Chemistry AS Textbook

AS Chemistry Cambridge International Examinations 2016 Syllabus

NES/Chemistry/AS Cambridge International Examinations AS Chemistry Syllabus Code 9701 2016 Syllabus Update Paper Type of Paper Duration Marks Weighting AS A2 1 Multiple Choice 1 Hour 40 31% 15.5% 46% 23% 2 Written paper 1 Hour and 15 Minutes 60 23% 11.5% 3 Practical 2 Hours 40 A Gru and minion production. 1

NES/Chemistry/AS Data Booklet You will get a copy of the data booklet for Paper 1 and Paper 2 (AS), as well as Paper 4 (A2). It also includes a copy of the AS/A2 Periodic Table. Contents: Tables of Chemical Data Page 1 Important values, constants and standards 3 2 Ionisation energies of selected elements in kJ mol–1 4 3 Bond energies 5 4 Standard electrode potential and redox potentials, E at 298K (25 °C) 7 5 Atomic and ionic radii 10 6 Typical proton (1H) chemical shift values (δ) relative to TMS = 0 12 7 Typical carbon (13C) chemical shift values (δ) relative to TMS = 0 13 8 Characteristic infra-red absorption frequencies for some selected bonds 14 9 The orientating effect of groups in aromatic substitution reactions 15 10 Names, structures and abbreviations of some amino acids 16 11 The Periodic Table of Elements 17 2

NES/Chemistry/AS 1 Important values, constants and standards molar gas constant R = 8.31 J K–1 mol–1 the Faraday constant F = 9.65 × 104 C mol–1 the Avogadro constant L = 6.02 × 1023 mol–1 the Planck constant h = 6.63 × 10–34 J s speed of light in a vacuum c = 3.00 × 108 m s–1 rest mass of proton, 1H mp = 1.67 × 10–27 kg 1 rest mass of neutron, 1n mn = 1.67 × 10–27 kg 0 rest mass of electron, 0e me = 9.11 × 10–31 kg −1 electronic charge e = –1.60 × 10–19 C molar volume of gas Vm = 22.4 dm3 mol–1 at s.t.p. Vm = 24.0 dm3 mol–1 under room conditions (where s.t.p. is expressed as 101 kPa, approximately, and 273 K [0 °C]) ionic product of water Kw = 1.00 × 10–14 mol2 dm–6 (at 298 K [25 °C]) specific heat capacity of water = 4.18 kJ kg–1 K–1 (= 4.18 J g–1 K–1) 3

NES/Chemistry/AS 2 Ionisation energies (1st, 2nd, 3rd and 4th) of selected elements, in kJ mol–1 Proton number First Second Third Fourth H 1 1310 – – – He 2 2370 5250 – – Li 3 519 7300 11800 – Be 4 900 1760 14800 21000 B 5 799 2420 3660 25000 C 6 1090 2350 4610 6220 N 7 1400 2860 4590 7480 O 8 1310 3390 5320 7450 F 9 1680 3370 6040 8410 Ne 10 2080 3950 6150 9290 Na 11 494 4560 6940 9540 Mg 12 736 1450 7740 10500 Al 13 577 1820 2740 11600 Si 14 786 1580 3230 4360 P 15 1060 1900 2920 4960 S 16 1000 2260 3390 4540 Cl 17 1260 2300 3850 5150 Ar 18 1520 2660 3950 5770 K 19 418 3070 4600 5860 Ca 20 590 1150 4940 6480 Sc 21 632 1240 2390 7110 Ti 22 661 1310 2720 4170 V 23 648 1370 2870 4600 Cr 24 653 1590 2990 4770 Mn 25 716 1510 3250 5190 Fe 26 762 1560 2960 5400 Co 27 757 1640 3230 5100 Ni 28 736 1750 3390 5400 Cu 29 745 1960 3350 5690 Zn 30 908 1730 3828 5980 Ga 31 577 1980 2960 6190 Br 35 1140 2080 3460 4850 Rb 37 403 4632 3900 5080 Sr 38 548 1060 4120 5440 Ag 47 731 2074 3361 – I 53 1010 1840 2040 4030 Cs 55 376 2420 3300 – Ba 56 502 966 3390 – 4

NES/Chemistry/AS 3 Bond Energies 3(a) Bond energies in diatomic molecules (these are exact values) Homonuclear Heteronuclear Bond Energy / kJ mol–1 Bond Energy / kJ mol–1 HH 436 HF 562 DD 442 HCl 431 N≡N 944 HBr 366 O=O 496 HI 299 P≡P 485 C≡O 1077 S=S 425 FF 158 Cl Cl 242 BrBr 193 II 151 5

NES/Chemistry/AS 3(b) Bond energies in polyatomic molecules (these are average values) Homonuclear Heteronuclear Bond Energy / kJ mol–1 Bond Energy / kJ mol–1 CC 350 CH 410 C=C 610 CCl 340 C≡C 840 CBr 280 C….C (benzene) 520 CI 240 NN 160 CN 305 N=N 410 C=N 610 OO 150 C≡N 890 SiSi 222 CO 360 PP 200 C=O 740 SS 264 C=O in CO2 805 NH 390 6 NCl 310 OH 460 SiCl 359 SiH 320 SiO (in SiO2(s)) 460 Si=O (in SiO2(g)) 640 PH 320 PCl 330 PO 340 P=O 540 SH 347 SCl 250 SO 360 S=O 500

