A-Level Chemistry Z Notes

TABLE OF CONTENTS 3 CHAPTER 1 3 CHAPTER 2 5 CHAPTER 3 7 CHAPTER 4 9 CHAPTER 5 10 CHAPTER 6 11 CHAPTER 7 12 CHAPTER 8 Atoms, Molecules & Stoichiometry Atomic Structure Chemical Bonding States of Matter Chemical Energetics Electrochemistry Equilibria Reaction Kinetics CHAPTER 9 14 Chemical Periodicity 16 CHAPTER 10 Group II – Alkaline Earth Metals CIE AS-LEVEL CHEMISTRY//9701 17 CHAPTER 11 18 CHAPTER 12 Group 17 – Halogens Nitrogen & Sulphur 18 CHAPTER 13 22 CHAPTER 14 25 CHAPTER 15 26 CHAPTER 16 Introduction to Organic Chemistry Hydrocarbons Halogen Derivatives Hydroxy Compounds CHAPTER 17 28 Carbonyl Compounds 29 CHAPTER 18 29 CHAPTER 19 Carboxylic Acids & Derivatives Analytical Techniques PAGE 2 OF 30 CIE AS-LEVEL CHEMISTRY//9701 1. ATOMS, MOLECULES AND STOICHIOMETRY 2. ATOMIC STRUCTURE 1.1 Relative Mass 2.1 Subatomic Particle Relative Atomic mass (Ar): Molecular mass (Mr): Formula mass: Isotopic mass: weighted average mass of an atom mass of a molecule mass of one formula unit of a compound mass of a particular isotope of an element compared with 12C where one atom of 12C has mass of exactly 12 units Subatomic Particle Protons (P) Neutrons (n) Electrons (e-) Relative Charge +1 0 -1 Relative mass/ a.m.u 1 1 1/1840 2.2 Behavior of a Beam of Subatomic Particles 1.2 The Mole  Mole: amount of substance that has the same number of particles (atoms, ions, molecules or electrons) as there are atoms in exactly 12g of the carbon-12 isotope.  Avogadro’s constant: number of atoms, ions, molecules or electrons in a mole = 6.02 10 1.3 Mass Spectra  Abundance of isotopes can be represented on a mass spectra diagram ℎ ℎ  Protons: positively charged ∴ deflected to -ve pole  Neutrons: no charge ∴ not deflected  Electrons: negatively charged ∴ deflected to +ve pole  e- lighter than P+ ∴ deflected at greater angle 100% 100 1.4 Empirical and Molecular Formulae 2.3 Protons, neutrons and electrons  Empirical formula: gives 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  Mass concentrated within centre; nucleus  An atom is electrically neutral; P+ = e Atomic no. or proton no. (Z) = no. of protons  Atomic mass or nucleon no. (A) = no. of P + N Where % . 100% 1.5 Calculations involving Mole Concept 24  Formula applies to gases at r.t.p.  Unit of volume is and 1000  Concentration unit 1  Isoelectronic ions: ions having same no. of e-s  Isotopes: are atoms of the same element with the same proton number but different number of neutrons 2.4 Electronic Configuration  Electrons are arranged in energy levels called shells  Each shell is described by a principle quantum no. (P.Q)  As the P.Q. increases, energy of shell increases  Inside the shell there are subshells: , , and  Orbital: region in space where there is a maximum probability of finding an electron  Each orbital can hold 2e-s in opposite directions PAGE 3 OF 30 CIE AS-LEVEL CHEMISTRY//9701 Orbitals 1 3 5 7 Max e s 2 6 10 14  Aufbau’s principle: method of showing how atomic orbitals are filled in a definite order to give lowest energy arrangement possible  Energy difference between 4 & 3 very small ∴ an e- from 4 can be promoted to half-fill or full-fill 3 orbital, to make atom more stable  When filling, fill 4s before 3d and when removing, also remove first from 4s 2.6 Shapes of Subshells -Subshell  Spherical shape  Increases in size as P.Q no. increases -Subshell Nuclear Charge Shielding Effect 2.5 Subshells 2.8 Factors affecting Ionization Energy Atomic Radius  When e s are placed in a set of orbital of equal energy, they occupy them singly and then pairing takes place  e-s placed in opposite direction: both -vely charge & if placed in same direction, they’d repel. In opposite direction they create a spin to reduce repulsion  Completely filled or half filled (i.e. one e- in each orbital) are more stable (reduced repulsion) Stable Config. -  +ve charge due to protons in nucleus  Greater nuclear charge  greater ionization energy  Inner shells of e-s repel outermost e-s, thus shielding them from +ve nucleus. The more eshells, the greater is the shielding effect  Greater effect  lower I.E because lesser attractive force between nucleus & outer e-s  Distance from the centre of the nucleus to the outermost orbit  Greater radius  lower I.E; distance of outermost e- to nucleus is large ∴ less energy needed to remove e High I.E needed to remove e-s from completely or half-filled orbitals 2.9 General 1st I.E Trends First Ionization Energy Trends Down a Group Across a Period DECREASES INCREASES  New shells added  Shell no. remains same  Attraction of nucleus to  Proton no. increases valence e s decreases  Effective nuclear charge  Shielding effect increases increases  Atomic radius decreases 2.10 Trend in 1st I.E across 3rd Period  Dumbbell shape 2.7 Ionization Energies (I.E)  1st I.E: energy needed to remove 1 mole of e-s from 1 mole of gaseous atom to form 1 mole of unipositive ions  Each successive I.E is higher than previous one because as e-s are removed, protons > e-s ∴ attraction between protons and remaining Big jump occurs electrons increases between I.E 1 & 2  Successive I.Es have ∴ part of 1st gp large jump in their value when e-s removed from lower energy shell  Deduce group no. by checking when 1st big jump occurs  I.E of Al lower than Mg: e- removed in Al is from higher energy 3p orbital which is further away from nucleus than 3s e- being removed from Mg. Nuclear attraction is less for 3p than 3s ∴ I.E of Al is lower than Mg  I.E of S lower than P: e- being removed in P is in a half filled, more stable 3p orbital whereas in S, the pairing of electrons in 3p results in increased repulsion ∴ less energy need to remove an e- PAGE 4 OF 30 CIE AS-LEVEL CHEMISTRY//9701 2.11 Ionic Radius 3.3 Coordinate/Dative Bonding  Ionic radius: describes the size of an ion  Positive ion: smaller radius than original neutral atom because shell no. decreases, screening effect decreases but the attraction of nucleus increases.  Negative ion: larger ionic radius than neutral atom because e-s added while nuclear charge remains same Groups 1 to 3 5 to 7 Ion Positive Negative No. of shells  Coordinate bond is a covalent bond where both electrons in the bond come from the same atom  Conditions: o An atom should have a lone pair of electrons o An atom should be in need of a pair of electrons  Donor: the atom that supplies the pair of electrons  Acceptor: the atom that accepts the pair of electrons  Coordinate bond is represented by an “” drawn from the atom donating to towards the atom accepting  Formation of Ammonium ion ( ):  Proton no. and effective nuclear charge increases  Ionic radius decreases  Negative ions always larger than positive ions in the same period as they have one more shell  Formation of dimer ( : 3. CHEMICAL BONDING 3.1 Ionic (Electrovalent) Bonding o Above 750oC, exists as vapor & covalent molecule AlCl3 o As vapor cools, exists as dimer Al2Cl6 o Bond angle as AlCl3 = 120o o Bond angle as Al2Cl6 = 109.5o  Ionic bond is the electrostatic attraction between oppositely charged ions.  