Pra-U STPM Chemistry Penggal 2 2019 CB039349b

Chemistry Term 2 STPM Chapter 9 Periodic Table: Periodicity 143 9 (iii) Explain why aluminium chloride is hydrolysed by water but not magnesium chloride. 8 A graph of first ionisation energy versus proton number for several elements in Period 3 of the Periodic Table is shown below: Proton number W I.E. 0 (a) Identify the element labelled W. (b) Explain why the first ionisation energy of W is lower than the preceding element. 9 The table below shows the formulae of ions with proton number 11 to 17 except proton number 14. Proton number 11 12 13 15 16 17 Formula of ion Na+ Mg2+ Al3+ P3– S2– Cl– (a) Sketch a graph of the ionic radius against proton number. (b) Explain your graph. 10 The graph below shows the variation of the melting point of the elements in Period 3 of the Periodic Table. Proton number Melting point 0 Based on the graph, explain the following. (a) The melting point increases from sodium to aluminium. (b) Silicon has an exceptionally high melting point. (c) Sulphur has a higher melting point than phosphorous. 11 The table below shows the melting point (M.p.) of some of the oxides of the Period 3 elements. Oxide Na2O MgO Al2O3 SiO2 P4O10 SO3 M.p./°C 1 275 2 852 2 067 1 610 580 17 (a) Explain the trend in the melting of the above oxides in terms of structure and bonding. (b) Classify the above oxides as acidic, basic or amphoteric. Illustrate their acid/base properties with reference to reactions with aqueous sodium hydroxide or hydrochloric acid.

CHAPTER 10 GROUP 2 Concept Map Learning earning Outcomes Group 2 Elements Physical Properties Anomalous Properties of Beryllium Diagonal Relationship Chemical Properties • General • Reactions with water Students should be able to: Selected Group 2 elements and their compounds • describe the trends in physical properties of Group 2 elements: Mg, Ca, Sr, Ba; • describe the reactions of Group 2 elements with oxygen and water; • describe the behaviour of the oxides of Group 2 elements with water; • explain qualitatively the thermal decomposition of the nitrates, carbonates and hydroxides of Group 2 elements in terms of the charge density and polarisability of large anions; • explain qualitatively the variation in solubility of sulphate of Group 2 elements in terms of the relative magnitudes of the enthalpy change of hydration for the relevant ions and the corresponding lattice energy. Anomalous behaviour of beryllium • explain the anomalous behaviour of beryllium as exemplified by the formation of covalent compounds; • describe the diagonal relationships between beryllium and aluminium; • explain the similarity of aqueous beryllium salts to aqueous aluminium salts in terms of their acidic property. Uses of Group 2 Compounds • state the uses of Group 2 compounds in agriculture, industry and medicine. Thermal Decomposition of the Carbonates, Nitrates and Hydroxides • General • Relative thermal stability of carbonates Solubility of Group 2 Sulphates • Lattice energy of Group 2 sulphates • Hydration energy of Group 2 cations • Relative solubility of Group 2 sulphates

Chemistry Term 2 STPM Chapter 10 Group 2 145 10 10.1 Selected Group 2 Elements and Their Compounds 1 2 Be Mg Ca Sr Ba 13 14 15 16 17 18 1 The Group 2 elements are beryllium, magnesium, calcium, strontium, barium and radium. 2 Some information of the Group 2 elements is listed in the table below. Element Be Mg Ca Sr Ba Name Beryllium Magnesium Calcium Strontium Barium Proton number 4 12 20 38 56 Electronic configuration 2.2 [He] 2s2 2.8.2 [Ne] 3s2 2.8.8.2 [Ar] 4s2 2.8.18.8.2 [Kr] 5s2 2.8.18.18.8.2 [Xe] 6s2 3 All have valence shell electronic configurations of ns2 . 4 They are all reactive metals and are not found in the free elemental states in nature. 5 In their pure state, they have a silvery colour but tarnish rapidly in air due to the formation of an oxide layer on the metals’ surface. For example, 2Mg(s) + O2(g) ⎯→ 2MgO(s) 6 They are soft and can be easily cut with a knife. 7 The Group 2 elements give characteristic flame test: Element Magnesium Calcium Strontium Barium Colour of the flame Brilliant white Brick red Crimson Apple green Physical Properties of Group 2 Elements Going down Group 2 from beryllium to barium: (a) Atomic radius increases due to the increase in the screening effect. Element Be Mg Ca Sr Ba Proton number 4 12 20 38 56 Electronic configuration 2.2 2.8.2 2.8.8.2 2.8.18.8.2 2.8.18.18.8.2 Effective nuclear charge 2 2 2 2 2 Atomic radius/nm 0.112 0.160 0.197 0.215 0.222 Info Chem The Group 2 elements are known as the alkaline earth metals. The oxidation of magnesium by air The flame test 2016/P2/Q20(a) Info Chem The Group 2 elements are extracted via the electrolysis of their molten chlorides. Refer to Section 9.1 for a more detailed discussion on the physical properties of the Group 2 elements. Exam Tips Exam Tips Group 2 elements are conductors due to the metallic bonds.

Chemistry Term 2 STPM Chapter 10 Group 2 146 10 (b) Ionic radius increases. Ion Proton number Electronic configuration Ionic radius/nm Be2+ 4 2 0.031 Mg2+ 12 2.8 0.065 Ca2+ 20 2.8.8 0.099 Sr2+ 38 2.8.18.8 0.113 Ba2+ 56 2.8.18.18.8 0.135 (c) Ionisation energy decreases. Element First ionisation energy/kJ mol–1 Second ionisation energy/kJ mol–1 Third ionisation energy/kJ mol–1 Be 900 1800 14 800 Mg 740 1450 7700 Ca 590 1150 4900 Sr 550 1060 4200 Ba 500 970 3390 Their third ionisation energies are exceptionally high because the third electron removed is from an inner shell compared to the first two electrons. As a result, Group 2 elements do not form the M3+ ions. Group 2 elements show a constant oxidation state of +2 in all their compounds. (d) Electronegativity decreases. Element Electronegativity Be 1.50 Mg 1.30 Ca 1.00 Sr 0.97 Ba 0.93 Chemical Properties of Group 2 Elements 1 Due to their relatively large atomic size and low ionisation energy, Group 2 elements react by way of losing two electrons to form the M2+ ions. Mg ⎯→ Mg2+ + 2e– ∆H° = +2190 kJ mol–1 Ba ⎯→ Ba2+ + 2e– ∆H° = +1470 kJ mol–1 2 They are all strong reducing agents as shown by their large negative standard electrode potentials (E°). Element Be Mg Ca Sr Ba E°/V –1.85 –2.37 –2.87 –2.89 –2.90 [Note: The more negative the E° value, the stronger is the reducing agent.] I.E. 1 2 Order of electron removed 3 2018/P2/Q9 They are all powerful reducing agents. Reducing strength increases down the group. Info Chem The size of atom increases. Effective nuclear charge does not change much. As a result, attraction for bonding electrons decreases.

Chemistry Term 2 STPM Chapter 10 Group 2 147 10 10 20 30 E°/V 40 50 –1.8 Proton number –1.9 –2.0 –2.1 –2.2 –2.3 –2.4 –2.5 –2.6 –2.7 –2.8 –2.9 Be Mg Ca Sr Ba 3 For example, magnesium is used in the extraction of titanium from titanium(IV) chloride. 2Mg(s) + TiCl4(s) ⎯→ Ti(s) + 2MgCl2(s) 4 The reducing strength of the elements increases with proton number. This is because of a decrease in the ionisation energy down the group as the atomic size increases while the effective nuclear charge remains almost constant. Quick Check 10.1 1 The standard electrode potentials of magnesium and barium are –1.85 V and –2.90 V respectively. (a) Write two balanced ionic equations to represent the standard electrode potentials of magnesium and barium. (b) Account for the fact that the standard electrode potential of barium is more negative than that of magnesium. 2 Magnesium is used in the extraction of titanium from titanium(IV) chloride. (a) Calculate the mass of magnesium required to react with 1000 kg of titanium(IV) chloride. (b) Calculate the mass of titanium produced assuming that the efficiency of the reaction is 65%.

Chemistry Term 2 STPM Chapter 10 Group 2 148 10 Reactions with Water 1 All Group 2 elements react with water to form their respective hydroxides with the liberation of hydrogen gas. M(s) + 2H2O ⎯→ M(OH)2 + H2(g) In the reaction, the Group 2 elements act as reducing agents and reduce water to hydrogen. 2H2O + 2e– ⎯→ H2 + 2OH– 2 Going down the group, the reactivity towards water increases as the reducing power of the elements increases. 3 Beryllium reacts slowly with steam: Be(s) + 2H2O(g) ⎯→ Be(OH)2(s) + H2(g) Δ 4 Magnesium reacts slowly with hot water, but rapidly with steam. Mg(s) + 2H2O(g) ⎯→ Mg(OH)2(s) + H2(g) Δ 5 Calcium reacts rapidly with hot water, but slowly with cold water. Ca(s) + 2H2O(l) ⎯→ Ca(OH)2(aq) + H2(g) Δ 6 Strontium and barium reacts vigorously with cold water. Sr(s) + 2H2O(l) ⎯→ Sr(OH)2(aq) + H2(g) Ba(s) + 2H2O(l) ⎯→ Ba(OH)2(aq) + H2(g) 7 All the hydroxides are bases except beryllium hydroxide, which is amphoteric. 8 For example, Mg(OH)2(s) + H2SO4(aq) ⎯→ MgSO4(aq) + 2H2O(l) Ba(OH)2(aq) + 2HCl(aq) ⎯→ BaCl2(aq) + 2H2O(l) 9 The solubility of the hydroxides increases down the group. Hydroxide Solubility Solubility/g per 1000 g H2O Be(OH)2 Insoluble Insoluble Mg(OH)2 Insoluble 0.012 Ca(OH)2 Sparingly soluble 1.2 Sr(OH)2 Soluble 10.0 Ba(OH)2 Soluble 47.0 10 An aqueous suspension of magnesium hydroxide (milk of magnesia) is used as antacid for the treatment of gastrointestinal discomfort by neutralising the acid in the stomach. 2018/P2/Q18(a) Reduction of water to hydrogen. Exam Tips Exam Tips Reactivity towards water increases as the reducing power of the elements increases. Beryllium hydroxide is amphoteric. Solubility Mr of M(OH)2 Quick Check 10.2 1 Explain why aqueous sodium hydroxide is not used as antacid? 2 An aqueous solution of barium hydroxide turns cloudy when exposed to air for sometimes. Explain with the aid of an equation what happens. 3 Explain why calcium hydroxide or calcium oxide are sometimes added to agricultural soils. Info Chem Barium hydroxide reacts with carbon dioxide to form insoluble barium carbonate. INFO Reactions of the Group 2 Elements with Water

