Summary Term Chemistry STPM Prepared by Crystal Goh AI Tuition Centre 017713136 Period 3 elements property Na Mg Al Si P (P 4 ) Type of element Metal Metalloid Non-metal Structure Giant metallic lattice Giant covalent Simple covalent molecule molecular lattice Bonding Strong Metallic bond Strong Covalent bond S (S 8 ) Cl Covalent bond in the molecule. Weak Van der Waals forces between molecules Melting and High High Low boiling points H vap High High Low Electric conductivity Good Increases when temperature increases; Poor Formulae of oxides P 4 O 10 SO 3 Cl O 7 Na O MgO Al O 3 SiO P 4 O 6 SO Cl O Lattice structure Giant ionic lattice Giant covalent molecular Simple covalent molecule lattice Physical state at 5 o C Solid Gas Liquid gas Oxidation state +1 + +3 +4 +5 Acid/ base properties Basic Amphoteric Acidic +3 +6 +4 +7 +1 Note : Na and chlorine reacts with water, other elements across period reacts with steam. Elements Reaction of elements with oxygen Reaction with water /steam Sodium 4Na + O Na O Na + H O(l) NaOH + H Magnesium Mg + O MgO Mg + H O(g) MgO + H Aluminum Al + O Al O 3 Al +3 H O(g) Al O 3 + 3H Silicon Si + O SiO Si + H O(g) SiO + H Phosphorous P 4 + 3O P 4 O 6 No reaction P 4 + 5O P 4 O 10 Sulphur S + O SO No reaction SO + O SO 3 Chlorine No reaction Cl + H O (l) HCl + HClO 1
Formula of Properties of oxides Reacts with Observation oxides alkali /acid Na O MgO Al O 3 SiO Strong Basic dissolves in water Basic Slightly soluble water Amphoteric Insoluble water Slightly Acidic Acid Na O + H O(l) Na + (aq) + OH (aq) Na O + HCl NaCl + H O (l) Acid MgO (s)+ H O (l) Mg(OH) MgO (s) + H + (aq) Mg + (aq) + H O (l) MgO + HCl MgCl + H O (l) Acid and alkali Al O 3 (s) + 6H + (aq) Al 3+ (aq) + 3 H O (l) Al O 3 (s) + 3H SO 4 (aq) Al (SO 4 ) 3 + 3H O Al O 3 (s) + OH - (aq) + 3 H O (l) [Al(OH) 4 ] - (aq) Al O 3 (s) + NaOH(aq) + 3 H O (l) Na[Al(OH) 4 ] Strong alkali SiO (s) + OH - (aq) SiO - 3 (aq) + H O (l) Insoluble water SiO (s) + NaOH (aq) Na SiO 3 (aq) + H O Oxides of phosphorous, sulphur and chlorine, are all acidic. They dissolve in water to give acids. P 4 O 6 P 4 O 10 Acidic Alkali P 4 O 6 (s) + 6 H O (l) 4 H 3 PO 3 (aq) P 4 O 10 (s) + 6 H O (l) 4 H 3 PO 4 (aq) P 4 O 6 (s) + 1NaOH (aq) P 4 O 10 (s) +1NaOH (aq) 4Na 3 PO 3 (aq) + 6 H O 4Na 3 PO 4 (aq) + 6 H O SO SO 3 Acidic Alkali SO (g) + H O (l) H SO 3 (aq) SO 3 (g) + H O (l) H SO 4 (aq) SO (g) + NaOH (aq) Na SO 3 + H O SO 3 (g) + NaOH (aq) Na SO 4 + H O Cl O Cl O 7 Acidic Alkali Cl O (l) + H O (l) HClO (aq) Cl O 7 (g) + H O (l) HClO 4 (aq) Cl O (l) + NaOH (aq) NaClO + H O Cl O 7 (g) + NaOH (aq) NaClO 4 + H O NaCl and MgCl are soluble salts dissolves fully in water. The rest chlorides are simple covalent molecule undergo hydrolysis form white fume HCl. Al Cl 6 + 6H O SiCl 4 (l) + 4 H O (l) Al(OH) 3 + 6HCl SiO.H O + 4HCl PCl 5 + 4 H O (l) H 3 PO 4 + 5HCl SCl 4 + 4 H O (l) H SO 4 + 4HCl
Reactions of the Group elements with air / oxygen gas. 1. Group elements burn in oxygen when heated to form oxides Equations : M (s) + O (g) MO(s) Mg (s) + O (g) Ca (s) + O (g) Sr (s) + O (g) Sr(s) + O (g) MgO(s) CaO(s) SrO(s) SrO (s) Strontium peroxide ( poisonous gas causes nose irritation) Ba (s) + O (g) BaO(s) (The solid s are white solid). Ba (s) + O (g) BaO (s) Barium peroxide Peroxides are not stable Solubility of oxide Group in water BeO(s) + H O Be(OH) (s) slightly soluble in water MgO (s)+ H O CaO (s)+ H O Quick lime Mg(OH) (s) Ca(OH) (s) slacked lime SrO (s)+ H O Sr(OH) (aq) very soluble in water Strong alkali ( ph=13) BaO + H SO 4 H O + BaSO 4 BaO (s)+ H O Ba(OH) (aq) Strong alkali ( ph =13) Beryllium chloride is hydrolysed in water exothermically to produce an acidic solution BeCl + H O Be(OH) + HCl Beryllium chloride does not conduct electricity in molten form. Beryllium oxide and beryllium hydroxide are amphoteric. a. BeO(s) + H + (aq) Be + (aq) + H O (l) BeO (s)+ OH - (aq) + H O(l) [Be(OH) 4 ] - (aq) b. Be(OH) (aq) + H + (aq) Be + (aq) + H O (l) Be(OH) (s)+ OH - (aq) [Be(OH) 4 ] - (aq) 3