Structure: giant ionic lattice, crystalline solids  Have high melting and boiling points 3.4 Orbital Overlap S–S Sigma Σ  Coordination number: number of oppositely charged ions that surround a particular ion in an ionic solid S–P Sigma Σ 3.2 Covalent Bonding P–P Sigma Σ P–P  For a covalent bond to form, atomic orbitals containing unpaired valence electrons must overlap each other Pi  Covalent bond is the bond formed by the sharing of pairs of electrons between two atoms.  Bonding electrons: e-s involved in bond formation  Non-bonding electrons or lone pair: pair of valence e-s that are not involved in bond formation  Covalent compounds are made of molecules which are held together by weak intermolecular forces  They have low melting and boiling points  Sigma bond has greater overlap ∴  Pi bond cannot exist without a Sigma bond. PAGE 5 OF 30 CIE AS-LEVEL CHEMISTRY//9701 3.5 Shapes of Covalent Molecules 3.7 Electronegativity  Shape and bond angles of molecules depend on: o The number of pairs of electrons around central atom o 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):  Ability of a particular atom, covalently bonded, to attract the bonded pair of e-s towards itself  Electronegativity depends on: o Radius of atom inversely ∝ electronegativity o Nuclear attraction directly ∝ electronegativity  Electronegativity increases across a period because atomic radius ↓ and nuclear attraction ↑  Electronegativity decreases down a group because atomic radius ↑ and nuclear attraction ↓  Dipole moment: slight charges on atoms in a covalent bond due to differences in electronegativity Lone -Lone Lone - Bonded Type Shape Angle 2 Pairs of e-s 2 bonded 180O Linear Bonded - Bonded Example CO2 3 Pairs of e-s 3 bonded Trigonal Planar 3.8 Bonds 120O BF3 - 4 Pairs of e s 4 bonded 109.5O Tetrahedral CH4 3 bonded 1 lone Pyramidal 2 bonded 2 lone Angular 107O NH3 104.5O H2O  Bond energy: energy needed to break one mole of a given bond in one mole of gaseous molecules  Bond length: distance between the centers 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  Bond length depends on radii of the two bonded atoms; larger the radius, longer the bond length  Strength of the bond depends on the length of the bond Longer bond - 5 Pairs of e s 5 bonded More reactive 3.9 Polar and Non-Polar Trigonal Bipyramid 90O PF5 6 Pair of e-s 6 bonded Weaker bond 90O Octahedral SF6 3.6 Hydrogen Bonding  Strongest type of intermolecular force in covalent bonds  For hydrogen bonding to occur, we need: o Molecule having a H atom bonded to F, O or N o Molecule having F, O or N atom with lone pair of e-s Ammonia Water .. .. .. Polar Covalent Bonds  Bonds with slight ionic character  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  Bond formed between: o Identical atoms: the electronegativity of both atoms is the same so pair of electron shared equally o Symmetrical polyatomic molecules: dipoles of bond exert equal & opposite effects hence cancel charge  Non-polar molecules have no overall charge PAGE 6 OF 30 CIE AS-LEVEL CHEMISTRY//9701 3.10 Intermolecular Forces 4. STATES OF MATTER  Intermolecular forces: weak forces present between two covalent molecules Induced Dipole (Van Der Waals’ Forces)  Very weak forces present between non-polar molecules  Due to constant motion of e-s, at an instant, a non-polar molecule develops poles due to distortion of electron density giving rise to instantaneous dipole, which is able to induce a dipole in the adjacent molecules  Van der Waals forces increase with: o increasing number of contact points between molecules; point where molecules come close together o increasing number of electrons (+ protons) in molecule ↑ no. of electron ↓ nuclear attraction Faster edistortion Stronger force Permanent Dipole-Dipole Forces  Weak forces present between polar molecules  Molecules always attracted to charged rod, whether +ve or –ve because molecules have +ve and –ve charges 3.11 Metallic Bonding  Strong electrostatic forces of attraction between metal cations and delocalized mobile electrons  Structure: lattice of +ve ions surrounded by mobile e-s  Strength of metallic bond increases with: o Increasing positive charge on the ions in the lattice o Decreasing size of metal ions in the lattice o Increasing number of mobile e-s per atom 3.12 Summary 4.1 Basic Assumptions of Kinetic Theory  Ideal gas: a gas whose volume varies in proportion to temperature and in inverse proportion to pressure.  Noble gases such as helium and neon approach ideal behavior because of their low intermolecular forces. Ideal Gas Laws:  Gas molecules move rapidly and randomly  Distance between gas molecules is greater than diameter of molecules ∴ volume is negligible  No forces of attraction/repulsion between molecules  All collisions between particles are elastic conserved  Temperature of gas related to average of molecules  Conditions at which gases behave ideally: o High temperature o Low pressure Limitations of Ideal Gas Laws:  Real gases do not obey kinetic theory in two ways: o There is not zero attraction between molecules o We cannot ignore volume of molecules themselves Deviations visible at low temp. and high pressure  Molecules are close to each other  Volume of molecules not negligible relative to container  VDW forces present, pulling molecules to each other  Pressure is lower than expected from ideal gas  Effective volume is less than expected from ideal gas 4.2 General Gas Equations Quantity Unit Conversion Pressure Pascal 1 1000 Volume m3 1 = 1000 =1 O O Temperature K C + 273  Standard Conditions: 101KPa and 273oK 10 cm3 Mole Fraction Partial Pressure of a Gas 4.3 Liquid State  Particles touching but may have gaps  Have to slide past each other in random motion  Enthalpy of fusion: heat energy required to change 1 mole of solid into a liquid at its melting point PAGE 7 OF 30 CIE AS-LEVEL CHEMISTRY//9701  Heating a solid (melting): o Energy transferred makes solid particles vibrate faster o Forces of attraction weaken & solid changes to liquid  Enthalpy of vaporization: heat energy required to change 1 mole of liquid into a gas at its boiling point  Heating a liquid (vaporization): o Energy transferred makes liquid particles move faster o Forces of attraction weaken o Highest energy particles escape first o Liquid starts to evaporate – temp. below b.p. o Forces weaken further – particles move faster & spread o Liquid boils – temp. at b.p.  