Chemistry AS Textbook

NES/Chemistry/AS Types of Catalyst A heterogeneous catalyst is in a different phase (state of matter) than the reagents. This is the most common type of catalyst in chemistry.  Example 1: The Haber process with a solid iron catalyst and gaseous nitrogen and hydrogen reagents. A homogeneous catalyst is in the same phase as the reagents. Most heterogeneous catalysts are enzymes.  Example 2: Protase is the enzyme that acts as a catalyst for the hydrolysis of protein. Enzymes are biological catalysts. They have specific features:  They are often more efficient that inorganic catalysts, speeding up reactions by a factor of 106  They are very specific and will usually only catalyse one reaction  They denature at temperatures above 400C  They usually work best at pH 7 150

NES/Chemistry/AS Topic 9 – The Periodic Table: Periodicity This topic illustrates the regular patterns in chemical and physical properties of the elements in the Periodic Table. 9.1 Periodicity of physical properties of the elements in the third period a) describe qualitatively (and indicate the periodicity in) the variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements (see the Data Booklet) b) explain qualitatively the variation in atomic radius and ionic radius c) interpret the variation in melting point and electrical conductivity in terms of the presence of simple molecular, giant molecular or metallic bonding in the elements d) explain the variation in first ionisation energy (see the Data Booklet) e) explain the strength, high melting point and electrical insulating properties of ceramics in terms of their giant structure; to include magnesium oxide, aluminium oxide and silicon dioxide 9.2 Periodicity of chemical properties of the elements in the third period a) describe the reactions, if any, of the elements with oxygen (to give Na2O, MgO, Al2O3, P4O10, SO2, SO3), chlorine (to give NaCl, MgCl2, Al2Cl6, SiCl4, PCl5) and water (Na and Mg only) b) state and explain the variation in oxidation number of the oxides (sodium to sulfur only) and chlorides (sodium to phosphorus only) in terms of their valence shell electrons c) describe the reactions of the oxides with water (treatment of peroxides and superoxides is not required) d) describe and explain the acid/base behaviour of oxides and hydroxides including, where relevant, amphoteric behaviour in reaction with acids and bases (sodium hydroxide only) e) describe and explain the reactions of the chlorides with water 151

NES/Chemistry/AS f) interpret the variations and trends in 9.2(b), (c), (d) and (e) in terms of bonding and electronegativity g) suggest the types of chemical bonding present in chlorides and oxides from observations of their chemical and physical properties 9.3 Chemical periodicity of other elements a) predict the characteristic properties of an element in a given Group by using knowledge of chemical periodicity b) deduce the nature, possible position in the Periodic Table and identity of unknown elements from given information about physical and chemical properties Ceramics are a class of material that is made of solid compounds. They are usually heat-resistant, corrosion-resistant, good insulators, hard and brittle. Coordination Number is how many particles surround a central particle in a crystal lattice – at IGCSE you learned that sodium chloride has a coordination number of 6. Physical Properties: Strong – can withstand large loads (weight) without deforming. Hard – resists shape change when a force is applied, such as scratch resistant. Brittle – breaks easily without deforming when subject to stress. 152

NES/Chemistry/AS 9.1 Periodicity – Physical Properties The Periodic Table is arranged in increasing atomic number, with each element in a Group having similar electronic configurations. Moving from left to right across a Period, the following physical properties change:  Atomic Radius  Ionic Radius  Melting Point  Electrical Conductivity  Ionisation Energy Period 3 Na Mg Al Si P S Cl Ar Element Structure giant metallic macromolecular simple molecular monatomic Bonding metallic covalent covalent / Atomic Decreases across the Period Radius Ionic Radius Two energy levels of electrons Three energy levels of electrons Decreases across the Period Decreases across the Period Melting low high high high low low low low Point Electrical high high high moderate low low low low Conductivity Ionisation Increases across the Period* Energy * For full details see Topic 2.3 153

NES/Chemistry/AS Atomic Radius atomic radius Na Mg Al Si P S Cl Ar The atomic radius decreases across Period 3 because the number of protons increases. The unusual value for Argon is due to how the atomic radius is measured, it is expected that Argon’s atomic radius is less than that of chlorine. Ionic Radius ionic radius Na Mg Al Si P S Cl Ar From sodium to silicon the elements give away valence electrons to form cations. This means that they will have only two energy levels of electrons. From phosphorus to chlorine the elements gain valence electrons to form anions. This means that they will have three energy levels of electrons making the anions significantly larger than the cations. 154

NES/Chemistry/AS Melting Point melting point Na Mg Al Si P S Cl Ar This pattern is more complex. Sodium to aluminium: The melting point increases as the number of valence electrons delocalised in the metallic lattice increases. There is a lower than expected increase from magnesium to aluminium. Sodium to magnesium: The increase is caused by an extra valence electron and the lattice structure changing from a coordination number of 8 to 12. Magnesium to aluminium: The increase is only caused by the extra valence electron. Both magnesium and aluminium have a coordination number of 12. Sodium Magnesium Aluminium 155