Chemistry Term 2 STPM Chapter 10 Group 2 149 10 Reactions with Oxygen 1 All Group 2 elements burn in oxygen (air), when heated, to form oxides. 2M(s) + O2(g) → 2MO(s) For example, 2Mg(s) + O2(g) → 2MgO(s) 2Ba(s) + O2(g) → 2BaO(s) 2 All Group 2 oxides are white solids that dissolve in water to form their respective hydroxides which are alkaline (except for beryllium hydroxide which is amphoteric). MO(s) + H2O(l) → M(OH)2(aq) For example, CaO(s) + H2O(l) → Ca(OH)2(aq) BaO(s) + H2O(l) → Ba(OH)2(aq) 3 Strontium and barium can also react with oxygen to form the peroxides. Sr(s) + O2(g) → SrO2(s) Ba(s) + O2(g) → BaO2(s) 4 The formula for the peroxide ion is O2 2-. Lewis diagram for the peroxide ion is as follows: O – – O 5 The peroxides dissolve in water to produce hydrogen peroxide. SrO2(s) + 2H2O(l) → Sr(OH)2(aq) + H2O2(aq) BaO2(s) + 2H2O(l) → Ba(OH)2(aq) + H2O2(aq) Thermal Decomposition of the Carbonates, Nitrates and Hydroxides 1 All Group 2 elements form carbonates, nitrates and hydroxides which are stable at room conditions. 2 All the above salts are white solids. 3 However, they decompose to their respective oxides when heated. For example, BeCO3(s) → BeO(s) + CO2(g) Δ 2Mg(NO3)2(s) → 2MgO(s) + 4NO2(g) + O2(g) Δ Brown fumes Ba(OH)2(s) → BaO(s) + 2H2O(g) Δ 4 All these show that the oxides are more stable than their respective carbonates, nitrates and hydroxides. 2011/P2/Q3(a)(iii) Info Chem Magnesium burns with a brilliant white flame. 2013/P2/Q11 2016/P2/Q9; Q20(b) 2014/P2/Q20(b) 2017/P2/Q8; Q17(a)

Chemistry Term 2 STPM Chapter 10 Group 2 150 10 5 This is because the O2– ion with its small size and higher charge can approach the cations more closely than the larger NO3 – , CO3 2– or OH– ions. This increases the strength of the ionic bonds in the oxides making them more stable. CO3 2– 2+ O2– 2+ 6 The lattice energy of the oxides is more exothermic than the corresponding nitrates, carbonates and hydroxides. 7 The energy profile for the thermal decomposition of magnesium hydroxide is shown below: Energy Mg(OH)2(s) MgO(s) + H2O(g) 8 However, no appreciable decomposition occurs at room conditions. This is because of the high activation energy involved. At room conditions, the particles do not have enough energy to overcome the high activation energy. 9 On heating, the kinetic energy of the particles increases. At high enough temperatures, they can overcome the activation energy and decomposition occurs. 10 For example, the decomposition of the carbonates. MCO3(s) → MO(s) + CO2(g) Δ 11 First of all, we have to know how the decomposition occurs. 12 The electron cloud of the large CO3 2– ion is polarised by the small and highly charged cation: O– O O– CO3 C 2– 2+ 2+ Polarisation ⎯⎯⎯→ The oxide is more stable than the carbonates, nitrates or hydroxides. Activation energy of decomposition is high. Polarisation of the carbonate ion by the small M2+ cation Weakening of the carbon-oxygen bond INFO Relative Thermal Stability of Group 2 Carbonates 2009/P1/Q20 2012/P2/Q8(a) 2011/P2/Q3(a)(i), (ii) 2014/P2/Q9

Chemistry Term 2 STPM Chapter 10 Group 2 151 10 13 The polarisation weakens the carbon-oxygen bonds in the CO3 2– ion. On heating, one of the carbon-oxygen bonds breaks to produce the O2– ion and carbon dioxide. 14 As a result, the relative stability of the carbonates depends on the extent of polarisation of the anion by the +2 cations. 15 Going down Group 2 from BeCO3 to BaCO3, the size of the cation increases. Ion Be2+ Mg2+ Ca2+ Sr2+ Ba2+ Ionic radius/nm 0.031 0.065 0.099 0.113 0.135 Charge density/charge nm–1 64.5 30.8 20.2 17.7 14.8 As a result, the charge density ( Charge Atomic radius ) and the polarising power decreases down the group. The extent of the polarisation of the CO3 2– ion decreases, making the carbonates more stable. 16 This is evident in the decomposition temperature of the carbonates given in the table below. Carbonate BeCO3 MgCO3 CaCO3 SrCO3 BaCO3 Temperature/°C 97 197 897 1277 1357 80 100 120 Mr of carbonates 140 160 Temperature/°C 1200 800 400 1400 1000 600 200 60 180 200 BeCO3 MgCO3 CaCO3 SrCO3 BaCO3 17 The same trend applies to the decomposition of the nitrates and hydroxides. 18 Beryllium carbonate is usually kept under an atmosphere of carbon dioxide to reduce its degree of decomposition by forcing the following equilibrium to the left-hand side. BeCO3(s) BeO(s) + CO2(g) Charge density of the cation decreases down the group. 2014/P2/Q19(b) VIDEO Thermal Decomposition of Calcium Nitrate

Chemistry Term 2 STPM Chapter 10 Group 2 152 10 19 The thermal stability of the anions also depends on the polarisibility of the anion. The larger the anion, the easier it is to be polarised because the electron cloud is not as tightly held by the nucleus. 20 The relative thermal stability of the salts of the same element increases in the order: CO3 2– < NO3 – < OH– Solubility of the Group 2 Sulphates 1 The dissolution of the Group 2 sulphates is represented by the general equation: MSO4(s) + aq → M2+(aq) + SO4 2–(aq) 2 The dissolution process can be thought of as comprising the following steps (although in the real sense it is not necessarily so). (a) Breaking of the ionic bonds holding the cations and anions in the solid structure MSO4(s) → M2+(g) + SO4 2–(g) ∆H1 = – Lattice energy The energy involved is the reverse of the lattice energy. (b) The gaseous ions dissolve in water to form hydrated ions. M2+(g) + aq → M2+(aq) ∆H2 = Hydration energy SO4 2–(g) + aq → SO4 2–(aq) The energy released is the hydration energies of the cation and the anion. (c) The total enthalpy change for the process, ∆H3: MSO4(s) + aq → M2+(aq) + SO4 2–(aq) ∆H3 = Heat of solution is given by Hess' law: – Lattice energy MSO4(s) → M2+(g) + SO4 2–(g) ∆H3 Hydration energy M2+(aq) + SO4 2–(aq) ∆H3 = (– Lattice energy) + (Hydration energy) or = (Lattice dissociation energy) + (Hydration energy) 3 The relative solubility of the sulphates depends on the magnitude of the enthalpy change of solution, ∆H3, which in turn is dependent on the lattice energy and the hydration energy. 4 Salts with negative values for ∆H3 are generally more soluble than those with positive ∆H3 values. The smaller the anion, the more difficult it is to be polarise. ∆H, is sometimes called lattice dissociation energy. It has a positive value. 2007/P1/Q20 2012/P2/Q8(b) 2015/P2/Q11 INFO Relative Solubility of Group 2 Sulphates

Chemistry Term 2 STPM Chapter 10 Group 2 153 10 Lattice Energy of the Group 2 Sulphates 1 Lattice energy of an ionic solid is defined as the heat released when one mole of the solid is formed from its gaseous ions under standard conditions. For example, Mg2+(g) + O2–(g) → MgO(s) ∆H = –3889 kJ mol–1 2 Lattice energy is a measure of the strength of the ionic bond in the solid. The stronger the ionic bond, the more exothermic is the lattice energy. 3 The lattice energy of an ionic compound is proportional to: Q+Q– (r+ + r– )2 Q+ = charge on the cation Q– = charge on the anion r+ = radius of the cation r– = radius of the anion 4 For a particular series of Group 2 compounds, the size of the cation increases down the group. This leads to an increase in the inter-ionic distance, (r+ + r– ) causing a decrease in the electrostatic attraction between the cations and the anions. Hence, the lattice energy decreases (becoming less exothermic). For example, Compound Lattice energy/kJ mol–1 BeO –4443 MgO –3889 CaO –3510 SrO –3310 BaO –3512 Hydration Energy of the Group 2 Cations 1 Hydration energy is the energy released when one mole of gaseous ions dissolve in water to form hydrated ions under standard conditions. For example, Mg2+(g) + aq → Mg2+(aq) ∆H = –1921 kJ mol–1 2 The energy released is the result of ion-dipole attraction between the cations and water molecules. 3 The stronger the ion-dipole attraction, the more exothermic is the hydration energy. 4 Going down Group 2, the size of the cations increases. As a result, the ion-dipole attraction decreases causing a decrease in the hydration energy. The lattice energies of the Group 2 sulphates are not available. Info Chem From BeO to BaO, the lattice energy decreases by about 21%.

Chemistry Term 2 STPM Chapter 10 Group 2 154 10 2+ δ+ δ– δ+ δ– δ+ δ– δ+ δ– δ+ δ– δ+ δ– δ δ– + δ+ δ– – – – – – – – – – – – – – – – – – – – – – – – – H2O molecule Ion-dipole attraction Ion Be2+ Mg2+ Ca2+ Sr2+ Ba2+ Ionic radius/nm 0.031 0.065 0.099 0.113 0.135 ∆Hhydration /kJ mol–1 –2494 –1921 –1577 –1443 –1305 Relative Solubility of the Group 2 Sulphates 1 Going down Group 2, the lattice energy and the hydration energy of the sulphates decreases. 2 However, due to the very large size of the SO4 2– ion, the increase in the interionic distance in the sulphates is slight. As a result, the lattice energy decreases only slightly. 3 On the other hand, the increase in the size of the cations causes a larger decrease in the hydration energy. Energy Lattice energy Hydration energy Group 2 sulphates 4 As a result, the enthalpy of solution changes gradually from exothermic to endothermic. This causes a decrease in the solubility of the sulphates down the group. Sulphate BeSO4 MgSO4 CaSO4 SrSO4 BaSO4 ∆Hsolution /kJ mol–1 – –91.2 –17.8 –8.7 +19.4 Solubility/g per 100 g H2O 41.0 36.4 0.21 0.010 0.00025 Hydration energy decreases more than the lattice energy. The hydration energy for BeSO4 is not available. Info Chem From Be2+ to Ba2+, the hydration energy decreases by about 48%. 2018/P2/Q18(a)

Chemistry Term 2 STPM Chapter 10 Group 2 155 10 120 140 160 Mr of sulphates 180 200 Solubility/(g/100 g H2O) 40 20 30 10 0 100 220 240 BeSO4 MgSO4 CaSO4 SrSO4 BaSO4 5 BeSO4 and MgSO4 are soluble in water. CaSO4 is sparingly soluble. SrSO4 and BaSO4 are insoluble. 6 The very low solubility of BaSO4 is employed as a test for the sulphate ions. Addition of aqueous barium chloride will cause an immediate precipitation of BaSO4, a white solid. 7 Aqueous Ba2+ ions are highly poisonous. However, due to the very low solubility of BaSO4, an aqueous suspension of BaSO4 can be safely ingested as ‘barium meal’ in the X-ray investigation of the stomach or other parts of the digestive system. Example 10.1 Draw a comparative energy profile for the dissolution of MgSO4 and BaSO4. SolutionEnergy Mg2+(g) + SO4 2–(g) Ba2+(g) + SO4 2–(g) Mg2+(aq) + SO4 2–(aq) Ba2+(aq) + SO4 2–(aq) MgSO4(s) –∆HLattice –∆HLattice ∆HHydration ∆HHydration BaSO4(s) –91.2 kJ mol–1 +19.4 kJ mol–1 Test for SO4 2– ion Exam Tips Exam Tips BaSO4 does not react with the dilute HCl in the stomach.

