IB DP1 Chemistry Bonding

1. IB DP1 ChemistryBonding 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 C60 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 Cl- Na+

9. Sodium and chloride ions formed when electrons transfer

10. Ions • Group 1: H+, Li+, Na+, K+, Rb+, Cs+, Fr+ • Group 2: Be2+, Mg2+, Ca2+, Sr2+, Ba2+ • Group 3?/13: B3+, Al3+, Ga3+ • Group 6?/16: O2-, S2-, • Group 7?/17: F-, Cl-, Br-, I-

11. Which is the smallest ion? Na+ Al+3 Cl- P3-

12. Different sized atoms give different mineral structures as they pack in a different way Two or more electrons can be transferred 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 examplephosphate, PO4-3 can be found in products of reactions of phosphoricacid

15. Some common polyatomic ions • Nitrate NO3- • Hydroxide OH- • Sulphate SO42- • Carbonate CO32- • Hydrogen carbonate HCO3- (Bicarbonate) • Phosphate PO43- • Ammonium NH4+

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 egsulphate, phosphate, etc. • different ions often have similar names... • nitrate NO3- • nitrite NO2- • 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 Electronegativityis 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 De-neg> 1.8 difference) • usually metal + non-metal

23. Which aluminium compounds will be ionic?

24. ‘Sharing’ electrons De-neg < 1,7covalent 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 CH3CH2CH2OH

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

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

33. Non-polar covalent bond In, H2 the two electrons in the bond are shared equally between the two hydrogen atoms. • H-H De-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 De-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. 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. VSEPRfor 5- and 6-negative charge centres • Some molecules have expanded valence shell with five or six negative centres. Shapes similar to trigonalbipyramid and octahedron.

43. Trigonalbipyramid Two types of electron rich region: • Equatorial: 3 bonds with 120obetween. • Axial: 2 bonds with 180o between • Equatorial to Axial: 90o. • Non-bonding orbitals always occupy equatorial positions. • Many of the compounds that form trigonalbipyramids and octahedrons are fluorides because only highly electronegative atoms can increase the number of valence electrons. Fluoride is also quite small, but there is not enough space for larger ions.

44. Octahedron • All positions are equal. 90oangle • If two non-bonding orbitals: they take place opposite each other plane square shape.

45. Examples • PCl5 • SF6 • XeF4 • PF6-

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

47. Intermolecular forces hold molecules together, affecting physical properties • Melting and boiling points • Strength • Flexibility • Viscosity • Deflection in electric field • Volatility (how easy a compound will convert to gas) • Electrical conductivity • Solubility

48. Intermolecular forces Hydrogen bond strong Dipole-dipole weaker van der Waal’s forces weakest

49. 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) A dipole is a overall charge imbalance in a molecule.

50. Induced dipoles in all molecules (van der Waal’s forces) Movements in electron cloud Temporary dipoles. Temporary dipole in one molecule can induce a temporary dipole in another. Image: http://www.uwec.edu/boulteje/Boulter103Notes/11December.htm