NES/Chemistry/AS Silicon Silicon has a giant covalent lattice, like diamond. Phosphorus to chlorine These are simple covalent molecules, so the melting point is dependent on the strength of van der Waals’ forces. Phosphorus Forms P4 molecules. Sulphur Forms S8 molecules. Chlorine Forms Cl2 molecules. Electrical Conductivity electrical conductivity Na Mg Al Si P S Cl Ar Sodium to aluminium: The electrical conductivity increases due to an increase in delocalised valence electrons. Silicon Silicon is a semi-conductor as it does have some delocalised electrons in its lattice structure. Phosphorus to chlorine These are poor conductors as they are simple covalent molecules without charged particles. 156

NES/Chemistry/AS First Ionisation Energy There is a general increase across the Period as the number of protons increases. There is a small drop after magnesium as the electron is being removed from the 3p orbital rather than the 3s orbital. There is a small drop after phosphorus as the electron being removed from sulphur has electron pair repulsion. See Topic 2.3 for more details. Ceramics Ceramics are a class of material with the following properties:  Strong  High Melting Point  Electrical Insulator This is due to their giant molecular structures. Common ceramics are made from:  Magnesium oxide  Aluminium oxide  Silicon dioxide 157

NES/Chemistry/AS 9.2 Periodicity – Chemical Properties This topic covers the chemical properties of:  Period 3 elements  Period 3 oxides  Period 3 chlorides Period 3 Elements Period 3 elements reacting with oxygen: Element Reaction Observation Bonding Na 4Na(s) + O2(g)  2Na2O(s) burns readily in air with an intense ionic yellow flame forming a white solid Mg 2Mg(s) + O2(g)  2MgO(s) burns with an intense white flame ionic on heating forming a white solid powdered aluminium burns with a Al 4Al(s) + 3O2(g)  2Al2O3(s) bright white flame forming a white ionic solid Si Si(s) + O2(g)  SiO2(s) slow reaction covalent white phosphorus catches fire spontaneously in air and burns P P4(s) + 5O2(g)  P4O10(s) with a bright white flame forming a white solid, but red phosphorus covalent burns on heating forming a white solid S(s) + O2(g)  SO2(g) burns with a blue flame in air forming a colourless gas. S covalent SO2(g) can then react under the 2SO2(g) + O2(g) ⇌ 2SO3(g) correct conditions forming sulphur trioxide Cl no reaction - - Ar 158

NES/Chemistry/AS Period 3 elements reacting with chlorine: Element Reaction Observation Bonding Na 2Na(s) + Cl2(g)  2NaCl(s) giant ionic reacts on heating with an giant ionic intense yellow flame forming covalent a white solid covalent covalent reacts on heating with an Mg Mg(s) + Cl2(g)  2MgCl2(s) intense white flame forming a white solid Al glows producing a pale Al 2Al(s) + 3Cl2(g)  2AlCl3(s) yellow solid (AlCl3 sublimes at 178oC to form the dimer Al2Cl6) Si Si(s) + 2Cl2(g)  SiCl4(l) reacts slowly on heating to form a colourless liquid in excess chlorine, burns with P P(s) + 5Cl2(g)  2PCl5(l) a white flame to give a colourless liquid Period 3 elements reacting with water: Element Reaction Observation Na 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g) violent reaction at room temperature Mg Mg(s) + 2H2O(l)  Mg(OH)2(aq) + H2(g) very slow fizzing with cold water Steam passed over heated Mg Mg(s) + H2O(g)  MgO(s) + H2(g) metal gives a vigorous exothermic reaction 159

NES/Chemistry/AS Period 3 Oxides Variation of Oxidation Number: Element Na Mg Al P S Na2O MgO Al2O3 P4O10 Oxide SO2 Oxidation +1 +2 +3 +5 SO3 Number +4 +6 Period 3 oxides reacting with water: Oxide Bonding Type of Reaction pH of Structure Oxide Na2O(s) + H2O(l)  2NaOH(aq) Solution MgO(s) + H2O(l)  Mg(OH)2(aq) Na2O giant ionic basic 14 MgO 9 Al2O3 giant ionic amphoteric No Reaction / Insoluble - with No Reaction / Insoluble - covalent 2 character 4 2 SiO2 giant weakly covalent acidic P4O10 covalent acidic P4O10(s) + 6H2O(l)  4H3PO4(aq) SO2 SO2(g) + H2O(l)  H2SO3(aq) SO3 simple SO3(g) + H2O(l)  H2SO4(aq) molecular 160

NES/Chemistry/AS Acid/Base behaviour of Period 3 oxides and hydroxides: Oxide/ Reaction Hydroxide Na2O(s) + 2HCl(aq)  2NaCl(aq) + H2O(l) Na2O NaOH NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l) MgO MgO(s) + 2HCl(aq)  MgCl2(aq) + H2O(l) Mg(OH)2 Mg(OH)2(aq) + 2HCl(aq)  MgCl2(aq) + 2H2O(l) Al2O3 Al2O3(s) + 6HCl(aq)  2AlCl3(aq) + 3H2O(l) Al2O3 With hot, concentrated NaOH: Al2O3(s) + 6NaOH(aq) + 3H2O(l)  2Na3Al(OH)6(aq) H3PO4(aq) The reaction takes place in 3 steps: H3PO4(aq) + NaOH(aq)  NaH2PO4(aq) + H2O(l) SO2 NaH2PO4(aq) + NaOH(aq)  Na2HPO4(aq) + H2O(l) SO3 Na2HPO4(aq) + NaOH(aq)  Na3PO4(aq) + H2O(l) SO2(g) + 2NaOH(aq)  Na2SO3(aq) + H2O(l) SO3(g) + 2NaOH(aq)  Na2SO4(aq) + H2O(l) 161