Thermal Decomposition MgCO 3 MgO + CO Mg(NO 3 ) MgO + NO + O Mg(OH) MgO + H O Thermal stability increase down the group. Explain : the atomic size increase, the charge density increases, higher polarizing power. Summary of the trend in reactivity The Group metals become more reactive towards water as you go down the Group. Element Reactivity reacts with water Equation Be Mg Ca Sr Ba No reaction Reacts slowly with water but react readily with steam to form MgO and H Reacts with cold water slowly to form Ca(OH) and H Reacts with cold water vigorously to form strong alkali Sr(OH) and H Reacts with cold water vigorously to form strong alkali Ba(OH) and H Mg + H O(g) MgO (s) + H MgO slightly soluble to form Mg(OH) (s) Ca(s) + H O(l) Ca(OH) (aq) + H Sr (s) + H O(l) Sr(OH) (aq) + H ph = 13 strong alkali Ba (s) + H O(l) Ba(OH) (aq) + H ph = 13 strong alkali Anomalous Behaviour of Beryllium (a) Due to 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 higher than expected compared to other group elements. (b) Formation of Be + is harder than the rest of Group elements. (c) Be + ion with high charge density able to polarizes any anion that is bonded to form covalent character to the compound. (d) Beryllium chloride is a covalent compound while the others chloride are ionic. BeCl is white solid that sublimes when heated c. Beryllium chloride is hydrolysed in water exothermically to produce an acidic solution BeCl + H O Be(OH) + HCl d. Beryllium chloride does not conduct electricity in molten form. 4
(e) Beryllium oxide and beryllium hydroxide are amphoteric. a. BeO(s) + H + (aq) Be + (aq) + H O (l) BeO (s)+ OH - (aq) + H O(l) [Be(OH) 4 ] -( aq) b. Be(OH) (aq) + H + (aq) Be + (aq) + H O (l) Be(OH) (s)+ OH - (aq) [Be(OH) 4 ] -( aq) (f) Beryllium can form complexes ions due to high charge and small size of Be + ion. (g) Aluminium oxide and aluminium hydroxide are amphoteric. a. Al O 3 (s) + 6H + (aq) Al 3+ (aq) + 3H O (l) Al O 3 (s)+ OH - (aq) + H O(l) [Al(OH) 4 ] - (aq) b. Al(OH) 3 (aq) + 6H + (aq) Al 3+ (aq) + 3H O (l) Al(OH) 3 (s)+ OH - (aq) [Al(OH) 4 ] - (aq) Group 14 elements The stability of oxidation state + increases down the group. Pb + is more stable than Pb 4+ due to inert electron pair effect. Pb 4+ is strong oxidising agent Chemical properties of group 14 Elements (a)the reaction with oxygen All the elements in Group 14, except lead react with oxygen when heated to produce a dioxide Lead only forms monoxide when heated with oxygen. C(s) + O CO Si(s) + O SiO Ge(s)+ O GeO Sn(s)+ O SnO Pb(s) + O PbO (yellow) 6 PbO + O Pb 3 O 4 (yellow) (orange) 5
Formula of oxides Acid /base Ionic Equation properties CO Acidic CO + OH CO - 3 + H O SiO Acidic SiO + OH SiO 3 - + H O GeO, SnO, PbO Amphoteric PbO + H + Pb + + H O PbO + OH - + H O [Pb(OH) 4 ] - GeO, SnO, PbO Amphoteric eg : PbO + 4H + Pb 4+ + H O Diluted any acid PbO + 4HCl PbCl + Cl + H O Concentrated HCl used PbO + OH - (aq) + H O(l) [Pb(OH) 6 ] - (aq) Thermal Stability CCl 4, SiCl 4 and GeCl 4 are ( simple covalent molecule )very stable even at high temperature. SnCl 4 decomposes when heated to form stanum(ii) chloride and chlorine, whereas PbCl 4 decomposes when it is warmed to form lead(ii) chloride and chlorine. SnCl 4 SnCl + Cl PbCl 4 PbCl + Cl ( bleach the blue litmus paper ) Hydrolysis All tetrachlorides except CCl 4 can be hydrolysed During hydrolysis, silicon, germanium, stanum and lead atoms of the respective chlorides make use of their empty d orbitals to form dative bond with water molecules. After the coordination, the reaction is followed by an elimination of hydrogen chloride molecule. Example: SiCl 4 (l) + 4 H O (l) SiO.H O + 4HCl 6
