1. IB DP1 Chemistry Bonding What makes atoms join together to make compounds?

2. Topic 4: Bonding (12.5 hours) 4.1 Ionic bonding 4.1.1 Describe the ionic bond as the electrostatic attraction between oppositely charged ions. 4.1.2 Describe how ions can be formed as a result of electron transfer. 4.1.3 Deduce which ions will be formed when elements in groups 1, 2 and 3 lose electrons. 4.1.4 Deduce which ions will be formed when elements in groups 5, 6 and 7 gain electrons. 4.1.5 State that transition elements can form more than one ion. 4.1.6 Predict whether a compound of two elements would be ionic from the position of the elements in the periodic table or from their electronegativity values. 4.1.7 State the formula of common polyatomic ions formed by non- metals in periods 2 and 3. 4.1.8 Describe the lattice structure of ionic compounds. 4.2 Covalent bonding 4.2.1 Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei. 4.2.2 Describe how the covalent bond is formed as a result of electron sharing. 4.2.3 Deduce the Lewis (electron dot) structures of molecules and ions for up to four electron pairs on each atom. 4.2.4 State and explain the relationship between the number of bonds, bond length and bond strength. 4.2.5 Predict whether a compound of two elements would be covalent from the position of the elements in the periodic table or from their electronegativity values. 4.2.6 Predict the relative polarity of bonds from electronegativity values 4.2.7 Predict the shape and bond angles for species with four, three and two negative charge centres on the central atom using the valence shell electron pair repulsion theory (VSEPR). 4.2.8 Predict whether or not a molecule is polar from its molecular shape and bond polarities. 4.2.9 Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C 60 fullerene). 4.2.10 Describe the structure of and bonding in silicon and silicon dioxide. 4.3 Intermolecular forces 4.3.1 Describe the types of intermolecular forces (attractions between molecules that have temporary dipoles, permanent dipoles or hydrogen bonding) and explain how they arise from the structural features of molecules. 4.3.2 Describe and explain how intermolecular forces affect the boiling points of substances. 4.4 Metallic bonding 4.4.1 Describe the metallic bond as the electrostatic attraction between a lattice of positive ions and delocalized electrons. 4.4.2 Explain the electrical conductivity and malleability of metals. 4.5 Physical properties 4.5.1 Compare and explain the properties of substances resulting from different types of bonding.

3. Ionic Bonding

4. Crystals: 7 ‘perfect’ crystal shapes

6. Halite- rock salt- sodium chloride

7. Sodium chloride is an ionic compound with ions arranged in a lattice

8. Ions charged particles with electrostatic attraction between them Na+ Cl-

9. Sodium and chloride ions formed when electrons transfer Na+Cl  Na + +Cl - 2,8,1 2,8,7 2,8 2,8,8

10. Ions  Group 1: H +, Li +, Na +, K +, Rb +, Cs +, Fr +  Group 2: Be 2+, Mg 2+, Ca 2+, Sr 2+, Ba 2+  Group 3?/13: B 3+, Al 3+, Ga 3+  Group 6?/16: O 2-, S 2-,  Group 7?/17: F -, Cl -, Br -, I -

11. Which is the smallest ion? Na + Al +3 Cl - P 3-

12. Two or more electrons can be transferred  Different sized atoms give different mineral structures as they pack in a different way Hexagonal Beryl crystal; Image Wikipedia

13. What is the formula of iron (III) oxide? Fe2O FeO Fe3O2 Fe2O3

14. Polyatomic ions: charge distributed over more than one atom For example phosphate, PO 4 -3 can be found in products of reactions of phosphoric acid

15. Some common polyatomic ions  Nitrate NO 3 -  Hydroxide OH -  Sulphate SO 4 2-  Carbonate CO 3 2-  Hydrogen carbonate HCO 3 - (Bicarbonate)  Phosphate PO 4 3-  Ammonium NH 4 +

16. Common Anions Common NameFormulaAlternative name Simple Anions ChlorideCl − FluorideF−F− BromideBr − OxideO 2− Polyatomic anions CarbonateCO 3 2- Hydrogen carbonate HCO 3 − bicarbonate HydroxideOH − NitrateNO 3 2- PhosphatePO 4 3- SulfateSO 4 2- Anions from Organic Acids EthanoateCH 3 COO − acetate MethanoateHCOO − formate EthandioateC 2 O 4 −2 oxalate CyanideCN - Common Cations Common NameFormulaAlternative name Simple Cations AluminiumAl 3+ CalciumCa 2+ Copper(II)Cu 2+ cupric HydrogenH+H+ Iron(II)Fe 2+ ferrous Iron(III)Fe 3+ ferric MagnesiumMg 2+ Mercury(II)Hg 2+ mercuric PotassiumK+K+ kalic SilverAg + SodiumNa + natric Polyatomic Cations AmmoniumNH 4 + HydroniumH3O+H3O+