NES/Chemistry/AS Period 3 Chlorides Variation of Oxidation Number: Element Na Mg Al Si P S NaCl MgCl2 AlCl3 SiCl4 SCl2 Chloride PCl3 +2 +1 +2 +3 +4 PCl5 Oxidation Number +3 +5 Period 3 chlorides reacting with water: Chloride Bonding Reaction Observations pH Structure NaCl(s) + (aq)  NaCl(aq) MgCl2(s) + (aq)  MgCl2(aq) NaCl giant ionic white solid 7 MgCl2 dissolves to 6.5 form colourless solution Al2Cl6* See below 3 SiCl4 covalent SiCl4(l) + 2H2O(l)  SiO2(s) + 4HCl(g) chlorides react 2 PCl5 PCl5(s) + 4H2O(l)  H3PO4(aq) + 5HCl(g) with water 2 simple molecular giving off misty white fumes of HCl(g) SCl2 2SCl2(s) + 2H2O(l)  4HCl(g) +S(s) + SO2(g) 2 *Anhydrous aluminium chloride exists as a dimer, Al2Cl6 - when added to water aluminium chloride breaks down to form Al3+ ions. The Al3+ ion has a high charge density and displaces a H+ ion from water, hence forming an acidic solution. [Al(H2O)6]3+(aq)  [Al(H2O)5OH]3+(aq) + H+(aq) The aluminium ions are shown as hydrated with water ligands. The central metal ion has vacant orbitals and can form dative covalent bonds with water molecules – this is covered in more detail at A2 level. 162

NES/Chemistry/AS Topic 10 – The Periodic Table: Group 2 The physical and chemical properties of the elements of Group 2 (the alkaline Earth metals) are introduced in this topic. 10.1 Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their compounds a) describe the reactions of the elements with oxygen, water and dilute acids b) describe the behaviour of the oxides, hydroxides and carbonates with water and dilute acids c) describe the thermal decomposition of the nitrates and carbonates d) interpret, and make predictions from, the trends in physical and chemical properties of the elements and their compounds e) state the variation in the solubilities of the hydroxides and sulfates f)--A2 Only-- g)--A2 Only-- 10.2 Some uses of Group 2 compounds a) describe and explain the use of calcium hydroxide and calcium carbonate (powdered limestone) in agriculture 163

NES/Chemistry/AS 10.1 Group 2 Metals Elements in Group 2 of the Periodic Table are called the Alkaline Earth Metals. They all have valence electrons in s-orbitals and form ions with a 2+ charge. Trends in Physical Properties Going down the Group:  Atomic radius increases  Ionic radius increases  Density increases  Melting point decreases  Ionisation energy decreases These trends are all due to the increasing size of the atom/cation and the decreasing charge density of the cations. Trends in Chemical Properties Going down the Group:  More reactive  Reducing power increases Chemical Properties Group 2 metal reactions:  With oxygen  With water  With acids  Thermal decomposition of Group 2 compounds 164

NES/Chemistry/AS Group 2 elements reacting with oxygen: Group 2 metals burn in air to form solid, white oxides. The metals get more reactive going down the Group as the valence electrons are further away from the nucleus, with greater shielding. In fact, barium and radium are so reactive that they are stored in oil like Group 1 metals. Element Reaction with Oxygen type of Be 2Be(s) + O2(g)  2BeO(s) oxide amphoteric Mg, Ca, Sr, Ba 2M(s) + O2(g)  2MO(s) basic Where M is a Group 2 metal. Group 2 elements reacting with water: Element Reaction with Water Be no reaction with water or steam Mg Mg(s) + 2H2O(l)  Mg(OH)2(aq) + H2(g) Mg Mg(s) + H2O(g)  MgO(s) + H2(g) Ca, Sr, Ba M(s) + 2H2O(l)  M(OH)2(aq) + H2(g) Where M is a Group 2 metal. 165

NES/Chemistry/AS Group 2 elements and compounds reacting with acids: Metal Reaction with Acid M(s) + H2SO4(aq)  MSO4(aq) + H2(g) Oxide MO(s) + H2SO4(aq)  MSO4(aq) + H2O(l) Hydroxide M(OH)2(aq) + H2SO4(aq)  MSO4(aq) + H2O(l) Carbonate MCO3(s) + H2SO4(aq)  MSO4(aq) + H2O(l) + CO2(g) Where M is a Group 2 metal. Thermal decomposition of Group 2 compounds: Nitrate Reaction M(NO3)2(s)  MO(s) + 2NO2(g) + ½O2(g) Carbonate MCO3(s)  MO(s) + CO2(g) Where M is a Group 2 metal. The high charge density on the metal cation causes the electron cloud around the anion to distort. The greater the distortion, the easier it is to break down the compound, so the less stable the compound is. Thermal stability increases down the Group and the decomposition reactions take longer and a higher temperature. This is because going down the Group, the charge density decreases, so the distortion of the anion decreases. 166

NES/Chemistry/AS Solubility of Group 2 Hydroxides and Sulphates Hydroxides: None of the hydroxides are very soluble. However, the solubility of Group 2 hydroxides increases going down Group 2. As the solubility increases, the pH of the solution increases. Aqueous calcium hydroxide is also called limewater. Sulphates: Magnesium sulphate is soluble in water. The rest of the Group 2 sulphates are sparingly soluble in water. The solubility of Group 2 sulphates decreases going down Group 2. Barium sulphate is so insoluble that it is used in analysis, Ba2+(aq) ions are used to test for SO42-(aq) ions as they form a white precipitate when mixed together. 167