NES/Chemistry/AS 4 Standard electrode potential and redox potentials, E at 298 K (25 oC) For ease of reference, two tables are given: (a) an extended list in alphabetical order; (b) a shorter list in decreasing order of magnitude, i.e. a redox series. (a) E in alphabetical order Electrode reaction E /V Ag+ + e– Ag +0.80 –1.66 Al 3+ + 3e– Al –2.90 Ba2+ + 2e– +1.07 Ba –2.87 Br2 + 2e– 2Br– +1.36 Ca2+ + 2e– +1.64 Ca +0.89 Cl2 + 2e– 2Cl – 2HOCl + 2H+ + 2e– –0.28 Cl2 + 2H2O +1.82 Cl O– + H2O + 2e– Cl – + 2OH– –0.43 –0.91 Co2+ + 2e– Co –0.74 Co2+ –0.41 Co3+ + e– +1.33 Co + 6NH3 +0.52 [Co(NH3)6]2+ + 2e– Cr +0.34 Cr2+ + 2e– +0.15 Cr –0.05 Cr3+ + 3e– Cr2+ +2.87 2Cr3+ + 7H2O –0.44 Cr3+ + e– Cu –0.04 +0.77 Cr2O72– + 14H+ + 6e– Cu +0.36 Cu+ + e– Cu+ –0.56 Cu2+ + 2e– Cu + 4NH3 0.00 2F– –0.83 Cu2+ + e– Fe +0.54 [Cu(NH3)4]2+ + 2e– –2.92 F2 + 2e– Fe –3.04 Fe2+ –2.38 Fe2+ + 2e– [Fe(CN)6]4– –1.18 Fe(OH)2 + OH– +1.49 Fe3+ + 3e– H2 +1.23 H2 + 2OH– +0.56 Fe3+ + e– 2I– +1.67 +1.52 [Fe(CN)6]3– + e– K +0.81 Fe(OH)3 + e– +0.94 2H+ + 2e– Li +0.87 2H2O + 2e– Mg I2 + 2e– Mn K+ + e– Mn2+ Mn2+ + 2H2O Li+ + e– MnO42– MnO2 + 2H2O Mg2+ + 2e– Mn2+ + 4H2O NO2 + H2O Mn2+ + 2e– HNO2 + H2O NH4+ + 3H2O Mn3+ + e– 7 MnO2 + 4H+ + 2e– MnO4– + e– MnO4– + 4H+ + 3e– MnO4– + 8H+ + 5e– NO3– + 2H+ + e– NO3– + 3H+ + 2e– NO3– + 10H+ + 8e–

NES/Chemistry/AS Electrode reaction E /V Na+ + e– Na –2.71 –0.25 Ni2+ + 2e– Ni –0.51 +1.77 [Ni(NH3)6]2+ + 2e– Ni + 6NH3 +0.88 H2O2 + 2H+ + 2e– +1.23 2H2O +0.40 HO2– + H2O + 2e– 3OH– +0.68 O2 + 4H+ + 4e– –0.08 2H2O –0.13 O2 + 2H2O + 4e– 4OH– +1.69 O2 + 2H+ + 2e– +1.47 H2O2 +0.17 O2 + H2O + 2e– HO2– + OH– +2.01 Pb2+ + 2e– +0.09 Pb –0.14 Pb4+ + 2e– Pb2+ +0.15 Pb2+ + 2H2O –1.20 PbO2 + 4H+ + 2e– –0.26 SO42– + 4H+ + 2e– SO2 + 2H2O +0.34 2SO42– +1.00 S2O82–+ 2e– 2S2O32– +1.00 S4O62–+ 2e– –0.76 Sn Sn2+ + 2e– Sn2+ Sn4+ + 2e– V V2+ V2+ + 2e– V3+ + H2O VO2+ + H2O V3+ + e– VO2+ + 2H2O VO2+ + 2H+ + e– Zn VO2+ + 2H+ + e– VO3– + 4H+ + e– Zn2+ + 2e– All ionic states refer to aqueous ions but other state symbols have been omitted. 8

NES/Chemistry/AS (b) E in decreasing order of oxidising power (a selection only – see also the extended alphabetical list on the previous pages) Electrode reaction E /V F2 + 2e– 2F– +2.87 S2O82–+ 2e– 2SO42– +2.01 H2O2 + 2H+ + 2e– 2H2O +1.77 MnO4– + 8H+ + 5e– Mn2+ + 4H2O +1.52 PbO2 + 4H+ + 2e– Pb2+ + 2H2O +1.47 2Cl – +1.36 Cl2 + 2e– 2Cr3+ + 7H2O +1.33 Cr2O72– + 14H+ + 6e– 2H2O +1.23 2Br– +1.07 O2 + 4H+ + 4e– Cl – + 2OH– +0.89 Br2 + 2e– NH4+ + 3H2O +0.87 ClO – + H2O + 2e– NO2 + H2O Ag +0.81 NO3– + 10H+ + 8e– Fe2+ +0.80 2I– +0.77 NO3– + 2H+ + e– 4OH– +0.54 Ag+ + e– +0.40 Cu +0.34 Fe3+ + e– +0.17 SO2 + 2H2O +0.15 I2 + 2e– Sn2+ +0.09 O2 + 2H2O + 4e– 2S2O32– H2 0.00 Cu2+ + 2e– Pb –0.13 –0.14 SO42– + 4H+ + 2e– Sn –0.44 Sn4+ + 2e– –0.76 Fe –0.83 S4O62–+ 2e– 2H+ + 2e– Zn –1.20 H2 + 2OH– Pb2+ + 2e– –2.38 V –2.87 Sn2+ + 2e– –2.92 Mg Fe2+ + 2e– Ca Zn2+ + 2e– K 2H2O + 2e– V2+ + 2e– Mg2+ + 2e– Ca2+ + 2e– K+ + e– 9

5 Atomic and ionic radii NES/Chemistry/AS (a) Period 1 atomic / nm ionic / nm C4– 0.260 single covalent H 0.037 H– 0.208 N3– 0.171 van der Waals He 0.140 O2– 0.140 atomic / nm ionic / nm F– 0.136 (b) Period 2 Li 0.152 Li+ 0.060 metallic Be 0.112 Be2+ 0.031 P3– 0.212 single covalent B 0.080 B3+ 0.020 S2– 0.184 C 0.077 C4+ 0.015 Cl – 0.181 van der Waals N 0.074 (c) Period 3 O 0.073 ionic / nm F 0.072 Na+ 0.095 metallic Ne 0.160 Mg2+ 0.065 atomic / nm Al 3+ 0.050 single covalent Na 0.186 Si4+ 0.041 Mg 0.160 van der Waals Al 0.143 ionic / nm (d) Group 2 Si 0.117 Be2+ 0.031 P 0.110 Mg2+ 0.065 metallic S 0.104 Ca2+ 0.099 Cl 0.099 Sr2+ 0.113 Ar 0.190 Ba2+ 0.135 atomic / nm Ra2+ 0.140 Be 0.112 Mg 0.160 Ca 0.197 Sr 0.215 Ba 0.217 Ra 0.220 10