17. Careful with...  name of atom can change when ion is formed chlorine atom (Cl)  chloride ion (Cl - )  -ate is often a polyatomic ion with oxygen eg sulphate, phosphate, etc.  different ions often have similar names...  nitrate NO 3 -  nitrite NO 2 -  nitride N -3

18. What is the formula of ammonium sulphate?  NH4SO4  (NH4)2SO4  NH4(SO4)2  SO4(NH4)2

19. d-block (transition elements) can have variable valencies Mn2+manganese(II) Mn3+manganese(III) Mn4+manganese(IV) Ni2+nickel(II)/nickelous Ni3+nickel(III)/nickelic Pb2+lead(II)/plumbous Pb4+lead(IV)/plumbic Cr2+chromium(II)/chromous Cr3+chromium(III)/chromic Cu1+copper(I)/cuprous Cu2+copper(II)/cupric Fe2+iron(II)/ferrous Fe3+iron(III)/ferric Hg2+mercury(I)/mercurous

20. Covalent bonding

21. Define electronegativity Electronegativity is the tendency of an atom to attract electrons towards itself. The atoms with higher values attract electrons more strongly. Highest flourine (and rest of groups 7,6,5) FONClBrISCH Wikipedia table

22. How ionic is an ionic compound?  bigger difference in electronegativity  more ionic  (‘ionic’ usually  e-neg> 1.8 difference)  usually metal + non-metal

23. Which aluminium compounds will be ionic? atomAlFOClBr electronegativity1.54.03.53.02.8 Formula of aluminium compound  e-neg ‘Ionic’ or ‘covalent’?

24. ‘Sharing’ electrons  e-neg < 1,7 covalent bonding forms molecules Often between non-metals

25. Covalent bond formation- valence electrons

26. 2, 4 or 6 electrons?  Single bond: the two atoms share two electrons (1 pair)  Double bond: the two atoms share four electrons (2 pairs)  Triple bond: the two atoms share six electrons (3 pairs)

27. Lewis structures (dot structures) show valence electrons in pairs as dots, crosses or lines

28. skeletal formula for complex organic molecules

29. Condensed formula propanol CH 3 CH 2 CH 2 OH

30. Coordinate covalent bond (dative bond) both electrons in the bond from the same atom once formed, is the same as any other covalent bond

31. Bond lengths and Bond strengths  As the number of shared electrons increases (single to triple) the bond lengths shortens and the bond energy increase BondBond typeLengths (pm)Energy (kJ/mol) CCSingle154347 CCDouble134614 CCTriple120839

32. Which bond has the highest bond polarity, δ H-H Cl-Cl Al-F Al-Br

33. Non-polar covalent bond In, H 2 the two electrons in the bond are shared equally between the two hydrogen atoms.  H-H  e-neg =0.  The electron distribution is symmetrical.

34. Polar covalent bond  If two different atoms form a covalent bond there will be a difference in  e-neg.  The atom with highest electronegativity will have the electrons closer; they don’t share equally.  Unsymmetrical electron distribution.

35. Bonds 100% Covalent bond  Polar covalent bond  Ionic bond % ionic character of a bond: 0-90% (there are no 100% ionic compounds)

36. Molecular shapes

37. What shape are molecules?  VSEPR theory (Valence shell electron pair repulsion)  pairs of electrons repel and sit as far away as possible from each other  double and triple bonds count as a pair

38. VSEPR: electron repulsion  molecular shape  Structure of molecule given by pairs of electrons arranging around an atom to be as far apart as possible  non-bonded pairs repel more than bonded pairs  double and triple bonds count as one

39. Build molecules from plasticine and straws  bond: 3cm length of straw  atom: 1cm diameter plasticine ball  unbonded pair of electrons 1cm straw length

40. Number of charge centres Name of shapeBond angles (s)Example 2linear180BeCl 2 3trigonal planar120BF 3 4tetrahedral109.5CH 4 5trigonal bipyramidal90, 120, 180 6octahedral90, 180 Shapes of simple molecules http://en.wikipedia.org/wiki/Phosphorus_pentafluoride http://en.wikipedia.org/wiki/Sulphur_hexafluoride http://en.wikipedia.org/wiki/Boron_triflouride

41. Methane, Water and Ammonia greater repulsion between non-bonding pairs  smaller bond angles than predicted

42. Intermolecular forces Why do molecules stick together to form liquids and solids?

43. Intermolecular forces hold molecules together, affecting physical properties  Melting and boiling points  Strength  Flexibility  Viscosity

44. Intermolecular forces Hydrogen bondstrong Dipole-dipoleweaker van der Waal’s forces weakest

45. Why do molecules attract each other to make liquids and gases? Intermolecular forces: electrostatic attraction between  permanent dipoles (polar molecules)  permanent dipole and a temporary dipole (induced polarity)  temporary diploes (induced polarity)

46. Why do molecules attract each other? electrostatic attraction between…  permanent dipoles (in polar molecules)  temporary diploes A dipole is a overall charge imbalance in a molecule. Which of the following molecules are polar?

