1. Covalent bonding

2. Learning Objectives Candidates should be able to:  describe, including the use of ‘dot-and-cross’ diagrams, covalent bonding, as in hydrogen; oxygen; chlorine; hydrogen chloride; carbon dioxide; methane; ethene.  describe covalent bonding in terms of orbital overlap.

3. Starter Activity

4. Forming a bond

5. Shared pair of electrons

6. Electron density maps for a hydrogen molecule Electron density map

7. This single σ covalent bond can be simply represented as: orH – H Representing a covalent bond

8. Overlap of atomic orbitals Overlap of one s and one p orbital – e.g. HF Overlap of two p-orbitals – e.g. F 2

9. Dot-cross diagrams Oxygen EtheneHydrogen chloride Carbon dioxide MethaneChlorine

10. Unexpected structures !! Breaking the octet rule

11. Bonding in CH 4 – promotion of an electron C 1s 2 2s 2 2p 2

12. Bonding in CH 4 – hybridisation

13. Hybridisation in PCl 5

14. Co-ordinate bonding

15. Learning Objectives Candidates should be able to describe, including the use of ‘dot-and-cross diagrams, co-ordinate (dative covalent) bonding, as in the formation of the ammonium ion and in the Al 2 Cl 6 molecule. A co-ordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom.

16. Starter Activity

17. Dot-cross diagrams H2SH2SSF 6 SiCl 4 CH 3 OHCO

18. Reaction between NH 3(g) and HCl (g)

20. Reaction between H 2 O and HCl

21. Electron-deficient BF 3

22. Aluminium chloride vapour

24. Electronegativity

25. Learning Objectives Candidates should be able to explain the origin of polar bonds, with reference to electronegativity differences between atoms. Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons.

26. Electronegativity values

27. Starter Activity

28. Why are bonds like bears? …’cos some of them are POLAR!

29. What happens if two atoms of equal electronegativity bond together? What happens if B is slightly more electronegative than A? Electronegativity differences

30. Representing polar bonds

31. MoleculeElectronegativity difference Dipole/debye HCl0.91.03 HBr0.70.78 HI0.40.38 Representing polar bonds

32. What happens if B is a lot more electronegative than A? An ionic bond is formed!!! Large electronegativity difference

33. Type of bond Electronegativity difference Non-polar Covalent  0.5 Polar covalentBetween 0.5 and 1.7 Ionic  1.7 As a rough guide:

34. Ionic with increasing covalent character

35. Positive cation Negative anion Polarisation of Anions Truly ionic Ionic with some covalent character

36. PropertyCation is the most powerful polarising agent when.... Anode is most easily polarised when... Charge Radius High Small High Large Polarisation of Anions

37. Do polar bonds make polar molecules? CCl 4 molecule is tetrahedral - the partial negative charges on the Cl atoms are distributed pretty symmetrically around the molecule. The partial positive charge on the C is buried in the center of the molecule.

38. The most electronegative element is fluorine. If you remember that fact, everything becomes easy, because electronegativity must always increase towards fluorine in the Periodic Table.

39. Trends in electronegativity

40. Ionic bonding

41. Learning Objectives Candidates should be able to describe ionic (electrovalent) bonding, as in sodium chloride and magnesium oxide, including the use of ‘dot-and-cross’ diagrams.

43. “The name’s Bond, Ionic Bond – taken, not shared!!!

44. Starter Activity

45. Approaching atoms

46. Ionic bonding

47. MgOCaCl 2 Al 2 O 3 K2OK2O

48. Why CaCl 2 and not CaCl or CaCl 3 ?

50. Positive cation Negative anion Polarisation of Anions Truly ionic Ionic with some covalent character

51. PropertyCation is the most powerful polarising agent when.... Anode is most easily polarised when... Charge Radius High Small High Large Polarisation of Anions

52. Shapes of molecules

53. Learning Objectives Candidates should be able to explain the shapes of, and bond angles in, molecules by using the model of electron-pair repulsion.

54. Starter Activity

55. 180 o Linear BeCl 2 Starter Activity

56. Balloon molecules 2 Balloons give a linear geometry 3 Balloons give a trigonal planar geometry 4 Balloons give a tetrahedral geometry 5 Balloons give a trigonal bipyramidal geometry 6 Balloons give an octahedral geometry

57. Electrons in outer shell of central atom Electrons added from other atoms (and any charge on an ion) No. of pairs of electrons No. of bonding pairs Diagram of molecule (including bond angles) Description of shape BF 3 3 333 Trigonal planar CCl 4 4444 Tetrahedral NH 3 5343 Trigonal pyramidal H2OH2O 6242 Bent (or V- shaped) Shapes of molecules

58. SF 6 6666 Octahedral CO 2 4422 Linear C2H6C2H6 4444 Tetrahedral C2H4C2H4 4433 Trigonal planar ClF 4 - 7564 Square planar Shapes of molecules

59. Metallic bonding

60. Lesson Objectives Candidates should be able to describe metallic bonding in terms of a lattice of positive ions surrounded by mobile electrons.