NES/Chemistry/AS 10. 2 Uses of Group 2 Compounds Calcium hydroxide (slaked lime):  Used to neutralise excess acidity in lakes and soil  Water purification  Glass manufacture  Neutralising acidic industrial waste products Calcium carbonate (limestone):  Used in the manufacture of iron and steel  Used in the manufacture of cement.  Used to neutralise excess acidity in lakes and soil 168

NES/Chemistry/AS Topic 11 – The Periodic Table: Group 17 The physical and chemical properties of the elements of Group 17 (the halogens) are introduced in this topic. At IGCSE this Group was numbered Group VII, however, at AS level it is numbered Group 17. 11.1 Physical properties of the Group 17 elements a) describe the colours and the trend in volatility of chlorine, bromine and iodine b) interpret the volatility of the elements in terms of van der Waals’ forces 11.2 The chemical properties of the elements and their hydrides a) describe the relative reactivity of the elements as oxidising agents (see also Section 6.3(f)) b) describe and explain the reactions of the elements with hydrogen c) (i) describe and explain the relative thermal stabilities of the hydrides (ii) interpret these relative stabilities in terms of bond energies 11.3 Some reactions of the halide ions a) describe and explain the reactions of halide ions with: (i) aqueous silver ions followed by aqueous ammonia (ii) concentrated sulfuric acid 11.4 The reactions of chlorine with aqueous sodium hydroxide a) describe and interpret, in terms of changes of oxidation number, the reaction of chlorine with cold and with hot aqueous sodium hydroxide 11.5 Some important uses of the halogens and of halogen compounds a) explain the use of chlorine in water purification b) state the industrial importance and environmental significance of the halogens and their compounds (e.g. for bleaches, PVC, halogenated hydrocarbons as solvents, refrigerants and in aerosols. See also Section 16.2) 169

NES/Chemistry/AS 11.1 Physical Properties of Group 17 Elements in Group 17 of the Periodic Table are called the Halogens. They all have valence electrons in p-orbitals and can form anions with a 1- charge. Chlorine, bromine and iodine can also form cations with a positive charge. Trends in Physical Properties Going down the Group:  Atomic radius increases  Ionic radius increases  Density increases  Colour gets darker Element Colour F pale yellow Cl green/yellow Br orange/brown I solid – grey/black gas - purple  Melting and boiling points increase Going down Group 17, the number of electrons in the atom increases, so the van der Waals’ forces of attraction increase and require more energy (heat) to overcome  Volatility decreases Volatility is when a liquid easily turns into a gas. This happens when the boiling point is only just above room temperature, so not much heating is required. As the van der Waals’ forces increase down the Group, the volatility decreases. 170

NES/Chemistry/AS 11.2 Chemical Properties Going down the Group:  Reactivity decreases  Oxidising power decreases When halogens react as oxidising agents, they gain an electron from another species to form a halide ion. The oxidising power decreases down the Group as electrons are being added to an energy level further away from the nucleus, with greater shielding. Group 2 elements reacting with hydrogen: Element Reaction with Hydrogen Observations F F2(g) + H2(g)  2HF(g) reacts explosively Cl Cl2(g) + H2(g)  2HCl(g) reacts explosively in Br Br2(l) + H2(g)  2HBr(g) sunlight I I2(s) + H2(g) ⇌ 2HI(g) reacts slowly on heating forms an equilibrium mixture on heating Stability of hydrogen halides: The stability of the hydrogen halides decreases down the Group. The ionic radius of the halide ion increases, so the bond gets longer and therefore weaker. The weaker the bond is, the easier it is to break and the less stable the hydrogen halide is. 171

NES/Chemistry/AS 11.3 Reactions of Halide Ions Group 17 halide reactions:  With aqueous silver ions, followed by aqueous ammonia  Concentrated sulphuric acid Group 17 halides reacting with silver: Halide Reaction with Ag+(aq) Colour of Precipitate Cl- Ag+(aq) + Cl-(aq)  AgCl(s) white cream Br- Ag+(aq) + Br-(aq)  AgBr(s) pale yellow I- Ag+(aq) + I-(aq)  AgI(s) Followed by aqueous ammonia: Precipitate Reaction with dilute NH3(aq) Reaction with concentrated NH3(aq) AgCl dissolves dissolves AgBr insoluble dissolves AgI insoluble insoluble 172

NES/Chemistry/AS Group 17 halide reacting with concentrated sulphuric acid: The halide ions act as reducing agents and get more powerful going down the group. You will see this effect illustrated in how far the sulphur in sulphuric acid is reduced. Halide Reaction NaCl NaCl(s) + H2SO4(l)  HCI(g) + NaHSO4(s) Sulphur 6+ to 6+ no reduction Halide Reactions NaBr NaBr(s) + H2SO4(l)  HBr(g) + NaHSO4(s) 2HBr(g) + H2SO4(I)  Br2(l) + 2H2O(l) + SO2(g) Sulphur 6+ to 4+ reduced by 2 Halide Reactions NaI Nal(s) + H2SO4(l)  HI(g) + NaHSO4(s) 2HI(g) + H2SO4(l)  I2(s) + SO2(g) + 2H2O(l) 6HI(g) + H2SO4(l)  3I2(s) + S(s) + 3H2O(l) 8HI(g) + H2SO4(l)  4I2(s) + H2S(g) + 4H2O(l) Sulphur 6+ to 2- reduced by 8 Observations: HCl(g) misty white fumes - true for all HX(g) S(s) yellow solid H2S(g) rotten egg smell I2(g) purple gas 173

