1. Periodicity Period 3 - elements

2. Can you write one fact about each of the following elements?  Na  Mg  Al  Si  P  S  Cl  Ar

3. Candidates should be able to: describe qualitatively (and indicate the periodicity in) the variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements (see the Data Booklet) explain qualitatively the variation in atomic radius and ionic radius interpret the variation in melting point and in electrical conductivity in terms of the presence of simple molecular, giant molecular or metallic bonding in the elements explain the variation in first ionisation energy.

4. Na[Ne] 3s 1 Mg[Ne] 3s 2 Al[Ne] 3s 2 3p x 1 Si[Ne] 3s 2 3p x 1 3p y 1 P[Ne] 3s 2 3p x 1 3p y 1 3p z 1 S[Ne] 3s 2 3p x 2 3p y 1 3p z 1 Cl[Ne] 3s 2 3p x 2 3p y 2 3p z 1 Ar[Ne] 3s 2 3p x 2 3p y 2 3p z 2

5. Atomic radius

6. www.chemguide.co.uk/atoms/properties/atradius.html Atomic radius – what about argon?

7. What is atomic radius?

8. www.rjclarkson.demon.co.uk/found/period3_atomrad.gif

9. An atomic radius is a measure of the distance from the nucleus to the bonding pair of electrons. From sodium to chlorine, the bonding electrons are all in the 3 rd shell being screened by the electrons in the first and second levels, i.e. the screening remains fairly constant. The increasing nuclear charge as you go across the period pulls the bonding electrons more tightly towards it. Explaining the trend

10. You have to ignore the noble gas at the end of each period. Because neon and argon don't form bonds, you can only measure their van der Waals radius - a case where the atom is pretty well "unsquashed". All the other atoms are being measured where their atomic radius is being lessened by strong attractions. You aren't comparing like with like if you include the noble gases. Leaving the noble gases out, atoms get smaller as you go across a period.

11. IONIC RADIUS Ions aren't the same size as the atoms they come from. Compare the sizes of sodium and chloride ions with the sizes of sodium and chlorine atoms. Positive ions Positive ions are smaller than the atoms they come from. Sodium is 2,8,1; Na + is 2,8. You've lost a whole layer of electrons, and the remaining 10 electrons are being pulled in by the full force of 11 protons. Negative ions Negative ions are bigger than the atoms they come from. Chlorine is 2,8,7; Cl - is 2,8,8. Although the electrons are still all in the 3-level, the extra repulsion produced by the incoming electron causes the atom to expand. There are still only 17 protons, but they are now having to hold 18 electrons.

12. Electronegativity

13. First ionisation energy

14. Electronegativity

15. NaMgAlSiPSClAr StructureGiant metallic Giant covalent Simple molecule Monatomic Type of element Metal Non-metal BondingMetallic Covalent FormulaP4P4 S8S8 Cl 2 Ar Type of force broken on melting/ boiling Metallic bond Covalent bond vdW Does the element conduct electricity ? Yes No Bonding, Structure and Properties

16. Structure and Properties

17. Electrical conductivity

18. Melting and boiling points

19. Periodicity Period 3 - oxides

20. Starter activity Complete these sketches to show how these properties change as you go along Period 3 from Na to Ar.

23. Candidates should be able to:  describe the reactions, if any, of the elements with oxygen to give Na 2 O, MgO, Al 2 O 3, P 4 O 10, SO 2 and SO 3.  state and explain the variation in oxidation number of the oxides.  describe the reactions of the oxides with water.  describe and explain the acid/base behaviour of oxides and hydroxides, including, where relevant, amphoteric behaviour in reaction with NaOH and acids.

24. Group number123456 Element in Period 3 NaMgAlSiPS Nuclear charge11+12+13+14+15+16+ [Ne] electronic configuration 3s 1 3s 2 3s 2 3p 1 3s 2 3p 2 3s 2 3p 3 3s 2 3p 4 Trend in Atomic radius decreases Trend in 1st ionisation energyincreases Trend in electronegativityincreases Formula of oxide/s Na 2 OMgOAl 2 O 3 SiO 2 P 4 O 10 SO 2 / SO 3 Oxidation state +1+2+3+4+5+4/+6 Reactions with oxygen

25. Sodium Sodium burns in oxygen with an orange flame to produce the white solid sodium oxide. Magnesium Magnesium burns in oxygen with an intense white flame to give white solid magnesium oxide. Reactions with oxygen

26. Aluminium Aluminium will burn in oxygen if it is powdered, otherwise the strong oxide layer on the aluminium tends to inhibit the reaction. If you sprinkle aluminium powder into a Bunsen flame, you get white sparkles. White aluminium oxide is formed. Silicon Silicon will burn in oxygen if heated strongly enough. Silicon dioxide is produced. Reactions with oxygen

27. Phosphorus White phosphorus catches fire spontaneously in air, burning with a white flame and producing clouds of white smoke - a mixture of phosphorus(III) oxide and phosphorus(V) oxide. The proportions of these depend on the amount of oxygen available. In an excess of oxygen, the product will be almost entirely phosphorus(V) oxide: Reactions with oxygen

28. Sulphur Sulphur burns in air or oxygen on gentle heating with a pale blue flame. It produces colourless sulphur dioxide gas. In an excess of pure oxygen, some SO 3 is also formed. This utilises the highest oxidation state of sulphur.

