1. Periodicity Mr Murphy

2. Main Menu Menu:  Lesson 1 – The Periodic Table Lesson 1 – The Periodic Table  Lesson 2 – Physical Properties Lesson 2 – Physical Properties  Lesson 3 – Chemical Trends Lesson 3 – Chemical Trends  Lesson 4 – The Period 3 Oxides Lesson 4 – The Period 3 Oxides  Lesson 5 – HL – The Period 3 Oxides and Chlorides Lesson 5 – HL – The Period 3 Oxides and Chlorides  Lesson 6 – HL – Transition Metals - Introduction Lesson 6 – HL – Transition Metals - Introduction  Lesson 7 – HL – Coloured Complexes and Catalysts Lesson 7 – HL – Coloured Complexes and Catalysts  Lesson 8 – Test Lesson 8 – Test  Lesson 9 – Test Debrief Lesson 9 – Test Debrief

3. Main Menu Lesson 1 The Periodic Table

4. Main Menu Overview  Copy this onto an A4 page. You should add to it as a regular review throughout the unit.

5. Main Menu We Are Here

6. Main Menu The Traditional Based on Mendeleev’s work. Easiest to use and display.

7. Main Menu Dmitri Mendeleev’s Periodic Table The one that started it all off.

8. Main Menu Wide Format Periodic Table Shows true position of the f-block (lanthanides and actinides)

9. Main Menu Janet Periodic Table Elements arranged in order of orbital filling. Used frequently by physicists.

10. Main Menu Benfey Periodic Table Spiral form shows the steady increase in atomic number.

11. Main Menu Stowe Periodic Table Emphasises the symmetrical nature of the increase in quantum numbers.

12. Main Menu Zymaczynski Periodic Table Another way to show the symmetry in the underlying quantum numbers.

13. Main Menu Giguere Periodic Table A 3D representation emphasising the s, p, d and f blocks

14. Main Menu The Structure of the Periodic Table PERIODS GROUPS

15. Main Menu Groups and Periods  Groups  Elements show similar chemical properties  Elements show similar trends in their chemical properties  Periods  As you move across periods, changes in the chemical and physical properties that are repeated in the next period  This is what ‘period’ and ‘periodic’ refers to

16. Main Menu The periodic table and electron configuration  How does an element’s position in the PT relate to its electron configuration?

17. Main Menu Being Mendeleev  The first widely accepted periodic table was produced by the Russian chemist Dmitri Mendeleev  It was a tremendous example of scientists as risk-takers as it was able to make a number of predictions thought unlikely at the time  Complete the exercise here in which you will use the information available to Mendeleev to construct your own periodic tablehere

18. Main Menu Key Points  The periodic table arranges the elements according to:  Their chemical properties  Their electronic structure

19. Main Menu Lesson 2 Physical Properties

20. Main Menu Refresh  Nitrogen and silicon belong to different groups in the periodic table. a) Distinguish in terms of electronic structure, between the terms group and period. b) State the maximum number of orbitals in the n = 2 energy level.

21. Main Menu We Are Here

22. Main Menu Lesson 2: Physical Properties  Objectives:  Identify and explain the trends in the physical properties of the first 20 elements including:  Atomic radius  Ionic radius  First ionisation energy  Electronegativity  Melting point  Use Microsoft Excel to produce a spreadsheet to graph the above physical data

23. Main Menu Atomic Radius  This is the ‘size’ of an atom  There is no simple measure as atoms do not have a well defined ‘edge’  We use the: covalent radius  This is half the distance between the nuclei of two atoms in a covalent bond  This means we don’t have values for the noble gases as they do not form bonds  Values range from are measured in picometres ( 1 pm = 1x10 -12 metres…a thousand-billionth of a metre) and range over:  270 picometres Francium  30 picometres for Hydrogen (helium would be smaller but does not form covalent bonds to be measured)  The main factors influencing atomic radius are:  Number of shells (the principal quantum number)  The charge in the nucleus

24. Main Menu Ionic Radius  This is the ‘size’ of an ion and is measured in a similar way to atomic radius  It is measured in a picometres with values ranging over:  272 pm for the Ge 4- ion  16 pm for the B 3+ ion  The main factors influencing ionic radius are:  Number of shells (the principal quantum number)…don’t forget this can be affected by the type of ion formed  The charge in the nucleus