NES/Chemistry/AS 11.4 The Reactions of Chlorine with Aqueous Sodium Hydroxide Chlorine reactions:  Cold aqueous sodium hydroxide (room temperature) Cl2(aq) + 2NaOH(aq)  NaClO(aq) + NaCI(aq) + H2O(l) The oxidation number of chlorine changes from 0 to 1- and 1+, this is a disproportionation reaction.  Hot aqueous sodium hydroxide (80oC) 3Cl2(aq) + 6NaOH(aq)  NaCIO3(aq) + 5NaCI(aq) + 3H2O(l) The oxidation number of chlorine changes from 0 to 5- and 1+, this is also a disproportionation reaction. The higher temperature allows chlorine to reach a higher oxidation number (5+). 174

NES/Chemistry/AS 11.5 Uses of Halogens and Halogen Compounds Uses of Chlorine:  Sterilise drinking water  Making bleach Uses of Halogens:  Making PVC  Making solvents  Making refrigerants  Making aerosols See also Topic 16.2 175

NES/Chemistry/AS Topic 13 – The Periodic Table: Nitrogen and Sulphur This topic introduces some of the chemistry associated with nitrogen and sulfur. 13.1 Nitrogen a) explain the lack of reactivity of nitrogen b) describe and explain: (i) the basicity of ammonia (see also Section 7.2) (ii) the structure of the ammonium ion and its formation by an acid-base reaction (iii) the displacement of ammonia from its salts c) state the industrial importance of ammonia and nitrogen compounds derived from ammonia d) state and explain the environmental consequences of the uncontrolled use of nitrate fertilisers e) state and explain the natural and man-made occurrences of oxides of nitrogen and their catalytic removal from the exhaust gases of internal combustion engines f) explain why atmospheric oxides of nitrogen are pollutants, including their catalytic role in the oxidation of atmospheric sulfur dioxide (see also Section 8.3(e)(iii)) 13.2 Sulfur: the formation of atmospheric sulfur dioxide, its role in acid rain a) describe the formation of atmospheric sulfur dioxide from the combustion of sulfur-contaminated fossil fuels b) state the role of sulfur dioxide in the formation of acid rain and describe the main environmental consequences of acid rain 176

NES/Chemistry/AS 13.1 Nitrogen Molecular nitrogen, N2, is relatively inert. It has a high activation energy because of the strong nitrogen – nitrogen triple bond. Reactions of Nitrogen: In the atmosphere nitrogen reacts with oxygen during thunderstorms. The lightning provides enough energy to overcome the high activation energy of nitrogen molecules. Step 1: N2(g) + O2(g)  2NO(g) Step 2: NO(g) + ½O2(g)  NO2(g) Step 3: 2NO2(g) + H2O(l) + ½O2(g)  2HNO3(aq) These reactions also happen from burning petrol due to the high temperature and pressure in a car engine. Oxides of nitrogen can be removed by a catalytic converter – the oxides of nitrogen are reduced to molecular nitrogen. Nitrogen also reacts with hydrogen in the Haber process to make ammonia (see also Topic 7.1), nitric acid, fertilisers and explosives. Ammonia Ammonia dissolves in water to form a weak alkaline solution. Aqueous ammonia is reacted with acids to make artificial fertilisers. However, over use of soluble nitrogen fertilisers can cause the eutrophication of rivers and lakes. This is where algae growth is promoted by the fertiliser and reduces levels of dissolved oxygen. 177

NES/Chemistry/AS 13.2 Sulphur Sulphur is an impurity in fossil fuels. When fossil fuels are burnt, the sulphur impurity also burns to form sulphur dioxide. S(s) + O2(g)  SO2(g) The sulphur dioxide can react further with oxygen and oxides of nitrogen in the atmosphere. SO2(g) + ½O2(g)  SO3(g) SO2(g) + NO2(g)  SO3(g) + NO(g) The sulphur trioxide can further react with water in the atmosphere. SO3(g) + H2O(l)  H2SO4(l) The sulphuric acid formed causes acid rain which:  Acidifies lakes and rivers  Corrodes limestone buildings  Harms trees and plants 178

NES/Chemistry/AS Topic 14 - An Introduction to Organic Chemistry Organic chemistry involves the study of a large group of chemical compounds containing carbon. This topic introduces naming conventions, organic reaction terminology and structures of organic molecules. 14.1 Formulae, functional groups and the naming of organic compounds a) interpret and use the general, structural, displayed and skeletal formulae of the following classes of compound: (i) alkanes and alkenes (ii) halogenoalkanes (iii) alcohols (including primary, secondary and tertiary) (iv) aldehydes and ketones (v) carboxylic acids and esters (vi) amines (primary only) and nitriles b) understand and use systematic nomenclature of simple aliphatic organic molecules with functional groups detailed in 14.1(a), up to six carbon atoms (six plus six for esters and amides, straight chains only) c) --A2 only-- d) deduce the possible isomers for an organic molecule of known molecular formula e) deduce the molecular formula of a compound, given its structural, displayed or skeletal formula 14.2 Characteristic organic reactions a) interpret and use the following terminology associated with organic reactions: (i) functional group (ii) homolytic and heterolytic fission 179