The evaporation of a liquid in a closed container o Constant evaporation from surface o Particles continue to break away from surface but are trapped in space above the liquid o As gaseous particles collide, some of them hit the surface of the liquid again, and become trapped there o An equilibrium is set up in which number of particles leaving surface is balanced by number rejoining it. liquid water molecules ⇌ vapor water molecules o In this equilibrium, there will be a fixed number of the gaseous particles in the space above liquid.  Vapor pressure: pressure exerted by a vapor in equilibrium with a liquid.  Vapor pressure increases as: Temp. ↑ Ek ↑ Forces weaken More vapor 4.4 Solid State Ionic Lattice Metallic Lattice Simple Molecular Macromolecular Lattice:  Diamond: o High m.p./b.p. - each carbon forms four covalent bonds o Hard - tetrahedral structure o Doesn’t conduct heat or electricity – no free eo Used for cutting as is strongest known substance and has sharp edges  Graphite: o Three strong (sp2) covalent bonds o Fourth e- in p orbital ∴ forms a pi bond, forming a cloud of delocalised electron above and below the planes o Layers kept together by weak Van der Waal’s forces o High m.p./b.p. - strong covalent bonds throughout o Soft – forces between layers are weak o Conducts electricity - has delocalized electrons  Silicon(IV) Oxide: o Each Si is bonded to 4 oxygen atoms, but each oxygen is bonded to 2 Si atoms o Sand is largely SiO2 o Similar properties to diamond 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 & 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  Effect of Hydrogen Bonding on Physical Properties: o Relatively high m.p./b.p.: many strong H-bonds o High viscosity: hydrogen bonding reduces ability of water molecules to slide over each other o High surface tension: hydrogen bonds in water exert a downward force on surface of liquid o Ice less dense than water: larger spaces between molecules in hexagonal structure Simple Molecular Lattice:  Iodine: o Dark grey crystalline solid; vaporizes into purple gas o m.p./b.p. are slightly higher than room temp o Slightly soluble in water; dissolves in organic solvents o Diatomic molecule formed due to covalent bond between individual atoms o Molecules have weak Van der Waals forces of attraction between them PAGE 8 OF 30 CIE AS-LEVEL CHEMISTRY//9701 Exothermic Reactions  Fullerenes: Buckminsterfullerenes(C60) Nanotubes C atoms in pentagonal and C atoms in hexagonal hexagonal rings rings only Spherical Cylindrical C60 molecules held Structure is rod like due together by VDWs to continuing rings Conducts heat and electricity Very strong and tough Insoluble in water High m.p./b.p.  Metal + Acid or Water  Thermal decomposition  Neutralization  Dissolving salts  Combustion  Photosynthesis  Condensation & Freezing  Boiling & Evaporation  Standard Enthalpy Conditions: o Temperature: 298 or 25 o Pressure: 101 or 1 o Solution Conc.: 1 5.2 Enthalpy Change Definitions Standard molar enthalpy change of Combustion ∆ Formation Solution ∆ 4.5 Ceramics ∆ Hydration ∆ Atomisation ∆ Neutralization ∆  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 1 mole of and combine to form 1 mole of 1 mole of gaseous atom formed from its element 1 mole of ions in the gas phase are dissolved in water 1 mole of a solute is dissolved in a solvent to form an infinitely dilute solution 1 mole of compound is formed from its elements 4.6 Recycling 1 mole of element or compound is completely combusted Enthalpy change when  Ceramic: an inorganic non-metallic solid prepared by heating one or a mixture of substance(s) to a high temp.  Most ceramic are giant molecular structures  Properties of ceramics: o High m.p./b.p. and hard – strong covalent bonds o Don’t conduct electricity/heat – no mobile ions or e-s o Chemically unreactive – e-s held in covalent bonds under standard conditions in their standard states 5.3 Bond Energy  Energy needed to break a specific covalent bond  Also how much energy is released when a bond forms 5.4 Calculating Enthalpy Changes ∆ ∆  When substance dissolved in water use & of water  ∆ is change in temp.: add –ve or +ve to show rise/fall 5. CHEMICAL ENERGETICS 5.1 Energy Change in Reactions Exothermic Reactions  Energy given out  Surrounding warmer  Bond making  ∆ negative Endothermic Reactions Endothermic Reactions  Energy taken in  Surrounding cooler  Bond breaking  ∆ positive PAGE 9 OF 30 CIE AS-LEVEL CHEMISTRY//9701 5.5 Hess’s Law 6. ELECTROCHEMISTRY  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: o Std. conditions hard to maintain (e.g. exo/endo) o Elements don’t always react directly 6.1 Calculating Oxidation Numbers 5.6 Calculating Enthalpy Change of… …Reaction from Formation …Formation from Combustion  Ionic Molecules: group number = valence electrons  Covalent molecules: o Rules:  Atoms in a diatomic molecule; oxidation number =0  Oxygen in a compound; oxidation number = -2  Oxygen as peroxide; oxidation number = -1  1st group elements & hydrogen; oxidation number = +1  H with highly reactive metal; oxidation number = -1 o Following these rules, all other atoms in a covalent bond must balance out the charge 6.2 Redox Reactions  Reaction where both oxidation and reduction occur  Can be shown with changes in oxidation numbers of elements from the product side to the reactant side  E.g. 6 →2 4 o 4 → 2 ⟹ gain of negative charge ∴ reduction o 4 → 4 ⟹ loss of negative charge ∴ oxidation …Hydration from Anhydrous Salt 6.3 Balancing Equations  Equation: →  Half ionic: 2 2 → 2 →  For every 2 iodines, there will be 1 nitrogen  Thus first put in correct ratio for iodine and nitrogen then balance hydrogens and oxygens  Balanced: 4 2 →2 3 …Reaction from Bond Energies PAGE 10 OF 30 CIE AS-LEVEL CHEMISTRY//9701 6.4 Electrolysis 7.2 Equilibrium Constants  Electrolysis: decomposition of an electrolyte by an electric current. Electrical energy is used to bring about a chemical reaction; endothermic  Electrolyte: an aqueous solution of an ionic substance or a molten ionic salt that conducts electricity due to mobile ions  Electrodes: o Rods which help current enter the electrolyte o Inert electrodes: do not take part in the reaction e.g. graphite or platinum. Steel/titanium used in industry. o Reactive electrodes: take part in the reaction 6.5 Products of Electrolysis At the cathode Equilibrium Constant Expressed in terms of concentration Only liquids and gases Only gases  Large value of / ⇒ equi. towards products side ⇒ equi. towards reactants side  Smaller value of /  / 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 7.3 Manufacture of Ammonia Metal HABER