Tetrachloride of carbon cannot be hydrolysed because central carbon atom can only hold 8 valence electrons ( maximum covalency is 4). The carbon atom cannot expand its octet because there is no empty orbital at period elements Group 17 elements halogens are diatomic molecules,simple covalent compound. The intensity colour of halogens getting darker as going down the group and physical state changes gas to liquid then to solid. The Solubility of Halogens in Water All halogens are slightly soluble in water because they cannot form hydrogen bonding with water. Chlorine and Bromine react with water to give HX and HOX whereas Iodine does not. Iodine is very soluble in aqueous Potassium Iodide to give triodide ion, I₃. Cl + H O Br + H O I₂ (s) + I (aq) HCl + HClO HBr + HBrO I₃ (aq) HClO decompose to form HCl and O Chemical Properties of Halogen (a) Reaction with hydrogen : Formation of Hydrogen Halides F₂(g) + H₂(g) HF (g) Even in dark and fluorine in solid state hv Cl₂(g) + H₂(g) HCl (g) Pt Br₂ (g) + H₂ (g) HBr (g) 50ºC Pt I₂ (g) + H₂ (g) HBr (g) 400ºC Stability of the Formation of product HF > HCl >HBr >HI ( hence enthalpy change of formation HF is the most exothermic) 7
Thermal Stability of Hydrogen Halides 1. From the enthalpy change of formation and the bond dissociates energy, the thermal stability of hydrogen halides decrease in the order HF > HCl > HBr > HI bond length increase All the hydrogen halides decompose to their elements on heating. HX(g) H (g) + X (g) Reaction of Halides Ions with Silver Ion in Aqueous Solution Halides ion react immediately with silver ion in aqueous solution to produce silver iodide precipitate. Ag+ + X AgX Silver Halide AgCl AgBr AgI Colour of Precipitate White Creamy Yellow Yellow Silver Chloride and Silver Bromide dissolve in aqueous ammonia whereas Silver Iodide does not. AgX (s) + NH₃ (aq) Ag(NH₃)₂+ + X Diammine silver(i) ion Reaction Between Halide Ions with Lead (II) Ions in Aqueous Solution Halides react immediately with Lead (II) ions to form Lead (II) Halides precipitates. Eg : Pb + + Cl PbCl₂ Lead (II) Halides PbCl₂ PbBr₂ PbI₂ Colour Precipitate White White Yellow Note : Formation of Precipitate is through double decomposition reaction. CaF₂(s) + H₂SO₄ HF + CaSO 4 NaCl(s) + H₂SO₄ HCl + Na SO 4 concentrated sulphuric acid is strong enough to oxidise the hydrogen bromide to bromine and hydrogen iodide to iodine. Example: Br + H₂SO₄ - HBr + HSO 4 HBr + H₂SO₄ Br + SO + H O I + H₂SO₄ HI + HSO 4 - HI + H₂SO₄ I + SO + H O 8
Note Combining these last two half-equations gives: (a) Redox Reaction 1. Br + Cl₂ Br₂ + Cl. I + Br₂ I₂ + Br 3. I₂ + S₂O₃² I + S₄O 6 ² Cl Br I Conc. H₂SO₄ Acid HCl HBr + Br₂ HI + I₂ Conc. H₂SO₄ Acid + MnO₂ Cl₂ Br₂ I₂ Conc H₃PO₄ Acid HCl HBr HI (b) Reactions which involve disproportion 1. Cl₂ + NaOH NaCl + NaClO + H₂O. 3Cl₂ + 6NaOH 5NaCl + NaClO₃ +3H₂O 3. 3OCl ClO₃ + Cl 4. Cl₂ + H₂O HCl + HOCl 5. Br₂ + H₂O HBr + HOBr 6. 4KClO₃ KCl + 3KClO₄ Homogeneous catalysis. (Reactants products and catalyst are in the same phase): Oxidation of iodide ion by peroxydisulphate ions in aqueous solution. Fe²+ (aq) / Fe 3+ (aq) I (aq) + S₂O₈² (aq) I₂ (aq) + SO₄² (aq) Without catalyst, this reaction is almost not feasible, as both the anions will repel each other. Mechanism: S₂O₈² (aq) + Fe²+ (aq) Fe³+ (aq) + I (aq) SO₄² (aq) + Fe³+ (aq) I₂ (aq) + Fe²+ (aq) 9