NES/Chemistry/AS (iii) free radical, initiation, propagation, termination (iv) nucleophile, electrophile (v) addition, substitution, elimination, hydrolysis, condensation (vi) oxidation and reduction 14.3 Shapes of organic molecules; σ and π bonds a) (i) describe and explain the shape of, and bond angles in, the ethane and ethene molecules in terms of σ and π bonds (ii) predict the shapes of, and bond angles in, other related molecules 14.4 Isomerism: structural and stereoisomerism a) describe structural isomerism and its division into chain, positional and functional group isomerism b) describe stereoisomerism and its division into geometrical (cis-trans) and optical isomerism c) describe cis-trans isomerism in alkenes, and explain its origin in terms of restricted rotation due to the presence of π bonds d) explain what is meant by a chiral centre and that such a centre normally gives rise to optical isomerism e) identify chiral centres and cis-trans isomerism in a molecule of given structural formula 180

NES/Chemistry/AS A Homologous Series is a group of chemicals with: the same general formula have the same functional group consecutive member differ by –CH2– and have a gradation in physical properties Structural Isomerism is said to occur when two or more compounds have the same molecular formula, but different structural formulae (have different physical or chemical properties). Stereoisomerism occurs when two or more compounds have the same molecular formula and the same structural formula, corresponding atoms being linked to the same atoms, but different spatial arrangements of their bonds. 181

NES/Chemistry/AS 14.1 Formula, Functional Groups and Naming Organic Compounds Types of Formula There are different types of formula used in organic chemistry and you will have to become familiar with each. 1. Molecular Formula This style of formula just gives the formula without showing any structure of the molecule.  Example 1: 1-chlorobut-2-ene - C4H7Cl Don't worry about naming molecules, it is coming later on in this topic. 2. Structural Formula This style of formula lists each carbon separately and also shows what atoms are bonded to the carbon.  Example 2: 1-chlorobut-2-ene - CH3CH=CHCH2Cl The = symbol is used to represent a double bond.  Example 3: 2-methylbutane - CH3CH(CH3)CH2CH3 or - (CH3)2CHCH2CH3 this style is much harder to see clearly. It might help to draw it as a display formula. The brackets are used to show side chains.  Example 4: pentane - CH3(CH2)3CH3 Here brackets are used to condense the formula, like factorising in maths. Only adjacent carbons can be condensed, which is why the two CH3 are not condensed. 182

NES/Chemistry/AS 3. Display Formula This style of formula shows every atom and every bond.  Example 5: 1-chloropent-2-ene H H H H Cl HCCCCCH HH H 4. Skeletal Formula This style of formula shows the carbon chain (and hydrogens) as a zig-zag line. It is an easier way of showing the display formula - once you get used to it.  Example 6: pentane  Example 7: 1-chloro-2-methylpent-2-ene Cl  Example 8: cyclobutane 183

NES/Chemistry/AS Functional Groups You will have to draw, identify and name the following functional groups:  Alkane carbon and hydrogen atoms carbon - carbon single bonds  Alkene carbon and hydrogen atoms at least one carbon = carbon double bond  Halogenoalkane carbon and hydrogen atoms at least one halogen atom  Alcohol (Primary, Secondary and Tertiary) carbon, hydrogen and oxygen atoms has at least one hydroxyl (OH) group  Aldehyde carbon, hydrogen and oxygen atoms has at least one CHO at the end of the carbon chain O C H  Ketone carbon, hydrogen and oxygen atoms has at least one C=O in the middle of the carbon chain O C 184

NES/Chemistry/AS  Carboxylic Acid carbon, hydrogen and oxygen atoms has at least one COOH at the end of the carbon chain O C OH Note - acids have their own properties and reactions, they are not a combination of alcohol and aldehyde.  Ester carbon, hydrogen and oxygen atoms has COO in the middle of the carbon chain O C O  Amine (Primary only) carbon, hydrogen and nitrogen atoms has NH2 at the end of the carbon chain H N H  Nitrile carbon, hydrogen and nitrogen atoms has CN at the end of the carbon chain CN 185

NES/Chemistry/AS Naming Organic Compounds There is a separate system used for naming organic molecules. The name consists of four main parts:  Principal Functional Group - this is the highest priority functional group and is used for the main name ending.  Main Chain Length - the longest chain of continuous carbons, including the principal functional group.  Substituents - other functional groups attached to the main chain.  Numbering - assigning numbers (locants) as to where on the chain the functional group and substituents are. Principal Functional Group The table is laid out in functional group priority, with the most important functional group at the top. So whichever functional group is the highest in the list forms the main part of the name. The other functional groups are usually then added to the front of the name as extra substituents. Functional Group Name if used Name as part as a of Main Name Carboxylic Acid Ester Substituent Nitrile carboxy- -oic acid Aldehyde Ketone alkyl alkanoate Alcohol Amine cyano- alkanenitrile Alkene Alkane oxo- -al oxo- -one hydroxy- -ol amino- alkylamine alkenyl- -en(e)* -an(e)* alkyl- *Alkane and alkene functional groups end in -ane or -ene if there is no other functional group present. If there is another functional group present the name is -an- (for alkane) or -en- (for alkene) with the second functional following. 186

