CAIE AS Level Chemistry Theory PDF

ZNOTES.ORG UPDATED TO 2023-2025 SYLLABUS CAIE AS LEVEL CHEMISTRY SUMMARIZED NOTES ON THE THEORY SYLLABUS Prepared for Aarav for personal use only. CAIE AS LEVEL CHEMISTRY 1. Atoms, Molecules and Stoichiometry 1.1. Relative Mass Atomic mass (Ar): weighted average mass of an atom Molecular mass (Mr): mass of a molecule Formula mass: mass of one formula unit of a compound Isotopic mass: mass of a particular isotope of an element Compared with 12C where one atom of 12C has mass of 1.5. Calculations Involving Mole exactly 12 units Unified atomic mass unit: u = 1.66 x 10-27kg Concept Mass 1.2. The Mole M oles = Molar Mass Mole: amount of substance that has the same number V olume of a Gas = M oles × 24 of particles (atoms, ions, molecules or electrons) as there The formula applies to gases at r.t.p. are atoms in exactly 12g of the carbon-12 isotope. Avogadro’s constant: number of atoms, ions, molecules Unit of volume is dm3 and 1000cm3 = 1dm3 or electrons in a mole = 6.02 × 1023 Moles Concentration = Volume 1.3. Mass Spectra Concentration unit = mol dm−3 Anhydrous: a compound in which all water molecules Abundance of isotopes can be represented on a mass spectra diagram are removed Hydrated: a compound which has a number of water Relative Abundance = Total Peak Height × 100% molecules associated with its crystalline structure Height Ar = ∑ M ass × Relative Abundance Water of Crystallisation: these water molecules in a 100 hydrated compound are called water of crystallisation 2. Chemical Bonding 2.1. Electronegativity and Bonding 1.4. Empirical and Molecular Formulae Empirical Formula: gives the simplest ratio of different atoms present in a molecule Molecular Formula: gives actual numbers of each type of atom in a molecule Molecular formula can be calculated using the Mr of a compound and its empirical formula M olecular F ormula = (Empirical Formula)n Molecular Mass Where n = Mass of Empirical Formula Atomic M ass × N o. of M oles % Composition = × Molar Mass of Compound WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Electronegativity: the power of an atom to attract electrons to itself The electronegativity depends on: The radius of atom (atomic size) inversely ∝ electronegativity Nuclear attraction directly ∝ electronegativity Electronegativity increases across a period because of atomic radius ↓ and nuclear attraction ↑, so polarity ↑ Electronegativity decreases down a group because of atomic radius ↑ and nuclear attraction ↓, so polarity ↓ Dipole Moment: slight charges on atoms in a covalent bond due to differences in electronegativity The difference between the electronegativity of two Have high melting and boiling points atoms in a compound determines the overall dipole moment and overall polarity of the compound. A large difference in electronegativity will make the bond more polar (more ionic in nature), but a small difference in electronegativity will make the bond less polar (more covalent in nature) Coordination Number: number of oppositely charged ions that surround a particular ion in an ionic solid E.g: NaCl, MgCl2 2.3. Dot and Cross Diagrams 2.2. Ionic (Electrovalent) Bonding An ionic bond is the electrostatic attraction between oppositely charged ions. Structure: giant ionic lattice, crystalline solids WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY A coordinate bond is a covalent bond where both 2.4. Metallic Bonding electrons in the bond come from the same atom Conditions: Strong electrostatic forces of attraction between metal An atom should have a lone pair of electrons cations and delocalized mobile electrons An atom should be in need of a pair of electrons Structure: lattice of +ve ions surrounded by mobile e-s Donor: the atom that supplies the pair of electrons Strength of metallic bond increases with: Acceptor: the atom that accepts the pair of electrons Increasing positive charge on the ions in the lattice Coordinate bond is represented by an “→” drawn from Decreasing size of metal ions in the lattice the atom donating to towards the atom accepting Increasing number of mobile e-s per atom Formation of Ammonium ion NH+ 4: Formation of AlCl3 dimer (Al2 Cl6 ): 2.5. Covalent Bonding Covalent Bond: the bond formed by the sharing of pairs of electrons between the nuclei of two atoms. Bonding Electrons: e-s involved in bond formation Above 750oC, exists as vapor & covalent molecule AlCl3 As vapor cools, exists as dimer Al2Cl6 Non-bonding electrons or lone pair: pair of valence e-s that are not involved in bond formation Bond angle as AlCl3 = 120o Covalent compounds are made of molecules which are Bond angle as Al2Cl6 = 109.5o held together by weak intermolecular forces Elements in period 3 can expand their octet, including the compounds sulfur dioxide, SO2, phosphorus 2.7. Orbital Overlap pentachloride, PCl5, and sulfur hexafluoride, SF6. For a covalent bond to form, atomic orbitals containing They have low melting and boiling points. unpaired valence electrons must overlap each other S – S (Sigma Σ) 2.6. Coordinate (Dative Covalent) Bonding S – P (Sigma Σ) WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY P – P (Sigma Σ) 3 pairs of e’s 3 bonded, 0 lone pair Trigonal planar P – P (Pi π ) 120O E.g. BF3 Sigma bond has greater overlap ∴ σ > π Pi bond cannot exist without a Sigma bond. Note: Elements in Period 3 can expand their octet by making use of the energetically accessible but lower lying d-subshell 4 pairs of e’s for bonding. This means that some elements of period 3 can 4 bonded, 0 lone pair bond with more than 4 electrons at once. (e.g: Sulfur, Tetrahedral Phosphorus etc) 109.5O E.g. CH4 2.8. Shapes of Molecules The shape and bond angles of