25. Main Menu First ionisation energy  This is the energy required to remove one mole of electrons from one mole of gaseous atoms to form positive ions i.e.: A(g)  A + (g) + e -  Values range over:  393 kJ mol-1 for Caesium  1681 kJ mol-1 for Helium  Values are positive because this is an endothermic process  Values are influenced by:  Number of inner electron shells (and their shielding)  Charge on the nucleus  HL: At the finest level – repulsion between electrons in their orbitals

26. Main Menu Electronegativity  This is a measure of the degree to which an element attracts the shared pair of electrons in a covalent bond  Again, this means there are no values for the noble gases  Values range over:  4.0 for Fluorine  0.7 for Francium  Values are influenced by:  Number of inner electron shells (and their shielding)  Charge on the nucleus  Values are unit-less as this is a relative measure

27. Main Menu Melting Point  This is the temperature (in Kelvin…i.e. Celsius + 273) at which an element melts  Values range over:  3935 K for Carbon  1 K for Helium  Values are influenced by:  Nature of bonding: giant covalent, giant ionic, metallic  Strength of bonding  Strength of intermolecular forces

28. Main Menu Trends in Physical Properties  You need to produce an Excel spreadsheet to help you analyse the physical data.  Use the blank here and follow the instructions on the instructions pageblank here  Once you have done this you need to use this to help you identify and explain the following trends:  Atomic and ionic radius, first ionisation energy, electronegativity and melting point  Down group I (alkali metals)  Down group VII (halogens)  Atomic and ionic radius, first ionisation energy, electronegativity  Across period 3  The general trend in electronegativity over the whole PT

29. Main Menu Key Points  Each of the following physical parameters follow trends and patterns in the PT  These patterns are generally explained by:  Charge in the nucleus  Number of electron shells  Electron shielding

30. Main Menu Lesson 3 Chemical Properties

31. Main Menu Refresh Which species has the largest radius? Do not use the data booklet…work it out! A. Cl – B. K C. Na + D. K +

32. Main Menu We Are Here

33. Main Menu Lesson 3: Chemical Properties  Objectives:  Understand the following trends in reactivity:  Alkali metals with water  Alkali metals with halogens  Halogens with halide ions  Complete an experiment to investigate the above

34. Main Menu Chemical Trends  Members of a group often have very similar reactivity.  You probably know that carbon will react with hydrogen to form methane, CH 4  You probably did not know that silicon will also react with hydrogen to form silane, SiH 4  Watch this demonstration to see some silane being madethis demonstration

35. Main Menu Three reactions to know  The Group I (alkali) metals react with water as follows:  Metal + Water  Metal Hydroxide + Hydrogen  The Group I (alkali) metals react with halogens (Group VII) as follows:  Metal + Halogen  Metal Halide  Halogens can react with halide ions as follows (using the example of bromide and chlorine):  Bromide + Chlorine  Chloride + Bromine

36. Main Menu Key Points  Alkali metals become more reactive down the group:  Due to the outer shell electron becoming increasingly easy to remove  Halogens become less reactive down the group:  Due to the increased numbers of electron shells (and thus shielding) causing them to attract electrons less strongly

37. Main Menu Lesson 4 Period 3 Oxides

38. Main Menu Refresh Which properties of the alkali metals decrease going down group 1? A. First ionization energy and reactivity B. Melting point and atomic radius C. Reactivity and electronegativity D. First ionization energy and melting point

39. Main Menu We Are Here

40. Main Menu Lesson 4: Period 3 Oxides  Objectives:  Understand and explain the trend in acid-base behaviour of the period 3 oxides  Complete an experiment to demonstrate the amphoteric nature of aluminium oxide

41. Main Menu The Period 3 Oxides ElementFormula of oxide Reaction of oxide with waterAcid/base nature Sodium*Na 2 ONa 2 O + H 2 O  2NaOHStrongly basic Magnesium*MgOSlight: MgO + H 2 O  Mg(OH) 2 Weakly basic AluminiumAl 2 O 3 Amphoteric SiliconSiO 2 Very weakly acidic Phosphorous*P 4 O 10 P 4 O 10 + 6 H 2 O  4 H 3 PO 4 Strongly acidic Sulphur*SO 2 SO 3 SO 3 + H 2 O  H 2 SO 4 Strongly acidic Chlorineno direct reaction but: Cl 2 O 7 Cl 2 O 7 + H 2 O  2 HClO 4 Strongly acidic Argonno oxides  There is a gradual transition from basic to acidic character, reflecting a gradual transition from metallic to non-metallic nature Note: you will only be tested on the elements marked with an asterisk, *