NES/Chemistry/AS (e) Group 14 atomic / nm ionic / nm C4+ 0.015 single covalent C 0.077 Si4+ 0.041 Ge2+ 0.093 Si 0.117 Sn2+ 0.112 Pb2+ 0.120 Ge 0.122 ionic / nm F– 0.136 metallic Sn 0.162 Cl – 0.181 Br– 0.195 Pb 0.175 I– 0.216 (f) Group 17 atomic / nm single covalent F 0.072 Cl 0.099 Br 0.114 I 0.133 At 0.140 (g) First row transition atomic / nm ionic / nm elements Ti2+ 0.090 metallic Sc 0.164 V2+ 0.079 Sc3+ 0.081 Cr2+ 0.073 Ti3+ 0.067 Ti 0.146 Mn2+ 0.067 V3+ 0.064 Fe2+ 0.061 Cr3+ 0.062 V 0.135 Co2+ 0.078 Mn3+ 0.062 Ni2+ 0.070 Fe3+ 0.055 Cr 0.129 Cu2+ 0.073 Co2+ 0.053 Zn2+ 0.075 Ni3+ 0.056 Mn 0.132 Fe 0.126 Co 0.125 Ni 0.124 Cu 0.128 Zn 0.135 11

NES/Chemistry/AS 6 Typical proton (1H) chemical shift values (δ) relative to TMS = 0 type of environment of proton example structures chemical shift proton –CH3, –CH2–, >CH– range (δ) alkane CH3–C=O, –CH2–C=O, >CH–C=O 0.9–1.7 C–H alkyl next to C=O CH3–Ar, –CH2–Ar, >CH–Ar 2.2–3.0 alkyl next to aromatic CH3–O, –CH2–O, –CH2–Cl, >CH–Br 2.3–3.0 ring ≡C–H alkyl next to =CH2, =CH– 3.2–4.0 electronegative atom attached to alkyne 1.8–3.1 attached to alkene 4.5–6.0 attached to aromatic ring H 6.0–9.0 aldehyde O 9.3–10.5 alcohol RC 0.5–6.0 H RO–H O-H phenol OH 4.5–7.0 (see carboxylic acid 9.0–13.0 note O 1.0–5.0 below) alkyl amine RC OH R–NH– N-H aryl amine NH2 3.0–6.0 (see amide 5.0–12.0 note O below) RC NH Note: δ values for –O-H and –N-H protons can vary depending on solvent and concentration 12

7 Typical carbon (13C) chemical shift values (δ) relative to TMS = 0 NES/Chemistry/AS hybridisation of environment of carbon example structures chemical the carbon atom CH3–, –CH2–, –CH< shift range atom alkyl (δ) sp3 0–50 sp3 next to alkene/arene CH2 10–40 25–50 –CH2–C=C, 30–65 sp3 next to carbonyl/carboxyl –CH2–COR, –CH2–CO2R 30–60 50–70 sp3 next to nitrogen –CH2–NH2, –CH2–NR2, 110–160 –CH2–NHCO 160–185 next to chlorine 190–220 sp3 (-CH2-Br and -CH2-I are in –CH2–Cl 65–85 the same range as alkyl) 100–125 sp3 next to oxygen –CH2–OH, –CH2–O–CO– CC sp2 alkene or arene CC >C=C<, CC sp2 carboxyl R–CO2H, R–CO2R sp2 carbonyl R–CHO, R–CO–R sp alkyne R–C≡C– sp nitrile R–C≡N 13

NES/Chemistry/AS 8 Characteristic infra-red absorption frequencies for some selected bonds bond functional groups absorption range (in appearance of peak containing the bond wavenumbers) /cm–1 (s = strong, w = weak) C–O alcohols, ethers, esters 1040–1300 s C=C aromatic compounds, 1500–1680 w unless conjugated alkenes amides, 1640–1690 s C=O ketones and aldehydes 1670–1740 s 1710–1750 s esters, 2150–2250 w unless conjugated C≡C alkynes 2200–2250 w C≡N nitriles 2850–2950 s C–H alkanes, CH2–H 3000–3100 w alkenes/arenes, =C–H 3300–3500 w N–H amines, amides 2500–3000 s and very broad O–H carboxylic acids, RCO2–H 3200–3600 s H–bonded alcohol, RO–H 3580–3650 s and sharp free alcohol, RO–H 14

NES/Chemistry/AS 9 The orientating effect of groups in aromatic substitution reactions. The position of the incoming group, Y, is determined by the nature of the group, X, already bonded to the ring, and not by the nature of the incoming group Y. XX 2 + H+ + Y+ Y 3 X- groups that direct the incoming 4 Y group to the 3- position –NO2 X- groups that direct the incoming –NH3 Y group to the 2- or 4- positions –CN –CHO, –COR –NH2, –NHR or –NR2 –CO2H, –CO2R –OH or –OR –NHCOR –CH3, –alkyl –Cl 15

NES/Chemistry/AS 10 Names, structures and abbreviations of some amino acids structure of side chain R- in name 3-letter abbreviation 1-letter symbol NH2 R CH alanine aspartic acid CO2H cysteine glutamic acid Ala A CH3– glycine lysine Asp D HO2CCH2– Cys C HSCH2– Glu E HO2CCH2CH2– Gly G H– Lys K H2NCH2CH2CH2CH2– phenylalanine Phe F CH2 serine Ser S HOCH2– tyrosine Tyr Y HO CH2 CH3 valine Val V CH CH3 16