molecules depend on: The number of pairs of electrons around the central atom Whether these pairs are lone pairs or bonded pairs Valence shell electrons are arranged in pairs to minimize repulsion between themselves Order of Repulsion Strength (VSEPR Theory): 3 bonded, 1 lone pair Pyramidal 107O 2 pairs of e’s E.g. NH3 2 bonded, 0 lone pair Linear 180O E.g. CO2 WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY 2 bonded, 2 lone pair Angular 2.9. Hybridisation 104.5O E.g. H2O Definition Mix atomic orbitals like s and p to form new hybrid orbitals like sp2 , sp3 , and sp. The resultant orbital has different energy levels, shapes, and properties. sp Hybridisation When 1 s orbital combines with 1 p orbital. They form linear molecules with 180o angles, e.g. CO2 5 pairs of e’s and all hydrocarbons containing 2 bonds like C2 H2 5 bonded, 0 lone pair (Carbon-Carbon triple bond). Trigonal Bipyramid Has an equal amount of s and p orbital properties. 90O and 120O E.g. PF5 sp2 Hybridisation 6 pairs of e’s When 1 s orbital combines with 2 p orbital. 6 bonded, 0 lone pair They form trigonal planar molecules with 120o angles, Octahedral e.g. BH3 and all hydrocarbons containing 3 bonds like 90O C2 H4 ( C = C ). E.g. SF6 Has twice the property of p orbital properties than s orbital. WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Bond Energy: the energy needed to break one mole of a given bond in one mole of gaseous molecules Bond Length: distance between the centres of two nuclei of two adjacent atoms Double bonds are shorter than single bonds because double bonds have a greater negative charge density between the two atomic nuclei hence greater attraction The bond length depends on the radii of the two bonded atoms; the larger the radius, longer the bond length The strength of the bond depends on the length of the bond sp3 Hybridisation When 1 s orbital combines with 3 p orbital. They form tetrahedral molecules with 109.25o angles, 2.11. Hydrogen Bonding e.g. SiCl4 , and all hydrocarbons containing 4 bonds like The strongest type of intermolecular force in covalent C2 H4 ( C − C ). bonds Has three times the property of p orbital properties than For hydrogen bonding to occur, we need: s orbital. A molecule having a H atom bonded to F, O or N Molecule having F, O or N atom with lone pair of e-s Note: Important properties of water (high melting/boiling points, high surface tension) are due to the strong hydrogen bonds present between water molecules. 2.12. Polar and Non-Polar 2.10. Bonds Polar Covalent Bonds Bonds with slight ionic character The bond formed with atoms of different electronegativity Bonding e-s attracted more towards atom with greater electronegativity ∴ unequal sharing of electrons ∴ molecule develops slight charges = Polar Molecule Polar molecules have dipoles, electric charges of equal magnitude and opposite sign The greater the difference in electronegativity of the two bonded atoms, the greater is the ionic character Non-Polar Covalent Bonds WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY The bond formed between: Identical Atoms: the electronegativity of both atoms 3. States of Matter is the same so a pair of electrons shared equally Symmetrical Polyatomic Molecules: dipoles of bond expert equal & opposite effects, hence cancel 3.1. Basic Assumptions of Kinetic charge Theory Non-polar molecules have no overall charge Ideal Gas: a gas whose volume varies in proportion to 2.13. Intermolecular Forces temperature and inverse to pressure. Noble gases such as helium and neon approach ideal Intermolecular Forces: weak forces present between behaviour because of their low intermolecular forces. two covalent molecules Ideal Gas Laws: Induced Dipole (Van Der Waals’ Forces) Gas molecules move rapidly and randomly Very weak forces present between non-polar molecules The distance between gas molecules is greater than the Due to constant motion of e-s, at an instant, a non-polar diameter of molecules ∴ volume is negligible molecule develops poles due to distortion of electron No forces of attraction/repulsion between molecules density giving rise to instantaneous dipole, which can All collisions between particles are elastic EK conserved induce a dipole in the adjacent molecules Temperature of gas related to average EK of molecules Van der Waals forces increase with: Conditions at which gases behave ideally: increasing the number of contact points between High temperature molecules; point where molecules come close Low pressure together increasing number of electrons (+ protons) in Limitations of Ideal Gas Laws: molecule Real gases do not obey kinetic theory in two ways: There is not zero attraction between molecules We cannot ignore the volume of molecules themselves Permanent Dipole-Dipole Forces Deviations visible at low temp. and high pressure Weak forces present between polar molecules Molecules are close to each other Molecules are always attracted to charged rods, whether The volume of molecules is not negligible relative to the +ve or –ve because molecules have +ve and –ve charges container VDW forces present, pulling molecules to each other 2.14. Summary Pressure is lower than expected from ideal gas The effective volume is less than expected from the ideal gas 3.2. General Gas Equations P V = nRT M ass × RT Mr = PV P 1 V1 P 2 V2 = T1 T2 WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Quantity Unit Conversion Pressure Pascal 1KPa = 1000Pa Constant evaporation from the surface Particles continue to break away from the surface but Volume m3 1m3 = 1000dm3 = 1x106cm3 are trapped in space above the liquid. Temperature K OC + 273 As gaseous particles collide, some of them hit the surface of the liquid again and become trapped there. Standard Conditions: 101KPa and 298 K An equilibrium is set up in which the number of particles Mols of One Gas leaving the surface is balanced by the number rejoining Mole Fraction = Total Mols of Gases it. Partial Pressure of a Gas = M ole F raction × Liquid water molecules ⇌ Vapor water molecules T otal P ressure A fixed number of gaseous particles will be in the space above the liquid in this equilibrium. 