Bonding

1. Bonding How do compounds benefit society? Which type of compound is the most useful to you and why?

2. Octet Rule • In compounds, atoms tend to achieve noble gas electron configurations • Metals lose to achieve octet in lower energy level • Non-metals gain to get 8 if combine with metal • or can share with other non-metals • Metal cation + non-metal anion = ionic bond (chapter 7) • 2 non-metals = covalent bond (chapter 8) • 2 metals = metallic bond (also chapter 7)

3. 8.1 Molecular Compounds • Covalent bonds • Sharing of electrons to hold atoms together • Between 2 non-metals • Neutral • Makes molecules • Bond dissociation energy is the amount of energy required to break a covalent bond between atoms • A large bond dissociation energy corresponds to a strong covalent bond • Double and triple bonds are stronger than single cov bonds

4. Molecular compounds • Chemical formula = molecular formula • Not lowest whole number, just actual number of atoms of each element • Properties • Comprised of two or more non-metals • Not good conductors of heat and electric current • Lower melting and boiling points than ionic • Diatomic molecules • Molecule of 2 identical atoms • H2 O2 N2 Cl2 Br2 I2 F2

5. Nomenclature • The same elements can combine in multiple ways via different numbers of atoms to form different molecules • Ex CO vs CO2 • –ide suffix on second element just like ionic • Use prefixes to indicate number of each • Mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca- • No mono- for first element • Ex: carbon monoxide and carbon dioxide

6. Naming molecular compounds • Practice • SiO2 • Silicon dioxide • SO3 • sulfur trioxide • XeF4 • Xenon tetrafluoride • Phosphorus pentachloride • PCl5 • Dinitrogen monoxide • N2O

7. 8.2 Nature of covalent bonding • Still obeys octet rule • Share as many electrons as needed for 8 • Single covalent bond • Share one pair of electrons (1 from each atom) • Ex H-H • Double covalent bond • Share two pairs of electrons (2 from each) • Ex O=O • Triple covalent bond • Share 3 pairs of electrons (3 from each) • Ex N≡N • Coordinate covalent bond • Share pair of electrons (usually only 1) but both from 1 atom • Ex :C: + :Ö:  :C≡O: (often drawn :C=O: )

8. Recall electron configurations • Valence electrons • Highest occupied energy level • Largely determine chemical properties • Corresponds to American group number • Exception = He • Used for chemical bonds • Shown in electron-dot structures/diagrams • Ex: H·

9. Electron Dot Diagrams • Practice: Na F He Cl Mg S B Li Ne C Kr N Si O Al P Ca Find a trend?

10. Structures • Dot structures can be made for covalent bonds • Place shared electrons between the 2 symbols either as pair of dots or dashes • Ex H-H or :Ö=Ö: or :N≡N: • Each dash represents a single covalent bond • Unshared pairs / lone pairs / non-bonding pairs = valence e- pairs not in cov bond • Structural formulas use dashes, but also show spatial arrangement of atoms

11. How Covalent Bonds Form • Draw the Lewis dot diagrams for all the atoms in CH4 • Draw 1 carbon diagram • Draw 4 hydrogen diagrams

12. Bonding in Methane C H H H H

13. Bonding in Methane C H H H H

14. H Bonding in Methane C H H H

15. H Bonding in Methane C H H H

16. H Bonding in Methane C H H H

17. H H Bonding in Methane C H H

18. H H Bonding in Methane C H H

19. H H Bonding in Methane C H H

20. H H H Bonding in Methane C H

21. H H H Bonding in Methane C H

22. H H H H Bonding in Methane C

23. Bonding in Methane H C H H H

24. Draw the following compounds: PCl3 H2O SF2 AlCl3

25. Do Now: Get out your homework to be stamped and be prepared to answer: • What is a covalent bond? • What types of elements do covalent bonds form between? • What are valence electrons? • How many electrons are in a bond?

26. Single Bond • Two electrons are shared between two atoms • 1 electron pair • Represented by 1 line between the atoms

27. Let’s Make F2 F F

28. Let’s Make F2 F F

29. Let’s Make F2 F F

30. Double Bond • Four electrons are shared between two atoms • 2 electron pairs • Represented by 2 lines between the atoms

31. Let’s Make O2 O O

32. Let’s Make O2 O O

33. Let’s Make O2 O O

34. Let’s Make O2 O O

35. Double bonds Try It! • Draw the Lewis dot diagram for • O3 • CO2 • CF2S

36. Triple Bond • Six electrons are shared between two atoms • 3 electron pairs • Represented by 3 lines between the atoms

37. Let’s Make N2 N N

38. Let’s Make N2 N N

39. Let’s Make N2 N N

40. Let’s Make N2 Let’s Make N2 N N

41. Let’s Make N2 N N

42. Triple bonds Try It! • Draw the Lewis Dot Diagram for • HCN • N2O • C2H2

43. Lewis Structures summary • Place shared electrons between the 2 symbols either as pair of dots or dashes • Ex H-H or :Ö=Ö: or :N≡N: • Each dash represents a single covalent bond • Unshared pairs / lone pairs / non-bonding pairs = valence e- pairs not in covalent bond • Structural formulas use dashes, but also show spatial arrangement of atoms • Oriented to maximize distance between e- pairs (bonded or unshared)

44. Polyatomic ions • Even though overall charged, polyatomic ions are covalently bonded together • Overall charge comes from extra electrons either gained or lost in order to achieve an octet for all constituent atoms • Ex: SO32- O S O O • Use brackets to denote total charge

45. Polyatomic Ions • Covalently bonded • With additional or removed electrons • Cations lose and anions gain in e- count NO31- SO32- PO43- CN1-

46. Making Sense

47. Making Sense

48. Resonance structures • Occur when possible to draw two or more valid electron dot structures for a molecule • Originally thought that electrons would flip back and forth (resonate) between the options • Actual bonding is a hybrid (mixture) of the options • All valid structures are separated by double headed arrow • Ex: O3 (ozone) :Ö = Ö – Ö: ↔ :Ö - Ö = Ö: ˙˙ ˙˙

49. Exceptions to Octet Rule • If total # of valence e- for compound is odd, the octet rule cannot be satisfied • Nitrogen dioxide • Phosphorus pentachloride • Sulfur hexafluoride • *Boron trifluoride • readily reacts with NH3

50. VSEPR Theory • Valence-shell electron-pair repulsion theory • States: the repulsion between electron pairs cause molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible • Non-bonding pairs are important because they will repel the other e- pairs more • No other atom stakes claim, therefore held closer to the central nucleus and will squeeze out the others