The Periodic Table of Elements Group 1 2 13 14 15 16 17 18 3 Key 1 9 2 Li 4 atomic number H 5 6 7 8 F He lithium Be atomic symbol hydrogen B C N O fluorine helium 6.9 beryllium name 1.0 boron carbon nitrogen oxygen 19.0 4.0 11 17 9.0 relative atomic mass 8 10.8 12.0 14.0 16.0 10 Na Cl 12 26 13 14 15 16 Ne sodium chlorine Mg Fe Al Si P S neon 23.0 35.5 19 magnesium 3 4 5 6 7 iron 9 10 11 12 aluminium silicon phosphorus sulfur 35 20.2 K 24.3 55.8 27.0 28.1 31.0 32.1 Br 18 44 potassium 20 21 22 23 24 25 27 28 29 30 31 32 33 34 bromine Ar Ru 39.1 Ca Sc Ti V Cr Mn Co Ni Cu Zn Ga Ge As Se 79.9 argon 37 ruthenium 53 calcium scandium titanium vanadium chromium manganese cobalt nickel copper zinc gallium germanium arsenic selenium 39.9 Rb 101.1 I 40.1 45.0 47.9 50.9 52.0 54.9 76 58.9 58.7 63.5 65.4 69.7 72.6 74.9 79.0 36 rubidium iodine 38 39 40 41 42 43 Os 45 46 47 48 49 50 51 52 Kr 85.5 126.9 55 Sr Y Zr Nb Mo Tc osmium Rh Pd Ag Cd In Sn Sb Te 85 krypton Cs strontium yttrium zirconium niobium molybdenum technetium 190.2 rhodium palladium silver cadmium indium tin antimony tellurium At 83.8 108 caesium 87.6 88.9 91.2 92.9 95.9 – 102.9 106.4 107.9 112.4 114.8 116.7 121.8 127.6 astatine 54 NES/Chemistry/ASHs 17132.9 56 57–71 72 73 74 75 77 78 79 80 81 82 83 84 – Xe 87 hassium Ba lanthanoids Hf Ta W Re Ir Pt Au Hg Tl Pb Bi Po xenon Fr – barium hafnium tantalum tungsten rhenium iridium platinum gold mercury thallium lead bismuth polonium 131.3 francium 137.3 178.5 180.9 183.8 186.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 – 86 – Rn radon – 88 89–103 104 105 106 107 109 110 111 112 114 116 Ra actinoids Rf Db Sg Bh Mt Ds Rg Cr Fl Lv radium rutherfordium dubnium seaborgium bohrium meitnerium darmstadtium roentgenium copernicium flerovium livermorium – – – – – –––– – – lanthanoids 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 actinoids La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu lanthanum cerium praseodymium neodymium promethium samarium europium gadolinium terbium dysprosium holmium erbium thulium ytterbium lutetium 138.9 140.1 140.9 144.4 – 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.1 175.0 89 90 91 92 93 94 95 96 97 98 99 100 101 102 103 Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr actinium thorium protactinium uranium neptunium plutonium americium curium berkelium californium einsteinium fermium mendelevium nobelium lawrencium – 232.0 231.0 238.0 – – – – – – – – – – –

NES/Chemistry/AS Table of Contents Section Topic Page Physical 1 Atoms, Molecules and Stoichiometry 19 Chemistry 2 Atomic Structure 31 3 Chemical Bonding 47 Inorganic 4 States of Matter 79 Chemistry 5 Chemical Energetics 96 6 Electrochemistry 109 Organic 7 Equilibria 115 Chemistry and 8 Reaction Kinetics 143 9 The Periodic Table (Periodicity) 151 Analysis 10 Group 2 163 11 Group 17 169 Practical 12 Transition Metals (A2 only) - 13 Nitrogen and Sulfur 176 14 Introduction to organic Chemistry 179 15 Hydrocarbons 202 16 Halogen Derivatives 220 17 Hydroxy Compounds 229 18 Carbonyl compounds 239 19 Carboxylic Acids 250 20 Nitrogen Compounds (A2 only) - 21 Polymerisation (A2 only) - 22 Analytical Techniques 255 23 Organic Synthesis (A2 only) - 24 Practical 258 18

NES/Chemistry/AS Topic 1 - Atoms, Molecules and Stoichiometry This topic illustrates how quantitative relationships can be established when different substances react. (The term relative formula mass or Mr will be used for all compounds including ionic.) 1.1 Relative masses of atoms and molecules a) define and use the terms relative atomic, isotopic, molecular and formula masses, based on the 12C scale 1.2 The mole and the Avogadro constant a) define and use the term mole in terms of the Avogadro constant 1.3 The determination of relative atomic masses, Ar a) analyse mass spectra in terms of isotopic abundances (knowledge of the working of the mass spectrometer is not required) b) calculate the relative atomic mass of an element given the relative abundances of its isotopes, or its mass spectrum 1.4 The calculation of empirical and molecular formulae a) define and use the terms empirical and molecular formula b) calculate empirical and molecular formulae, using combustion data or composition by mass 1.5 Reacting masses and volumes (of solutions and gases) a) write and construct balanced equations b) perform calculations, including use of the mole concept, involving: (i) reacting masses (from formulae and equations) (ii) volumes of gases (e.g. in the burning of hydrocarbons) (iii) volumes and concentrations of solutions 19

NES/Chemistry/AS c) deduce stoichiometric relationships from calculations such as those in 1.5(b) The relative atomic mass (Ar) of an element is the weighted average of the masses of its isotopes relative to 1/12 of the mass of a carbon-12 atom. -This is for monatomic elements. The relative molecular mass (Mr) is the weighted average of the masses of the molecules relative to 1/12 of the mass of a carbon-12 atom. -This is for covalent molecules. The relative isotopic mass is the mass of the isotope relative to 1/12 of the mass of a carbon-12 atom. -This is for isotopes of an atom. The relative formula mass (Mr) is the weighted average of the masses of the formula units relative to 1/12 of the mass of a carbon-12 atom. -This is for any compound (ionic, or covalent). A mole is the amount of substance which contains the same number of particles as there are atoms in exactly 12g of carbon-12. Avogadro's Constant is the number of atoms in one mole of 12C. The limiting reagent in a particular experiment is the reactant that governs the maximum amount of product that can be formed. -The smallest number of moles on the reagent side, using the mole ratio. The percentage yield of a product is the percentage of its theoretical yield achieved in practice. -This is how much product is made in an incomplete reaction, compared to how much could be made if the reaction was 100% complete. The percentage purity of a reactant is the actual mass reacted divided by the initial mass used expressed as a percentage. -This is how pure a reagent is. Empirical formula shows the simplest whole number ratio of atoms present in a substance. Molecular formula shows the actual number of atoms of each element present in a substance. 20