Chemistry Term 2 STPM Chapter 10 Group 2 156 10 10.2 Anomalous Behaviour of Beryllium 1 Due to the very small size of the beryllium atom coupled with the fact that it has only two inner shell electrons, the ionisation energy of beryllium is exceptionally high. Element Be Mg Ca Sr Ba M → M2+/kJ mol–1 2700 2190 1740 1610 1470 10 20 30 Mr of element 40 50 M(g) M2+(g)/kJ mol–1 2200 1800 2000 1600 2800 2400 2600 1400 0 Be Mg Ca Sr Ba 2 As a result, most of the compounds of beryllium are covalent, or are ionic with a significant amount of covalent character. 3 The Be2+ ion, with its high charge density is able to polarise any anion that it is bonded to, to give covalent character to the compound. 4 For example, beryllium chloride is a covalent compound that sublimes when heated, while chlorides of the other members of the group are ionic. • fi • Cl Be fi Cl fifi fifi fifi fifi fi fi fi fi 2+ Anion In the solid state, beryllium chloride exists as a polymeric chain. Cl Cl Be Cl Cl Be Cl Cl Be Be The beryllium atom: Nucleus 2015/P2/Q9 2016/P2/Q8 2017/P2/Q9 INFO Beryllium Chloride

Chemistry Term 2 STPM Chapter 10 Group 2 157 10 When added to water, beryllium chloride undergoes hydrolysis to produce an acidic solution. BeCl2(s) + 2H2O(l) Be(OH)2(aq) + 2H+(aq) + 2Cl– (aq) 5 Beryllium oxide and beryllium hydroxide are amphoteric. BeO(s) + 2HCl(aq) → BeCl2(aq) + H2O(l) BeO(s) + 2NaOH(aq) + H2O(l) → Na2Be(OH)4(aq) Sodium beryllate Similarly, Be(OH)2(s) + H2SO4(aq) → BeSO4(aq) + 2H2O(l) Be(OH)2(s) + 2NaOH(aq) → Na2Be(OH)4(aq) 6 Due to the high charge density of the Be2+, it can form complexes, a property not shared by other members of the group. Examples of the complexes are: BeCl2•2NH3; BeF4 2–; Be(OH)4 2– Cl Cl NH3 NH3 Be2+ F F F F Be2+ Be2+ OH OH HO OH Diagonal Relationship 1 Certain pairs of diagonally adjacent elements in Period 2 and Period 3 of the Periodic Table exhibit similar chemical properties. This is called diagonal relationship. 2 Examples are between lithium and magnesium; between beryllium and aluminium and, between boron and silicon. Period 3 Mg Be 2 Al B 13 Si 14 Li 1 Period 2 3 The similarities between beryllium and aluminium are as below: (a) Beryllium oxide and hydroxide are amphoteric. Aluminium oxide and hydroxide are also amphoteric. (b) Beryllium chloride and aluminium chloride are covalent compounds that sublime on heating. (c) Beryllium and aluminium form complexes. (d) Aqueous solutions of beryllium salts and aluminium salts are acidic. Complex formation. Info Chem Info Chem Info Chem In the vapour state, beryllium chloride forms dimer (Be2Cl4) at 410 °C and monomer at 900 °C. Be2Cl4(g)  2BeCl2 Cl Cl – Be Be – Cl  2BeCl2 Cl Beryllium and aluminium dissolve in alkalis with the release of hydrogen. Be + 2OH– + 2H2O → Be(OH)4 2– + H2 2Al + 2OH– + 6H2O → 2Al(OH)4 – + 3H2 Aqueous beryllium salts and aqueous aluminium salts are acidic due to the hydrolysis of the Be2+ and Al3+ aqueous ion: [Be(H2O)4] 2+(aq)  [Be(OH)(H2O)3] +(aq) + H+(aq) [Al(H2O)6] 3+(aq)  [Al(OH)(H2O)5] 2+(aq) + H+(aq) 2016/P2/Q10 2017/P2/Q17(b); Q20(c) 2013/P2/Q17

Chemistry Term 2 STPM Chapter 10 Group 2 158 10 4 These similarities come about because both beryllium and aluminium have the same electronegativity and their ions have similar charge densities. Element Aspect Aluminium Beryllium Electronegativity 1.5 1.5 Size of ion/nm 0.050 0.033 Charge density of ion/charge nm–1 60 60.6 5 Going across a period in the Periodic Table, the atomic radius decreases while the electronegativity increases. 6 Going down a group, the atomic radius increases while the electronegativity decreases. 7 Hence, going from beryllium to aluminium, the effects cancel out one another causing both elements to have the same electronegativity and charge density. Mg Be 2 Al B 13 Si 14 Li 1 Atomic radius decreases Electronegativity increases Electronegativity decreases Atomic radius increases 8 The table below lists the similarities between the pair, lithium and magnesium. Lithium Magnesium Lithium forms lithium nitride by direct combination with nitrogen at high temperatures. 6Li + N2 → 2Li3N Magnesium forms magnesium nitride by direct combination with nitrogen at high temperatures. 4Mg + 3N2 → 2Mg2N3 Lithium carbonate and nitrate decompose to lithium oxide on heating. Li2CO3 → Li2O + CO2 4LiNO3 → 2Li2O + 4NO2 + O2 Magnesium carbonate and nitrate decompose to magnesium oxide on heating. MgCO3 → MgO + CO2 2Mg(NO3)2 → 2MgO + 4NO2 + O2 INFO Anomalous Properties of Beryllium

Chemistry Term 2 STPM Chapter 10 Group 2 159 10 SUMMARY SUMMARY 10.3 Uses of Group 2 Compounds Compound Uses Magnesium oxide Refractory material in high temperature furnace Magnesium hydroxide (Milk of magnesia) As an antacid to relieve indigestion Magnesium sulphate As a laxative Calcium sulphate ‘Plaster of Paris’ Calcium hydroxide To neutralise acidic in soil Calcium carbonate Making writing chalks Barium carbonate Rat poison Barium sulphate As ‘barium meal’ in the X-ray of the digestive system Objective Questions 1 Which of the following statements regarding the Group 2 elements and their compounds is not true? A They are known as the alkaline metals. B Beryllium oxide is the only amphoteric oxide in Group 2. C The melting point of calcium is higher than that of magnesium. D Barium oxide is soluble in water. 2 Going down Group 2 of the Periodic Table, A the maximum oxidation state of the elements increases. B the polarising power of the cations decreases. C the electronegativity of the elements decreases. D the standard electrode potential becomes more negative. 3 Beryllium oxide is amphoteric. This is because A beryllium is a metalloid B Be2+ is unstable C Be2+ has a high charge density D Be2+ forms the [Be(OH)4]2- complex ion 1 The Group 2 elements are also known as the ‘alkaline earth metals’. 2 They have two electrons in their valence shell. 3 They exhibit a constant valence of +2 in all their compounds. 4 They are powerful reducing agents. The reducing strength increases down the group. 5 Their salts are white except when in combination with the oxo-anions of transition elements. 6 The carbonates, nitrates and hydroxides decompose when heated to form the corresponding oxides. 7 The thermal stability of the salts increases down the group. 8 The solubility of the sulphates decreases down the group. 9 Beryllium shows some properties that are not typical of the Group 2 elements: • Beryllium oxide and beryllium hydroxide are amphoteric. • Beryllium chloride is covalent. • Beryllium forms complexes. STPM PRACTICE 10

Chemistry Term 2 STPM Chapter 10 Group 2 160 10 4 Which of the following is not true? A All nitrates of Group 2 elements decompose on heating to release brown fumes of nitrogen dioxide. B The carbonates of Group 2 elements are more stable to heat compared to their corresponding hydroxides. C Beryllium carbonate is usually kept under an atmosphere of carbon dioxide to prevent its decomposition. D Barium hydroxide is soluble in water but barium sulphate is not. 5 Radium is below barium in Group 2 of the Periodic Table. Which statement about radium and its compounds is incorrect? A Radium is an alkaline earth metal. B Radium carbonate decomposes spontaneously at room temperature. C Radium sulphate is insoluble in water. D Radium reacts with dilute hydrochloric acid to liberate hydrogen gas. 6 Which of the following best explains why CaCO3 decomposes at a higher temperature than MgCO3. A MgO is less stable than CaO. B Ca2+ has more electrons than Mg2+. C Ca2+ is larger than Mg2+. D Ca is more reactive than Mg. 7 Among the carbonates of Group 2, beryllium carbonate is the least stable towards heat. This is because A the carbonate ion is polarised by the beryllium(II) ion B beryllium is amphoteric C beryllium oxide is more stable than beryllium carbonate D beryllium(II) ion has the smallest size 8 Aqueous barium(II) ions are highly poisonous. However, barium sulphate can be safely ingested (as in barium meal for X-ray purpose) but not barium carbonate. This is because A barium sulphate is not soluble in water. B barium sulphate is covalent but barium carbonate is ionic. C barium carbonate dissolves in the stomach but not barium sulphate. D barium carbonate is soluble in water. 9 Which of the following is true of the Group 2 elements? A All the sulphate are insoluble in water. B All salts of Group 2 elements decompose on heating. C All the elements are powerful reducing agents. D The lattice energy of the compounds increases down the group. 10 The standard enthalpy of hydration of the Group 2 cations decreases down the group. This is because A the size of the atoms increases. B the ionisation energy decreases. C the electronegativity of the elements decreases. D the charge density of the cations decreases. 11 Magnesium burns in air with a brilliant white flame. When the white residue is warmed with water, a gas that turns moist red litmus blue is liberated. What conclusion can be drawn from this observation? What is the composition of the white residue? A Magnesium oxide and magnesium nitride B Magnesium oxide and magnesium nitrate C Magnesium nitride only D Magnesium oxide only 12 Which of the following is not true about the Group 2 compounds? A Magnesium nitrate is thermally more stable than magnesium carbonate. B Barium hydroxide is more soluble than calcium hydroxide. C Strontium sulphate is less soluble than magnesium sulphate. D An aqueous solution of beryllium chloride is acidic. 13 The carbonates, nitrates and hydroxides of Group 2 elements decompose to their respective oxides when heated. Hence, it can be concluded that A all compounds of Group 2 elements are unstable to heat B the oxide is energetically more stable than the carbonates, nitrates and hydroxides C Group 2 elements are strong reducing agents D the metallic properties of the elements increases down the group 14 Which of the following is not true about beryllium chloride? A It has a bent structure. B It is covalent. C It exists as a dimer in the solid state. D It dissolves in water to form Be2+ and Cl- ions.

Chemistry Term 2 STPM Chapter 10 Group 2 161 10 15 Which of the following elements is most likely to form some covalent compounds? A Beryllium C Potassium B Calcium D Magnesium 16 Which of the following statements is not true? A All carbonates of Group 2 elements are decomposed by heat to give carbon dioxide. B Beryllium hydroxide is amphoteric. C Beryllium reacts with alkali to release hydrogen gas. D All sulphates of Group 2 elements are insoluble in water. 17 Which of the following properties of the Group 2 elements and their compounds increases with increasing proton number? A The stability of the nitrate to heat B The tendency to form a complex C The magnitude of the hydration energy D The solubility of the sulphates 18 The Group 2 elements are powerful reducing agents. Which property accounts for the increasing reducing power of the elements down the group? A Standard reduction potential B Ionisation energy C Electronegativity D Enthalpy of atomisation 19 Beryllium carbonate, BeCO3, decomposes when heated to 97 °C; whereas barium carbonate, BaCO3, decomposes at 1357 °C. Which statement best explains why beryllium carbonate is thermally less stable than barium carbonate? A Ba2+ ion is smaller than Be2+ ion. B BeCO3 is ionic while BaCO3 is covalent. C The C—O bond in BeCO3 is weaker than the C—O bond in BaCO3. D The lattice energy of BeCO3 is more negative than that of BaCO3. 20 Strontium is situated above barium in Group 2 of the Periodic Table. Which of the following statements is true? A The strontium ion has a lower charge density than the barium ion. B Strontium hydroxide is more stable to heat compared to barium hydroxide. C Strontium sulphate is more soluble than barium sulphate. D Strontium is a more powerful reducing agent than barium. 21 Which statement best explains why the thermal stability of the nitrates of Group 2 elements increases down the group? A The melting point of the nitrates increases. B The size of the Group 2 metal ions increases. C The first ionisation decreases. D The size of the nitrate ion decreases. 22 Which of the following factors best explain the trend in the decomposition temperature of the Group 2 nitrates? A Charge density of the cations B Size of the anion C Lattice energy D Size of the atom Structured and Essay Questions 1 (a) Describe and explain the reactivity of the Group 2 elements (Be → Ba) with water. (b) The solubility of four Group 2 sulphates is given below: Salt MgSO4 CaSO4 SrSO4 BaSO4 Solubility / g per 100 g water 36.4 0.210 0.0100 0.000254 Explain the trend of the solubility of the sulphates.