3.3. Liquid State Vapour Pressure: pressure exerted by a vapour in Particles touching but may have gaps equilibrium with a liquid. Have EK slide past each other in random motion Vapour pressure increases as: Enthalpy of Fusion: heat energy required to change 1 mole of solid into a liquid at its melting point Heating a Solid (melting): Energy transferred makes solid particles vibrate 3.4. Crystallisation faster Forces of attraction weaken & solid changes to liquid Anhydrous Enthalpy of Vaporisation: heat energy required to change 1 mole of liquid into a gas at its boiling point Salt containing no water molecules. Heating a Liquid (vaporisation): Energy transferred makes liquid particles move faster Hydrous Forces of attraction weaken The highest energy particles escape first Salt containing water molecules. Liquid starts to evaporate – temp. Below b.p. Forces weaken further – particles move faster & Water of Crystallisation Spread Liquid boils – temp. At b.p. The number of water molecules present in the Hydrous salt. The evaporation of a liquid in a closed container For example, CuSO4 ⋅ 5(H2 O) has 5 water of crystalisation. 3.5. Recycling Finite resource: resource which doesn't get replaced at the same rate that it is used up. Examples of finite resources: copper, aluminium, glass Advantage of Recycling: ○ Saves energy ○ Reduces environmental issues ○ Conserves ore supplies ○ Less wastage ○ Cheaper than extracting 3.6. States of Matter and Properties Solubility For solids, generally, solubility increases with increasing temperature as the increase in temperature facilitates the overcoming of intermolecular bonds, making it easier for the solid to dissolve. For gases, generally, solubility decreases with increasing temperature as the pressure of the gas increases (pressure only affects the solubility of gases) WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Conductivity Diamond: High m.p./b.p. - each carbon forms four covalent Solids (metals) are generally the best conductors of bonds electricity, while gases are the worst. Hard-tetrahedral structure Solids (metals) are generally better conductors of heat Doesn’t conduct heat or electricity – no free e- than liquids, while liquids are better thermal conductors Used for cutting as is the strongest known substance than gases. This is because of the proximity of molecules and has sharp edges in solids, allowing heat to be transferred rapidly through vibrations of neighbouring molecules. 3.7. Solid State Ionic lattice Graphite: Three strong (sp2) covalent bonds Fourth e- in p orbital ∴ forms a pi bond, forming a cloud of delocalised electrons above and below the Metallic lattice planes Layers kept together by weak Van der Waal’s forces High m.p./b.p. - strong covalent bonds throughout Soft – forces between layers are weak Conducts electricity - has delocalized electrons Simple molecular Macromolecular Lattice: WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Silicon(IV) Oxide: Effect of Hydrogen Bonding on Physical Properties: Each Si is bonded to 4 oxygen atoms, but each Relatively high m.p./b.p.: many strong H-bonds oxygen is bonded to 2 Si atoms High viscosity: hydrogen bonding reduces the ability Sand is largely SiO2 of water molecules to slide over each other Similar properties to diamond High surface tension: hydrogen bonds in water exert a downward force on the surface of the liquid Ice is less dense than water: larger spaces between molecules in a hexagonal structure Simple Molecular Lattice: Iodine: Dark grey crystalline solid; vaporizes into purple gas m.p./b.p. are slightly higher than room temp Slightly soluble in water; dissolves in organic solvents Diatomic molecules formed due to covalent bonds between individual atoms Molecules have weak Van der Waals forces of attraction between them Hydrogen-Bonded Lattice: In ice form, water molecules slow down and come closer together Due to polarity, molecules form hydrogen bonds between lone pairs of oxygen & δ + charge of hydrogens Each water molecule has 2 H-bonds They arrange themselves into an open crystalline, hexagonal structure Due to large spaces, ice is less dense than water Fullerenes: Buckminsterfullerenes(C60) C atoms in pentagonal and hexagonal rings Spherical C60 molecules held together by Van der Waals forces Can conduct heat and electricity Very strong and tough Insoluble in water Low m.p./b.p. WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Energy Change is measured in 1mol dm−3 4.2. Reaction Pathway Diagrams Exothermic Reactions 4.3. Nanotubes C atoms in hexagonal rings only Cylindrical The structure is rod-like due to continuing rings Conducts heat and electricity Very strong and tough Insoluble in water High m.p./b.p. Endothermic Reactions 4. Chemical Energetics 4.4. Enthalpy Change Definitions 4.1. Energy Change in Reactions Standard Conditions (this syllabus assumes that these Exothermic Reactions Endothermic Reactions are 298K and 101 kPa) shown by ⦵ Energy given out Energy taken in Standard Enthalpy Change of: Surrounding warmer Surrounding cooler Bond making Bond breaking ΔH negative ΔH positive EReactants > EProducts EReactants < EProducts WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY 1. Reaction ΔH 1. Enthalpy changes when 1 mole of element or compound is completely reacted under standard conditions in their standard states. 