NES/Chemistry/AS 1.1 Relative Masses of Atoms and Molecules This is the same as your IGCSE moles. The only real difference is that the questions are more complex, especially on Paper 1. When performing calculations, candidates’ answers should reflect the number of significant figures given or asked for in the question. When rounding up or down, candidates should ensure that significant figures are neither lost unnecessarily nor used beyond what is justified Relative Atomic/Isotopic Mass (Ar) The relative atomic mass (Ar), in grams, of any element contains 1 mole of atoms. Ar is the same as the element's mass number and can be found on the Periodic Table and is quoted to 1 decimal place. It has the units gmol-1, which just means how many grams there are in 1 mole. Different isotopes of an element will have different Ar values as the number of neutrons in the nucleus will vary.  Example 1: Lithium has two different isotopes 6Li and 7Li. Relative Molecular/Formula Mass (Mr) The relative molecular mass, or relative formula mass (Mr) in grams, of any compound contains 1 mole of molecules. Mr is worked out by adding the Ar values for all the elements in a molecule, or compound. It also has the units gmol-1.  Example 2: The Mr of Cu(NO3)2.2H2O = 63.5 + 2(14 + (3 x 16)) + 2((2 x 1) + 16) = 223.5 21

NES/Chemistry/AS 1.2 The Mole and the Avogadro Constant A mole is the amount of substance which contains 6.02 x 1023 atoms, ions or molecules. Avogadro's Constant (L) This is quoted in the data booklet.  Example 3: How many particles of Al3+ are there in 0.2 moles of Al(OH)3? Each Al(OH)3 has 1 particle of Al3+ So number of particles = 0.2 x 6.02x1023 = 1.20 x1023  Example 4: How many particles of OH- are there in 0.01 moles of Al(OH)3? Each Al(OH)3 has 3 particles of OH- so number of particles = 0.03 x 6.02x1023 = 1.80 x1022 22

NES/Chemistry/AS 1.3 The Determination of Relative Atomic Masses, Ar Mass Spectrometry Relative atomic mass can be determined by mass spectrometry. It measures the relative mass of a sample, when compared to a standard (carbon-12). ionisation sample vapour Step 1 - the sample is vapourised chamber (turned into a gas). electron gun detector accelerator Step 2 - the sample is ionised by electric field collision with a high speed electron. X(g) + e–  X+(g) + 2e– Step 3 - the sample is accelerated by an electric field. Step 4 - the sample is deflected by a magnetic field. magnetic field Step 5 - a mass spectra graph is plotted from the results produced by the detector to vacuum pump More massive particles need a stronger magnetic field to follow the curved path. Therefore, we can work out the mass by the magnetic field strength, assuming all ions are unipositive. The spectrum is plotted as relative abundance of a unipositive (1+) cation (y-axis) against mass/charge ratio (x-axis). The x-axis is normally labelled m/e or m/z. At AS Level, the charge is usually 1+ so the x-axis is effectively the same as the mass. More complex questions are asked at A2 Level. 23

NES/Chemistry/AS  Example 5: The mass spectrum for neon, Ne 114 abundance 11.2 0.20 20 21 22 m/e Species responsible for each peak are: 20 m/e = 20Ne+ 21 m/e = 21Ne+ 22 m/e = 22Ne+ 24

NES/Chemistry/AS  Example 6: The mass spectrum for chlorine, Cl2 abundance 35 37 70 72 74 m/e There are two sets of peaks. The 35 and 37 peaks are from individual chlorine atoms, which got broken apart in the spectrometer. The peaks at 70, 72 and 74 are from chlorine molecules which did not break apart in the machine. Species responsible for each peak are: 35 m/e = 35Cl+ 37 m/e = 37Cl+ 70 m/e = [35Cl–35Cl]+ 72 m/e = [35Cl–37Cl]+ and [37Cl–35Cl]+ 74 m/e = [37Cl–37Cl]+ As the chlorine-35 and chlorine 37 atoms are in a 3:1 ratio, then the peaks at 70, 72 and 74 are in a 9:6:1 ratio. 25

NES/Chemistry/AS 1.4 The Calculation of Empirical and Molecular Formulae Empirical Formula The empirical formula of a molecule is the smallest, whole number ratio of atoms present in the molecule. For ionic compounds and macromolecular compounds - the molecular formula will be the same as the compound's formula. It is already in the simplest, whole number ratio. For covalent molecules - the molecular formula can sometimes be simplified and this becomes the empirical formula.  Example 7: The formula and empirical formula of silicon (IV) oxide, a macromolecular compound, is SiO2. It is already in its simplest ratio of one silicon atom to two oxygen atoms.  Example 8: The molecular formula for glucose is C6H12O6 whereas the empirical formula is simplified to just CH2O. The empirical formula can be calculated from the mass, or percentage composition of elements in a compound. This is the same as the calculations at IGCSE level. Step 1: Write out the elements in the compound Step 2: Divide by the mass (or %) by the elements Ar Step 3: Divide by the smallest number to convert ratio to whole numbers This will give you the ratio of elements. Sometimes you have to multiply the numbers by 2, or 3 to get whole numbers. 26

NES/Chemistry/AS  Example 9: A phosphide of arsenic contains 40.8% phosphorus by mass. Calculate the empirical formula of the compound. Step 1 As P Step 2 Step 3 1 1.67 Ratio 3 5 Whole number ratio So the formula of the compound is As3P5 27

NES/Chemistry/AS Molecular Formula For covalent compounds (such as organic compounds) the empirical formula is often different to the molecular formula. It can be calculated if you know the molecular mass, just like at IGCSE level. Step 1: Calculate the empirical mass Step 2: Divide the molecular mass by the empirical mass to find the number of formula units. Step 3: Multiply the empirical formula by the number of formula units.  Example 10: A carboxylic acid has the empirical formula CHO2 and a molecular mass of 90.0. Calculate the molecular formula. Step 1 Empirical mass = (12.0 + 1.0 + (2 x 16.0)) = 45.0 Step 2 = = 2 formula units Step 3 Molecular formula = CHO2 x 2 = C2H2O4 This is ethandioic acid, HOOCCOOH 28

NES/Chemistry/AS 1.5 Reacting masses and volumes (of solutions and gases) The equations for these calculations are the same as at IGCSE level, although the calculations are often more complex, involving several calculations combined together.  Equation for mass The mass must be in g  Equation for solutions The concentration is in moldm-3 The volume is in dm3  Equation for gases The volume must be in dm3 29