PROCESS FLOWCHART Methane and steam Fractional distillation of ⇌ 3 liquid air Molten Electrolyte At the anode At the cathode Non-metal Higher in reactivity then H+ H2 formed Lower in reactivity then H+ Metal formed Aqueous Electrolyte Concentrated Halogen fomed Dilute Oxygen formed Halide present At the anode Halide not present Expressed in terms of partial pressure Oxygen Formed 7. EQUILIBRIA  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 7.1 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 PAGE 11 OF 30 Hydrogen 3 Nitrogen Compressed and Heated  Catalyst: Iron beds ⇌2  Temp: 500oC  Pressure: 200atm Ammonia CIE AS-LEVEL CHEMISTRY//9701 7.4 Manufacture of Sulphuric Acid 7.6 Conjugate Pairs CONTACT PROCESS FLOWCHART Burning fossil fuels Fractional distillation of air Sulphur  When acid-base reacts, an equilibrium mixture is formed Oxygen ⟶ Filter sulphur dioxide 2 3  Catalyst: V2O5  Temp: 500oC  Pressure: 1atm (slight) ⇌2 ⟶ [Oleum] ⟶2 Sulphuric Acid  SO3 not dissolved directly into water because reaction explosive and causes H2SO4 to vaporize  Forward reaction exothermic ∴ temp. not too high so rate of backward reaction doesn’t increase & not too low so particles have for collision and catalyst works  Since reaction highly exothermic, gases must be cooled  No impurities otherwise catalyst will be poisoned  Atmospheric pressure enough because equilibrium already favours product side     is a conjugate acid of base is acid-I and is base-I is a conjugate acid of base is acid-II and is base-II & vice versa & vice versa 7.7 Strong and Weak Acids and Bases  Strong acids/bases: acids/bases which dissociate almost completely in solutions  Weak acids/bases: acids/bases which are only partially dissociated in solutions  Strong and weak acids and bases can be distinguished by the pH value of their aqueous solutions  Monoprotic acids: donate one H+ proton per molecule  Diprotic acids: donate two H+ protons per molecule 8. REACTION KINETICS  Rate of a reaction: change in concentration of reactants or products per unit time  Activation energy: minimum energy colliding particles must possess for a successful collision to take place  Catalysis: acceleration of a chemical reaction by catalyst 8.1 Effect of Concentration Changes 7.5 Acid-Base Equilbria  Brønsted-Lowry Theory: o An acid is a proton (H+) donor o A bases is a proton (H+) acceptor  Amphoteric: substances that can act like bases or acids  Increasing conc. of reactants increases rate of reaction: more particles per unit volume, collision rate between reacting particles increases, ∴ rate of successful collision increases, resulting in increased rate of reaction. PAGE 12 OF 30 CIE AS-LEVEL CHEMISTRY//9701 8.2 Maxwell-Boltzmann Theory  Explains effect of temp. & catalyst on rate of reaction  Based on distribution of energy among reacting molecules under different conditions 8.3 Effect of Temperature  Number of collisions and chance of success will increase 8.4 Effect of Catalyst  Catalyst: a substance that increases rate of reaction but remains chemically unchanged itself at the end  Does not alter the chemical composition of substances and only lowers the activation energy  It provides a new route or mechanism to follow for reactants that requires less energy  Curve unchanged, only activation energy changes  Homogeneous catalysts: reactant and catalyst are in the same physical state  Heterogeneous catalysts: reactant and catalyst are in different physical states  Enzymes: a protein molecule that is a biological catalyst. Most are specific to a substrate & function as lock-key PAGE 13 OF 30 MgO(s) Al2O3(s) SiO2(s) P2O3(s) P2O5(s) SO2(g) SO3(g) Mg Al Si Burns blue flame Reaction Burns yellow flame Burns blinding white flame Coating Coating Burns yellow flame Simple molecular Giant covalent Giant ionic lattice Structure 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g) Mg(s) + 2H2O(l)  Mg(OH)2(aq) + H2(g) Mg(s) + H2O(g)  MgO(s) + H2(g) Na & Water Mg & Water Mg & Steam 9.3 Reaction of Na & Mg with Water S P Na2O(s) Na Formulae 9.1 Reaction of Elements with Oxygen Diagram S. acidic S. Acidic Amphoteric W. acidic Basic Basic Nature +3 +5 +4 +6 S. Acidic S. Acidic PAGE 14 OF 30  Aluminium oxide is amphoteric ∴ reacts with acid and base Al2O3 + H2SO4  Al2(SO4)3 + H2O | Al2O3 + NaOH  NaAlO2 + H2O  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 P2O3(s) + 3H2O(l) 2H3PO3(aq) P2O5(s) + 3H2O(l)  2H3PO4(aq) SO2(g) + H2O(l)  H2SO3(aq) SO3(g) + H2O(l)  H2SO4(aq) NO REACTION Reaction Na2O(s) + H2O(l)  2NaOH(aq) MgO(s) + H2O(l)  Mg(OH)2(aq) 9.4 Acid-Base Reactions Na2O(s) MgO(s) Al2O3(s) SiO2(s) P2O3(s) P2O5(s) SO2(g) SO3(g) Simple atoms Atoms held by VDWs 0 Argon Oxid. Nature +1 S. Alkaline +2 W. Alkaline 6 7 Sulphur Chlorine Non-metals Simple molecular covalent Intra = covalent Inter = weak VDWs 5 Phosphorous 9.2 Reaction of Oxides with Water 4 Silicon Metalloid Macromolecular Covalent bonds between atoms Very fast, floats, forms ball & dissolves Very slow Very fast +3 +4 +3 +5 +4 +6 +2 +1 Oxid. Metallic bond between cations and delocalized e- 3 Aluminium Bonding 2 Magnesium Metal Giant metallic lattice 1 Sodium Group Element Character Structure 9. CHEMICAL PERIODICITY CIE AS-LEVEL CHEMISTRY//9701 CIE AS-LEVEL CHEMISTRY//9701 9.8 Ionic Radius 9.5 Reactions of Elements with Chlorine Na Mg Al Si P Formula NaCl(s) MgCl2(s) AlCl3(s) SiCl4(l) PCl3(l) PCl5(l) Structure Giant ionic Simple molecular Oxid. +1 +2 +3 +4 +3 +5 Nature Neutral Neutral Acidic S. Acidic S. Acidic Oxide + HCl(g) 9.6 Reactions of Chloride with Water Reaction Nature NaCl(s) NaCl(s) + H2O(l)  NaCl(aq) Neutral MgCl2(s) MgCl2(s) + H2O(l) MgCl2(aq) W. Acidic AlCl3(s) AlCl3(s) + H2O(l) Al2O3(s) + HCl(g) Acidic SiCl4(l) SiCl4(l) + H2O(l) SiO2(s) + HCl(g) S. Acidic PCl3(l) PCl3(l) + H2O(l) H3PO3(aq) + HCl(g) S. Acidic PCl5(l) PCl5(l) + H2O(l)  H3PO4(aq) + HCl(g)  Sodium chloride simply dissolves in water. Water is polar ∴ positive Na+ attracted to OH- while Cl- attracted to H+  MgCl2 slightly acidic because Mg ion has smaller radius & higher charge ∴ attraction to water is so strong that H2O loses a proton and solution becomes slightly acidic AlCl3  Ionic radius decreases across a period however, since non-metals gain electrons, they have one more shell than metals therefore they always have a larger radius than metal ions 9.9 Melting Point anhydrous covalent; exists as dimer Al2Cl6 Hydrated - exists as ions; Al3+ and Cl- AlCl3(s) + H2O(l) Al2O3(s) + HCl(g) [Al(OH)(H2O)5]2+ + HCl 9.7 Atomic Radius  Na  Al m.p. increases because delocalized e- per atom increases making metallic bond stronger  Si has highest m.p. due to giant covalent structure  The larger the molecule size, the stronger the VDW forces ∴ S8 > P4 > Cl2 > Ar 9.10 Electrical Conductivity  P+ in nucleus increases so nuclear charge increases  There are more e-, but increase in shielding is negligible because each extra e- enters same principal energy level  ∴ force of attraction between nucleus & e- increases ...  