Chemistry 9701/22 May/June 2025: question paper and mark scheme

Question

(a) Diamond and graphite are both crystalline solids made from carbon atoms. Graphite conducts electricity. Diamond does not conduct electricity. (i) Name the type of lattice structure present in the crystalline solids diamond and graphite. …… [1] (ii) Explain how graphite conducts electricity. …… [1] (iii) Explain why diamond does not conduct electricity. …… [1] (b) Separate samples of phosphorus(V) chloride and silicon(IV) chloride are each added to an excess of cold water. (i) Write an equation for each reaction. …… [2] (ii) Describe the appearance of phosphorus(V) chloride and silicon(IV) chloride at room temperature. phosphorus(V) chloride …… silicon(IV) chloride …… [2] (iii) Compare the appearance of the mixtures produced when each reaction is complete. …… [2] (c) Some oxides are amphoteric. (i) Describe what is meant by amphoteric. …… [1] (ii) Identify the formula of a Period 3 oxide that is amphoteric. …… [1] (d) The melting points of different oxides are shown in Table 1.1. Table 1.1 oxide melting point / °C force of attraction broken during melting SO2 –73 H2 O 0 SiO2 1610 Na2O 1132 MgO 2852 (i) Complete Table 1.1 by identifying the strongest force of attraction in each oxide that is broken during melting. Use the abbreviations below. i.d. = instantaneous dipole–induced dipole p.d. = permanent dipole–permanent dipole H = hydrogen bond C = covalent bond I = ionic bond [3] (ii) A student suggests the following hypothesis. The stronger the covalent bond between atoms in non‑metal oxides, the higher the melting point. Use Table 1.1 to deduce if this hypothesis is true or false or if there is not enough information to make a conclusion. Explain your answer. …… [2] [Total: 16]

Question

A sample of iron contains three different isotopes and has a relative atomic mass, Ar, of 55.8. (a) (i) Define relative atomic mass. …… [2] (ii) Table 2.1 shows the abundances of two of the isotopes present in the sample of iron. Table 2.1 isotope relative isotopic mass abundance / % 54Fe 53.9 6.0 56Fe 55.9 91.9 Use Table 2.1 to calculate the relative isotopic mass of the third isotope of iron in the sample. Show your working. relative isotopic mass = …… [2] (b) Deduce the number of pairs of electrons in the 3d sub‑shell in an iron(II) ion. …… [1] (c) Sketch the shape of the lowest energy orbital in the shell with principal quantum number n = 2. [1] (d) Complete Table 2.2 to show information about particles in one ion of 56Fe3+. Table 2.2 number of particles present in particle one ion of 56Fe3+ electrons 30 [2] (e) The atomic radius of iron is 1.26 × 10–10 m. Suggest the change to the radius, if any, after an iron atom reacts to produce an Fe3+ ion. Explain your answer. …… [2] [Total: 10]