2. Combustion ΔHC 1. Enthalpy changes when 1 mole of element or compound is completely combusted under standard conditions in their standard states. 3. Formation ΔHF 1. Enthalpy changes when 1 mole of a compound is formed from its elements under standard conditions in their standard state. Formation from Combustion 4. Neutralization ΔHn 1. Enthalpy change when 1 mole of H + and OH − combine to form 1 mole of H2 O under standard conditions in their standard states 4.5. Bond Energy Energy needed to break a specific covalent bond Also how much energy is released when a bond forms 4.6. Calculating Enthalpy Changes Q = −mcΔT Hydration from Anhydrous Salt Q ΔH = no. of moles When substance dissolved in water use c & m of water ΔT is a change in temp.: add –ve or +ve to show rise/fall 4.7. Hess’s Law The total enthalpy change in a chemical reaction is independent of the route by which the chemical reaction takes place as long as the initial and final conditions are the same. Reason to Use Hess’s Law: Reaction from Bond Energies Standard conditions hard to maintain (e.g. Exo/endo) Elements don’t always react directly 4.8. Calculation of Enthalpy Change, ΔH Reaction from Formation WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Source: Alt Academy 5. Equilibria 5.1. Introduction Reversible reaction: a reaction in which products can be changed back to reactants by reversing the conditions Dynamic Equilibrium: the state of a reversible reaction carried out in a closed container where the rates of forward and backward reactions are equal and constant 4.9. Worked Examples - Hess Law 5.2. Le Chatelier’s Principle When a chemical system in dynamic equilibrium is disturbed (conditions changed) it tends to respond in such a way so as to oppose the change and a new equilibrium is set up 5.3. Equilibrium Constants Equilibrium constant expressed in terms of concentration Product mols KC = [[Reactant]]mols Only liquids and gases Equilibrium constant expressed in terms of partial pressure p(P roduct)mols KP = p(Reactant) mols Only gases Large value of KC /KP ⇒ equi. towards products side Smaller value of KC /KP ⇒ equi. towards reactants side KC /KP changes only with changes in temperature The amount of reactants that disappear will always appear in the products in the same ratio as present in a balanced equation 5.4. Brønsted–Lowry Theory of Acids and Bases WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Group 1 2 3 4 Brønsted-Lowry Theory: Element Sodium Magnesium Aluminium Silicon An acid is a proton (H+) donor Character Metal Metalloid A bases is a proton (H+) acceptor Structure Giant metallic lattice Macromolecular Amphoteric: substances that can act like bases or acids Metallic bond between cations Covalent bonds Bonding Strong acid: an acid that dissociates completely in and delocalized e- between atoms solution (e.g., HCl) Weak acid: an acid that dissociates partially in solution (e.g., ethanoic acid) Strong base: a base that dissociates completely in Diagram solution (e.g., NaOH) Weak base: a base that dissociates partially in solution (e.g., ammonia) Strong acids/bases react more vigorously than weak acids/bases. 6.2. Reaction of Elements with Oxygen Strong acids have lower pH values than weak acids. Strong bases have higher pH values than weak bases. Formulae Reaction Structure Oxid. Nature When an acid reacts with a base, salt & water are Burns yellow Na Na2O(s) +1 Basic formed. The pH changes in this neutralisation reaction flame can be graphed as shown in the image below: \n Burns Bright Giant ionic Mg MgO(s) +2 Basic White Flame lattice Burns Bright Al Al2O3(s) +3 Amphoteric White Flame SiO2(s) Giant Si Coating +4 W. acidic covalent P2O3(s) Burns yellow +3 P S. Acidic P2O5(s) flame Simple +5 SO2(g) Burns blue molecular +4 S S. acidic SO3(g) flame +6 P2 O5 empirical formula of P4 O10 6.3. Reaction of Na & Mg with Water 6. The Periodic Table: Na & 2Na(s) + 2H2O(l) Very fast, floats, forms Chemical Periodicity Water 2NaOH(aq) + H2(g) Mg(s) + 2H2O(l) ball & dissolves Mg & Very slow Water Mg(OH)2(aq) + H2(g) 6.1. Introduction Mg & Mg(s) + H2O(g) MgO(s) + Very fast Steam H2(g) 6.4. Reaction of Oxides with Water WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Reaction Oxid. Nature Na2O(s) Na2O(s) + H2O(l) 2NaOH(aq) +1 S. Alkaline Sodium chloride simply dissolves in water. Water is polar MgO(s) MgO(s) + H2O(l) Mg(OH)2(aq) +2 W. Alkaline ∴ positive Na+ attracted to OH- while Cl- attracted to H+ MgCl2 slightly acidic because Mg ion has smaller radius & Al2O3(s) NO REACTION higher charge ∴ attraction to water is so strong that H2O SiO2(s) loses a proton and solution becomes slightly acidic P2O3(s) P2O3(s) + 3H2O(l) 2H3PO3(aq) +3 S. Acidic P2O5(s) P2O5(s) + 3H2O(l) 2H3PO4(aq) +5 SO2(g) SO2(g) + H2O(l) H2SO3(aq) +4 S. Acidic SO3(g) SO3(g) + H2O(l) H2SO4(aq) +6 6.5. Acid-Base Reactions Aluminum oxide is amphoteric ∴ reacts with the acid 6.8. Atomic Radius and base Al2O3 + H2SO4 → Al2(SO4)3 + H2O | Al2O3 + NaOH →NaAlO2 + H 2O Silicon dioxide is acidic: SiO2 + NaOH (hot & conc.) Na2SiO3 Sulphur dioxide and trioxide are strongly acidic With Produces SO2(g) NaOH NaHSO3(aq) SO2(g) Excess NaOH Na2SO3(aq) + H2O SO3(g) NaOH NaHSO4(aq) SO3(g) Excess NaOH Na2SO4(aq) + H2O P+ in nucleus increases so nuclear charge increases There are more e-, but increase in shielding is negligible 6.6. Reactions of Elements with because each extra e- enters same principal energy level Chlorine ∴ force of attraction between nucleus & e- increases... So atomic radius decreases. Formula Structure Oxid. Nature Na NaCl(s) Giant ionic +1 Neutral 6.9. Ionic Radius