So atomic radius decreases.  Na < Mg < Al because no. of delocalized electrons which can carry charge increases  Silicon is a semi-conductor  Non-metals – covalent ∴ no charge PAGE 15 OF 30 CIE AS-LEVEL CHEMISTRY//9701 9.11 Electronegativity 10.2 Reaction of Gp. 2 Metals with Oxygen M O2 g →2MO s s  All gp. 2 metals tarnish in air forming oxide coatings  Burn vigorously in oxygen forming white solids 10.3 Reactions with Water Metals: Ms Metal Oxides: MO s H2 O l →M OH 2 aq H2 H2 O l →M OH 2 aq g  Increases across period because the bonded e- are in the same energy level but are attracted more strongly as no. of protons increases  Solubility of M, MO and M(OH)2 increases down group  Alkalinity of solution increases down the group  Solubility of M and MO increases do+wn the group  Solubility of M(OH)2 and MSO4 decreases down group 9.12 First Ionization Energy 10.4 Reaction with Acid MCO3 s M s Acid aq → Salt Hydrogen MO s Acid aq → Salt Water M OH x s Acid aq → Salt Water Acid aq → Salt Water Carbon Dioxide 10.5 Thermal Decomposition of Gp. 2 Metals MCO3 2M NO3  Generally increases as no. of protons increases  Decrease Mg  Al: more distant and less effective nuclear charge on 3p orbital  Decrease P  S: in S, one electron paired ∴ causing repulsion and easier to lose electron 2 s s ∆H ∆H MO 2MO s s CO2 g 2NO2 g O2 g  NO2: thick brown, acidic and soluble gas  Thermal stability increases down the group ∴ decomposition becomes more difficult. 10.6 Uses of Group II Metals 10. GROUP II – ALKALINE EARTH METALS  m.p./b.p. decreases down group: atoms/ions get larger, distance between nuclei & e-s increases ∴ bonds weaker  m.p./b.p. higher in gp. 2 than 1: 2e-s per atom donated into delocalized system ∴ metallic bonding stronger  density increases down group: mass of atoms increases faster than their size (volume) as atomic no. increases  Calcium compounds: 10.1 Reactivity of Alkaline Earth Metals  Atomic radius increases down group  Ionisation energy (I.E) decreases down the group  The lower the I.E, easier to remove electrons  Hence metals more reactive down the group  Gp. 2 less reactive than gp. 1 since they need to lose two e-s ∴ total I.E = 1st I.E. and 2nd I.E.  Gp. 2 metals form ionic compounds  Calcium oxide (lime): basic oxide used to neutralize acidic soil and used as a drying agent for drying ammonia  Calcium carbonate (limestone): used as building material (cement, concrete) etc., for extraction of iron, glass industry, neutralize soil or chemical waste PAGE 16 OF 30 CIE AS-LEVEL CHEMISTRY//9701 (Sub) Halide ions and aq. Sulphuric Acid 11. GROUP 17 – HALOGENS 11.1 Trends in Colour and Volatility m.p. & b.p. increases  Volatility decreases  Fluorine Yellow Gas Chlorine Yellow-Green Bromine Orange-Brown Liquid Iodine Grey-Blue Solid Astatine Black  As atomic number increases, the number of electrons increases, this increases VDW forces so stronger bonds thus m.p./b.p. increases and volatility decreases 11.2 Oxidising Ability  Halogens have high electron affinity (they gain electrons easily) hence they are good oxidising agents  Oxidising ability decreases down the group because electron affinity decreases as atomic size increases. 11.3 Some reactions of the halide ions X2(g) + H2(g)  2HX(g) Product Reaction Description HF Reacts explosively in all conditions HCl Reacts explosively in sunlight HBr Reacts slowly on heating HI Forms an equilibrium mixture on heating  Thermal stability of halogen hydrides decreases down the group because: o Size of halogen atom increases o ∴ nuclear attraction decreases o The H – X bond becomes longer and weaker o Thus less energy needed to break the bond  Bond energies decrease down the group (Sub) Halide ions and aq. Silver Ions Ag+(aq) + X-(aq)  AgX(s) Halide With Silver With dilute With conc. Ion Nitrate aq. ammonia aq. ammonia ClWhite ppt. ppt. dissolves Br Cream ppt. ppt. dissolves IYellow ppt.  The solubility of these ppts. are tested with dilute and conc. aq. ammonia to confirm presence of ion.  If ppt. dissolves, it forms a complex ion: AgX(s) + 2NH3(aq)  [Ag(NH3)2]+(aq) + X The complex ion formed is called Diamine Silver(I) ion ∶⟶ ←∶ Metal Halide + Conc. H2SO4(aq)  Hydrogen Halide  Conc. H2SO4(aq) is an oxidising agent  This reaction is used for preparation of hydrogen halides Chlorine NaCl(s) + H2SO4(aq)  HCl(g) + NaHSO4(aq) NaBr(s) + H2SO4(aq)  HBr(g) + NaHSO4(aq) Bromine HBr(g) + H2SO4(aq)  Br2(g) + SO2(g) + H2O(l) NaI(s) + H2SO4(aq)  HI(g) + NaHSO4(aq) Iodine HI(g) + H2SO4(aq)  I2(g) + SO2(g) + H2O(l) HI(g) + H2SO4(aq)  I2(g) + H2S(g) + H2O(l) 11.4 The reactions of chlorine with aqueous sodium hydroxide  Disproportionation: a reaction in which the same substance is oxidized and reduced simultaneously producing two different products  When chlorine reacts with a solution of cold aqueous sodium hydroxide, the disproportionation goes to lower oxidation states 2NaOH → NaCl NaClO H O Cl  With a hot solution, the oxidation state of chlorine goes up to V 3Cl 6NaOH → NaClO 5NaCl 3H O  This happens as the chlorate is formed by disproportionation of hypochlorite and hypochlorous acid ClO 2HClO → ClO 2HCl Higher temperatures promotes the formation of hypochlorous acid through hydrolysis of hypochlorite and therefore speeds up the reaction 11.5 Some important uses of halogens and of halogen compounds  Fluorine: o To make chlorofluorocarbon (CFCs) o As fluoride in toothpaste o To make polytetrafluoroethylene (PTFE) – non sticking coating in pots and pans  Bromine and Iodine: manufacture of photographic films  Chlorine: o In bleaches o To make PVC and chlorofluorocarbon (CFCs) o As solvents PAGE 17 OF 30 CIE AS-LEVEL CHEMISTRY//9701  Use of chlorine in water purification: o The oxidising power of chlorine is used in treatment of water to kill bacteria Cl2(aq) + H2O(l)  HCl(aq) + HClO(aq) HClO(aq)  HCl(aq) + O o This disproportionation reaction produces reactive oxygen atoms which kill bacteria 12.6 Pollution 12. NITROGEN AND SULPHUR 12.1 Lack of Reactivity of Nitrogen  Nitrogen molecule has three strong covalent bonds  Bond is very strong and requires high energy for splitting the two nitrogen atoms of a molecule.  It reacts only under extreme temperature or pressure or in presence of catalyst. 