42. Main Menu Key Points  The oxides of period 3 display a gradual transition basic to acidic character  This reflects a gradual transition from metallic to non-metallic nature of the elements

43. Main Menu Lesson 5 Period 3 Oxides and Chlorides

44. Main Menu Refresh  Which oxides produce an acidic solution when added to water? I. P 4 O 10 II. MgO III. SO 3 A. I and II only B. I and III only C. II and III only D. I, II and III

45. Main Menu We Are Here

46. Main Menu Lesson 6 Transition Metal Complexes - Introduction

47. Main Menu Refresh  By reference to the structure and bonding in NaCl and SiCl 4 : a) State and explain the differences in electrical conductivity in the liquid state. b) Predict an approximate pH value for the solutions formed by adding each compound separately to water. Explain your answer.

48. Main Menu We Are Here

49. Main Menu Lesson 6: Transition Metal Complexes - Introduction  Objectives:  Describe the properties of transition metals  Understand the term ligands  Understand and explain the formation of transition metal complexes

50. Main Menu The Traditional Based on Mendeleev’s work. Easiest to use and display.

51. Main Menu The Transition Metals  A transition metal is an element in which at least one ion has a partially filled d-orbital  For example, Cu 2+ : 1s 2 2s 2 2p 6 3s 2 3p6 (4s 0 ) 3d 9  Properties of the transition metals include:  Variable oxidation states (for example iron: Fe 2+, Fe 3+, Fe 6+ )  Formation of coloured compounds (more later)  Catalytic properties (more later)  Formation of complex ions (much more later)

52. Main Menu Scandium and Zinc  Although in the first row of the d-block, these are not transition metals.  To understand why, write the full electron configuration for:  Sc and Sc 3+  Zn and Zn 2+

53. Main Menu Variable oxidation numbers (ions)  Transition metals have large numbers of electrons in d-orbitals,  This means the amount of energy required to remove the second electron is not much different to that required to remove the first and so on.  Some common oxidation states we need to know:  All of them in the +2 oxidation state  Cr(III), Cr(VI)  Mn(IV), Mn(VII)  Fe(III)  Cu(I)  Task: select 4 of these and write the electron configuration

54. Main Menu Ligands  A ligand is a species with a lone pair  Often negative ions  Common ligands include:  Water, H 2 O  Ammonia, NH 3  Chloride, Cl -  Hydroxide, OH -  Cyanide, CN -  Thiocyanate, SCN -

55. Main Menu Transition Metal Complexes  The lone pair on a ligand can form a dative covalent bond to a metal ion to form a transition metal complex.  This involves the ligands donating charge into the empty 4d and 4s orbitals (at least for the first-row of transition elements), not the partially occupied 3d orbitals.  [Fe(H 2 O) 6 ] 3+ [Fe(CN) 6 ] 3-  [Cu(Cl) 4 ] 2- [Ag(NH 3 ) 2 ] +

56. Main Menu Making transition metal complexes  Add potassium thiocyanate solution to a solution of iron (III)  Add conc. HCl (fume hood!) to 1 cm 3 of a strong solution of cobalt (II). Repeat but use conc. NH 3 instead (fume hood!).  Add dilute NH 3 to a copper (II) solution until no further change occurs  Record all observations  Suggest possible structures for the complexes you have formed and possible reaction equations

57. Main Menu Key Points  Transition metals form ions with partially filled d-orbitals  Ligands are species with lone pairs  Ligands will form dative covalent bonds to transition metals forming ‘complex ions’

58. Main Menu Lesson 7 Complex Colours and Catalysts

59. Main Menu Refresh  By reference to the structure and bonding in NaCl and SiCl 4 : a) State and explain the differences in electrical conductivity in the liquid state. b) Predict an approximate pH value for the solutions formed by adding each compound separately to water. Explain your answer.

60. Main Menu We Are Here

61. Main Menu Lesson 7: Complex Colours and Catalysts  Objectives:  Complete a group exercise to:  Understand the origin of colour in transition metal complexes  Understand the uses of transition metals as catalysts

62. Main Menu Key Points  The formation of complexes causes d-orbitals to split into two energy levels  Electron transitions between these energy levels give rise to their colour  Transition metals are hugely important for their catalytic properties