61. Starter Activity

62. This is sometimes described as "an array of positive ions in a sea of electrons". Metallic Bonding

63. Close packed structures? Dense metals Group 1 metals

64. Malleability

65. Metal grains

66. Intermolecular forces

67. Lesson Objectives Candidates should be able to describe intermolecular forces (van der Waals’ forces), based on permanent and induced dipoles, as in CHCl 3(l), Br 2(l) and the liquid noble gases.

68. Starter Activity

69. Particles in solids, liquids and gases In a liquid or a solid there must be forces between the molecules causing them to be attracted to one another, otherwise they would move apart from each other and become a gas.

70. Intermolecular attractions are attractions between one molecule and a neighbouring molecule. Intermolecular forces

71. The lozenge-shaped diagram represents a small symmetrical molecule - H 2, perhaps, or Br 2. The even shading shows that on average there is no electrical distortion (i.e. the molecule is non-polar). How do intermolecular (or van der Waals) forces arise? Temporary or instantaneous dipoles

72. Electrons are mobile. The constant "sloshing around" of the electrons in the molecule causes rapidly fluctuating dipoles even in the most symmetrical molecule. It even happens in monatomic molecules - molecules of noble gases, like helium, which consist of a single atom.

73. Temporary dipole - induced dipole interaction

74. How temporary dipoles give rise to intermolecular attractions

75. This diagram shows how a whole lattice of molecules could be held together in a solid using van der Waals’ forces. An instant later, of course, you would have to draw a quite different arrangement of the distribution of the electrons as they shifted around - but always in synchronisation. van der Waals’ forces

76. van der Waals lattice

77. helium -269°C neon -246°C argon -186°C krypton -152°C xenon -108°C radon-62°C How molecular size affects the strength of the dispersion forces The boiling points of the noble gases are:

78. How molecular size affects the strength of the dispersion forces There is a gradual increase in the very low boiling temperatures of the noble gases with increasing atomic size. As the size of the atoms increases the number of electrons increases and the magnitude of the van der Waals forces increases.

79. How molecular shape affects the strength of the temporary dipole interactions Butane has a higher boiling point because the intermolecular forces are greater. The molecules are longer and can lie closer together than the shorter, fatter 2- methylpropane molecules.

80. Permanent dipoles

82. Hydrogen bonding

83. Lesson Objectives Candidates should be able to describe hydrogen bonding, using ammonia and water as simple examples of molecules containing N-H and O-H groups.

84. Starter Activity

85. The increase in boiling point happens because the molecules are getting larger with more electrons, and so van der Waals forces become greater. Boiling points of the Group 4 hydrides

86. Boiling points of the hydrides in Groups 5, 6 and 7.

87. The origin of hydrogen bonding

88. Hydrogen bonding is a particularly strong intermolecular force that involves three features:  a large dipole between an H atom and the highly electronegative atoms N, O or F;  the small H atom which can get very close to other atoms;  a lone pair of electrons on another N, O or F, with which the positively charge H atom can line up. The origin of hydrogen bonding

89. Drawing hydrogen bonds 1 mark for indicating bond polarity 1 mark for showing lone pair 1 mark for showing H-bond

90. Hydrogen bonding accounts for many of the other unusual properties of water including:  its high specific heat capacity  its very high surface tension  its high viscosity and  the low density of ice compared to water Hydrogen bonding in water

91. Which type of intermolecular force?

92. A summary B onding, structure and properties B onding, structure and properties

93. Lesson Objectives Candidates should be able to:  describe, interpret and/or predict the effect of different types of bonding on the physical properties of substances.  describe, in simple terms, the lattice structure of a crystalline solid which is ionic, simple molecular, giant molecular, hydrogen-bonded and metallic.  suggest from quoted physical data the type of structure and bonding present in a substance.

94. Starter Activity

95. The properties of substances are decided by their bonding and structure.  Bonding means the way the particles are held together: ionic, covalent, metallic or weak intermolecular bonds.  Structure means the way the particles are arranged relative to one another. You have already met the major types of structure at IGCSE. Bonding, structure and properties

97. Bond Average bond enthalpy/ kJmol -1 Bond length/nm C – C+3470.154 C = C+6120.134 C ≡ C +8380.120 C – H+4130.108 O – H+4640.096 C – O+3580.143 C = O+8050.116 Bond energies

