carbons): Number of C atoms (n) Name Molecular Formula Structural Formula 2 Ethene $C_2H_4$ $CH_2=CH_2$ 3 Propene $C_3H_6$ $CH_2=CH-CH_3$ 4 Butene $C_4H_8$ Complete 5 Pentene $C_5H_{10}$ Complete Isomers: Positional isomerism (changing double bond position) begins with butene. Chemical Properties (Addition Reactions): Combustion: Burn with luminous flame (excess air: $CO_2 + H_2O$; limited air: $CO + H_2O$). Hydrogenation (Hydrogen): Alkene + $H_2 \rightarrow$ Alkane. Halogenation (Chlorine/Bromine): Alkene + Halogen $\rightarrow$ Dihaloalkane. Hydrohalogenation (Hydrogen Halide): Alkene + HX $\rightarrow$ Haloalkane. Hydration (Steam): Alkene + $H_2O \rightarrow$ Alcohol. Polymerization: Alkenes react with themselves to form polyalkenes. Uses: Ethene (polythene, ripening fruits), propene (polypropene, ropes). 18.5. Alcohols (Alkanols) Organic compounds with hydroxyl group (-OH). General formula $C_nH_{2n+1}OH$. Names end in -ol. Formation: Hydration of Alkene: Alkene + $H_2SO_4 \rightarrow$ alkyl hydrogen sulphate, then hydrolyze with $H_2O \rightarrow$ alcohol. Fermentation of Carbohydrates: Starch $\rightarrow$ maltose $\rightarrow$ glucose $\rightarrow$ ethanol (yeast). Chemical Properties: Combustion: Burn in excess oxygen to form $CO_2$ and $H_2O$. Reaction with Sodium Metal: Alcohol + Na $\rightarrow$ Sodium alkoxide + $H_2$. Dehydration: Excess conc. $H_2SO_4$ at $170^\circ C \rightarrow$ alkene + water. Esterification: Alcohol + Carboxylic acid $\rightarrow$ Ester + water. Oxidation: Primary alcohols $\rightarrow$ alkanal $\rightarrow$ alkanoic acid. Uses: Organic solvent, fuel, antiseptic, alcoholic beverages, thermometric liquid. 18.6. Carboxylic Acids (Alkanoic Acids) Organic compounds with carboxyl group (-COOH). General formula $C_nH_{2n+1}COOH$. Names end in -oic acid. Formation: Oxidation of alcohols, hydrolysis of esters. Chemical Properties: React with bases $\rightarrow$ salt + water. Liberate hydrogen with active metals. React with carbonates/hydrogen carbonates $\rightarrow$ salt + water + $CO_2$. Esterification (with alcohols) $\rightarrow$ ester + water. Uses: Preservative, tenderizer, flavoring, solvent. 18.7. Esters (Alkanoates) Organic compounds, general formula RCOOR (R = alkyl radicals). Formation: Alcohol + organic acid (esterification). Chemical Properties: Combustion ($CO_2 + H_2O$), hydrolysis (alcohol + organic acid). Uses: Perfumes, food flavorants, varnishes. 18.8. Homologous Series Collection of organic compounds with same general formula, similar chemical properties, adjacent members differ by -$CH_2$ group. Gradual change in physical properties (melting/boiling points, density, solubility) as molecular mass changes. 18.9. Macromolecules (Polymers) Giant (large) molecules formed by joining small molecules (monomers). Synthetic Macromolecules (Polymers): Human-made. Examples: Polythene, polystyrene, PVC. Made by addition polymerization. Nylon: Formed by condensation polymerization of diamines and dioic acids, forming amide (-NHCO-) linkages. Terylene: Formed by condensation polymerization of diols and dioic acids, forming ester (-COO-) linkages. Plastics: Used for carrier bags, buckets, pipes. Synthetic Fibres (Nylon, Terylene): Clothes, tents, strings, ropes. Biodegradability: Non-biodegradable (litter problems, air pollution on burning). Natural Macromolecules: Occur naturally. Examples: Cellulose (wood, cotton) from glucose monomers. Natural rubber (isoprene). Proteins (amino acids). Fats/oils (fatty acids + glycerol). Starch: Glucose monomers joined by oxygen bonds (glycosidic linkages) via condensation polymerization. Proteins: Amino acid monomers joined by amide (peptide) linkages via condensation polymerization. Fats/Oils: Glycerol + fatty acids joined by ester linkages via esterification. Biodegradability: Biodegradable (rot). Hydrolysis of Fats: Reaction with water/steam (in presence of alkali) to produce glycerol and fatty acids. Saponification: reaction with alkali to produce soap and glycerol. Hydrolysis of Starch: Produces simple sugars (monosaccharides) like glucose. Hydrolysis of Proteins: Produces amino acids.