Entry Task: Oct 16 th Block 1

1. Entry Task: Oct 16th Block 1 Provide the name or formula for the following NO DiPhosphoruspentaoxide C4H8 Boron monocarbide

2. Agenda: Discuss Lewis Dot ws Ch. 9 FLASH LECTURE Start VSEPR Activity

4. The objective to this activity is to familiarize you with Lewis dot structures for both atoms and ions. From this information, you can use Lewis dot structures to see how ionic bonds form to create compounds.

6. Now you try:

7. 2 6 +2 -2 Ca Ca S S Calcium sulfide Ca S

8. 3 5 +3 -3 Al Al P P Aluminum phosphide Al P

9. 1 6 +1 +1 -2 K K O K O K Potassium oxide K K O

10. 3 7 -3 +3 Fe Fe F F F F F F Iron III fluoride F Fe F F

11. Questions 1. What happened to the metal’s electron(s)? They were given (transferred) to the nonmetal. 2. How did the action from question 1 affect the nonmetal? The nonmetal becomes a negative charge 3. What is the end result of the actions from questions 1 &2? The transferring of electrons created a charge on the metal and nonmetal therefore created a bond between the two.

13. From the previous activity, you learned how ionic compounds bond by transferring electrons to create an octet. In this activityyou will use nonmetals to create molecules by sharing their valence electrons to create an octet.

14. Guidelines for Drawing Lewis Dot Structures for Molecules: • Find out how many valence electrons are on each atom involved. • Put the atom which can share the most electrons in the middle. • Arrange dots so that unshared (loners) are “facing” other loners from another atom. • Some atoms can share two or three pairs of electrons (usually it is • either C, N, or O) • Hydrogen’s are always on the ends and only bond with 1 atom • (sharing only a pair of electrons) • 6. Circle the shared pair of electrons.

17. H H O O H H

18. H H H H H C C H H H

19. H H H H P P H H

20. O C O O C O O C O

21. O O O O

22. H H H H N N H H

23. Questions • How does covalent bonding differ from ionic bonding- • THINK ON THIS ONE!! (4 differences) Ionic Covalent Transfer electrons Shared electrons Metal/nonmetals Nonmetals Has a charge No charge Single bonds only Can have double/triple bonds

24. Ch. 9 FLASH LECTURE

25. I can… • Explain the properties of molecules that coincide with their molecular shape and geometries- VSEPR theory.

26. Molecular Shapes • The shape of a molecule plays an important role in its reactivity. • By noting the number of bonding and nonbonding electron pairs, we can easily predict the shape of the molecule.

27. What Determines the Shape of a Molecule? • Simply put, electron pairs, whether they be bonding or nonbonding, repel each other. • By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

28. Electron Domains • We can refer to the electron pairs as electron domains. • In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom; therefore, they count as one electron domain. • The central atom in this molecule, A, has four electron domains.

29. Valence-Shell Electron-Pair Repulsion Theory (VSEPR) “The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

30. Electron-Domain Geometries Table 9.1 contains the electron-domain geometries for two through six electron domains around a central atom.

31. Electron-Domain Geometries • All one must do is count the number of electron domains in the Lewis structure. • The geometry will be that which corresponds to the number of electron domains.

32. Molecular Geometries • The electron-domain geometry is often not the shape of the molecule, however. • The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs.

33. Molecular Geometries Within each electron domain, then, there might be more than one molecular geometry.

34. Linear Electron Domain • In the linear domain, there is only one molecular geometry: linear. • NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.

35. Trigonal Planar Electron Domain • There are two molecular geometries: • Trigonal planar, if all the electron domains are bonding, • Bent, if one of the domains is a nonbonding pair.

36. Nonbonding Pairs and Bond Angle • Nonbonding pairs are physically larger than bonding pairs. • Therefore, their repulsions are greater; this tends to decrease bond angles in a molecule.

37. Multiple Bonds and Bond Angles • Double and triple bonds place greater electron density on one side of the central atom than do single bonds. • Therefore, they also affect bond angles.

38. Tetrahedral Electron Domain • There are three molecular geometries: • Tetrahedral, if all are bonding pairs, • Trigonal pyramidal, if one is a nonbonding pair, • Bent, if there are two nonbonding pairs.

39. Trigonal Bipyramidal Electron Domain • There are two distinct positions in this geometry: • Axial • Equatorial

40. Trigonal Bipyramidal Electron Domain Lower-energy conformations result from having nonbonding electron pairs in equatorial, rather than axial, positions in this geometry.

41. Trigonal Bipyramidal Electron Domain • There are four distinct molecular geometries in this domain: • Trigonal bipyramidal • Seesaw • T-shaped • Linear

42. Octahedral Electron Domain • All positions are equivalent in the octahedral domain. • There are three molecular geometries: • Octahedral • Square pyramidal • Square planar

43. Larger Molecules In larger molecules, it makes more sense to talk about the geometry about a particular atom rather than the geometry of the molecule as a whole.

44. Polarity • In Chapter 8, we discussed bond dipoles. • But just because a molecule possesses polar bonds does not mean the molecule as a whole will be polar.

45. Polarity By adding the individual bond dipoles, one can determine the overall dipole moment for the molecule.

46. Polarity

47. Overlap and Bonding • We think of covalent bonds forming through the sharing of electrons by adjacent atoms. • In such an approach this can only occur when orbitals on the two atoms overlap.

48. Overlap and Bonding • Increased overlap brings the electrons and nuclei closer together while simultaneously decreasing electron– electron repulsion. • However, if atoms get too close, the internuclear repulsion greatly raises the energy.

49. Hybrid Orbitals • Consider beryllium: • In its ground electronic state, beryllium would not be able to form bonds, because it has no singly occupied orbitals.

50. Hybrid Orbitals But if it absorbs the small amount of energy needed to promote an electron from the 2s to the 2p orbital, it can form two bonds.