Question

(a) C3H6 reacts with HBr(g) in an addition reaction. (i) Define addition reaction. …… [1] (ii) Complete Fig. 3.1 to show the mechanism for the addition reaction between C3H6 and HBr to produce 2‑bromopropane. Include charges, dipoles, lone pairs of electrons and curly arrows, as appropriate. H3C H CH3 C C H C CH3 H H Br H Br Fig. 3.1 [4] (iii) Explain why the major product of this reaction is 2‑bromopropane rather than 1‑bromopropane. …… [2] (b) H2(g) and nickel are added to alkene X. Fig. 3.2 shows how the concentration of X changes with time. 1.0 0.8 concentration 0.6 of X / mol dm–3 0.4 0.2 0.0 0 200 400 600 time / s Fig. 3.2 (i) Use Fig. 3.2 to identify the limiting reagent in this reaction. Explain your answer. …… [1] (ii) Use Fig. 3.2 to describe how the gradient changes as the reaction proceeds. State what this shows about the rate during this reaction. …… [1] (c) Hydrocarbon Y contains two C=C groups. There are no other functional groups present. Y reacts with an excess of H2(g) to produce hexane, C6H14. Table 3.1 bond energy / kJ mol–1 C−C 350 C=C 610 C≡C 840 H−H 436 C−H 410 Use Table 3.1 to calculate the enthalpy change per mole of C6H14 produced in this reaction. enthalpy change = …… kJ mol–1 [2] (d) Hydrocarbon Y reacts with H2(g). Fig. 3.3 shows the distribution of energies of H2(g) at temperature T. Area A represents the number of molecules with energy greater than or equal to the activation energy, EA, at temperature T. number of molecules area A EA energy Fig. 3.3 (i) Annotate Fig. 3.3 to show the effect of adding nickel to Y and H2(g) at temperature T. [1] (ii) Area B (not labelled on Fig. 3.3) represents the number of molecules with energy greater than or equal to the activation energy when nickel is added at temperature T. State the difference, if any, between areas A and B. Explain the significance of your answer on the rate of hydrogenation of Y. Give your answer in terms of collisions. …… [2] (e) Alkene Z contains two C=C bonds. Z reacts with an excess of hot concentrated acidified KMnO4 to produce only CH3COCH3, HOOCCH2COOH, CO2 and H2O. Suggest the structure of Z. [2] [Total: 16]

Question

(a) Propanoic acid, CH3CH2COOH, reacts with reducing agent Q to produce propan‑1‑ol, CH3CH2CH2OH. (i) Suggest the formula of reducing agent Q. …… [1] (ii) Complete the equation to show the reduction of propanoic acid to propan‑1‑ol. Use [H] to represent one atom of hydrogen from Q. ……CH3CH2COOH + ……[H] …… [1] (b) Propan‑1‑ol is converted to compound T in a three‑step synthesis. In step 1, SOCl 2 is added to propan‑1‑ol to produce compound E. In step 2, E reacts with a suitable reagent to produce butanenitrile. In step 3, butanenitrile is heated with NaOH(aq). (i) Name the type of reaction that occurs in step 1. …… [1] (ii) Identify the reagent and conditions required in step 2. …… [1] (iii) Construct an equation to describe the reaction of butanenitrile with NaOH(aq) in step 3. …… [2] (iv) Name the type of reaction that occurs in step 3. …… [1] (c) Compounds A, B and C belong to the alcohol homologous series. Each molecule of A, B and C contains four saturated carbon atoms. (i) Identify the type of hybridisation shown in the saturated carbon atoms of all alcohols. …… [1] (ii) Identify the gas produced when Na(s) is added to separate samples of each alcohol. …… [1] (iii) Describe the role of Na(s) when it reacts with alcohols. …… [1] (d) Table 4.1 shows the results of two tests on separate samples of A, B and C. Table 4.1 heat under reflux with warm with compound acidified K2Cr2O7 alkaline I2(aq) A remains orange no visible change B orange to green no visible change C orange to green pale yellow precipitate (i) Deduce the structure of A. [1] (ii) Deduce a possible name for B. …… [1] (iii) Deduce the identity of C. …… [1] [Total: 13]