Chemistry Term 2 STPM Chapter 10 Group 2 162 10 2 Explain the following observations. (a) The thermal decomposition temperature of the Group 2 carbonates decreases with the proton number of the Group 2 elements. (b) Calcium hydroxide decomposes at a higher temperature than calcium carbonate. (c) Magnesium sulphate is soluble in water, but barium sulphate is not. 3 (a) Describe and explain the trend in the ease of thermal decomposition of the Group 2 carbonates. (b) Which is more stable to heat, magnesium carbonate or magnesium hydroxide? Explain your answer. 4 (a) Beryllium has certain properties not shared by other members of the Group 2. For example, beryllium chloride is a white solid that fumes in moist air. Write a balanced equation for the reaction of beryllium chloride with moist air. (b) When beryllium chloride dissolves in cyclohexane, its relative molecular mass was found to be 160. (i) What is the nature of the intermolecular forces between beryllium chloride and cyclohexane? (ii) Determine the molecular formula of beryllium chloride in cyclohexane. (iii) Draw the structure for beryllium chloride in cyclohexane. 5 Beryllium chloride is prepared by heating beryllium with chlorine gas. In the solid state, anhydrous beryllium chloride exists as a polymeric chain with the following structure: Cl Cl Cl Be Be Be Cl Cl Cl (a) Write an equation for the production of beryllium chloride. (b) In the vapour state, beryllium chloride exists in the dimeric form at 405 °C and in the monomer form at temperature exceeding 900 °C. Draw the structures for the monomer and dimer of beryllium chloride and state their shapes. (c) Beryllium chloride reacts vigorously with water at room temperature. (i) State two observations for the reaction. (ii) Write an equation for the hydrolysis of beryllium chloride. (d) Explain why beryllium chloride is covalent whereas the other Group 2 chlorides are ionic. 6 (a) When magnesium nitrate is heated to 200 °C, a white residue formed and a brown gas is evolved. (i) Write an equation for the decomposition of magnesium nitrate and name the white residue and the brown gas formed. (ii) Would you expect the decomposition temperature of barium nitrate to be higher or lower than 200 °C? Explain your answer. (b) When heated using bunsen flame, magnesium carbonate gives out a gas that turns limewater chalky, whereas barium carbonate does not. Explain this observation as fully as you can.

CHAPTER 11 GROUP 14 Concept Map Learning earning Outcomes Students should be able to: Physical properties of Group 14 elements • explain the trends in physical properties (melting points and electrical conductivity) of Group 14 elements: C, Si, Ge, Sn, Pb. Tetrachlorides and oxides of Group 14 elements • explain the bonding and molecular shapes of the tetrachlorides of group 14 elements; • explain the volatility, thermal stability and hydrolysis of tetrachlorides in terms of structure and bonding; • explain the bonding, acid-base nature and the thermal stability of the oxides of oxidation states +2 and +4. Relative stability of +2 and +4 oxidation states of Group 14 elements • explain the relative stability of +2 and +4 oxidation states of the elements in their oxides, chlorides and aqueous cations. Silicon, silicone and silicates • describe the structures of silicone and silicates (pyroxenes and amphiboles), sheets (mica) and framework structure (quartz) (general formulae are not required); • explain the uses of silicon as a semiconductor and silicone as a fluid, elastomer and resin; • describe the uses of silicates as basic materials for cement, glass, ceramics and zeolites. Tin alloys • describe the uses of tin in solder and pewter. Group 14 Elements Variation in Physical Properties Tetrachlorides of Group 14 Elements • General • Physical properties • Hydrolysis Uses of Silicon and its Compounds Silicates • Chain • Sheet • Framework and giant structure Oxides of Group 14 Elements Chemical Properties Zeolites Glass Ceramics Uses of Tin

Chemistry Term 2 STPM Chapter 11 Group 14 164 11 11.1 Physical Properties of Group 14 Elements 1 The Group 14 elements consists of carbon, silicon, germanium, tin and lead. 1 2 Ge C Si Sn Pb 13 14 15 16 17 18 2 Some basic informations of the Group 14 elements are listed in the table below. Element Carbon Silicon Germanium Tin Lead Symbol C Si Ge Sn Pb Proton number 6 14 32 50 82 Electronic configuration 2.4 [He]2s2 2p2 2.8.4 [Ne]3s2 3p2 2.8.18.4 [Ar]4s2 4p2 2.8.18.18.4 [Kr]5s2 5p2 2.8.18.32.18.4 [Xe]6s2 6p2 3 They are all p-block elements with the valence shell electronic configuration of ns2 np2 . ns np 4 Group 14 consists of a mixture of metals (tin and lead), non-metal (carbon) and metalloids (silicon and germanium). Hence, unlike those of Group 2 elements, there are significant differences in their properties. Antomic Radius and Ionisation Energy 1 The table below shows the atomic radius and ionisation energy of Group 14 elements. Element C Si Ge Sn Pb Atomic radius/nm 0.077 0.114 0.122 0.140 0.154 1 st ionisation energy/kJ mol–1 1090 970 760 710 720 Info Chem Tin is also known as stanum. Lead is also known as plumbum. Info Chem Carbon is the only element in the group to exhibit catenation.

Chemistry Term 2 STPM Chapter 11 Group 14 165 11 10 20 30 Proton number 40 50 Atomic radius (fi 10–3)/nm 100 90 80 70 60 140 130 120 110 150 0 60 70 C Si Ge Sn Pb 80 700 800 900 1000 1100 10 20 30 40 50 60 70 80 Proton number I.E. / kJ mol–1 Info Chem The first ionisation energy of lead is slightly higher than that of tin.

Chemistry Term 2 STPM Chapter 11 Group 14 166 11 2 Going down the group, the effective nuclear charge remains almost constant, but each successive element has one shell extra than the preceding ones. Element C Si Ge Sn Pb No. of protons 6 14 32 50 82 Electronic configuration 2.4 2.8.4 2.8.18.4 2.8.18.18.4 2.8.18.32.18.4 Effective nuclear charge 4 4 4 4 4 3 This causes the attraction between the nucleus and the electron cloud to get progressively weaker. As a result, the atomic radius increases. 4 As the atomic radius increases, the first ionisation energy of the elements generally decreases. 5 However, the first ionisation energy of lead is slightly higher than that of tin although the lead atom is larger than the tin atom. This is because the screening effect does not increase as much as expected when going from tin to lead. 6 Going from tin to lead, there is an increase of 32 protons and 32 electrons. Sn: [Ar] 3d10 4s 2 4p6 4d10 5s 2 5p2 Pb: [Ar] 3d10 4s 2 4p6 4d10 4f 14 5s 2 5p6 5d10 6s 2 6p2 7 However, the screening effect provided by the 4f electrons is less than expected. As a result, the increase in the nuclear charge is not totally being balanced by the increase of additional 32 electron. 8 As a result, the effective nuclear charge increases slightly from tin to lead leading to an increase in the first ionisation energy. Electrical Conductivity Electrical conductivity Proton number (Diamond) Pb Sn Ge Si C 1 Going down the group, the properties of the elements change from non-metallic to metallic. As a result, there is a corresponding increase in the electrical conductivity. Info Chem Effective nuclear charge does not change much but the number of shells increases. The 4f electrons have poor screening effect. 2012/P2/Q7(a)

Chemistry Term 2 STPM Chapter 11 Group 14 167 11 2 Carbon (in the form of diamond) is a non-conductor because there are no delocalised electrons in the giant covalent structure. 3 Silicon and germanium are metalloids. They are semiconductors. 4 Tin and lead have delocalised electrons in their giant metallic structures. They are good conductors of electricity. Melting Point 1 Carbon (diamond) has a giant covalent structure with strong covalent bonds holding the individual atoms together in a three dimension array. A lot of energy is required to break the covalent bonds. As a result, it has a very high melting point. Element C (diamond) Si Ge Sn Pb Structure Giant covalent Giant metallic Melting point/°C 3730 1410 937 232 327 500 1000 1500 2000 2500 3000 3500 4000 10 20 30 40 50 60 70 80 Proton number Melting point / °C C Si Ge Sn Pb 2 Silicon and germanium also have a structure like diamond. But due to their larger size, the covalent bonds are significantly weaker than those in diamond. Their melting points are lower than that of diamond. Due to its larger size, the melting point of germanium is lower than that of silicon. 2015/P2/Q19(c) 2016/P2/Q11 2017/P2/Q10 C, Si or Ge Info Chem Another allotrope of carbon, graphite, is a conductor. Info Chem The melting point of Pb is slightly higher than Sn.

Chemistry Term 2 STPM Chapter 11 Group 14 168 11 3 Tin and lead have giant metallic structures. Due to their large size, the metallic bonds are relatively weak. This accounts for their low melting points compared to those of carbon, silicon and germanium. 4 Solid lead has a close-packed structure while that of tin is more open. That is why the melting point of lead is higher than that of tin even though the atomic radius of lead is larger than that of tin. 11.2 Tetrachlorides and Oxides of Group 14 Elements 1 All Group 14 elements form tetrachlorides which are simple molecules with the general formula of MCl4. 2 They are all colourless liquids at room conditions. 3 The Lewis diagram and shape of the tetrachloride are as shown: Cl Cl Cl Cl x x M x x Cl Cl Cl M Cl The central atom undergoes sp3 hybridisation to give a bond angle of 109.5°. 4 The tetrachlorides of silicon, germanium, tin and lead can be prepared by direct combination, by heating the element in a stream of dry chlorine gas. Si(s) + 2Cl2(g) ⎯→ SiCl4(l) ∆ Ge(s) + 2Cl2(g) ⎯→ GeCl4(l) ∆ Sn(s) + Cl2(g) ⎯→ SnCl4(l) ∆ 5 Carbon tetrachloride is prepared by passing dry chlorine gas into liquid carbon disulphide boiling under reflux in the presence of a little iodine as catalyst. 3Cl2 + CS2(l) ⎯→ CCl4(l) + S2Cl2(l) ∆ 6 Lead(IV) chloride is prepared by reacting lead dioxide with concentrated hydrochloric acid at 5 °C. PbO2(s) + 4HCl(aq) ⎯→ PbCl4(l) + 2H2O(l) 2012/P2/Q3(b) 2015/P2/Q10 2017/P2/Q11 2013/P2/Q12 2016/P2/Q12 2018/P2/Q10 Direct combination Preparation of CCl4 Preparation of PbCl4 Info Chem When lead is heated with chlorine, lead(II) chloride is formed. Pb(s) + Cl2(g) → PbCl2(s)

Chemistry Term 2 STPM Chapter 11 Group 14 169 11 Boiling Point of the Tetrachlorides 1 The bonds that hold the atoms together in the tetrachloride molecule are covalent bonds, but the intermolecular forces are weak van der Waals forces. van der Waals force MCl4 MCl4 2 All of them are colourless liquids at room conditions. 3 Going down the group, the size and the total number of electrons in the molecule increase. The van der Waals forces get stronger. As a result, the boiling points of the tetrachlorides increase down the group. Tetrachloride CCl4 SiCl4 GeCl4 SnCl4 PbCl4 No. of electrons 74 82 100 118 150 Melting point/°C –23* –70 –50 –33 –15 Boiling point/°C 77* 59 86 114 Decomposes [* exceptional behaviour of CCl4 that has no apparent explanation] 50 60 70 80 90 100 110 120 CCl4 GeCl4 SnCl4 SiCl4 150 160 170 180 190 200 210 220 Relative molecular mass of MCl4 Boiling point/ o C 4 There is no normal boiling point for lead tetrachloride because it decomposes on heating. PbCl4(l) ⎯→ PbCl2(s) + Cl2(g) ∆ 2011/P1/Q22 Info Chem The melting point and boiling point of CCl4 are higher than expected.