29. Reactions with oxygen Which is which?

30. Melting points of oxides Na 2 OMgOAl 2 O 3 SiO 2 P 4 O 10 SO 2 T m /K1548312523451883573200 BondingIonic Covalent Structure Giant lattice Simple molecular Giant ionic and covalent solids contain only strong bonds and have high melting points. Simple molecules have weak vdW forces between molecules and have low melting points.

31. Reaction with waterpH Na 2 O + H 2 O  2Na + + 2OH - 14 MgO + H 2 O  Mg 2+ + 2OH - 9 Al 2 O 3 insoluble – no reaction7 SiO 2 insoluble – no reaction7 P 4 O 10 + 6H 2 O  4H 3 PO 4 0 SO 2 + H 2 O  H 2 SO 3 3 SO 3 + H 2 O  H 2 SO 4 0 Acid/base properties of the oxides

32. Acid/base behaviour of aluminium oxide BASE: Al 2 O 3 + 3H 2 SO 4 → Al 2 (SO 4 ) 3 + 3H 2 O ACID: Al 2 O 3 + 2NaOH + 3H 2 O → 2NaAl(OH) 4

33. Periodicity Period 3 - chlorides

34. Candidates should be able to: describe the reactions, if any, of the elements with chlorine to give NaCl, MgCl 2, Al 2 Cl 6, SiCl 4, and PCl 5. describe and explain the reactions of the chlorides with water.

35. ElementNaMgAlSiP Description of reaction with chlorine Very vigorousVigorous Slow Formula of chloride/s NaClMgCl 2 Al 2 Cl 6 SiCl 4 PCl 3 / PCl 5 Oxidation state of period 3 element +1+2+3+4+3 / +5 State of chloride at r.t.p. Solid Liquidliquid / solid b.pt. of chloride ( o C) 146514184235774 / 164 Structure of chloride Giant latticeSimple molecular Bonding in chloride IonicCovalent Reaction with chlorine

36. Structure of Al 2 Cl 6

37. Reaction with waterpH NaCl (s)  Na + (aq) + Cl - (aq) 7 MgCl 2(s)  Mg 2+ (aq) + 2Cl - (aq) 6/7 Al 2 Cl 6(s) + 12H 2 O (l)  2[Al(H 2 O) 6 ] 3+ 6Cl - (aq) 3 SiCl 4(l) + 2H 2 O (l)  SiO 2(s) + 4HCl (g) 0 PCl 5(l) + 4H 2 O (l)  H 3 PO 4(aq) + 5HCl (g) 0 Acid/base properties of the chlorides

38. Periodicity Group 7 - Halogens

39. Starter activity Can you complete task 1 – IGCSE revision on the halogens

40. Candidates should be able to: describe the trends in volatility and colour of chlorine, bromine and iodine. interpret the volatility of the elements in terms of van der Waals’ forces. describe the relative reactivity of the elements as oxidising agents. describe and explain the reactions of the elements with hydrogen. describe and explain the relative thermal stabilities of the hydrides. interpret these relative stabilities in terms of bond energies.

41. The Halogens Ionic Covalent

42. The Halogens Trend in colour and physical state

43. Like dissolves like!

44. Periodicity Group 7 – the halides

45. Starter activity Can you complete the table ‘Reducing power of the halides’?

46. Trend in reducing ability NaXObservationsProductsType of reaction NaFsteamy fumes HF acid-base (F - acting as a base) NaClsteamy fumes HCl acid-base (Cl - acting as a base) NaBrsteamy fumes HBr acid-base (Br - acting as a base) colourless gas SO 2 redox (reduction product of H 2 SO 4 ) brown fumes Br 2 redox (oxidation product of Br - ) NaIsteamy fumes HI acid-base (I - acting as a base) colourless gas SO 2 redox (reduction product of H 2 SO 4 ) yellow solid S redox (reduction product of H 2 SO 4 ) smell of bad eggs H2SH2S redox (reduction product of H 2 SO 4 ) Grey solid, purple fumes I2I2 redox (oxidation product of I - )

47. Candidates should be able to: describe and explain the reactions of halide ions with aqueous silver ions followed by aqueous ammonia concentrated sulphuric acid. describe and interpret in terms of changes of oxidation number the reaction of chlorine with cold, and with hot, aqueous sodium hydroxide. explain the use of chlorine in water purification. recognise the industrial importance and environmental significance of the halogens and their compounds, (e.g. for bleaches; PVC; halogenated hydrocarbons as solvents, refrigerants and in aerosols).

48. Fluoride Chloride Bromide Iodide Relative reducing power Summary of reducing power

49. Testing for halide ions ion presentobservation F-F- No precipitate Cl - White precipitate Br - Cream precipitate I-I- Yellow precipitate Silver fluoride is soluble, and so you don't get a precipitate.