47. Induced dipoles in all molecules (van der Waal’s forces) Image: http://www.uwec.edu/boulteje/Boulter103N otes/11December.htm Movements in electron cloud  Temporary dipoles. Temporary dipole in one molecule can induce a temporary dipole in another.

48. van der Waals forces  The strength increases with molar mass of the molecule. E.g. He b.p 4 K : Xe b.p. 165 K.  Only effective over short range so the molecule “area” is also important. E.g: Pentane, C 5 H 12, b.p. 309 K Dimethylpropane, (CH 3 ) 4 C b.p. 283 K

49. Is a molecule polar? A polar molecule  Has polar covalent bonds. Look at the difference in electronegativity ( FONClBrISCH) AND  Unsymmetrical shape according to charge distribution. Otherwise it will be a non-polar molecule.

50. Molecular polarity Images: http://en.wikipedia.org/wiki/Molecular_polarity HF H2O NH3

51. Molecular polarity  http://phet.colorado.edu/en/simulation/molecule-polarity http://phet.colorado.edu/en/simulation/molecule-polarity

52. Dipole-dipole Electrostatic attraction between molecules with permanent dipoles. Stronger than vdW. Hydrogen chloride M= 36,5 g/mol b.p. 188 K Fluorine M= 38 g/mol b.p. 85K

53. Induced dipole Image: http://www.uwec.edu/boulteje/Boulter103N otes/11December.htm

54. Polar and non-polar liquids are immiscible Image: http://en.wikipedia.org/wiki/Petroleum

55. Hydrogen bonding  H bonded to a highly electronegative element eg F, O or N  proton strongly attracts electronegative element in another molecule  important in water Image: http://en.wikipedia.org/wiki/Induced_dipole#Debye_.28induced_dipole.29_force

56. Hydrogen bond  In molecules that contain Hydrogen bonded to Oxygen, Nitrogen or Fluorine (high electronegativity and non- bonding electron pair).  Interaction of the non-bonding electron pair in one molecule and hydrogen (with high positive charge) in another molecule.

57. Examples  H 2 O b.p.=100 o CH 2 S b.p.= -61 o C  NH 3 b.p.= -33 o CPH 3 b.p.= -88 o C  C 3 H 8 b.p. CH 3 CHO C 2 H 5 OH b.p. 20 o C 42 o C 78 o C

58. Examples  H 2 O b.p.=100 o CH 2 S b.p.= -61 o C  NH 3 b.p.= -33 o CPH 3 b.p.= -88 o C  C 3 H 8 b.p. CH 3 CHO C 2 H 5 OH b.p. 20 o C 42 o C 78 o C

59. Ice Image: http://en.wikipedia.org/wiki/Ice

60. Trends in physical properties

61. How strong are the forces between molecules? Bond typeDissociation energy (kJ/mol) Covalent1600 Hydrogen bonds50–70 Permanent dipoles2–8 Induced dipoles<4 Data: http://en.wikipedia.org/wiki/Induced_dipole#Debye_.28induced_dipole.29_force

62. Trends in physical properties melting point /Cboiling point /C Flourine-220-188 Chlorine-102-34 Bromine-759 Iodine114184 Astatine302337 Plot one graph showing melting point and boiling point (in Kelvin) against molar mass for the halogens Describe the pattern (2 sentences) Explain the pattern (2 sentences) Data: http://en.wikipedia.org/wiki/Halogen

63. How strong are the forces between molecules? Bond typeDissociation energy (kJ/mol) Covalent1600 Hydrogen bonds50–70 Permanent dipoles2–8 Induced dipoles<4 Data: http://en.wikipedia.org/wiki/Induced_dipole#Debye_.28induced_dipole.29_force

64. Allotropes: different structural forms of the same element http://catalog.flatworldknowledge.com/bookhub/4309?e=averill_1.0-ch18_s04 Oxygen O2 diatomic oxygen O3 ozone

65. Allotropes of Carbon

66. Diamond  Hard, colourless, insulator  Tetrahedral, giant structure  Covalent bond => sp 3 orbitals.

67. Graphite  Slippery, black, conductor  Layers of fused six-membered rings. Each carbon surrounded by three others in a planar trigonal arrangement => sp 2 + p- orbital  The p-orbital is perpendicular to the layer and give close packed p-orbitals  stabilise the layers  Delocalisation of electrons => electrical conductivity