Chemistry Term 2 STPM Chapter 11 Group 14 170 11 5 Other than CCl4, the boiling points of the tetrachlorides show a gradual increase down the group. 6 Even though carbon tetrachloride has the smallest size, its melting point and boiling point are higher than expected. This is an abnormal behaviour with no suitable explanation. Thermal Stability of the Tetrachlorides 1 The strength of a covalent bond to a large extent depends on the bond length. The longer the covalent bond, the weaker it is and is easier to break. 2 Going down Group 14, the M—Cl bonds become longer and weaker as the atomic radius of M increases. 3 As a result, the thermal stability of the tetrachlorides decreases down the group. 4 CCl4, SiCl4 and GeCl4 are stable to heat even at high temperatures. 5 SnCl4 undergoes decomposition on strong heating to produce tin(II) chloride with the liberation of chlorine gas. SnCl4(l) ⎯→ SnCl2(s) + Cl2(g) ∆ 6 In contrast, PbCl4 is so unstable that it undergoes partial decomposition even at room temperature. PbCl4(l) ⎯→ PbCl2(s) + Cl2 Hydrolysis of the Tetrachlorides 1 Hydrolysis is the reaction of a compound with water. 2 All the Group 14 tetrachlorides, except carbon tetrachloride, undergo hydrolysis with water to produce hydrochloric acid or hydrogen chloride gas depending on the amount of water used. 3 For example, SiCl4(l) + 2H2O(l) ⎯→ SiO2(s) + 4HCl(aq) SnCl4(l) + 2H2O(l) ⎯→ SnO2(s) + 4HCl(aq) 4 During hydrolysis, the silicon, germanium, tin and lead atoms of the respectively chlorides make use of their empty d-orbitals in their valence shells to form coordinate bonds with water molecules. 5 Using silicon tetrachloride as an example, the valence shell electronic configuration of silicon in SiCl4 is: Si atom in SiCl4 3s 3p 3d orbitals used to form coordinate bonds with water molecules O H 2 H : : Exam Tips The van der Waals forces between CCl4 melecules are stronger than that between SiCl4 molecules. MCl4 ⎯→ MCl2 + Cl2 ∆ 2009/P1/Q22 2008/P1/Q22 This is a convenient way to prepare pure silicon dioxide. 2014/P2/Q11

Chemistry Term 2 STPM Chapter 11 Group 14 171 11 6 The silicon atom makes use of two of the empty 3d orbitals to accept lone pair electrons from two water molecules to form a hexa-valency intermediate. 7 The intermediate then decomposes to produce silicon(IV) oxide and hydrochloric acid. O H H Cl Cl Cl Cl Si O H H ⎯⎯→ SiO2 + 4HCl 8 The mechanism can be summarised as: SiCl4 + 2H2O ⎯→ SiCl4(H2O)2 SiCl4(H2O)2 ⎯→ SiO2 + 4HCl 9 On the other hand, the valence shell electronic configuration of carbon in CCl4 is: C atom in CCl4 2s 2p There are no empty d-orbitals in its valence shell to form coordinate bonds with water and is not hydrolysed by water. 10 Silicon, germanium, tin and lead can also make use of their empty d orbitals to form complexes such as [SnCl6]2– and [PbCl6]2–, while carbon cannot. OH2 OH2 Cl Cl Cl Cl Si •• •• Carbon has no d-orbitals of comparable energy. 2011/P2/Q8(c) Quick Check 11.1 1 Explain why carbon tetraiodide, CI4 is unstable. 2 In the preparation of lead(IV) chloride from lead(IV) oxide, the temperature of the reaction mixture must be maintain at about 5 °C. Explain why it is so. 3 When liquid lead(IV) chloride is added to cold water, a white precipitate is formed. Name the white precipitate and write a balanced equation for the reaction taking place. Thermal Stability of the Oxides 1 Group 14 elements form two series of oxides with oxygen. Monoxide, with the general formula of MO, and dioxide, with the general formula of MO2. 2 The oxidation state of the elements in the monoxide is +2, while that in the dioxide is +4. Info Chem SiO2 has a giant covalent structure. 2010/P2/Q8(a) 2014/P2/Q12, Q20(a) 2015/P2/Q19(a)

Chemistry Term 2 STPM Chapter 11 Group 14 172 11 3 The table below lists some of the properties of the monoxides of Group 14. Monoxide CO SiO GeO SnO PbO Physical state Gas Solid Structure Simple molecule Predominantly Ionic Thermal stability Form dioxide on heating in air Stable Acid/base nature Neutral Amphoteric 4 Monoxides of C, Si, Ge and Sn are converted to the dioxides on heating in air. 2CO(g) + O2(g) ⎯→ 2CO2(g) ∆ 2GeO(s) + O2(g) ⎯→ 2GeO2(s) ∆ 2SnO(s) + O2(g) ⎯→ 2SnO2(s) ∆ 5 Lead monoxide is stable to heat. 6 The table below lists some of the properties of the dioxides of Group 14. Dioxide CO2 SiO2 GeO2 SnO2 PbO2 Physical state Gas Solid Structure Simple molecule Giant covalent Intermediate between giant covalent and giant ionic Thermal stability Stable Unstable Acid/base nature Acidic Amphoteric 7 Dioxides of carbon, silicon, germanium and tin are stable to heat. 8 On the other hand, the brown lead dioxide decomposes to yellow lead monoxide on heating. 2PbO2(s) ⎯→ 2PbO(s) + O2(g) Brown ∆ Yellow Acid-base Nature of the Oxides 1 CO and SiO are neutral oxides. Monoxides of germanium, tin and lead are amphoteric. 2 This is because the metallic character of the elements increases with increasing proton number. 3 The monoxides dissolve in hot, dilute mineral acids. MO(s) + 2H+(aq) ⎯→ M2+(aq) + H2O(l) ∆ For example, GeO(s) + 2H+(aq) ⎯→ Ge2+(aq) + H2O(l) ∆ SnO(s) + 2H+(aq) ⎯→ Sn2+(aq) + H2O(l) ∆ PbO(s) + 2H+(aq) ⎯→ Pb2+(aq) + H2O(l) ∆ 2010/P2/Q8(a) Another oxide of lead is Pb3O4. It is called ‘red lead’. It is a mixture of Pb(II) and Pb(IV) oxide, PbO2 .2PbO. The Lewis diagram of CO: C O 2007/P2/Q8(b) 2013/P2/Q19(a) 2011/P2/Q8(a) 2014/P2/Q12 2016/P2/Q8

Chemistry Term 2 STPM Chapter 11 Group 14 173 11 4 Monoxides of germanium, tin and lead also dissolve in hot, concentrated alkali to form salts. MO(s) + 2OH– (aq) + H2O(l) ⎯→ M(OH)4 2–(aq) ∆ An alternative equation is: MO(s) + 2OH– (aq) ⎯→ MO2 2–(aq) + H2O(l) ∆ For example, GeO(s) + 2OH– (aq) + H2O(l) ⎯→ Ge(OH)4 2–(aq) ∆ Germanate(II) SnO(s) + 2OH– (aq) + H2O(l) ⎯→ Sn(OH)4 2–(aq) ∆ Stanate(II) PbO(s) + 2OH– (aq) + H2O(l) ⎯→ Pb(OH)4 2–(aq) ∆ Plumbate(II) 5 CO2 and SiO2 are acidic dioxides. They are covalent oxides. 6 Carbon dioxide dissolves in cold, dilute alkali to form carbonates. CO2(g) + 2NaOH(aq) ⎯→ Na2CO3(aq) + H2O(l) ∆ 7 Silicon dioxide dissolves slowly in hot, caustic alkali to form silicates. SiO2(s) + 2NaOH(aq) ⎯→ Na2SiO3(aq) + H2O(l) ∆ 8 The dioxides of germanium, tin and lead are amphoteric. This shows that the bonds in these dioxides are intermediate between that of ionic and covalent. 9 These amphoteric oxides dissolve in hot, dilute mineral acids. (a) Germanium dioxide and tin dioxide react with hot, concentrated hydrochloric acid to form the tetrachlorides. GeO2(s) + 4HCl(aq) ⎯→ GeCl4(l) + 2H2O(l) ∆ SnO2(s) + 4HCl(aq) ⎯→ SnCl4(l) + 2H2O(l) ∆ (Although the tetrachloride formed will undergo slight hydrolysis.) (b) Lead dioxide reacts with hot, concentrated hydrochloric acid to produce lead(II) chloride with the liberation of chlorine gas due to the decomposition of the unstable lead tetrachloride. PbO2(s) + 4HCl(aq) → PbCl2(s) + Cl2(g) + H2O(l) ∆ (c) Lead dioxide reacts with cold, concentrated hydrochloric acid to produce lead tetrachloride. PbO2(s) + 4HCl(aq) → PbCl4(l) + 2H2O(l) Info Chem The presence of hydrochloric acid suppresses the hydrolysis of the tetrachloride due to the presence of the common ion Cl– . For example, SnCl4(l) + 2H2O(l)  SnO2(s) + 4HCl(aq)

Chemistry Term 2 STPM Chapter 11 Group 14 174 11 10 The dioxides of germanium, tin and lead react with hot, concentrated alkali to form germanate(IV), stanate(IV) and plumbate(IV) respectively. MO2(s) + 2OH– (aq) + 2H2O(l) → M(OH)6 2–(aq) ∆ Alternative equation: MO2(s) + 2OH– (aq) → MO3 2–(aq) + H2O(l) ∆ For example, GeO2(s) + 2NaOH(aq) + 2H2O(l) → Na2Ge(OH)6(aq) ∆ SnO2(s) + 2NaOH(aq) + 2H2O(l) → Na2Sn(OH)6(aq) ∆ PbO2(s) + 2NaOH(aq) + 2H2O(l) → Na2Pb(OH)6(aq) ∆ 11.3 Relative Stability of +2 and +4 Oxidation States of Group 14 Elements 1 Group 14 elements, with the valence shell configuration of ns2 np2 can exhibit two oxidation states: +2 and +4. 2 The +2 oxidation state involves only the s 2 electrons, while the +4 oxidation state involves both the s 2 and p2 electrons. 3 The energy required to form the +2 cations is given below: Element C Si Ge Sn Pb M(g) → M2+(g) + 2e– /kJ mol–1 3440 2366 2302 2117 2166 The +2 oxidation states of carbon and silicon are covalent, while that of germanium is covalent with slight ionic characteristic. On the other hand, the +2 oxidations of tin and lead are predominantly ionic. 4 The energy required to form the +4 ions is given below: Element C Si Ge Sn Pb M(g) → M4+(g) + 4e– /kJ mol–1 14 270 9956 9992 8987 9326 The energies involved are very high. As a result, none of the Group 14 elements form simple +4 cations. The +4 oxidation states are all covalent. All +4 states are covalent. 2009/P2/Q7(a) 2012/P1/Q23 2014/P2/Q20(a)