NES/Chemistry/AS  Example 11: A moles question from a Cambridge Paper 1 Number of moles of NaOH = 2 x 100/1000 = 0.2 mol Using the mole ratio in the first equation: Number of moles of Al2O3 = 0.1 mol Number of moles of HCl = 2 x 800/1000 = 1.6 mol So the correct answer must be either C, or D. Using the mole ratio in the second equation: Number of moles of HCl used by Al2O3 = 0.6 mol Moles of acid left = 1.6 - 0.6 = 1.0 mol Using the mole ratio in the third equation: Number of moles of MgO = 0.5 mol So the correct answer is D. 30

NES/Chemistry/AS Topic 2 - Atomic Structure This topic describes the type, number and distribution of the fundamental particles which make up an atom and the impact of this on some atomic properties. 2.1 Particles in the atom a) identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses b) deduce the behaviour of beams of protons, neutrons and electrons in electric fields c) describe the distribution of mass and charge within an atom d) deduce the numbers of protons, neutrons and electrons present in both atoms and ions given proton and nucleon numbers and charge 2.2 The nucleus of the atom a) describe the contribution of protons and neutrons to atomic nuclei in terms of proton number and nucleon number b) distinguish between isotopes on the basis of different numbers of neutrons present c) recognise and use the symbolism xyA for isotopes, where x is the nucleon number and y is the proton number 2.3 Electrons: energy levels, atomic orbitals, ionisation energy a) describe the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3 and also the 4s and 4p orbitals b) describe and sketch the shapes of s and p orbitals c) state the electronic configuration of atoms and ions given the proton number and charge, using the convention 1s22s22p6, etc. d) (i) explain and use the term ionisation energy 31

NES/Chemistry/AS (ii) explain the factors influencing the ionisation energies of elements (iii) explain the trends in ionisation energies across a Period and down a Group of the Periodic Table (see also Section 9.1) e) deduce the electronic configurations of elements from successive ionisation energy data f) interpret successive ionisation energy data of an element in terms of the position of that element within the Periodic Table Proton Number or Atomic Number (Z) is the number of protons in the nucleus of an atom and indicates the element's position in the periodic table. Nucleon Number or Mass Number (A) is the number of protons and neutrons found in the nucleus of an atom of an element. An Orbital is the volume of space in which there is a 95% chance of finding an electron. Isotopes are Atoms of the same element with the same number of protons, but different number of neutrons. The First Ionisation Energy is the enthalpy change required to remove one mole of electrons from one mole of gaseous atoms of the element to form one mole of gaseous unipositive cations. The Second Ionisation Energy is the enthalpy change required to remove one mole of electrons from one mole of gaseous univalent cations of the element to form one mole of gaseous dipositive cations. 32

NES/Chemistry/AS 2.1 Particles in the Atom Atoms are made up of protons, neutrons and electrons. Hydrogen is unusual as it is the only atom that has no neutrons. Charge and Mass of Sub-Atomic Particles in an Atom Particle Relative charge Relative mass Proton 1+ Neutron 0 1 Electron 1– 1 1 1837 Atoms are always neutral overall, as the number of electrons (negative) is equal to the number of protons (positive). Ions An ion is an atom with a charge. They are formed when atoms lose, or gain electrons. Calculating the Number of Sub-Atomic Particles in an Atom, or Ion Number of Protons = Proton Number Number of Electrons in an Atom = Proton Number Note - when calculating the number of electrons in an ion, you must also take into account the charge on the ion. Number of Electrons in an Ion = Proton Number - Ion Charge Number of Neutrons = Nucleon Number - Proton Number 33

NES/Chemistry/AS Behaviour of Sub-Atomic Particles in an Electric Field  Protons protons + - Protons are deflected in an electric field towards the negative pole, as they have a positive charge.  Electrons electrons + - Electrons are deflected in an electric field towards the positive pole, as they have a negative charge. The angle of deflection for electrons is greater than that of protons because electrons has a smaller mass.  Neutrons neutrons + - Neutrons are not deflected in an electric field as they have no charge. 34

NES/Chemistry/AS increasing energy Structure of an Atom Above this level the electron has enough energy to escape; the atom is ionised Energy Level 5, 6 and 7 4 3 2 1 Nucleus  Nucleus The central part of an atom is called the nucleus and contains nucleons (protons and neutrons). This is where most of the mass of the atom is and it has an overall positive charge.  Energy Levels (called shells at IGCSE level) Electrons surround the nucleus at different energy levels. The combined mass of all the electrons is negligible compared to the mass of the nucleus and the electrons have a negative charge. 35

NES/Chemistry/AS 2.2 The Nucleus of the Atom Isotope Atoms of the same element can have different numbers of neutrons - these are called isotopes. This happens naturally with all elements and causes the mass number to change. Calculating the average mass of isotopes is covered in Topic 1. The mass numbers quote on the periodic table are average mass values rounded to 1 decimal place, or 3 significant figures. We do not focus much on the nucleus of the atom in Chemistry. However, the next section on electrons is quite detailed and significantly more complex than IGCSE. 36

NES/Chemistry/AS 2.3 Electrons: Energy Levels, Atomic Orbitals, Ionisation Energy Electron affinity is not part of the AS syllabus anymore, it has moved to A2 level. There are four quantum numbers that determine the likely location and properties of an electron in an atom:  Principal Quantum Number - this is the energy level the electron is in. It ranges from 1 (closest/lowest energy) to 7 (furthest away/highest energy).  Orbital Angular Momentum Quantum Number - this determines the shape of the orbital.  Magnetic Quantum Number - determines the number and orientation of the orbitals.  Electron Spin Quantum Number - electrons spin either up, or down Energy Levels Energy levels can be divided into energy sub-levels, which are made up of orbitals. Each orbital can hold a maximum of 2 electrons, one with spin up and the other with spin down. This is because there cannot be 2 electrons in an atom with the exact same four quantum numbers. We only need to know the arrangement of electrons in detail up to the fourth energy level. Energy Sub- Number of Maximum Number Maximum Number of Level Level Orbitals in of Electrons in Electrons in Energy Sub-Level Sub-Level 1 1s Level 2 2s 1 2 1 2 2 2p 3 6 8 3s 1 2 3 3p 3 6 18 3d 5 10 4s 1 2 32 4p 3 6 4 4d 5 10 4f 7 14 37