98. GIANT LATTICECOVALENT MOLECULAR IonicCovalent networkMetallicSimple molecularMacromolecular What substances have this type of structure? Compounds of metals with non-metals. Some elements in Group 4 and some of their compounds. Metals Some non-metal elements and usually some non-metal/non- metal compounds. Polymers Examples NaClSiO 2 CuH2OH2OPoly(ethene) What type of particle does it contain? ionsatoms positive ions and delocalised electrons molecules How are the particles bonded together? Strong ionic bonds; attraction between oppositely charged ions Strong covalent bonds; attraction of atoms’ nuclei for shared electrons Strong metallic bonds; attraction of atoms’ nuclei for delocalised electrons Weak intermolecular bonds between molecules; strong covalent bonds between atoms within each molecule. What are the typical properties? M. pt and b.pt. highvery highgenerally highlow moderate (often decompose on heating) Hardness hard but brittle very hard (if 3D) hard but malleable softvariable Electrical conductivity conduct when (l) or (aq) do not normally conduct conduct when (s) or (l) do not conduct do not normally conduct Solubility in water often solubleinsoluble insoluble (but some react) usually insoluble (but some H-bond) sometimes soluble Solubility in non-polar solvents (e.g. hexane) insoluble usually soluble sometimes soluble Structure table

99. The modern use of materials

100. Lesson Objectives Candidates should be able to:  Explain the strength, high melting point and insulating properties of ceramics in terms of their giant molecular structure.  Relate the uses of ceramics to their properties.  Describe and interpret the uses of the metals aluminium and copper (and their alloys) in terms of their physical properties.  Understand that materials are a finite resource and the importance of recycling processes.

101. Starter Activity

102. Five most common metals Aluminium Copper Zinc Steel Brass

103.  Low density, corrosion resistance and strength make it ideal for construction of aircraft, lightweight vehicles, and ladders.  Malleability, low density, corrosion resistance and good thermal conduction make it a good material for food packaging.  Good electrical conduction, corrosion resistance and low density leads to its use for overhead power cables hung from pylons (low density gives it an advantage over copper). Uses of aluminium

104. Uses of copper  Copper is an excellent conductor of electricity and heat.  Copper is soft and malleable.  Copper is very unreactive and therefore corrosion resistant.

105. Copper Alloy Other metal it contains Main properties Uses Brass Zinc Fairly soft and malleable Screws and hinges Bronze TinStrong Propellors and bearings Alloys of copper

106. A mineral is a naturally occurring solid formed through geological processes that has a characteristic chemical composition, a highly ordered atomic structure, and specific physical properties. Ceramics Ceramic: Any of various hard, brittle, heat-resistant and corrosion-resistant materials made by shaping and then firing a nonmetallic mineral.

107. Ceramics furnace: an enclosed chamber in which heat is produced to heat buildings, destroy refuse, smelt or refine ores, etc. heat shieldsglass and crockery furnace linings brake pads electrical insulators

108. Raw materials extracted (removed by chemical means) from the Earth cannot last forever. Although some materials are more (present in great quantity) abundant than others, they are all finite (have a limit) resources. Increasing demand for raw materials (items used to produce something else), coupled with ever growing problems of waste disposal, have led to considerable interest in recycling (processing for reuse) waste. Recycling has a number of possible advantages (beneficial factors):  It leads to reduced demand for new raw materials;  It leads to a reduction in environmental damage (harm to the surroundings);  It reduces the demand for landfill sites (a place for burying waste) to dump waste;  It reduces the cost of waste disposal;  It may reduce energy costs. Recycling

109. The kinetic-molecular model of liquids

110. Lesson Objectives Candidates should be able to describe using a kinetic- molecular model, the liquid state; melting; vaporisation and vapour pressure.

111. Starter Activity

112. SolidLiquidGas Arrangement of particles very orderly short-range order, longer range disorder almost complete disorder Movement of particles vibrate about fixed positions some movement from place to place continuous, rapid, random movement Proximity of particles close (~10 -10 m) far apart (~10 -8 m) Compressibility of substance very low high Conduction of heat poor except metals and graphite metals very good; others poor very poor The Kinetic-Molecular Model of Liquids

113. Liquids do not have a fixed __________ because the particles can move about. However, they remain very __________ together. This shows that the inter-particle forces have not been __________ broken. If sufficient __________ is supplied, the particles overcome the inter- particle forces almost completely and __________ from the liquid. This is called __________ or boiling. The energy required to boil a liquid is always __________ than that required to melt the same substance and is a better __________ of the strength of inter-particle forces.

114. Vapour pressure Vapour pressure is the pressure of a vapour over a liquid at equilibrium.

115. Vapour pressure Even at low temperature there are particles with high energy.

116. Vapour pressure At equilibrium, the rate at which molecules leave the liquid equals the rate at which molecules join the liquid.

117. Measuring vapour pressure

118. More on ideal gases

119. Lesson Objectives Candidates should be able to explain qualitatively in terms of intermolecular forces and molecular size the limitations of ideality at very high pressures and very low temperatures.

120. Starter Activity

121. More on ideal gases

122. Y is hydrogen. It is closest to ideal under all conditions. Hydrogen has the weakest intermolecular forces and is the smallest molecule. Z is ammonia. It is the least ideal at lower pressures. Ammonia molecules can hydrogen bond. X is nitrogen. Deviates greatly from ideality at high pressures where its larger molecular volume becomes important.