Chemistry Term 2 STPM Chapter 11 Group 14 175 11 5 The +4 oxidation state involves the promotion of one of the paired s 2 electrons to one of the empty p orbitals. +2: s p Excitation +4: s p Energy Ground state Excited state Excitation ⎯⎯⎯→ This process needs an input of energy. 6 Hence, the relative stability of the +4 state depends on the amount of energy released when the atom forms four covalent bonds. 7 If the energy released is sufficient to compensate for the energy required for the excitation, then the +4 oxidation state will be stable. Otherwise, the +4 state will be unstable. Energy (+4 is stable) Excitation of electron Formation of 4 covalent bonds (+4 is unstable) +4 +4 +4 compound +4 compound +2 +2 8 Going down the group, the atomic radius increases and the strength of the covalent bonds formed gets weaker and less energy is released. As a result, the +4 oxidation state becomes less stable and the +2 oxidation state becomes more stable. Info Chem Going down the group: +2 becomes more stable +4 becomes less stable INFO Inert − Pair Effect in Group 14

Chemistry Term 2 STPM Chapter 11 Group 14 176 11 Relative stability +2 +4 C Si Ge Sn Pb 9 For C, Si and Ge, the +2 oxidation state is much more stable than the +4 oxidation state. For Sn, the +2 is slightly more stable than the +4 oxidation state. For Pb, the +4 oxidation state is more stable than the +2 oxidation state. 10 The tendency of the latter members of a group to exhibit a valency 2 units less than the group valency due to the reluctance of the s electrons to participate in the reaction is called the inert pair effect. 11 The increasing stability of the +2 oxidation state can be seen from the standard electrode potential values below: Ge4+(aq) + 2e– Ge2+(aq) E° = –1.60 V Sn4+(aq) + 2e– Sn2+(aq) E° = +0.15 V Pb4+(aq) + 2e– Pb2+(aq) E° = +1.69 V The more positive the E° value, the less stable the +4 oxidation state. The more negative the E° value, the more stable the +4 oxidation state. 12 For lead, the +4 state is strongly oxidising. Its tendency to accept electrons and be converted to Pb2+ is very high. For example, lead(IV) oxide can oxidise hot, concentrated hydrochloric acid to chlorine and itself reduced to lead(II) chloride. PbO2(s) + 4HCl(aq) → PbCl2(s) + Cl2(g) + 2H2O(l) ∆ 13 In the presence of concentrated nitric acid, lead(IV) oxide oxidises manganese(II) aqueous ions to the purple permanganate ion. 5PbO2(s) + 2Mn2+(aq) + 4H+(aq) →∆ 2MnO4 – (aq) + 5Pb2+(aq) + 2H2O(l) 14 On the other hand, germanium(II) and tin(II) aqueous ions reduce iron(III) to iron(II) by way of donating electrons. Ge2+(aq) + 2Fe3+(aq) → Ge4+(aq) + 2Fe2+(aq) ∆ Sn2+(aq) + 2Fe3+(aq) → Sn4+(aq) + 2Fe2+(aq) ∆ Pb4+ is a powerful oxidising agent. 2007/P2/Q8(b) Ge2+ is a powerful reducing agent. Germanium(IV) and tin(IV) are written in the ionic form to balance the charges. The inert pair effect in lead is due to the ineffective screening provided by the inner d and f electrons. The inert pair effect increases down the group. 2012/P1/Q23

Chemistry Term 2 STPM Chapter 11 Group 14 177 11 11.4 Silicon, Silicone and Silicates Silicon and Silicone 1 Silicon, being a semiconductor, is used extensively in the electronic sectors to build transistors, microchips, integrated circuits and computer components. 2 Silicon is also used to make silicone, an organosilicon polymer. 3 The chlorine atoms in silicon tetrachloride can be replaced by alkyl groups to produce chlorosilanes such as CH3SiCl3 and (CH3)2SiCl2. 4 Chlorosilanes can be prepared by passing the vapour of haloalkanes or haloarenes over a copper-silicon mixture at 330 °C. For example, 2CH3Cl + Si → (CH3)2SiCl2 ∆ 5 Chlorosilanes can also be prepared using Grignard’s reagents. Grignard’s reagents are organomagnesium compounds such as methyl magnesium chloride, CH3MgCl. It is prepared by adding magnesium turnings to a mixture of methyl chloride in dry ether. CH3Cl + Mg → CH3MgCl 6 Grignard’s reagents react with silicon tetrachloride to produce chlorosilanes. 2CH3MgCl + SiCl4 → (CH3)2SiCl2 + 2MgCl2 The structure of the chlorosilane is shown below: CH3 CH3 Cl Cl Si 7 On hydrolysis, the chlorine atoms are replaced by ⎯OH groups. CH3 CH3 Cl Cl Si CH3 CH3 OH OH !!!: Si OH– 2008/P2/Q3(b) 2015/P2/Q16 2010/P2/Q23 2018/P2/Q12 INFO Silicone

Chemistry Term 2 STPM Chapter 11 Group 14 178 11 8 The resulting compound then undergoes polymerisation by losing water. HO Si CH3 CH3 OH HO Si CH3 CH3 O Si CH3 CH3 O Si CH3 CH3 O Si CH3 CH3 OH O Si CH3 CH3 H OH 9 By varying the nature of the alkyl group and the extent of hydrolysis, straight chains, ring or cross-linked polymers can be obtained. For example: !! Si !! R R !! !!O !! Si !! O O !! !!O !! Si !! R R !! !! !! Si !! R R !! !!O !!!! Si !! O !! !!O !! Si !! R R !! !! 10 By varying the degree of cross-links and chain length, the silicones can be obtained in the form of oils, grease or rubber-like solids. 11 Silicones are chemically inert and are good water repellents. 12 Silicones are used as lubricants, hydraulic fluids, electric insulators, elastomers, resins, electrical condensers, car polishes, implant in plastic surgery and to make water-proof fabric. 13 Silicones are also used in moulds to prevent the casting from sticking to the mould. Uses of silicone INFO Uses of Silicone Exam Tips Exam Tips The strong Si ─ O bonds are responsible for the inertness and water-proofing property of silicone.

Chemistry Term 2 STPM Chapter 11 Group 14 179 11 Silicates 1 Silicates are compounds containing silicon and oxygen. 2 Almost 75% of the Earth’s crust consists of silicates. Examples of silicates are sand, quartz, asbestos, granite, mica, clay and feldspar. 3 The building block (or unit cell) of all silicates is the tetrahedral SiO4 4– anion. This anion is sometimes called the orthosilicate ion. – O O– O– O– Si oxygen Top view silicon at the back of oxygen 4 The orthosilicate ions can share different numbers of terminal oxygen atoms with other orthosilicate units to give rise to a variety of structures. 5 In this section, we shall study the chain silicates (pyroxenes and amphiboles), sheet silicates (mica and talcum) and three-dimensional giant structures called framework silicates (quartz). Chain Silicates There are two main types of chain silicates. Single-chain silicates known as pyroxenes, and double-chain silicates known as amphiboles. Pyroxenes (Single-chain Silicates) 1 Pyroxenes are examples of single-chain silicates. 2 In pyroxenes, each SiO4 4– unit linked with other SiO4 4– units by sharing two oxygen atoms to form long linear chains. The repeating unit (empirical formula) of the single-chain silicate is the SiO3 2– ion. 3 Examples of pyroxenes are sodium silicate, Na2SiO3, magnesium silicate, MgSiO3 and calcium magnesium silicate, CaMg(SiO3)2. – O O– O O Si Si O – O Si O– Si O O– – O – O O O– – – – – – – – – – – Examples of silicates The orthosilicate ion Si O– O– O– O– Sharing of two oxygen atoms 2012/P1/Q22 The general formulae of the silicates are not required. VIDEO Silicates Exam Tips Exam Tips The negative charges in the silicate chains are balanced by the inclusion of positive metal ions such as Na+, Ca+ or Mg2+.

Chemistry Term 2 STPM Chapter 11 Group 14 180 11 Amphiboles (Double-chain Silicates) 1 When two single-chain silicates join by sharing one oxygen atom, a double-chain silicate called amphibole is formed. – – – – – – – – – – – – – – – – – 2 In amphiboles, some of the SiO4 2– units share two oxygen atoms each with other SiO4 2– units, while others share three oxygen atoms with other SiO4 2– units. 3 The repeating unit (empirical formula) in amphiboles is the [Si4O11]6– ions. 4 The bonds holding the atoms together in the chain are strong covalent bonds. However, the forces holding the individual chains together are weak. As a result, pyroxenes and amphiboles are brittle and breaks easily. 5 An example of amphiboles is asbestos. Asbestos cleaves easily in the plane parallel to the chain to form fibers. 6 Asbestos is used extensively as heat insulator and brake-pads. Asbestos dust is extremely hazardous and is thought to be a main cause of lung infection such as tuberculosis and lung cancer. Sheet Silicates 1 When each SiO4 4– units share three oxygen atoms with other SiO4 4– units, a sheet or layer silicate is formed. 2 In general, their structures are similar to those of amphiboles, except the hexagonal ring structures extend in both directions to give an infinite two-dimensional sheet structure. – – – – – – – – – – – – – – – – – – – The repeating unit has an empirical formula of [Si2O5]2–. Sharing of two and three oxygen atoms Asbestos 2007/P2/Q8(c) Sharing of three oxygen atoms

Chemistry Term 2 STPM Chapter 11 Group 14 181 11 3 Like that of the pyroxenes and amphiboles, strong covalent bonds hold the atoms together in the sheet, but weak intermolecular forces hold the individual sheets together. 4 The sheets can slide over one another. As a result, sheet silicates are soft and slippery. 5 Examples of sheet silicates are mica and talcum. Framework and Giant Sructure Silicates 1 When each SiO4 4– ion shares four terminal oxygen with other SiO4 4– units, a giant three-dimensional structure consisting of an array of silicon atoms and oxygen atoms is formed. 2 In the giant structure, each silicon atom is bonded to four oxygen atoms, and each oxygen atom is bonded to two silicon atoms. The arrangement repeats itself indefinitely. 3 The empirical formula of the framework silicate is SiO2. 4 A two-dimensional representation of the framework silicate is shown below: !! !! ! ! Si !! O ! O ! Si !! ! O ! Si !! O !! ! ! Si !! O ! O ! Si !! ! O ! Si !! O !! ! ! Si !! O ! O ! Si !! ! O ! Si !! O !! ! ! Si !! O ! O ! Si !! ! O ! Si !! O !!O!! !!O!! !!O!! !!O ! !! !!O!! !!O!! !!O!! !!O ! !! !!O!! !!O!! !!O!! !!O ! 5 An example of framework silicate is quartz (silicon dioxide). Due to the strong covalent bonds in the giant structure, quartz is a hard solid with high melting point. 6 Silicates are used as basic materials for cement, glass, ceramics and zeolites. The structure is almost similar to that of graphite. Sharing of 4 oxygen atoms silicon oxygen The geometry around each silicon atom is tetrahedral. Sand is an impure form of quartz.