Mg MgCl2(s) +2 Neutral Al AlCl3(s) +3 Acidic Si SiCl4(l) +4 S. Acidic PCl3(l) Simple molecular +3 P S. Acidic PCl5(l) +5 6.7. Reactions of Chloride with Water WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Ionic radius decreases across a period however, since non-metals gain electrons, they have one more shell Na < Mg < Al because no. of delocalized electrons which than metals therefore they always have a larger radius can carry charge increases than metal ions Silicon is a semi-conductor Non-metals – covalent ∴ no charge 6.10. Melting Point 6.12. Electronegativity Na Al m.p. increases because delocalized e- per atom increases making metallic bond stronger Increases across period because the bonded e- are in the Si has highest m.p. due to giant covalent structure same energy level but are attracted more strongly as no. The larger the molecule size, the stronger the Van Der of protons increases Waals forces ∴ S8 > P4 > Cl2 > Ar 6.13. First Ionisation Energy 6.11. Electrical Conductivity WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY All group 2 metals tarnish in air, forming oxide coatings Burn vigorously in oxygen-forming white solids 7.3. Reactions with Water Metals: M(s) + H2 O(l) → M (OH)2(aq) + H2(g) Metal Oxides: MO(s) + H2 O(l) → M(OH)2(aq) Solubility of MO (Metal Group 2 Oxide) and M(OH)2(Metal Generally, increases as no. of protons increases Group 2 Hydroxides) increases down group Decrease Mg → Al: more distant and less effective Alkalinity of solution increases down the group nuclear charge on 3p orbital Solubility of MSO4 decreases down group Decrease P → S: in S, one electron paired ∴ causing M(CO3) don’t react with water repulsion and easier to lose an electron 7.4. Reaction with Acids 7. Group 2 M(s) + Acid(aq) → Salt + Hydrogen 7.1. Introduction MO(s) + Acid(aq) → Salt + Water M(OH)x(s) + Acid(aq) → Salt + Water MCO3(s) + Acid(aq) → Salt + Water + Carbon Dioxi 7.5. Thermal Decomposition of Group 2 Metals MCO3(s) MO(s) +CO2(g) 2M(NO3 )2(s) 2MO(s) +4NO2(g) + O2(g) m.p./b.p. decreases down group: atoms/ions get larger, the distance between nuclei & e-s increases ∴ bonds NO2: thick brown, acidic and soluble gas weaker Thermal stability increases down the group ∴ m.p./b.p. higher in gp. 2 than 1: 2e-s per atom donated decomposition becomes more difficult. into delocalized system ∴ metallic bonding stronger density increases down group: mass of atoms 7.6. Uses of Group 2 Metals increases faster than their size (volume) as atomic no. increases Calcium compounds: 7.2. Reaction of Group 2 Metals with Oxygen WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Halide With Silver With dilute aq. With conc. aq. Calcium oxide (lime): basic oxide used to neutralize Ion Nitrate ammonia ammonia acidic soil and used as a drying agent for drying White ppt. ppt. dissolves Cl‑ ammonia Calcium carbonate (limestone): used as building Br- Cream ppt. X ppt. dissolves material (cement, concrete) etc., for extraction of iron, I- Yellow ppt. X X glass industry, neutralize soil or chemical waste AgX(s) + 2NH3(aq) → [Ag(NH3)2]+(aq) + X- 8. Group 17 The complex ion formed is called Diamine Silver(I) ion [H3 N :⟶ Ag ←: N H3 ]+ 8.1. Trends in Colour and Volatility Fluorine Yellow 8.5. (Sub) Halide Ions and Aqueous Chlorine Yellow- Gas The Sulfuric Acid Green m.p. & b.p. Volatility colour Orange- Metal Halide + Conc. H2SO4(aq) Hydrogen Halide Bromine Liquid increases decreases gets Brown down the down the darker Grey, group ↓ group ↓ down the Conc. H2SO4(aq) is an oxidising agent (except for chloride Iodine Black, group ↓ and fluorides as it is not strong enough) Violet Solid This reaction is used for the preparation of hydrogen Astatine Black halides 8.2. Oxidising Ability Chlorine NaCl(s) + H2SO4(aq) HCl(g) + NaHSO4(aq) NaBr(s) + H2SO4(aq) HBr(g) + NaHSO4(aq) Halogens have high electron affinity (they gain electrons Bromine HBr(g) + H2SO4(aq) Br2(g) + SO2(g) + H2O(l) easily) hence, they are good oxidising agents NaI(s) + H2SO4(aq) HI(g) + NaHSO4(aq) Oxidising ability decreases down the group because electron affinity decreases as atomic size increases. HI(g) + H2SO4(aq) I2(g) + SO2(g) + H2O(l) Iodine HI(g) + H2SO4(aq) I2(g) + H2S(g) + H2O(l) 8.3. Reactions of Halide Ions 6HI(g) + H2SO4(aq) 3I2(g) + S(s) + 4H2O(l) X2(g) + H2(g) 2HX(g) 8.6. The Reactions of Chlorine with Product Reaction Description HF Reacts explosively in all conditions Aqueous Sodium Hydroxide HCl Reacts explosively in sunlight Disproportionation: a reaction in which the same HBr Reacts slowly on heating substance is oxidized and reduced simultaneously, HI Forms an equilibrium mixture on heating producing two different products When chlorine reacts with a solution of cold aqueous The thermal stability of halogen hydrides decreases sodium hydroxide, the disproportionation goes to lower down the group because: oxidation states The size of halogen atom increases Nuclear attraction decreases Cl2 + 2N aOH → N aCl + N aClO + H2 O The H – X bond becomes longer and weaker Thus, less energy is needed to break the bond With a hot solution, the oxidation state of chlorine goes Bond energies decrease down the group up to +V 3Cl2 + 6N aOH → N aClO3 + 5N aCl + 3H2 O 8.4. (Sub) Halide ions and aq. Silver Ions This happens as the chlorate is formed by Ag+(aq) + X-(aq) AgX(s) disproportionation of hypochlorite and hypochlorous acid ClO− + 2HClO → ClO3− + 2HCl WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Higher temperatures promote the formation of hypochlorous acid through the hydrolysis of hypochlorite and, therefore, speed up the reaction 8.7. Important Uses of Halogens and Halogen Compounds Fluorine: To make chlorofluorocarbon (CFCs) As fluoride in toothpaste