12.2 Ammonium  Lone pair of e-s of nitrogen forms a coordinate bond with the H+ ion  Formation: NH3(g) + H+  NH4+  Shape: tetrahedral  Bond angle: 109.5o  Bond length: equal lengths Displacement of ammonia from its salts: Any Ammonium Salt + Any Base warm → Ammonia Gas  Catalytic convertors: exhaust gases passed through catalytic convertors containing a catalyst (platinum/ palladium/nickel) helping to reduce oxides to nitrogen.  Catalytic role in oxidation of sulphur dioxide: + Salt + Water 12.3 Uses of Ammonia & its Compounds  Used in the production of nitric acid  Used in the production of inorganic fertilizers  Used in the production of nylon  Used in the production of explosives 12.4 Eutrophication  Nitrate fertilisers leach into rivers and lakes after rain  Water plants grow more than usual  They block sunlight and kill plants underneath  Bacteria/fungi decompose remains using the O2  Fish and other creatures die from oxygen starvation 12.5 Oxides of Nitrogen N2(g) + O2(g)  2NO(g) or N2(g) + ½O2(g)  NO2(g)  Naturally: during lightning, EA provided for N2 to react  Man-made: in car engine, high temp. and pressure Acid Rain: SO3 + H2O  H2SO4 2NO2 + H2O  HNO3 + HNO2 or NO2 + H2O + ½O2  HNO3  Damages trees & plants, kills fish and other river life, buildings, statues and metal structures Combustion Pollutants:  Nitrogen oxide (NO): formed by reaction of N2 and O2 in the engine, forms acid rain and respiratory problems  Carbon monoxide (CO): source: incomplete combustion of hydrocarbon fuel, toxic effect on haemoglobin 12.7 Food Preservation  SO2 is used by itself or as a sulphite to preserve food SO2 + H2O  H2SO3(aq)  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. 13. INTRODUCTION TO ORGANIC CHEMISTRY  Organic chemistry: study of hydrocarbons and their derivatives  Carbon can form a variety of compounds because: o Carbon is tetravalent o Carbon-carbon bonds can be single, double or triple o Atoms can be arranged in chains, branches and rings  Homologous series: a series of compounds of similar structures in which: o contain the same functional group o all share same general formula o formula of homologue differs from neighbour by CH2 o similar chemical properties o gradual change in physical properties as Mr increases  Functional group: an atom or group of atoms in an organic molecule that determine the characteristic reactions of a homologous series.  Alkyl group: a reactive group which is alkane minus 1 H PAGE 18 OF 30 CIE AS-LEVEL CHEMISTRY//9701 13.1 Hybridization  Hybridisation: mixing up of different atomic orbitals resulting in new orbitals of equal energy.  Carbon’s electron configuration: Ground State Excited State 2s 2p 2s 2p ↿⇂ ↑ ↑ ↑ ↑ ↑ ↑ sp3 All orbitals mix 4 sp3 orbitals 1:3 sp2 sp 2s, 2px, 2py mix 2s and 2px mix 3 sp2 orbitals 2 sp orbitals 1 pure p orbital 2 pure p orbitals Ratio of characteristics s : p 1:2 1:1 PAGE 19 OF 30 CIE AS-LEVEL CHEMISTRY//9701 13.2 Classes of Compound Organic Family Suffix Example -ane Methane -ene Ethene Alkanes Alkenes halo- … -ane Chloroethane Halogenoalkanes -ol Methanol -al Methanal -one Propanone -oic Methanoic acid Alcohols Aldehydes Ketones Carboxylic Acid -oate Methyl ethanoate Esters -amine Methylamine -nitrile Ethyl nitrile Amines Nitriles PAGE 20 OF 30 CIE AS-LEVEL CHEMISTRY//9701 13.3 Types of Formulae Hexane Displayed Formula Structural Formula CH3-CH2CH2CH2CH2-CH3 or CH3(CH2)4CH3 Skeletal Formula Molecular Formula C6H14 13.4 Nomenclature  Select longest chain as the main chain  Other carbon chains as substituent alkyl groups  Give lowest number C in main chain to substituent  If different alkyl groups present on identical position, give simpler alkyl smaller number  Two or more alkyl groups present, order alphabetically  If same substituent repeated use di, tri, tetra prefix  If ring of carbon present, use prefix cyclo  Write position of double bond in alkene e.g. but-1-ene 13.5 Breaking of Covalent Bonds Homolytic Fission:  Two atoms sharing e- pair of similar electro-tivity  When bond breaks, each atom takes one e- from pair of electrons forming free radicals  Free radicals: electrically neutral atoms or group of atoms with unpaired electrons  very reactive  Free radical reaction catalysed by heat or light Heterolytic Fission:  Two atoms sharing e- pair are of different electro-tivity  When bond breaks, one of the bonded atoms takes both bonding e-s  Results in formation of +ve and –ve ions  If +ve charge on C, its called carbocation or carbonium  If –ve charge on C, its called carbanion  +ve ions or e- deficient molecules  Attack regions of high e- density  Examples: Br+, CH3+, AlCl3 13.7 Types of Reaction  Addition reaction: single product formed o Electrophilic addition (alkenes) o Nucleophilic addition (carbonyl compounds)  Substitution reaction: two products formed o Nucleophilic substitution (halogenoalkanes) o Free radical substitution (alkanes)  Elimination reaction: more than one product formed, small molecule removed from reactant (alcohols and halogenoalkanes)  Hydrolysis reaction: breaking down of molecule by water, sped up by acid or alkali (esters and alkenes) 13.8 Oxidation and Reduction  Oxidation: addition of oxygen or removal of hydrogen  Reduction: addition of hydrogen or removal of oxygen 13.9 Shapes of Ethane and Ethene Ethane sp3 bonds all sigma bonds Ethene Planar Shape H – C – H bond = 120o Benzene 13.10 Isomerism  Existence of two or more compounds with the same molecular formula but different structural formula Note: homolytic fission require less energy than heterolytic ISOMERS 13.6 Types of Reagents Structural: molecules that have the Nucleophilic reagent (nucleophile): donator of pair of esame molecular formula, but have a different arrangement of the atoms  Must have lone pair of e-s  Attack centre of +ve charge (positive pole) Chain Position Functional  Reaction with nucleophile called nucleophilic reactions  Examples: CH-, Cl-, NH3, H2O, CNElectrophilic reagent (electrophile): acceptor of pair of ePAGE 21 OF 30 Stereo: molecules that have same structure, but different spatial arrangement of atoms in 3D space Geometric Optical CIE AS-LEVEL CHEMISTRY//9701 Note:  Straight chain alkanes have higher b.p. than branched  Branching makes molecule more spherical  reduces contact points  VDW forces decreases 13.11 Chain Isomers  Isomers have different carbon chain length  Same chemical properties but slightly different physical  Example: 13.12 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 But-2-ene 13.13 Functional Isomers  Isomers have different functional groups, belong to different homologous series  Have different physical and chemical properties Ratio of C : H 1:3 1:2 1:2 Must have O2 Functional Gps. Alcohol & Ether Aldehyde & Ketone Carboxylic acid & Ester 13.15 Optical Isomers  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 chiral carbons 2 14. HYDROCARBONS 14.1 Properties