NES/Chemistry/AS Main Chain Length This also forms part of the main name with the principal functional group. Name Number of Carbons Meth- 1 Eth- 2 Prop- 3 But- 4 Pent- 5 Hex- 6 Hept- 7 Oct- 8 Non- 9 Dec- 10 Substituents This part of the name always goes at the front of the name, even if it is the only functional group present. Functional Group Substituent Name Halogenoalkane Fluoro- Chloro- Bromo- Iodo- Numbering The position of a functional group is located by using numbers to count which carbon the functional group is on. Numbering can be done from the left, or the right to get the lowest number. Once you have decided which side to count from, you must use the same side for the entire name of the organic molecule. Numbers are separated from words using a dash (-) and numbers are separated from other numbers using a comma (,). 187

NES/Chemistry/AS Other Cyclo - this is used when a chain form a cyclic (round) chain. It comes at the very beginning of the name. di / tri / tetra / penta / hexa / hepta / octa - are used to show there is more than one of that functional group present. Primary, Secondary and Tertiary Carbon Atoms Carbons can also be classified as primary, secondary, or tertiary. This refers to the number of alkyl groups that are joined to the carbon atom. An alkyl group, is just a chain of carbons. Primary Carbon These are carbon atoms at the end of a carbon chain.  Example 9: Propan-1-ol CH3CH2CH2OH The OH group is attached to a carbon atom. That carbon atom is connected to one alkyl group (one chain). Secondary Carbon These are carbon atoms in the middle of a carbon chain.  Example 10: Propan-2-ol CH3CH(OH)CH3 The OH group is attached to a carbon atom. That carbon atom is connected to two alkyl groups (two chains). Tertiary Carbon These are carbon atoms in the middle of a carbon chain, with a side chain as well.  Example 11: 2-methylpropan-2-ol (CH3)3COH The OH group is attached to a carbon atom. That carbon atom is connected to three alkyl groups (three chains). It might be a good ideas to draw this display diagram of this molecule so you can see it clearly. 188

Naming Summary NES/Chemistry/AS Substituent Chain Functional Functional Group 1 Group 2 bromo meth butyl eth an(e) ol chloro prop en(e) one ethyl but al fluoro pent oic acid hydroxy hex iodo help methyl oct propyl non dec  Example 12: Naming CH3CH2CH3 propane The only functional group is an alkane and there are three carbons in the chain, so the name is propane.  Example 13: Naming CH3CH2CH2CH=CHCH2CH3 hept-3-ene The only functional group is an alkene and there are seven carbons in the chain. The double bond starts on carbon three (counting from the right-hand side) so the name is hept-3-ene.  Example 14: Naming CH2=CHCH2CH=CHCH3 hex-1,4-diene The only functional group is an alkene and there are six carbons in the chain. The double bonds start on carbon one and four (counting from the left-hand side) so the name is hex-1,4-diene. 189

NES/Chemistry/AS  Example 15: Naming CH3CH2CHO propanal There is an alkane and an aldehyde functional group. The alkane will be named -an- and the aldehyde will be named -al. There are three carbons in the chain so the name is propanal. There is no need to use a number as aldehydes are at the end of the chain and must be numbered 1.  Example 16: Naming CH3CH2COCH3 butanone There is an alkane and an ketone functional group. The alkane will be named -an- and the ketone will be named -one. There are four carbons in the chain so the name is butanone. There is no need to use a number as it must be 2. Longer ketones will require a number.  Example 17: Naming CH2=CHCH2CH2COCH3 hex-5-en-2-one There is an alkene and an ketone functional group. The alkene will be named -en- and the ketone will be named -one. There are six carbons in the chain so the name is hex-5-en-2-one. The molecule is numbered from the right-hand side as this gives the ketone (the more important functional group) the lowest number.  Example 18: Naming CH2ClCHFCF(CH3)CH3 The are chlorine, fluorine and methyl substituents on a four carbon main chain. The substituents are ordered alphabetically and the first substituent is given the lowest possible number, so the name is 1-chloro-2,3-difluoro-3-methylbutane.  Example 19: Naming CH3NH2 methylamine There is an alkane and amine functional group with one carbon in the main chain, so the name is methylamine. 190

NES/Chemistry/AS  Example 20: Naming CH3CH2CH(OH)COOH 2-hydroxybutanoic acid There is an alkane, alcohol and carboxylic acid. Normally the alcohol would be named at the end as the main functional group; however, the carboxylic acid is more important, so the alcohol becomes a substituent. There are four carbons in the main chain, so the name is 2-hydroxybutanoic acid. 191

NES/Chemistry/AS 14.2 Characteristic Organic Reactions Organic reactions depend on the functional group(s) present in the molecule. Each functional group is looked at in more detail in Topics 15-19. The types of reactions that occur are categorised into the following types:  homolytic and heterolytic fission  free radical, nucleophile, electrophile  addition, substitution, elimination, hydrolysis, condensation  oxidation and reduction During a chemical reaction, reactants break bonds. There are two main ways for this to happen, homolytic fission and heterolytic fission. Each type of bond breaking gives rise to a different mechanism for the reaction. Remember each bond consists of two electrons shared between two atoms. Homolytic Fission Each atom takes one electron from the bond, effectively making atoms again. This is how free radicals are made.  Example 21: When UV light shines on chlorine Cl Cl ● Cl + ● Cl Free Radicals These are made by homolytic fission. Free radicals are species which have a single unpaired electron. You will see free radicals in Topic 15.1 - Alkanes. This is the only functional group that we study that has free radical reactions. The mechanism for the reaction is split into 3 parts: 1. Initiation - the creation of two free radicals. 2. Propagation - one free radical as reagent, producing one free radical as product. 3. Termination - when two free radicals react together. 192