To make polytetrafluoroethylene (PTFE) – non- sticking coating in pots and pans Bromine and Iodine: manufacture of photographic films Displacement of ammonia from its salts: Chlorine: Any Ammonium Salt + Any Base warm In bleaches Ammonia Gas + salt + water To make PVC and chlorofluorocarbon (CFCs) As solvents Use of chlorine in water purification: 9.3. Uses of Ammonia & its Compounds The oxidising power of chlorine is used in the treatment of water to kill bacteria Used in the production of nitric acid Used in the production of inorganic fertilizers Cl2(aq) + H2O(l)→ HCl(aq) + HClO(aq) Used in the production of nylon HClO(aq)→ HCl(aq) + O Used in the production of explosives This disproportionation reaction produces reactive oxygen atoms which kill bacteria 9.4. Oxides of Nitrogen N2(g) + O2(g)→ 2NO(g) or ½N2(g) + O2(g)→ NO2(g) 9. Nitrogen and Sulfur Naturally: during lightning, EA provided for N2 to react Man-made: in car engine, high temp. and pressure 9.1. Lack of Reactivity of Nitrogen Catalytic convertors: exhaust gases passed through catalytic convertors containing a catalyst (platinum/ Nitrogen molecule has three strong covalent bonds. palladium/nickel) helping to reduce oxides to nitrogen. The bond is very strong and requires high energy to split Catalytic role in oxidation of sulphur dioxide: the two nitrogen atoms of a molecule. It reacts only under extreme temperature or pressure or in the presence of a catalyst. 9.2. Ammonium Lone pair of e-s of nitrogen forms a coordinate bond with the H+ ion 9.5. Pollution Formation: NH3(g) + H+ NH4+ Shape: tetrahedral Acid Rain: SO3 + H2O→ H2SO4 Bond angle: 109.5o 2NO2 + H2O → HNO3 + HNO2 or 4NO2 + 2H2O + O2 → Bond length: equal lengths 4HNO3 Damages trees & plants, kills fish and other river life, buildings, statues and metal structures WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Combustion Pollutants: Hybridisation: mixing up of different atomic orbitals Nitrogen oxide (NO): formed by reaction of N2 and O2 resulting in new orbitals of equal energy. in the engine, forms acid rain and respiratory problems Carbon’s Electron Configuration: Atmospheric oxides of nitrogen (NO & NO2) can react with unburned hydrocarbons to form peroxyacetyl nitrate (PAN) which is a component of photochemical smog Carbon monoxide (CO): source: incomplete combustion of hydrocarbon fuel, toxic effect on haemoglobin sp3 sp2 sp All orbitals mix 2s, 2px, 2py mix 2s and 2px mix 9.6. Food Preservation 3 sp2 orbitals 2 sp orbitals 4 sp3 orbitals 2 pure p orbitals SO2 is used by itself or as a sulphite to preserve food 1 pure p orbital Ratio of characteristics s : p SO2 + H2O → H2SO3(aq) 1:3 1:2 1:1 SO2 & sulphites inhibit growth of bacteria, yeasts, etc. & are reducing agents, so reduce rate of oxidation of food. Used to prevent spoilage of dried fruit, dehydrated vegetables and fruit juices. 10. An Introduction to AS Level Organic Chemistry 10.3. Classes of Compound 10.1. Introduction Organic Chemistry: study of hydrocarbons and their derivatives Carbon can form a variety of compounds because: Carbon is tetravalent Carbon-carbon bonds can be single, double or triple Atoms can be arranged in chains, branches and rings Homologous series: a series of compounds of similar structures In which: contain the same functional group all share the same general formula the formula of homologue differs from neighbour by CH2 similar chemical properties gradual change in physical properties as Mr increases Functional group: an atom or group of atoms in an organic molecule that determines the characteristic reactions of a homologous series. Alkyl group: a reactive group which is alkane minus 1 H 10.2. Hybridisation WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY 10.4. Types of Formulae Two atoms sharing e- pair are of different electronegativity Displayed formula (Hexane) When the bond breaks, one of the bonded atoms takes both bonding e-s Results in the formation of +ve and –ve ions If +ve charged on C, it is called carbocation or carbonium If –ve charged on C, it is called carbanion Note: homolytic fission requires less energy than heterolytic Structural formula (Hexane) 10.7. Types of Reagents CH3-CH2CH2CH2CH2-CH3 or CH3(CH2)4CH3 Skeletal formula (Hexane) Nucleophilic reagent (nucleophile): donator of pair of e- Must have lone pair of e-s Attack centre of +ve charge (positive pole) Reaction with nucleophile called nucleophilic reactions Examples: CH-, Cl-, NH3, H2O, CN- Molecular formula (Hexane) C6H14 Electrophilic reagent (electrophile): acceptor of pair of e- +ve ions or e- deficient molecules 10.5. IUPAC Nomenclature Attack regions of high e- density Select the longest chain as the main chain Examples: Br+, CH3+, AlCl3 Other carbon chains as substituent alkyl groups Give the lowest number C in the main chain to a 10.8. Types of Reaction substituent If different alkyl groups are present on identical Addition reaction: single product formed positions, give simpler alkyl smaller number Electrophilic addition (alkenes) Two or more alkyl groups present, order alphabetically Nucleophilic addition (carbonyl compounds) If the same substituent is repeated use the di, tri, or tetra Substitution reaction: two products formed prefix Nucleophilic substitution (halogenoalkanes) If the ring of carbon is present, use the prefix “cyclo” Free radical substitution (alkanes) Write the position of the double bond in alkene, e.g. but- Elimination reaction: more than one product formed, 1-ene small molecule removed from reactant (alcohols and halogenoalkanes) 10.6. Breaking of Covalent Bonds Hydrolysis reaction: breaking down of molecule by water, sped up by acid or alkali (esters and alkenes) Homolytic Fission: 