68. Fullerene, C 60  Spherical molecule. Looks like a football. 12 pentagons and 20 hexagons.  Bonds: C 60 –hydration  C 60 H 60 (C 2 H 4 + H 2  C 2 H 6 ; 1 H 2 / double bond) Each carbon has a double bond

69. Silicon  Metalloid, Semiconductors, non-metallic structure  Similar structure as diamond.

70. Silicon dioxide  SO 2 Silica, giant structure similar to diamond  Silicates, SiO 4, tetrahedrical, silicon-oxygen single bond

71. Physical properties  Melting points (impurities lower the melting point)  Boiling points  Volatility (how easy a compound will convert to gas)  Electrical conductivity  Solubility

72. Properties Structure type Property Giant Metallic Giant Ionic Giant Covalent Molecular Covalent Hardness and malleability Variable hard-ness, malleable rather than brittle Hard and brittle Usually soft and malleable unless hydrogen bonded Melting and boiling points Variable dep. On No of valence e - HighVery HighLow Electrical and thermal conductivity Good in all states Not as solids, conduct in (aq) or (l) No Solubility Insoluble, except as alloys In Water mostlyInsoluble Often more soluble in other than water except if H-bonded ExamplesIron, copperNaCl, Na 2 SO 4 Diamond, SiO 2 (Sand) CO 2, Cl 2, ethanol, sugar

73. Ionic salts  Typical properties  Hard, brittle,  Conduct electricity in solution or melted.  High melting points => Strong bonds  Hydration of Ion in Water solution

74. Metallic bond  Metals have low electronegativity.  The atoms are packed close together in a lattice.  The valence electrons are delocalised among all atoms.  The valence electron have no “home”  The atoms can be seen as positive ions in a see of electrons that keep them together.

75. This can explain the metallic properties  Electrical conductivity: electrons float around. If you put in one, one will fall out.  Malleability (smidbarhet) and Ductility (sträckbarhet): if the atom is pushed from its location the electron will follow. The bond is between the ion and the electrons not between the ions.

76. Investigate a physical property of a mixture related to intermolecular forces Quantitative independent variable (cause) Quantitative dependent variable (effect) viscosity, deflection by charged object, or other physical property

77. Links  Ionic bonding http://www.teachersdomain.org/asset/lsps07_int_ionicbonding / http://www.teachersdomain.org/asset/lsps07_int_ionicbonding /  Covalent bonding http://www.teachersdomain.org/asset/lsps07_int_covalentbon d/ http://www.teachersdomain.org/asset/lsps07_int_covalentbon d/

78. Polarity links  http://phet.colorado.edu/en/simulation/molecule-polarity http://phet.colorado.edu/en/simulation/molecule-polarity  Viscosity http://www.youtube.com/watch?v=3KU_skfdZVQhttp://www.youtube.com/watch?v=3KU_skfdZVQ  States of matter http://phet.colorado.edu/en/simulation/states- of-matterhttp://phet.colorado.edu/en/simulation/states- of-matter

79. Polarity links  http://phet.colorado.edu/en/simulation/molecule-polarity http://phet.colorado.edu/en/simulation/molecule-polarity  http://antoine.frostburg.edu/chem/senese/101/liquids/faq/h-bonding-vs-london- forces.shtml http://antoine.frostburg.edu/chem/senese/101/liquids/faq/h-bonding-vs-london- forces.shtml  States of matter http://phet.colorado.edu/en/simulation/states-of-matterhttp://phet.colorado.edu/en/simulation/states-of-matter  http://employees.oneonta.edu/viningwj/modules/CI_dipoleinduced_dipole_forces _13_5a.html http://employees.oneonta.edu/viningwj/modules/CI_dipoleinduced_dipole_forces _13_5a.html  Notes: http://www.uwec.edu/boulteje/Boulter103Notes/11December.htmhttp://www.uwec.edu/boulteje/Boulter103Notes/11December.htm  Snowflakes: http://www.its.caltech.edu/~atomic/snowcrystals/class/class.htmhttp://www.its.caltech.edu/~atomic/snowcrystals/class/class.htm  Ice crystals http://www.edinformatics.com/interactive_molecules/ice.htmhttp://www.edinformatics.com/interactive_molecules/ice.htm

80. Links  http://phet.colorado.edu/en/simulation/molecule-shapes http://phet.colorado.edu/en/simulation/molecule-shapes  http://en.wikipedia.org/wiki/Phosphorus_pentafluoride http://en.wikipedia.org/wiki/Phosphorus_pentafluoride  http://en.wikipedia.org/wiki/Sulphur_hexafluoride http://en.wikipedia.org/wiki/Sulphur_hexafluoride  http://en.wikipedia.org/wiki/Boron_triflouride http://en.wikipedia.org/wiki/Boron_triflouride

81. Teaching notes