NES/Chemistry/AS Orbitals An orbital is an area of space in an atom where you are likely to find an electron 95% of the time. Each orbital can hold a maximum of 2 electron as long as they have opposite spin. As we move up the energy levels, the orbitals get larger in size. Orbitals usually fill from the lowest energy level, which is the closest to the nucleus of the atom. This is called the Ground State for the atom and is the arrangement that requires the least amount of energy. Orbitals also usually fill the sub-levels in the order s, p, d, then f. There are exceptions we will come to later on. Each type of orbital has its own shape which you must be able to recognise and draw. Orbital Shapes  s-orbitals These are spherical in shape and get larger as the energy level increases. Below is a scatter diagram of where the electron is most likely to be found and from this the spherical shape can be seen. The diagrams are just draw as spheres. There is only ever one s-orbital in an energy level. This corresponds to Groups I and II on the Periodic Table. 38

NES/Chemistry/AS  p-orbitals These are figure-8 shaped/hourglass shaped/dumb bell shaped and the shape gets larger as the energy level increases. p-orbitals are only found in energy levels 2-7. There are three types of p-orbital that lie on the x, y and z axis around the nucleus of an atom. You will be expected to identify and draw each type of orbital individually, but not all together on a single atom as this would get far too messy. Electronic Configuration At IGCSE level, you were shown a simplified model of how the electrons are arranged around the nucleus. We are now going to look at a more complex model of the arrangement of electrons. The electron can either be draw out as up and down arrows in boxes (called electrons-in-boxes) where the arrows represent the spin of the electron and the boxes show what energy level and orbital the electron is in. Alternatively, the electrons can be written out stating energy level, orbital and number of electrons present.  Electrons-In-Boxes There is a maximum of 2 electrons per box. Electrons are draw as up, or down arrows. In the p and d orbitals the boxes are filled with single electrons first, before pairing up the electrons, due to electron pair repulsion. 1s 2s 2p 3s 3p 39

NES/Chemistry/AS  Example 1: Carbon ↑↓ ↑↓ ↑ ↑ 1s 2s 2p  Example 2: Chlorine ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ 1s 2s 2p 3s 3p However, after the 3p energy sub-level, the orbitals start to fill out of order. The 4s energy sub-level fills before the 3d energy sub-level. This is because the 4s orbital requires less energy for the electron to occupy, despite the fact it is further away than the 3d orbital. Filling order is affected by the shapes of the orbitals, but we only have to go as far as the 4p orbital; rather than learn the entire sequence.  Example 3: Nickel ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ 1s 2s 2p 3s 3p 4s 3d Also, the 3d orbital is more stable when half-full, or full; so chromium and copper fill in a slightly different sequence.  Example 4: Chromium ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑ ↑ ↑ ↑ 1s 2s 2p 3s 3p 4s 3d You can see that there is only 1 electron in the 4s orbital, this is because the electron is found in the 3d orbital instead, half filling the energy sub-level. This results in a more stable arrangement for the electrons. 40

NES/Chemistry/AS These electronic configurations are all called 'ground state' because it is the arrangement that requires the least amount of energy. Later on in Topic 3 you will see electrons being 'promoted' to higher energy levels which is called an 'excited state'.  Stating Energy Levels This is similar to the electrons-in-boxes notation, but the configuration is simply stated. 3p4 The first number is the energy level The letter is the orbital The second number superscript is how many electrons there are in that energy sub-level.  Example 5: Calcium 1s22s22p63s23p64s2 Note - if you look at the last energy sub-level (here 4s2) it tells you that the element is in Period 4 and Group 2. Note - if you count the electrons of an atom (superscript numbers) then it will equal the proton number of the element.  Example 6: Copper 1s22s22p63s23p64s13d10 - filling order 1s22s22p63s23p63d104s1 - electronic configuration Remember - copper, like chromium, fills in a way that gets it a full d energy sub-level. Note - the filling order and electronic configuration are written in a different order. The filling order is written in the order that the electrons fill up the orbitals. The electronic configuration puts the energy sub-levels in energy level order. 41

NES/Chemistry/AS  Example 7: Nitride, N3- 1s22s22p6 Note - the electronic configuration of an ion is different from an atom, as it has either gained, or lost electrons. Note - ions are usually isoelectric with Noble Gases, as they tend to fill, or empty their valence energy level. Isoelectric means it has the same electronic configuration. It does not mean it has become a new element. Valence Energy Level is the outermost energy level.  Example 8: Titanium(II), Ti2+ 1s22s22p63s23p63d2 Note - with Transition Metals, they give away their 4s orbital electrons, before their 3d orbital electrons as they are further away from the nucleus. This is the opposite of the filling order. Note - the trends in the following sections are also linked to Topic 9.1. This topic is not taught until later, but it will help to look at the graphs, trends an explanations for this topic as well. 42

NES/Chemistry/AS Ionisation Energy Electrons can rise up the energy levels in an atom and if they have enough energy, leave the atom completely. This will result in the formation of a unipositive cation. The energy required to overcome the force of attraction between the nucleus (+) and the electron (-) is called first ionisation energy. Further electrons can be removed from the cation one at a time, called second ionisation energy, third ionisation energy, etc. Electrons in the outermost (valence) energy level are removed first, followed by each energy level in sequence. It is possible to remove all electrons from an atom, although it is unlikely to be stable. The exception to this is the order that electrons are removed from the Transition Metals (see example 8).  First Ionisation Energy as an equation: X(g) → X+(g) + e-  Second Ionisation Energy as an equation: X+(g) → X2+(g) + e- Note - any element can form a positive ion, including non-metals and Noble Gases. Factors Affecting Ionisation Energy The three main factors are: 1. Nuclear Charge 2. Distance of the electron from the nucleus 3. Shielding from other electrons There are also two more factors which also have an effect: 4. Sub-Energy levels 5. Electron Pair repulsion 43