Chemistry Term 2 STPM Chapter 11 Group 14 182 11 Zeolites 1 Aluminosilicates are compounds formed when some silicon atoms in the silicate structure are replaced by aluminium atoms. Some aluminosilicates lose water when heated to give an open structure that has large surface area and are porous. These structures are called zeolites. 2 Aluminosilicates are used as cation exchangers, drying agents, heterogeneous catalysts and as molecular sieves. 3 An example of a zeolite that is used as cation exchanger to ‘soften hard water’ is sodium aluminium silicate, NaAlSiO6. 4 Hard water is water that does not form lather with soaps due to the presence of dissolved magnesium or calcium salts. These cations, Mg2+ and Ca2+ forms insoluble precipitate with soap and covers the surface of the soap. 5 The soluble Mg2+ and Ca2+ can be removed from the water (this process is called softening of water) by passing the water through a column packed with zeolite. 6 The Ca2+ or Mg2+ will displace Na+ from the zeolite structure and thus are removed from the water. 7 The simplified equation for the softening process can be written as: 2NaAlSiO6(s) + Mg2+(aq) Mg(AlSiO6)2(s) + 2Na+(aq) 8 During use, the zeolite slowly loses its activity as more and more Ca2+ or Mg2+ are adsorbed. It can be regenerated by percolating with a concentrated solution of sodium chloride, which displaces the calcium ion and magnesium ion and replaces them with sodium ion. 9 Zeolites are also used as molecular sieve. The structure retains ions/ molecules that fit into their cavities while allowing smaller ion/ molecules to pass through. 10 Due to their anhydrous nature, zeolites can absorb water or moisture. They are used as drying agents. 11 In the petrochemical industry, zeolite is used as a catalyst during the cracking process to produce smaller hydrocarbon fragments from crude oil. Glass 1 When silica (silicon dioxide) is heated strongly, it melts at about 1710 °C when the silicon-oxygen bonds break. 2 When molten silica is cooled, the silicon-oxygen bonds are reformed but not in the regular crystalline structure as before. A transparent amorphous solid called quartz glass is formed. 3 Glass is a hard, non-crystalline transparent substance with an internal structure of a liquid. Glass is not a true solid in the sense that it flows slightly over time. Zeolites are microporous aluminosilicate minerals. Water containing Ca2+ and Mg2+ Cation Zeolite exchanger Water + Na+ Ca2+ and Mg2+ being adsorbed Molecular sieve Glass is a supercooled liquid. The Ca2+ and Mg2+ ions are adsorbed on the surface of zeolite. 2012/P2/Q7(b) VIDEO Zeolites 2014/P2/Q10 2009/P1/Q21

Chemistry Term 2 STPM Chapter 11 Group 14 183 11 4 The random ‘liquid-like’ molecular arrangement accounts for the fact that glass breaks irregularly instead of splitting along a plane like that of a crystal. 5 Glass has many uses especially as containers for liquids. It is transparent, chemically inert (except to concentrated hydrofluoric acid and caustic alkalis), non-toxic and easily recyclable. 6 Sometimes, other substances are added to quartz glass to modify their properties for particular uses. 7 Soda-lime glass which has sodium oxide and calcium oxide added is used for making window panes, bottles and dishes. 8 If a part of the silica is replaced by boron oxide, the glass produced has less tendency to crack with changes in temperature. This is borosilicate glass. The famous Pyrex glass used mainly as kitchen wares and as laboratory apparatus. 9 Lead glass with the ability to stop X-ray is made from a mixture of silica and lead(II) oxide. 10 Flint glass which has a brilliant appearance and high clarity and is suitable for optical instruments is made from silica, lead(II) oxide and sodium carbonate. 11 Coloured glass are obtained by adding compounds of transition elements during the manufacturing process. The table below lists some substances used to colour glass. Substance FeO Fe2O3 Cu2O SnO2 MnO2 CaF2 CoO Colour Green Yellow Red, green Opaque Violet Milky white Blue 12 In the manufacture of glass, proper annealing is very important. Annealing is the process of cooling the molten mixture to form a solid. If the glass is cooled too quickly, small, uneven areas of crystallinity will develop, this results in internal strain that causes the glass to shatter when knocked or when sudden temperature change occurs. 13 High quality optical glass must be annealed very carefully. For example, the huge Mount Palomar (in California, U.S.A.) observatory mirror was annealed from 500 °C to 300 °C over a period of nine months! Ceramics 1 Ceramics, which contain clay (silicates and aluminosilicates) and metal oxides are hard, brittle and stable even at high temperatures. They are used to make electrical insulators, glass-wares, pottery, water containers, bathroom tiles and as supports for the precious metal catalysts in catalytic converters. Glass dissolves slowly in concentrated hydrofluoric acid and concentrated strong alkali: SiO2(s) + 6HF(aq) → 2H+(aq) + SiF6 2–(aq) + 2H2O(l) SiO2(s) + 2NaOH(aq) → Na2SiO3(aq) + H2O(l) Info Chem Ceramics are inorganic, non-metallic solids prepared by the action of heat followed by cooling.

Chemistry Term 2 STPM Chapter 11 Group 14 184 11 SUMMARY SUMMARY 2 Ceramics are light weight and retain their properties and strength at temperatures above 1000 °C and are chemically inert (due to the strong bonds in the structure). 3 The one main problem with ceramic materials is their brittleness. Ceramics deform very little under stress, until they suddenly shatter without warning. 4 The brittleness of ceramics result from a weak point in the bonding within the ceramic matrix. The occurrence of the weak points is random. Hence, predictability of failure is impossible. 11.5 Tin Alloys 1 Tin is used in the plating of iron/steel container to form ‘tin can’. The iron/steel container is either dipped in molten tin, or the tin can be electroplated. As long as the tin layer is not scratched, the iron beneath it will not rust. 2 Tin is also used widely in the making of alloys. Examples are: Alloy Composition Bronze 70% Cu, 30% Sn Pewter 95% Sn, 3% Cu, 1% Sb Solder 30% Sn, 70% Pb Disadvantages of ceramics 1 Group 14 consists of a mixture of metals, metalloids and non-metals. 2 Going down Group 14: • the atomic radius increases • first ionisation energy decreases (although the 1st I.E. of Pb is slightly higher than Sn) • melting point decreases (although the melting point of Pb is slightly higher than Sn) • electric conductivity increases • the relative stability of the +2 oxidation state increases, while that of +4 state decreases • the inert pair effect increases • the metallic property increases 3 The tetrachlorides are all liquid at room conditions. 4 All tetrachlorides, except carbon tetrachloride, are hydrolysed by water to form acidic solutions. 5 The thermal stability of the tetrachlorides decreases down the group. 6 CO and SiO are neutral. GeO, SnO and PbO are amphoteric. 7 All monoxides, except lead monoxide, are converted to the dioxides on heating in air. 8 CO2, SiO2 are acidic. GeO2, SnO2 and PbO2 are amphoteric. 9 All dioxides are stable to heat except PbO which converts to PbO2 on heating in air. 10 The basic building block for all silicates is the orthosilicate ion, SiO4 4–. 11 Pyroxenes are single-chain silicates. 12 Amphiboles are double-chain silicates. 13 Quartz is an example of framework or giant structure silicate. 14 Silicones are organosilicon polymers. INFO Ceramics Exam Tips Exam Tips In solder, tin is added to lower the melting point of lead. 2007/P1/Q22 2013/P2/Q19(b)(ii)

Chemistry Term 2 STPM Chapter 11 Group 14 185 11 energy STPM PRACTICE 11 Objective Questions 1 Which of the following statements best explains the decreasing stability of the +4 oxidation states of the Group 14 elements with increasing proton number? A The metallic properties of the elements increases B The increase in the size of the atoms C The decrease in the ionisation energy D Electronegativity decreases 2 Which of the following tetrachlorides is the most unstable towards heat? A CCl4 C PbCl4 B PbCl2 D SiCl4 3 Which of the following statements is incorrect regarding the Group 14 elements and their compounds? A It is a group consisting of metals and nonmetals. B All the elements have giant structures. C All tetrachlorides are hydrolysed by water. D All the dioxides react with sodium hydroxide under suitable conditions. 4 Which of the following oxides of Group 14 elements is the most stable towards heat? A CO C SnO2 B SiO2 D PbO2 5 Which of the following statements is true regarding the Group 14 elements? A The oxides are either neutral, amphoteric or acidic. B Carbon is the only element with a giant structure. C The stability of the +4 oxidation state increases down the group. D Only tin and lead are conductors. 6 Which of the following statements about silicon tetrachloride (SiCl4) and lead tetrachloride (PbCl4) is true? A SiCl4 is a liquid, PbCl4 is a solid. B PbCl4 is more stable than SiCl4. C SiCl4 is more volatile than PbCl4. D SiCl4 is hydrolysed by water but PbCl4 is not. 7 Which statement about silicone (a.k.a. polysiloxanes) is not correct? A It is an inorganic polymer containing silicon atoms as the backbone of the polymer chain. B The repeating unit is siloxane. C It is used as a catalyst in many industrial processes. D It is an electrical insulator. 8 A tetrachloride of Group 14, MCl4 is insoluble in water and is stable to heat. M could be A carbon C tin B silicon D none of the above 9 What is the nature of the bond that holds the carbon atom and oxygen atom together in carbon monoxide? A van der Waals force B Ionic bond C Single covalent bond D Triple covalent bond 10 Lead reacts with cold concentrated hydrochloric acid to form lead(IV) chloride. Which statement is not correct about the property of lead(IV) chloride? A It is a yellow oily liquid. B It dissolves in water to give a reddish brown precipitate. C It is less stable towards heat than carbon tetrachloride. D It is ionic. 11 What causes lead to have a higher melting point than tin? A Lead atom is smaller. B The metallic bond in lead is weaker. C The inert-pair effect is more prominent in lead. D Lead is a metal while tin is a metalloid. 12 Which of the following statements is/are true when descending Group 14 in the Periodic Table from carbon to lead? I The atomic radius decreases. II The stability of the +4 oxidation state decreases. III The first ionisation energy decreases. A II only C II and III B I and III D I, II and III

Chemistry Term 2 STPM Chapter 11 Group 14 186 11 13 The first ionisation energy of lead is higher than that of tin because A the lead atom is smaller than the tin atom. B of the presence of 4f electrons in lead. C lead has more protons than tin. D the inert pair effect of lead is more profound than that in tin. 14 Which of the following graphs is correct with regards to the relative stability of the +2 and +4 oxidation states of the Group 14 elements? A C Si Ge Sn Pb Relative stability +4 +2 B +4 +2 C Si Ge Sn Pb Relative stability C C Si Ge Sn Pb Relative stability +4 +2 D C Si Ge Sn Pb Relative stability +4 +2 15 The boiling point of CCl4 is higher than that of SiCl4 because A the size of CCl4 is larger B the intermolecular forces in CCl4 are stronger C CCl4 is polar D carbon is a non-metal whereas silicon is a metalloid 16 In the preparation of lead(IV) chloride from lead(IV) oxide and concentrated hydrochloric acid, the temperature must be maintained at about 2 °C. This is A to prevent the mixture from boiling. B because lead(IV) oxide is unstable at high temperatures. C because concentrated hydrochloric acid will decompose to chlorine at higher temperatures. D to prevent the decomposition of lead(IV) chloride. 17 Group 14 elements (from carbon to lead) form tetrachlorides with the general formula of MCl4. What is not true about these chlorides? A Non-polar molecules B Liquid at room conditions C Thermally stable D Hydrolysed by water except for CCl4 18 Which of the following is not a silicate material? A Fullerene C Granite B Clay D Pyroxene 19 Solder is an alloy of tin and A zinc C copper B lead D carbon 20 Which property is not true of the Group 14 elements on descending down the group? A The inert pair effect increases. B The catenation of the elements decreases. C The thermal stability of the tetrachloride increases. D The electronegativity of the element decreases. 21 GeCl4 dissolves in water to form an acidic solution whereas CCl4 is insoluble in water because A the Ge—Cl bond is weaker than the C—Cl bond B GeCl4 is ionic but CCl4 is covalent C germanium atom has empty orbitals in its valence shell while carbon does not D GeCl4 has a lower melting point than CCl4 22 Which of the following is not a use for zeolites? A As a dehydrating agent B As a molecular sieve C As an ion exchanger D As a superconductor 23 Silicone is an inorganic polymer containing silicon. Silicone is used as I water-proof material II lubricant III electric insulator A I and II C II and III B I and III D I, II and III