50. Confirmatory tests original precipitateobservation AgClSoluble in dilute NH 3(aq) AgBrSoluble in concentrated NH 3(aq) AgIInsoluble in concentrated NH 3(aq) E.g. AgCl (s) + 2NH 3(aq) → [Ag(NH 3 ) 2 ] + (aq) + Cl - (aq)

51. Cl 2(g) + 2NaOH (aq) → NaClO (aq) + NaCl (aq) + H 2 O (l) 0+1-1 3Cl 2(g) + 6NaOH (aq) → NaClO 3 (aq) + 5NaCl (aq) + 3H 2 O (l) Reactions of chlorine with NaOH

52. Uses of halogens and their compounds bleach refrigerants solvents Aerosol propellants PVC

53. Periodicity Nitrogen and Sulphur

54. Starter activity Can you use information in your textbooks to complete task 9?

55. Candidates should be able to: explain the lack of reactivity of nitrogen. describe the o formation, and structure of, the ammonium ion o the displacement of ammonia from its salts. understand the environmental consequences of the uncontrolled use of nitrate fertilisers. understand and explain the occurrence, and catalytic removal, of oxides of nitrogen. explain why atmospheric oxides of nitrogen are pollutants, including their catalytic role in the oxidation of atmospheric sulphur dioxide. describe the formation of atmospheric sulphur dioxide from the combustion of sulphur contaminated carbonaceous fuels. state the role of sulphur dioxide in the formation of acid-rain and describe the main environmental consequences of acid-rain. understand the industrial importance of sulphuric acid. describe the use of sulphur dioxide in food preservation.

56. Gases in the air

57. Structure of nitrogen Bond dissociation enthalpy = +946 kJ/mol

58. High energy Enzymes Nitrogen

59. The Haber Process

60. The ammonium ion AmmoniumNH 4 + Add sodium hydroxide solution to a solution of the substance and gently heat. Ammonia gas is given off.

61. n. Any of a large number of natural and synthetic materials, including manure and nitrogen, phosphorus, and potassium compounds, spread on or worked into soil to increase its capacity to support plant growth. Fertiliser

62. Fertilisers Ammonium nitrate Potassium chloride Ammonium phosphate

63. Nitric acid NH 3(g) + 2O 2(g)  HNO 3(l) + H 2 O (l) Pt/Rh catalyst 900 o C

64. The environment and fertilisers fertilisers applied to farm land too much used, at the wrong time of year, during wet weather, if Excessive growth of aquatic plants. The bacteria which live on dead plants thrive and use up the oxygen in the water. The lack of oxygen causes death of fish. This is called eutrophication. Harm to infants - called ‘blue baby’ syndrome washed into rivers and lakes causes excess contaminates underground drinking water supplies causes

65. Natural sources of SO 2 and NO x

66. Sulphur dioxide - SO 2 Used as a food preservative:  Inhibits growth of moulds, yeasts and aerobic bacteria.  Acts as a reducing agent and retards the oxidation of foodstuffs.

67. Acid rain SO 2(g) + NO 2(g)  SO 3(g) + NO (g) SO 3(g) + H 2 O (l)  H 2 SO 4(l)

68. Uses of sulphuric acid Paints, pigments and dyestuffs Detergents Fertilisers

69. Periodicity Group 2 AS Chemistry

70. Candidates should be able to interpret, and make predictions from, the trends in physical and chemical properties of the elements and their compounds.

71. Trends in Atomic Radius

72. The radius of an atom is governed by:  the distance of the outer electrons from the nucleus  the nuclear charge, and  the amount of shielding

73. Trends in First Ionisation Energy

74. Ionisation energy is also governed by:  the charge on the nucleus,  the amount of shielding by the inner electrons,  the distance between the outer electrons and the nucleus. Trends in First Ionisation Energy Increased nuclear charge ‘off-set’ by increased shielding. Outermost electron increasingly distant from pull of nucleus, less energy needed to remove it.

75. Trends in Electronegativity

78.  As the metal atoms get bigger, any bonding pair gets further and further away from the metal nucleus, and so is less strongly attracted towards it.  Increasing nuclear charge ‘off-set’ by more shielding.  In other words, as you go down the Group, the elements become less electronegative. Trend in Electronegativity

79. Trend in Melting Point

80. Going down Group 2: the number of delocalised electrons remains the same... the charge on each metal cation stays the same at 2+, but... the ionic radius increases... so the attraction between the delocalised electrons and the metal cations decreases. Trend in Melting point

81. Periodicity Group 2 – chemical properties AS Chemistry

82. Candidates should be able to: describe the reactions of the elements with oxygen and water. describe the behaviour of the oxides with water. describe the thermal decomposition of the nitrates and carbonates. interpret, and make predictions from, the trends in chemical properties of the elements and their compounds. explain the use of magnesium oxide as a refractory lining material and calcium carbonate as a building material. describe the use of lime in agriculture.