NES/Chemistry/AS 1. Nuclear Charge The more protons there are in the nucleus of an atom, the greater the nuclear charge is. So, as we move across a Period on the Periodic Table, the proton number increases by one, so the nuclear charge increases. This makes it more difficult to remove the electron as the force of attraction between the electron and the nucleus is stronger. 2. Distance of the Electron from the Nucleus Electrons in higher energy levels are further away from the nucleus, so the force of attraction between the electron and the nucleus decreases as we go down a Group on the Periodic Table. 3. Shielding from other Electrons Valence electrons are shielded by any electrons at a lower energy level as they have the same negative charge and this causes repulsion, rather than attraction. So as we go down a Group on the Periodic Table shielding increases and the force of attraction between the valence electron and the nucleus is weaker. When working out ionisation energy, or comparing different ionisation energies, all three factors must be taken into account and explained. Trend in Ionisation Energy Across a Period Going from left to right across a Period, the ionisation energy usually increases because:  Nuclear charge increases, as there is one more proton in the nucleus each time  Distance is the same, as electrons are added to the same energy level  Shielding is the same, as electrons are added to the same energy level There are some exceptions to this trend, which we will discuss later on. 44

First Ionisation Energy kJ/mol NES/Chemistry/AS Trend in Ionisation Energy Down a Group Going down a Group the ionisation energy decreases because:  Despite the fact that the nuclear charge increases  The distance is greater, as the electrons are being removed from a higher energy level  The shielding is greater, as the electrons are being removed from a higher energy level Graph of First Ionisation Energy He Ne N Ar H O Be B Li Na Atomic Number 45

NES/Chemistry/AS 4. Sub-Energy Levels You will notice that from Beryllium to Boron in Period 2 and from Magnesium to Aluminium in Period 3 there is a decrease in first ionisation energy instead of the expected increase. This is because the electron in the Group III element (Boron, or Aluminium) is being removed from a p-orbital; rather than an s-orbital. The p-orbital is at a slightly higher energy level than the s-orbital, so it requires a little more energy to remove the electron from the atom. 5. Electron Pair Repulsion You will also notice that from Nitrogen to Oxygen in Period 2 and from Phosphorus to Sulphur in Period 3 there is a slight decrease in first ionisation energy instead of the expected increase. This is because in Oxygen and Sulphur the electron being removed is paired up in a p-orbital and there is a slight repulsive force between the electrons, making them slightly easier to remove. 46

NES/Chemistry/AS Topic 3 - Chemical Bonding This topic introduces the different ways by which chemical bonding occurs and the effect this can have on physical properties. 3.1 Ionic bonding a) describe ionic bonding, as in sodium chloride, magnesium oxide and calcium fluoride, including the use of ‘dot-and-cross’ diagrams 3.2 Covalent bonding and co-ordinate (dative covalent) bonding including shapes of simple molecules a) describe, including the use of ‘dot-and-cross’ diagrams: (i) covalent bonding, in molecules such as hydrogen, oxygen, chlorine, hydrogen chloride, carbon dioxide, methane, ethene (ii) co-ordinate (dative covalent) bonding, such as in the formation of the ammonium ion and in the Al2Cl6 molecule b) describe covalent bonding in terms of orbital overlap, giving σ and π bonds, including the concept of hybridisation to form sp, sp2 and sp3 orbitals (see also Section 14.3) c) explain the shapes of, and bond angles in, molecules by using the qualitative model of electron-pair repulsion (including lone pairs), using as simple examples: BF3 (trigonal), CO2 (linear), CH4 (tetrahedral), NH3 (pyramidal), H2O (non-linear), SF6 (octahedral), PF5 (trigonal bipyramidal) d) predict the shapes of, and bond angles in, molecules and ions analogous to those specified in 3.2(b) (see also Section 14.3) 3.3 Intermolecular forces, electronegativity and bond properties a) describe hydrogen bonding, using ammonia and water as simple examples of molecules containing N–H and O–H groups b) understand, in simple terms, the concept of electronegativity and apply it to explain the properties of molecules such as bond polarity (see also Section 3.3(c)), the dipole moments of molecules (3.3(d)) and the behaviour of oxides with water (9.2(c)) 47

NES/Chemistry/AS c) explain the terms bond energy, bond length and bond polarity and use them to compare the reactivities of covalent bonds (see also Section 5.1(b)(ii)) d) describe intermolecular forces (van der Waals’ forces), based on permanent and induced dipoles, as in, for example, CHCl3(l); Br2(l) and the liquid Group 18 elements 3.4 Metallic bonding a) describe metallic bonding in terms of a lattice of positive ions surrounded by delocalised electrons 3.5 Bonding and physical properties a) describe, interpret and predict the effect of different types of bonding (ionic bonding, covalent bonding, hydrogen bonding, other intermolecular interactions, metallic bonding) on the physical properties of substances b) deduce the type of bonding present from given information c) show understanding of chemical reactions in terms of energy transfers associated with the breaking and making of chemical bonds 48

NES/Chemistry/AS Hybridisation happens when atomic orbitals mix to form new atomic orbitals. The new orbitals have the same total electron capacity as the old ones. The properties and energies of the new, hybridized orbitals are an 'average' of the original unhybridized orbitals. Orbital Overlap is the concentration of orbitals on adjacent atoms in the same regions of space. -This leads to the formation of bonds in a molecule. Ionic Bonds are a strong non-directional electrostatic force of attraction between the alternating cations and anions formed due to the transfer of electrons. -This leads to the formation of ionic lattices. Covalent Bonds are a directional bond formed between two atoms by sharing one or more pairs of electrons in the overlap, one electron being contributed by each atom in order to have a full outer shell. -This leads to the formation of molecules and macromolecules. Co-ordinate (Dative Covalent) Bonds are formed when a pair of electrons is shared between two atoms, where both the shared electrons in the covalent bond come from a filled valence orbital of one of the atoms joining together. -This leads to the formation of molecules. σ Bonds are formed when atomic orbitals overlap linearly along the line joining the two bonded atoms forming molecular orbitals. -This is the type of bonding in all single covalent bonds π Bonds are formed due to the lateral/sideways overlap of two p orbitals, creating electron density above and below the plane of the molecule. -Double bonds have one π bond and one σ bond -Triple bonds have two π bonds and one σ bond Metallic Bonds are the strong non-directional electrostatic force of attraction between the lattice of positive ions and the mobile ‘sea’ of delocalised valence electrons in a giant 3-dimensional structure. -This leads to the formation of metallic lattices. Delocalised Electrons are electrons in a atom, ion or molecule not associated with any single atom, or a single covalent bond. - Metallic lattices and graphite have delocalised electrons. 49