10.9. Oxidation and Reduction Two atoms sharing e- pair of similar electronegativity When bond breaks, each atom takes one e- from pair of Oxidation: addition of oxygen or removal of hydrogen electrons forming free radicals Reduction: addition of hydrogen or removal of oxygen Free radicals: electrically neutral atoms or groups of atoms with unpaired electrons are very reactive 10.10. Shapes of Ethane and Ethene Free radical reaction catalysed by heat or light Heterolytic Fission: Ethane: sp3 bonds, all sigma bonds WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Note: Ethene: Planar Shape, H – C – H bond = 120o Straight chain alkanes have higher b.p. than branched Branching makes molecule more spherical reduces contact points VDW forces decreases 10.12. Chain Isomers Isomers have different carbon chain length Same chemical properties but slightly different physical Example: Benzene 10.13. Position Isomers Isomers differ in position of substituent atoms or group or the functional group Same chemical properties but slightly different physical Example: But-1-ene 10.11. Isomerism Existence of two or more compounds with the same molecular formula but different structural formula Example: But-2-ene WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Arises from different arrangement of atoms or groups in 3D space resulting in two isomers Have effect on polarised light Chiral carbon: a carbon having 4 single bonds and 4 different atoms or groups Isomers non-super-imposable images of each other Have same physical and chemical properties No. of optical isomers in a molecule containing n chiral carbons = 2n 10.14. Functional Isomers Isomers have different functional groups, belong to different homologous series Have different physical and chemical properties Ratio of C : H Functional Gps. Example 1:3 Alcohol & Ether C2H6O 1:2 Aldehyde & Ketone C3 H 6 O 1:2 Must have O2 Carboxylic acid & Ester C3 H 6 O 2 11. Hydrocarbons 10.15. Geometric (cis/trans) Isomers 11.1. Properties Shown only by alkenes Generally unreactive: Arises due to restriction of double bond Only possible when each carbon has 2 different groups All C–C bonds single; alkanes = saturated hydrocarbons cis-trans isomers have different b.p. Non-polar ∴ no center of charge to act as either cis isomers have higher dipole nucleophile or electrophile ∴ cannot attract polar trans isomer of symmetrical alkene has zero dipole reagents like acids, bases, metals or oxidizing agents Physical properties: The volatility of the alkanes decreases and m.p/b.p increases as number of carbon atoms increases Reason: increasing Van der Waals forces 11.2. Combustion Used as fuel because they burn in oxygen to given out 10.16. Optical Isomers large amounts of energy Alkanes kinetically stable in presence of O2; combustion occurs when necessary amount of Ea supplied Reaction occurs only in gas phase Complete: carbon dioxide + water Incomplete: carbon monoxide + carbon (soot) + water General Equation of Hydrocarbon Combustion: CxHy + (x + y4 ) O2 → xCO2 + y2 H2 O WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY 11.3. Substitution Hint: CH3o and Clo represents radicals which is neutral atom Alkanes react with halogens: Cl2 and Br2 and an unpaired electron. Example: Chlorination of Methane Reagent Condition Reaction 11.7. Cracking Type Mechanism Cl2(g) UV light Substitution Free Radical Breaking of large less useful alkanes into useful, more energy value smaller products using heat & catalyst Products: smaller alkanes and alkenes or smaller alkenes and hydrogen gas Thermal cracking: high temp. & pressure Catalytic cracking: high temp. & catalyst 11.8. Hydrocarbons as Fuels Source of alkanes: crude oil Steady change in b.p. of alkanes allows crude oil to be separated by fractional distillation Catalytic conversion of CO and NOx: Reactants are Halogens and Alkane. 2NO2 + 4CO → N2 + 4CO2 Involves 3 steps: initiation, propagation, and termination. 2NO + 2CO → N2 + 2CO2 requires the action of UV light or Heat. 11.9. Alkenes 11.4. Initiation Unsaturated hydrocarbons Breakdown of Chlorine into radicals. Contain at least one C=C double bond Uses the action of UV light. General formula: CnH2n (like cycloalkanes) Starts the reaction. Source of alkenes: Creates radicals. Cracking alkanes Cl2(g) → 2Cl(g)o Dehydration of alcohols More reactive than alkanes due to presence of double bond; pi electrons loosely and more susceptible to 11.5. Propagation attacks by e- deficient groups like electrophiles Chlorine radical reacts with the alkyl/alkane. Alkenes combust completely carbon dioxide + water Helps the reaction to propagate (chain reaction). Give energy but not used as fuels; have other uses Maintains balance of radicals. CH4 + Clo → CH3o + HCl 11.10. Electrophilic Addition Mechanism CH3o + Cl2 → CH3 Cl + Clo 11.6. Termination Chlorine radical and alkyl radical reacts. Stops propagation of reaction. Reduces the number of radicals. CH3o + Clo → CH3 Cl 2Clo → Cl2 CH3o + CH3o → CH3 C H3 WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Electrophile forms by heterolytic fission Hydrogenation (Alkene + H2 Alkane) Electrophile attacks double bond Reagent: H2(g) Pair of e-s from double bond migrate to electrophile and Condition: π bond breaks Catalyst: Nickel Carbocation formed which attacks the nucleophile Temp.: 100oC Press.: 2 atm. 11.11. Carbocations Use: convert liquid oils to saturated solid fats Halogenation (Alkene + X2 Dihaloalkane) Reagent: Halogen(aq) Condition: r.t.p./dark Halogenation (Alkene + Hydrohalogen Halogenoalkane) Reagent: Hydrohalogen(g) Condition: r.t.p. Hydration (Alkene + H2O(g) Alcohol) Reagent: steam Condition: Markovnikov’s principle: an electrophile adds to an