SCH4U – UNIT 1 STRUCTURE AND PROPERTIES

1. SCH4U – UNIT 1STRUCTURE AND PROPERTIES CHAPTER 4 – CHEMICAL BONDING

2. Activity • With a partner discuss everything you remember about chemical bonding • Eg. Types of bonds? why? What happens?

3. 4.1 Types of Chemical Bonds • What are the two main types of chemical bonds? • Ionic: chemical bond between oppositely charged ions • Electrostatic attraction • Covalent: a chemical bond in which atoms share bonding electrons • Bonding Electron Pair: electron pair that is involved in bonding

4. Bond type depends on the attraction for electrons of the atoms involved • i.e. electronegativity

5. Ionic Compounds

6. How do these work? Metal + Non-Metal  Metal+ + Non-Metal- Low IE High IE Isoelectronic with noble gases Low EA High EA Opposites attract in no particular direction, considered non-directional Ions cling together in clusters known as crystals

7. Get a lattice structure • Lattice energy: energy change when one mole of an ionic substance is formed from its gaseous ions • Depends on: • Charge on the ions • Size of the ions

8. Ionic Compounds and Bonding • Properties – WHY? • Do not conduct electric current in the solid state • Conduct electric current in the liquid state • When soluble in water, form good electrolyte • Relatively high MP and B • Brittle, easily broken under stress

9. Covalent Bonds Balance of attractive and repulsive forces What are the forces acting here?

12. Octet Rule • Atoms share electrons so that they are surrounded by 8 electorns • # bonds = 8 - # valence electorn • Example: Carbon, Oxygen, Nitrogen • Two covalent bonds = double bond • Three covalent bonds = triple bond

13. Lewis Structures • Atoms and ions are stable if they have a full valence shell • Electrons are most stable when they are pair • Atoms form chemical bonds to achieve full valence shells of electrons • Full valence shell may occur by an exchange or by sharing electrons • Sharing – covalent; exchange - ionic

16. Polar Covalent Bonds • When electrons are shared unevenly in a covalent bond • Example: HF, H2O

17. Coordinate Covalent Bonds • Both electrons are contributed by one atom • Example: • NH4+ • H3O+ • CO • N2O • NHO3

18. Resonance Structures • Single bonds are longer than double bonds, which are longer than triple bonds • Example: SO3 • Resonance Structure: Electron pair is shared over three bond evenly • Delocalized electrons

19. Less than 8 • BeH2 • BCl3

20. More than 8 • Octet rule only applies to the first two periods • After that, can have expanded octets • Example: • PF5 • BrF5 • SiF63-

21. Practice - Worksheet • H2 • F2 • OF2 • O2F2

22. Valence Bond Theory and Quantum Mechanics • Covalent bonds occur when orbitals overlap and two electrons occupy the same region of space • Example: H2

23. HF • What are the electron configurations for H and F? • How would the orbitals interact

24. H2O • What are the electron configurations for H and O? • How would the orbitals interact

25. Problem • We know from experiments in atomic structure that the bond angle in H2O is 104.5°… not 90° as predicted by valence bond theory • True for CH4 (109.5°) and NH3(107.5°) – VBT always predicts bond angles of 90° • So, we need a better theory…

26. Hybrization • Two problems still exist from Lewis Bonding Theory • Carbon atoms form 4 EQUAL C-H bonds in CH4 (or any other molecule) • Not predicted due to electron configuration of C • Recall: s orbitals have lower energy than p orbitals, therefore the bond length would be different • Existence of double and triple bonds

27. Hybridization of Carbon Orbitals • An s electron gets promoted to the empty p-orbital • This stabilizes the p- and s- orbitals and gives them all the same energy; • Half-filled subshells • Called sp3 orbitals (HYBRID ORBITALS) • Each sp3 orbital lies at 109.5°

29. Additional Hybrid Orbitals – sp - LINEAR

30. Additional Hybrid Orbitals – sp2 - PLANAR

31. Additional Hybrid Orbitals – sp3 - Tetrahedral

32. Double and Triple Bonds • Two types of orbital overlap exist • What we have seen so far is one type • Sigma bonds: σ-bonds • End-on-end overlap of orbitals • Pi bonds:π-bonds • Sideways overlap of orbitals

33. Sigma Bonds • Occur in single bonds and account for the FIRST bond in a double or triple bond • Examples:

34. Pi Bonds • Occur when p-orbitals not on the bonding axis (py or pz) overlap with each other

35. Making Double Bonds • Example: C2H4 • Draw a Lewis Structure • What occurs with the C atoms hybridization?

36. For double bonds, there must be one σ-bond from overlapping hybrid orbitals and one π-bond from overlapping py or pzorbitals • Come from sp2hybridized orbitals and result in trigonal planar structures

37. Making Triple Bonds • Triple bonds have one σ-bond and two π-bonds; come from sp-hybridized orbitals, and result in linear structures • Central atom has two un-hybridized p-orbitals

39. Practice • Explain the structure of the following molecules using electron configurations, orbital hybridization and VBT. • C2Cl4 • C2Cl2 • CO2

40. VSEPR Theory - Valence Shell Electron Pair Repulsion Theory

41. Work through VSEPR Chart • Fun times with molecular structure…

42. Practice Problems • Use Lewis Theory and VSEPR Theory to predict the structure of the following molecules: • Homework/Practice - Worksheet

43. Polar Molecules • Polar molecules are molecules where the electron charge is not distributed evenly

44. Electronegativity and Polar Covalent Bonds • Ionic Bond: ΔEN = >1.7 • Electron transfer • Polar Covalent Bond: ΔEN = 0.5-1.7 • Electrons shared unevenly • Pure Covalent Bond: ΔEN = 0.0-0.5 • Electrons shared evenly • Remember: Think of electrons as electron probabilities, electron cloud density is greater around one atom or another, therefore one gets a slight negative, the other slight positive charge

45. Think of the scale as a continuum

46. Polar Molecules • Cannot exist if there are no polar bonds! • Bond dipole: electronegativity difference of two atoms represented by an arrow pointing from the positive to the negative end (lower to higher EN) • Non-polar molecule: either perfectly symmetrical so the bond dipoles cancel out, or when no polar bonds exist • Polar molecule: occur when bond dipoles do not cancelout

47. Example: • Determine the polarity of the following molecules • H2O, CCl4, NH3, PCl5 • Practice: • CH3Cl, BeCl2, SiO2, BrF4 • CHF2Cl • CH3NH2

48. Intermolecular Forces • Forces that exist between molecules • Three types: • Dipole-Dipole • Hydrogen Bonds • London Dispersion • In order to determine the Intermolecular Forces (IMF), you need to first determine the polarity of the molecule • Much weaker than covalent bonds

49. Dipole-Dipole Forces • Occur in polar molecules • The slightly negative end on one molecule is attracted to the slightly positive end on another molecule • Strength depends on the size of the dipole

50. London Dispersion Forces • Simultaneous attraction of the electrons in one molecule to the nuclei in the surrounding molecules • Increase as the number of electrons and protons in a molecule increase • Exist in ALL molecules • Weakest Force