Chemistry Term 2 STPM Chapter 11 Group 14 187 11 24 Which of the following Group 14 oxides undergoes thermal decomposition most readily? A CO C SnO2 B SiO2 D PbO2 25 Which of the following properties of the elements of Group 14 (C to Pb) shows a general decrease with increasing proton number? A First ionisation energy B Basic character of the oxides C Stability of the +2 oxidation state D Ease of hydrolysis of the tetrachlorides 26 Which of the following features of graphite best accounts for its use as a lubricant? A Strong covalent bond within the layer B The layered structure C Delocalised electrons D van der Waals forces between layers 27 Which of the following is a property of lead tetrachloride, PbCl4? A High boiling point B Stable to heat C Fumes in moist air D It disproportionate on heating. 28 Which of the following compounds is not stable towards heat? A CO2 C CCl4 B PbCl2 D SiO 29 Which statement about the chlorides of the Group 14 elements is correct? A CCl4, SiCl4 and GeCl4 are covalent, whereas SnCl4 and PbCl4 are ionic. B All, except CCl4, dissolve in water to form an acidic solution. C All the tetrachlorides decompose to form the dichlorides when heated. D All can be prepared by heating the elements directly with chlorine. 30 Silicon tetrachloride undergoes hydrolysis with water. However, carbon tetrachloride does not react with water. This is because A the Si—Cl bond is weaker than the C—Cl bond. B CCl4 is ionic but SiCl4 is covalent. C of the presence of empty orbitals in the valence shell of silicon. D the SiCl4 molecule is larger. 31 Zeolite is a type of aluminosilicate. Which of the following is/are a use for zeolite? I Catalyst in petrochemical industry II Ion exchanger III Molecular sieve IV Lubricant A I and II C II, III and IV B I, II and III D I, II, III and IV 32 Silicone is an inorganic polymer. It is used as A catalyst B water-proof material C superconductor D semiconductor 33 The structure of an element is represented by the diagram below. Which property is not true of the material? A It is unstable at room conditions. B It is soft and slippery. C It conducts electricity. D It is a constituent of all organic compounds. 34 Zeolite is used extensively as a catalyst, a dehydrating agent or in ion exchangers. It is an example of a/ an A silicate B silicone C aluminosilicate D ceramic 35 Which of the following chlorides is the least stable to heat? A CCl4 B SnCl2 C PbCl2 D PbCl4 36 On going down Group 14 (carbon to lead) of the Periodic Table, A the physical state of the elements changes from gas to solid B the +2 oxidation state of the elements increases C the thermal stability of the dioxide increases D the acid-base properties of the monoxide change from acidic to amphoteric

Chemistry Term 2 STPM Chapter 11 Group 14 188 11 Structured and Essay Questions 1 (a) Describe and explain the relative thermal stability of tin(IV) chloride and lead(IV) chloride. (b) Silicon and diamond have similar crystal lattice structure. However, the melting point of silicon is 1410°C while that of diamond is 3550°C. Explain the difference in their melting points. (c) When aqueous lead(II) nitrate is heated with aqueous sodium chlorate(I) under alkaline condition, lead(IV) oxide is formed. (i) Construct an ionic equation for the above reaction. (ii) Calculate the maximum mass of lead(IV) oxide produced when 50.0 cm3 of 0.110 mol dm-3 sodium chlorate(I) solution is heated with excess of lead(II) nitrate solution. 2 (a) (i) Complete the table below for the acid/base properties of the dioxides of Group 14 elements. Oxide CO2 SiO2 GeO2 SnO2 PbO2 Acid/base property (ii) Using SiO2 and PbO2, write equations to illustrate the acid/base properties of the dioxides. (b) When PbO2 is heated with an aqueous solution of manganese(II) oxide, in the presence of concentrated nitric acid, a purple solution is obtained. (i) Name the species that gives rise to the purple colour. (ii) Write a balanced equation for the reaction. (c) Very pure silicon dioxide (silica) can be obtained by reacting silicon tetrachloride with water. (i) Write an equation for the hydrolysis of silicon tetrachloride. (ii) State one use of silica. 3 Consider the following reactions: C(s) + CO2(g) → 2CO(g) ∆H = +170 kJ mol–1 Pb(s) + PbO2(s) → 2PbO(s) ∆H = –160 kJ mol–1 With reference to the above enthalpies of reactions, explain the relative stability of the +2 and +4 oxidation states for carbon and lead. 4 (a) Silicon tetrachloride is a colourless liquid at room conditions. (i) Draw the shape of silicon tetrachloride. (ii) Using an appropriate equation (with state symbols) to illustrate the standard enthalpy of atomisation of silicon tetrachloride. (iii) Calculate the standard enthalpy of formation of silicon tetrachloride. Process ∆H°/kJ mol–1 Atomisation of silicon +368 Cl–Cl bond energy +242 Atomisation of silicon tetrachloride +1 492 (iv) Explain why the standard enthalpy of atomisation of silicon tetrachloride is not equal to four times the Si–Cl bond energy. 5 (a) The Group 14 elements (carbon to lead) form two series of oxides with oxidation states of +2 and +4 respectively. Describe the acid/base properties of the oxides. Write equations where appropriate. (b) When lead dioxide, a brown solid, is heated with aqueous manganese(II) nitrate in the presence of excess concentrated nitric acid, a purple solution is obtained. (i) Give the name and formula of the species responsible for the purple solution. (ii) What is the role of lead dioxide in the reaction? (iii) Write a balanced ionic equation for the reaction taking place.

CHAPTER 12 GROUP 17 Concept Map Group 17 Elements Reaction of Chlorine with Sodium Hydroxide Reaction of Halide Ions with Silver Nitrate Reaction of Halides with Concentrated Sulphuric Acid Black-and-white Photography Uses of Halogens and their Compounds • Uses of – Chlorine – Bromine – Iodine Chemical Properties • Bond energy • Oxidising agent • Reaction with hydrogen • Hydrogen halides Physical Properties • Volatility • Atomic radius • Solubility in water Learning earning Outcomes Students should be able to: Physical properties of selected Group 17 elements • state that the colour intensity of Group 17 elements: Cl2, Br2, I2, increase down the group; • explain how the volatility of Group 17 elements decreases down the group. Reactions of selected Group 17 elements • deduce and explain the relative reactivities of Group 17 elements as oxidising agents from E° values; • explain the order of reactivity of F2, Cl2, Br2, I2 with hydrogen, and compare the relative thermal stabilities of the hydrides; • explain the reactions of chlorine with cold and hot aqueous sodium hydroxide. Reactions of selected halide ions • explain and write equations for reactions of Group 17 ions with aqueous silver ions followed by aqueous ammonia; • explain and write equations for reactions of Group 17 ions with concentrated sulphuric acid. Industrial applications of halogens and their compounds • describe the industrial uses of the halogens and their compounds as antiseptic, bleaching agent and in black-and-white photography; • explain the use of chlorine in water treatment.

Chemistry Term 2 STPM Chapter 12 Group 17 190 12 12.1 Physical Properties of Selected Group 17 Elements 1 The Group 17 elements (also known as halogens) are fluorine, chlorine, bromine, iodine and astatine. 1 2 13 14 15 16 17 18 Br F Cl I At 2 All have seven electrons in their valence shells with the configuration of ns2 np5 . ns np 3 Under normal conditions, they exist as simple diatomic molecules of F2, Cl2, Br2, I2 and At2 through electron sharing. 4 All are reactive non-metals and are not found in the elemental state in nature. 5 The colour intensity of Group 17 elements increases down the group as the metallic properties of the element increases. Atomic Radius and Ionic Radius 1 Going down Group 17, the effective nuclear charge remains almost constant but the number of electronic shells increases. Element Fluorine Chlorine Bromine Iodine Proton number 9 17 35 53 Electronic configuration 2.7 2.8.7 2.8.18.7 2.8.18.18.7 Effective nuclear charge 7 7 7 7 Atomic radius/nm 0.072 0.099 0.196 0.213 Ionic radius/nm 0.135 0.180 0.195 0.215 2 As a result, the atomic radius and ionic radius increases with proton number. Info Chem Astatine is radioactive and one of the rarest element on Earth. • • • • • • • Cl Cl fifi fi fi fi fi fi 2012/P1/Q25 2015/P2/Q20(a) 2013/P2/Q13 2017/P2/Q12

Chemistry Term 2 STPM Chapter 12 Group 17 191 12 10 20 30 Proton number 40 50 Radius (fi 10–3)/nm 140 120 100 80 60 220 200 180 160 60 F Cl Br I F – Cl– Br– I – Volatility of Group 17 Elements Element Melting point /°C Boiling point/°C Physical state at 298 K Fluorine –219 –188 Pale yellow gas Chlorine –101 –34.7 Pale yellowish green gas Bromine –7.2 58.8 Reddish brown liquid Iodine 114 184 Black solid 1 Melting point and boiling point are measure of the strength of the intermolecular forces that hold the molecules together in the solid and liquid state respectively. The stronger the force, the higher is the melting point and boiling point. 2 The intermolecular force between halogen molecules is the van der Waals force. X2 X2 van der Waals force 3 Going down Group 17, the size and the total number of electrons in the halogen molecule increase. This causes an increase in the intermolecular van der Waals force. As a result, the melting point and boiling point increase. The colour of the elements gets darker down the group. INFO General Chemistry of Halogens

Chemistry Term 2 STPM Chapter 12 Group 17 192 12 Element F2 Cl2 Br2 I2 Atomic radius/nm 0.064 0.099 0.114 0.133 No. of electrons 18 34 70 106 4 The halogens get less volatile down the group and the colour of the elements gets darker with increasing proton number. Solubility in Water 1 The halogens are only sparingly soluble in water because they cannot form hydrogen bonds with water molecules. For example, Cl2(g) + H2O(l) HCl(aq) + HOCl(aq) 2 However, iodine is completely soluble in aqueous potassium iodide. This is due to the formation of a water-soluble complex ion, I3 – . I2(aq) + I– (aq) ⎯→ I3 – (aq) Exam Tips Exam Tips • F2 reacts violently with water to liberate oxygen gas. 2F2 + 2H2O → 2HF + O2 • When exposed to strong sunlight, chloric(I) acid decomposes according to the equation: 2HOCl(aq) → 2HCl(aq) + O2 (g) Quick Check 12.1 When a sample of chlorine water is exposed to strong sunlight, a gas that supports combustion is liberated. Explain the observation. Example 12.1 Draw the Lewis structure for the I3 – ion and predict its shape. Solution • • • • • • • I I I – fifi fifi fifi fifi fi fi fi fi fi fi fi The central iodine atom is surrounded by 5 electron pairs (2 bondpairs and 3 lone-pairs). As a result, its shape is linear. I I I – 2012/P2/Q3(a) 2010/P2/Q8(b)