Catalyst: H3PO4 – phosphoric acid unsymmetrical alkene so that the most stable Temp.: 300oC carbocation is formed as an intermediate Press.: 70atm Hydrogen binds to carbon that is more stable Inductive effect of alkyl groups: 11.13. Oxidation of Alkenes Alkyl groups donate e- to the ring Producing a positive inductive effect Both oxidation and addition to double bond involved A larger alkyl group has a weaker inductive effect KMnO4 changes from purple/pink to colourless With Cold Dilute Acidified KMnO4/H+ 11.12. Addition Reactions Diol is formed With Hot Concentrated Acidified KMnO4/H+ This leads to the rupture of the double bond Two compounds are formed Products formed depend on alkene WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY General conditions: high pressure, 11.14. Polymerisation high temperature and catalyst Disadvantages: Repeated addition of 1000s of alkene molecules Non-biodegradable: does not break down so (monomer) to each other forming a macromolecule increases the amount of volume needed for landfill Polyethene: sites LDPE: cling wrap Combustion produces harmful gases which HDPE: water pipes, wire insulation contribute to global warming e.g. SO2, CO2 and HCl from PVCs Disposal of Polymers: Recycle existing plastic Make polymers biodegradable by adding starch units 12. Halogen Compounds 12.1. Types of Halogenoalkanes Polychloroethene (PVC): Water pipes Insulation of wires WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Primary 1o (SN2) The C-X bond is a polar bond and has partial charges due to the high electronegativity of halogen. The δ+ carbocation is easily susceptible to attack by a nucleophile SN1 Mechanism: Secondary 2o (SN2 and SN1) Unimolecular – only one molecule involved in 1st step Secondary and Tertiary halogenoalkanes SN2 Mechanism: Tertiary 3o (SN1) Bimolecular – two molecules involved in 1st step Primary and secondary halogenoalkanes 12.4. Nucleophilic Substitution Reaction 12.2. Strength of C – Hal Bond Polar Nature Bond Energy Reactivity Fluoro Chloro Decrease Decrease Increases Bromo Iodo Electronegativity Bond length increases, bond decreases energy decreases, lower EA so down group more reactive 12.3. Nucleophilic Substitution Mechanism WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Hydrolysis (R – X + OH- R – OH + X-) 12.6. Elimination Reaction Reagent: strong alkali; NaOH(aq) or KOH(aq) Condition: heat/reflux R – X + OH- Alkene + X- + H2O Fluoroalkanes are not hydrolysed because the C – F bond is too strong Mechanism: Ease of hydrolysis increases: Primary < Secondary < Tertiary Tertiary halogenoalkanes can be hydrolysed without alkali Note: if any Cl- or Br- ions present in NaOH(aq), these ions will interfere with reaction Reagent: ethanolic NaOH or KOH Nitrile (cyanide) (R – X + CN- RCN + X-) Conditions: temp. 60oC, reflux Reagent: KCN or NaCN in ethanol OH- acts as a proton acceptor; it accepts the H+ loss from Condition: the halogenoalkanes during elimination Solvent: Ethanol Elimination becomes progressively easier Heat/Reflux Reaction forms a C – C bond; therefore number of Primary < Secondary < Tertiary Carbon increases; name has one more carbon Primary Amines (R – X + NH3 RNH2(l) + HX(g)) Note: the carbon atom adjacent to carbon with halide Reagent: Ammonia (NH3) must have at least one hydrogen attached to it. Condition: Ammonia in alcohol under pressure in a sealed container 12.7. Uses of Halogenoalkanes Note: If excess concentration of ammonia used, HX reacts with it forming NH4X CFCs are inert and can be liquefied easily: Strength of C – X bond is very high, hence do not decompose easily 12.5. Reflux and are not flammable. Uses: Many organic reactions proceed slowly As propellants in aerosol cans Heating done under reflux to prevent volatile organic As solvents in dry-cleaning solvents to evaporate As refrigerant for freezers and fridges Mechanism similar to simple distillation Fire extinguishers, insecticides and pesticides 12.8. CFCs Effect on Ozone Layer Destroys the ozone layer CFCs escape the atmosphere and, because of their inertness, remain without further reaction until they reach the stratosphere and ozone layer. In the stratosphere, high energy U.V causes the Cl atom to split CFC molecule forming Cl⋅, which reacts with ozone This is a catalytic cycle where one Cl⋅ can react with many O3 thus causing destruction of ozone layer: Cl⋅ + O3(g) ⋅OCl(g) + O2(g) ⋅OCl(g) + O(g) Cl⋅ + O2(g) Can react and breakdown another O3 molecule WWW.ZNOTES.ORG Copyright © 2025 ZNotes Education & Foundation. All Rights Reserved. This document is authorised for personal use only by Aarav at The Sanskaar Valley School on 02/02/25. CAIE AS LEVEL CHEMISTRY Boiling Point: Note: the alternative is using HCFCs (replace Cl with H or b.p. decreases→ more F atoms) as they break down more easily and do not release Cl → less effect on the ozone layer 13. Hydroxy Compounds Because: 13.1. Types of Alcohols b.p. of alcohols > alkenes as they have hydrogen bonds Primary 1o Solubility of Alcohols in Water: Smaller alcohols mix completely with water since strong hydrogen bonds occur between alcohol and water As hydrocarbon nature increases (i.e. more C-C… bonds), the non-polar character outweighs the ability of the OH to form hydrogen bonds and ∴ solubility decreases Small alcohols (e.g. ethanol) are good solvents for both Secondary 2o polar and non-polar compounds as they have polar and non-polar components 13.3. Reaction with Sodium R – OH + Na(l) RO- Na+ + ½ H2(g) Type of reaction: acid-base Reagent used: liquid sodium metal Tertiary 3o Reactivity of alcohols decreases with increasing chain lengths of hydrocarbon Reaction les

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