NES/Chemistry/AS Heterolytic Fission This is when the bond breaks, one species takes both electrons from the bond and the other species takes no electrons. This is how nucleophiles and electrophiles are made.  Example 22: When chlorine breaks to react with an alkene Cl Cl Cl+ + :Cl- Nucleophiles and Electrophiles These are made by heterolytic fission. The species that has both electrons from the bond usually also has a negative charge. This is called a nucleophile (attracted to other species with positive charges). The species that has no electrons from the bond will have a positive charge. This is called an electrophile (attracted to other species with a negative charge). Types of Reaction Reactions are also classified by what happens during the reaction. Type of Reaction What Happens Addition This is when there are two reagents that join together to form a single product Elimination This is when small molecules (such as H2O, or HCl) are removed from the organic molecule Condensation This is a type of elimination reaction, where water is specifically removed from the organic molecule Hydrolysis This is when water is added (with acid, or alkali) causing a larger organic molecule to break up into smaller parts Substitution This is when an atom is removed from an organic molecule and replaced by a different atom Oxidation The gain of oxygen, or loss of hydrogen from an organic molecule Reduction The loss of oxygen, or the gain of hydrogen from an organic molecule 193

NES/Chemistry/AS 14.3 Shapes of Organic Molecules; σ and π Bonds A quick review of bond angles and molecule shapes (Topic 3.2) will help for this section. Remember:  σ bonds are all single bonds formed by the end-to-end overlap of orbitals.  π bonds are the second bond in a double bond formed by the side-to-side overlap of p orbitals. σ Bonds - Ethane Ethane, like all alkanes forms single, covalent bonds between the atoms within the molecule. This results in each carbon atom having four single σ bonds. HH H C CH HH Molecule Shape = Tetrahedral Bond Angle = 109.50 π Bonds - Ethene Ethene has a double bond between the carbon atoms. This results in each carbon atom having three σ bonds and a π bond. The π bond is one of the two bonds in the double bond (the other bond in the double bond is a σ bond). H H C C H H Molecule Shape = Trigonal Planar Bond Angle = 1200 194

NES/Chemistry/AS 14.4 Isomerism: Structural and Stereoisomerism Structural Isomerism Stereoisomerism Chain Position Functional Geometric Optical Group The two main categories of isomerism:  Structural isomerism  Stereoisomerism Structural Isomerism These are molecules with the same molecular formula, but a different structure. It is split into: 1. Chain 2. Position 3. Functional Group 1. Chain Isomerism The molecules have a different structure because of side chains. The total number of carbon atoms stays the same, so the main chain gets shorter and some carbon atoms are attached as side chains.  Example 23: The chain isomers of pentane, C5H12 are: pentane 2-methylbutane 2,2-dimethylpropane 195

NES/Chemistry/AS 2. Position Isomerism The molecules have a different structure because the position of the functional group has move.  Example 24: The position isomers of chloropentane, C5H11Cl are: 1-chloropentane 2-chloropentane 3-chloropentane 3. Functional Group Isomerism The molecules have a different structure because there is a different functional group.  Example 25: The functional group isomers of C4H8 are: butene cyclobutane  Example 26: The functional group isomers of C3H6O are: propanal propanone Stereoisomerism These are molecules with the same molecular formula and the same structure. They have a different arrangement of bonds. It is split into: It is split into: 4. Geometric 5. Optical 196

NES/Chemistry/AS 4. Geometric Isomerism Geometric isomers (also called cis/trans) must have a carbon-carbon double bond, as well as different atoms (or groups of atoms) bonded to the carbons. As the double bond cannot rotate this gives rise to geometric isomers.  Example 23: The geometric isomers of 1,2-dichloroethene, C2H2Cl2 cis-1,2-dichloroethene HH CC Cl Cl trans-1,2-dichloroethene H Cl CC Cl H As the double bond cannot rotate, this is actually a different geometric isomer. 197

NES/Chemistry/AS When looking for geometric isomers only look for atoms, or groups of atoms above/below the double bond. Never look to the left/right of the double bond like the diagram below: 1,1-dichloroethene, C2H2Cl2 H Cl CC H Cl This is not a geometric isomer. It is actually a position isomer, as the chlorine has move from position 2 to position 1. Cis - means on the same side. Trans - means on the opposite side. 5. Optical Isomerism Optical isomers contain a carbon atom that is bonded to four different atoms, or groups of atoms. This makes the molecule have mirror images that are non-super imposable. The isomers are simply identified as Left and Right. A chiral centre is the asymmetric carbon atom that has four different functional groups.  Example 27: 2-chlorobutane chiral centres mirror plane C2H5 C2H5 C C H H3C H CH3 Cl Cl Left Right 198

NES/Chemistry/AS  Example 28: Testosterone has 6 chiral centres CH2OH CO H3C  OH O H3C     O 199