Generally unreactive:  All C–C bonds single; alkanes = saturated hydrocarbons  Non-polar ∴ no center of charge to act as either nucleophile or electrophile ∴ cannot attract polar 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 14.2 Combustion Example C2H6O C3H6O C3H6O2 13.14 Geometric (cis/trans) Isomers  Shown only by alkenes  Arises due to restriction of double bond  Only possible when each carbon has 2 different groups  cis-trans isomers have different b.p.  cis isomers have higher dipole  trans isomer of symmetrical alkene has zero dipole  Used as fuel because they burn in oxygen to given out 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: PAGE 22 OF 30 4 → 2 14.3 Substitution 14.6 Alkenes  Alkanes react with halogens: and Example: Chlorination of Methane  Unsaturated hydrocarbons  Contain at least one C=C double bond  General formula: CnH2n (like cycloalkanes)  Source of alkenes: o Cracking alkanes o Dehydration of alcohols  More reactive than alkanes due to presence of double bond; pi electrons loosely and more susceptible to attacks by e- deficient groups like electrophiles  Alkenes combust completely  carbon dioxide + water  Give energy but not used as fuels; have other uses Reagent Condition Reaction Type Mechanism Substitution Free Radical Cl2(g) UV light  Initiation: o Energy of a photon of light absorbed o bond breaks homolytically  Propagation: o Highly reactive ∙ collides with a forming a new free radical; ∙ ⋅→ ⋅ ⋅ molecule → o This can then react with another and process repeats if sufficient present until all are replaced  Termination: o Reaction ends when 2 free radicals collide & combine ⋅ ⋅→ ⋅ ⋅→  Products: forms large amounts of CH3Cl and HCl and small amount C2H6; separated by fractional distillation  Products and free radicals differ due to: o Halogen used: bromine requires more light o Alkane used: ↑ no. of C = ↑ variety of products 14.4 Cracking  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 ⋅ 14.7 Electrophilic Addition Mechanism  Electrophile forms by heterolytic fission  Electrophile attacks double bond  Pair of e-s from double bond migrate to electrophile and bond breaks  Carbocation formed which attacks the nucleophile 14.8 Carbocations 14.5 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: o 2NO2 + 4CO  N2 + 4CO2 o 2NO + 2CO  N2 + 2CO2  Markovnikov’s principle: an electrophile adds to an unsymmetrical alkene so that the most stable carbocation is formed as an intermediate  Hydrogen binds to carbon with more hydrogens  Inductive effect of alkyl groups: o Alkyl groups donate e- to the ring o Producing a positive inductive effect o A larger alkyl group has a weaker inductive effect PAGE 23 OF 30 Hydro- CIE AS-LEVEL CHEMISTRY//9701 CIE AS-LEVEL CHEMISTRY//9701 Halogenation Hydrogenation 14.9 Addition Reactions Alkene + H2  Alkane  Reagent: H2(g)  Condition: o Catalyst: Nickel o Temp.: 100oC o Press.: 2 atm.  Use: convert liquid oils to saturated solid fats Alkene + X2  Dihaloalkane  Reagent: Halogen(aq)  Condition: r.t.p./dark Hydration Alkene + Hydrohalogen  Halogenoalkane  Reagent: Hydrohalogen(g)  Condition: r.t.p. Alkene + H2O(g)  Alcohol  Reagent: steam  Condition: o Catalyst: H3PO4 – phosphoric acid o Temp.: 300oC o Press.: 70atm 14.10 Oxidation of Alkenes  Both oxidation and addition to double bond involved  KMnO4 changes from pink to colourless With Cold Dil. Acidified KMnO4/H+  Diol is formed With Hot Conc. Acidified KMnO4/H+  Leads to the rupture of the double bond  Two compounds are formed  Products formed depend on alkene 14.11 Polymerization  Repeated addition of 1000s of alkene molecules (monomer) to each other forming a macromolecule  Polyethene: o LDPE: cling wrap o HDPE: water pipes, wire insulation  Polychloroethene (PVC): o Water pipes o Insulation of wires  General conditions: high pressure, high temperature and catalyst  Disadvantages: o Non-biodegradable: does not break down so increases amount of volume needed for landfill sites o Combustion produces harmful gases which contribute to global warming e.g. SO2, CO2 and HCl from PVCs  Disposal of Polymers: o Recycle existing plastic o Make polymers biodegradable by adding starch units PAGE 24 OF 30 CIE AS-LEVEL CHEMISTRY//9701 15.4 Nucleophilic Substitution Reaction 15. HALOGEN DERIVATIVES 15.1 Types of Halogenoalkanes Secondary 2o Tertiary 3o Hydrolysis Primary 1o SN2 SN1 Bond length increases, bond energy decreases, lower EA so more reactive 15.3 Nucleophilic Substitution Mechanism  The C – X bond is a polar bond, has partial charges due to high electro-tivity of halogen.  The carbocation is easily susceptible to attack by a nucleophile SN1 Mechanism: Nitrile (cyanide) Reactivity Primary Amines Electro-tivity decreases down group Bond Energy Increases  Fluoro Chloro Bromo Iodo Decrease  Polar Nature Decrease  15.2 Strength of C – Hal Bond R – X + OH-  R – OH + X Reagent: strong alkali; NaOH(aq) or KOH(aq)  Condition: heat/reflux  Fluoroalkanes are not hydrolysed because the C – F bond is too strong  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 R – X + CN-  RCN + X Reagent: KCN or NaCN in ethanol  Condition: o Solvent: Ethanol o Heat/Reflux  Reaction forms a C – C bond therefore no. of C increases; name has one more carbon R – X + NH3  RNH2(l) + HX(g)  Reagent: Ammonia (NH3)  Condition: ammonia in alcohol under pressure in sealed container  Note: if excess conc. ammonia used, HX reacts with it forming NH4X 15.5 Reflux  Many organic reactions proceed slowly  Heating done under reflux to prevent volatile organic solvents to evaporate  Mechanism similar to simple distillation  Unimolecular – only one molecule involved in 1st step  Tertiary halogenoalkanes SN2 Mechanism:  Bimolecular – two molecules involved in 1st step  Primary and secondary halogenoalkanes 15.6 Elimination Reaction R – X + OH-  Alkene + X- + H2O Mechanism:  Reagent: ethanolic NaOH or KOH  Conditions: temp. 60oC, reflux  OH- acts as a proton acceptor; it accepts the H+ loss from the halogenoalkanes during elimination  Elimination become progressively more easier Primary < Secondary < Tertiary  Note: the carbon atom adjacent to carbon with halide must have at least one hydrogen attached to it. PAGE 25 OF 30 CIE AS-LEVEL CHEMISTRY//9701 15.7 Uses of Halogenoalkanes 16.2 Properties  CFCs are inert and can be liquefied easily: Strength of C – X bond is very high, hence do not decompose easily and are not flammable.  