Question

The structure of vitamin C is shown in Fig. 5.1. vitamin C HO O O HO HO OH Fig. 5.1 (a) Deduce the empirical formula of vitamin C. …… [1] (b) The mass of vitamin C present in 150.0 g of lemon is found in an experiment. stage 1 All the vitamin C in 150.0 g of lemon is extracted and dissolved in water to make 100.0 cm3 of solution L. stage 2 A 25.0 cm3 sample of solution L is titrated with 5.00 × 10–3 mol dm–3 I2(aq). vitamin C M HO HO O O O O HO HO I2 + + 2I– + 2H+ HO OH O O Fig. 5.2 Exactly 36.65 cm3 of I2(aq) reacts with the 25.0 cm3 sample of solution L. [Mr: vitamin C, 176] (i) Calculate the amount, in mol, of I2(aq) added in the titration. amount of I2(aq) = …… mol [1] (ii) Use your answer to (b)(i) to calculate the percentage by mass of vitamin C present in 150.0 g of lemon. (If you were unable to calculate a value for the amount of I2(aq) in (b)(i), use the value 7.65 × 10–4 mol. This is not the correct value.) percentage by mass = ……% [2] (iii) The progress of the reaction of vitamin C with I2(aq) to produce M is monitored using infrared spectroscopy. Table 5.1 indicates the presence of some absorptions in the infrared spectrum of vitamin C. Complete Table 5.1 to predict which of these absorptions, if any, are present in the infrared spectrum of M. Table 5.1 present in spectrum present in spectrum absorption / cm–1 of vitamin C of M 1500–1680 ✓ 2850–2950 ✓ 3200–3650 ✓ [1] Table 5.2 bond functional groups containing the bond characteristic infrared absorption range (in wavenumbers) / cm–1 C–O hydroxy, ester 1040–1300 C=C aromatic compound, alkene 1500–1680 C=O amide 1640–1690 carbonyl, carboxyl 1670–1740 ester 1710–1750 C≡N nitrile 2200–2250 C–H alkane 2850–2950 N–H amine, amide 3300–3500 O–H carboxyl 2500–3000 hydroxy 3200–3650 [Total: 5] Important values, constants and standards molar gas constant R = 8.31 J K–1 mol–1 Faraday constant F = 9.65 × 104 C mol–1 Avogadro constant L = 6.02 × 1023 mol–1 electronic charge e = –1.60 × 10–19 C molar volume of gas Vm = 22.4 dm3 mol–1 at s.t.p. (101 kPa and 273 K) Vm = 24.0 dm3 mol–1 at room conditions ionic product of water Kw = 1.00 × 10–14 mol2 dm–6 (at 298 K (25 °C)) specific heat capacity of water c = 4.18 kJ kg–1 K–1 (4.18 J g–1 K–1) The Periodic Table of Elements Group 1 2 13 14 15 16 17 18 1 2 H He hydrogen helium Key 1.0 4.0 3 4 atomic number 5 6 7 8 9 10 Li Be atomic symbol B C N O F Ne lithium beryllium name boron carbon nitrogen oxygen fluorine neon 6.9 9.0 relative atomic mass 10.8 12.0 14.0 16.0 19.0 20.2 11 12 13 14 15 16 17 18 Na Mg Al Si P S Cl Ar sodium magnesium aluminium silicon phosphorus sulfur chlorine argon 23.0 24.3 3 4 5 6 7 8 9 10 11 12 27.0 28.1 31.0 32.1 35.5 39.9 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 at www.cambridgeinternational.org after the live examination series. K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr potassium calcium scandium titanium vanadium chromium manganese iron cobalt nickel copper zinc gallium germanium arsenic selenium bromine krypton 39.1 40.1 45.0 47.9 50.9 52.0 54.9 55.8 58.9 58.7 63.5 65.4 69.7 72.6 74.9 79.0 79.9 83.8 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 16 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe rubidium strontium yttrium zirconium niobium molybdenum technetium ruthenium rhodium palladium silver cadmium indium tin antimony tellurium iodine xenon 85.5 87.6 88.9 91.2 92.9 95.9 – 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 56 57–71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 lanthanoids Cs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn caesium barium hafnium tantalum tungsten rhenium osmium iridium platinum gold mercury thallium lead bismuth polonium astatine radon 132.9 137.3 178.5 180.9 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 – – – 87 88 89–103 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 actinoids Fr Ra Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og francium radium rutherfordium dubnium seaborgium bohrium hassium meitnerium darmstadtium roentgenium copernicium nihonium flerovium moscovium livermorium tennessine oganesson – – – – – – – – – – – – – – – – – 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 lanthanoids 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 138.9 140.1 140.9 144.2 – 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 actinoids 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 – – – – – – – – – – – Assessment International Education Copyright Acknowledgements Booklet. This is produced for each series of examinations and is freely available to download To avoid the issue of disclosure of answer‑related information to candidates, all copyright acknowledgements are reproduced online in the Cambridge