Uses: o As propellants  in aerosol cans o As solvents  in dry-cleaning o As refrigerant  for freezers and fridges o Fire extinguishers, insecticides and pesticides Physical Properties:  Colourless liquids at r.t.p  b.p. and density increases with increasing C atoms and also with increasing OH groups Boiling Point: b.p. decreases  15.8 CFCs Effect on Ozone Layer 16. HYDROXY COMPOUNDS 16.1 Types of Alcohols Secondary 2o Secondary Tertiary 3o Tertiary because: branching increases  Causes the destruction of the ozone layer  CFCs escape in atmosphere and because of their inertness, remain without further reaction until they reach the stratosphere and ozone layer.  In stratosphere, high energy U.V causes Cl atom to split of 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  Note: alternative is using HCFCs (replace Cl with H or more F atoms) as they break down more easily and do not release Cl  less effect on ozone layer Primary 1o Primary VDWs decrease b.p. decreases  b.p. of alcohols > alkenes as they have hydrogen bonds Solubility of Alcohols in Water:  Smaller alcohols mix completely with water since strong hydrogen bonds occur between alcohols and water  As hydrocarbon nature increase (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 polar and non-polar compounds as they have polar and non-polar components 16.3 Reaction with Sodium R – OH + Na(l)  RO- Na+ + ½ H2(g)  Type of reaction: acid-base  Reagent used: liquid sodium metal  Reactivity of alcohols decreases with increasing chain lengths of hydrocarbon  Reaction less vigorous than that of Na and water which shows water is a stronger acid than alcohol 16.4 Reaction with Carboxylic Acids  Source of Alcohols: o Hydration of alkenes o Fermentation Reagent R-COOH PAGE 26 OF 30 Condition Heat-reflux Conc. H2SO4 Type of Reaction Esterification CIE AS-LEVEL CHEMISTRY//9701 Naming esters: Properties of Esters:  Esters are volatile compounds – no H-bonds so low m.p.  Polar molecules – soluble in organic solvents  Sweet, fruity smelling liquids  Many occur naturally e.g. as fats, oils & flavours in fruits  Used in food flavourings and perfumes and as solvents 16.5 Hydrolysis of Esters 16.6 Dehydration of Alcohols Alcohol(l)  Alkene + H2O(l) Condition Type of Reaction or Conc. H2SO4 H3PO4 at 180oC or Elimination Al2O3 at 300oC Mechanism:  Adjacent carbon to carbon with OH must have at least one hydrogen (tertiary cannot undergo dehydration) PAGE 27 OF 30 CIE AS-LEVEL CHEMISTRY//9701 16.7 Halogenation  Type of Reaction: Nucleophilic Substitution R – OH  R – X Producing: Alkyl Chlorides NaBr + H2SO4(aq)  HBr P + Br2 –warm PBr3 P + I2 –warm PI3 Alkyl Bromides Alkyl Iodide R – OH + Forming Reagent Reactions Conc. HCl  RCl(l) + H2O SOCl2  RCl(l) + SO2(g) + HCl(g) PCl5  RCl(l) + POCl3(aq) + HCl(g) PCl3  RCl(l) + H3PO3(aq) + HCl(g) HBr(g)  RBr(l) + H2O PBr3(g)  RBr(l) + H3PO3(aq) PI3(g)  RI(l) + H3PO3(aq) 16.8 Oxidation of Alcohols 17. CARBONYL COMPOUNDS  Reagent: Oxidising agents Reagent Type of Reaction Acidified K2Cr2O4 Acidified KMnO4 Oxidation Orange to Green Pink to Colourless Boiling Point: Primary Alcohol R-CH2-OH Secondary Alcohol R-CH(OH)-R Condition Zn + Heat/Reflux r.t.p Heat/Reflux r.t.p r.t.p Solubility:  Smaller carbonyl compounds: completely soluble as they form hydrogen bonds with water molecules; are good solvents for polar & non-polar solutes  Larger carbonyl compounds: polar nature decreases and non-polar nature increases; ability to form hydrogen bonds decreases 17.1 Nucleophilic Addition with HCN Reagent  If aldehyde needed, must distil as soon as it forms  Tertiary alcohols not oxidised because no hydrogens attached to carbon with OH group so oxidising agent colour does not change 16.9 Tests for Alcohols Reagent Result with: Secondary Bubble of H2 Gas Green Colourless Primary Na metal K2Cr2O4/H+ KMnO4/H+ Condition Type of Reaction HCN w/alkali or HCN Nucleophilic Addition HCN w/KCN  Since HCN added, carbon chain increases  Product formed is hydroxynitrile or cyanohydrine  Aldehydes are more susceptible to nucleophilic attacks than ketones  Smaller carbonyl compounds more reactive  Product has a chiral carbon ∴ exhibits optical isomerism Mechanism: Tertiary PAGE 28 OF 30 CIE AS-LEVEL CHEMISTRY//9701  Note: HCN is a poor nucleophile and with few CN- ions, the reaction is slow. To increase CN- conc.:  Make HCN react in presence of alkali HCN + OH-  H2O + CN Addition of KCN and dilute H2SO4 can provide HCN and more CN- ions 17.2 Reduction of Carbonyl Compounds  Type of Reaction: nucleophilic addition (H- ions)  Reducing agents: o NaBH4 – sodium tetrahydrioborate o LiAlH4 – lithium aluminium hydride o H2/Pt or Ni Aldehydes ⟹ 1o Alcohols Ketones ⟹ 2o Alcohols R-CHO + 2[H]  RCH2OH R-CO-R + 2[H]  R-CH(OH)-R 18. CARBOXYLIC ACIDS AND DERIVATIVES  Weak acids; don’t dissociate completely  Forms hydrogen bonds: o High m.p./b.p. o High solubility of smaller carboxylic acids  Forms hydrogen bonded dimers when pure vapour, liquid or solid & when dissolved in non-polar organic solvents 18.1 Formation of Carboxylic Acids  From alcohols: complete oxidation of primary alcohols  From aldehydes: oxidation of aldehydes  From nitriles: acid/base hydrolysis of a nitrile 17.3 Testing Carbonyl Compounds 2,4,- dinitrophenylhydrazine: 18.2 Formation of Salts  Heterolytic fission of the hydroxyl bond (-OH)  Salts called carboxylates  It is a nucleophilic addition & condensation/elimination  Forms: red/orange ppt.  The m.p. of the ppt. can be used to identify individual aldehydes and ketones Tests Given only by Aldehydes: Tollen’s Reagent Solution of AgNO3 + aq. NH3  excess [Ag(NH3)2]+ Aldehyde + Tollen’s Reagnet  Silver Mirror Ag+ reduced to Silver and –CHO oxidised to acid 2Ag+ + RCHO  2Ag + RCOOH- + H+ Fehling’s Solution CuSO4 in ammonia solution Aldehyde + Fehling’s Solution  Red ppt. Cu2+ reduced to Cu(I) oxide and –CHO oxidised to acid 2Cu2+ + RCHO  2Cu+ + RCOOH- + H+ 19. ANALYTICAL TECHNIQUES 19.1 Infra-red Spectroscopy  This is when a sample being analysed is irradiated with electromagnetic waves in the infra-red region of the electromagnetic spectrum.  Machine used is spectrophotometer and it detects intensity of wavelengths of infra-red that pass through the sample  The energy absorbed corresponds to changes in vibration of bonds leading to the bond being to stretch, bend and twist PAGE 29 OF 30 CIE AS-LEVEL CHEMISTRY//9701  At a specific frequency, the resonance frequency, the largest vibrations are obtained  Each type of vibration will absorb characteristic wavelengths of infra-red radiation  We can hence identify the presence (or absence) of different functional groups from the absorbance pattern on an infra-red spectrum 19.2 Monitoring Air Pollution  IR spectroscopy identifies particular bonds in a molecule, and so each pollutant will show a different pattern of absorptions – this allows the identification of the pollution  It is also possible to measure the concentration of each pollutant with the different amounts of absorption PAGE 30 OF 30