Honors Chemistry

1. HonorsChemistry Unit 7 – Chapter 6 (p.175-217) Bonding and Lewis Structures

2. Bond force that holds groups of 2 or more atoms together and makes them function as a unit Bond energy energy required to break a bond

3. Types of Bonds • Ionic- transfer of electrons; held together by the attraction of opposite charges; usually between a metal and nonmetal • Covalent- electrons are shared by two nuclei; held by the attraction of positive nuclei to negative shared electrons • Non-Polar Covalent - electrons are shared equally • Polar Covalent - electrons are shared unequally

5. Ionic Bonds - Ex. LiF

6. Ex. LiF continued When soluble ionic compounds are placed in water, the aqueous ions separate from one another. Ex. LiF If placed in water, this compound will separate into aqueous Li+1 and F-1 ions. Sketch:

7. Polyatomic Salts Compounds with polyatomic ions are held together by the attraction of opposite charges. However, the individual polyatomic ion is held together by covalent bonds and does not break apart.

8. Polyatomic Salts – continued Ex. NaNO3 if placed in water will separate into aqueous Na+1 and NO3-1 ions Note: the ionic bond is broken, but NOT the covalent bond holding the nitrate ion together. Sketch:

9. Isoelectronic- when two atoms have the same number of electrons in the same energy levels, sublevels and orbitals Ex. Cl(7 valence electrons) = [Ne] 3s23p5 Cl-1is isoelectronic to [Ne] 3s23p6 – the same as [Ar]

10. Covalent Bonds In covalent bonds, electrons are shared in order to achieve a noble gas configuration. That is why hydrogen is diatomic!

12. Understanding: Oxygen’s two unpaired electrons in two 2p orbitals share separately with one electron in a 1s orbital from each hydrogen. This sharing effectively puts _2_ electrons in each orbital. This results in each atom having the configuration of a noble gas: Oxygen is now isoelectronic with _____Ne____ Hydrogen is now isoelectronic with ______He____

13. Formation of Covalent Bonds

14. Electronegativity and Bonds The difference in electronegativity values between atoms can determine the type of bond they will form. Non-Polar Polar Covalent Covalent (has a dipole) Ionic ←――――――│――――――――│――――――→ DEN < or = 0.4 DEN 0.5 – 1.6 DEN > or =1.7 Note: Do NOT MULTIPLY electronegativity values by the SUBSCRIPTS in a chemical formula. The electronegativity difference is PER bond. VERY IMPORTANT

15. Example Problems Ex. O2 = 3.5 – 3.5 = 0 non-polar covalent Ex. CO = 3.5 – 2.5 = 1.0 polar covalent Ex. LiF = 4.0 – 1.0 = 3.0 ionic

16. Polarity / dipole moment- a molecule containing a partial positive and partial negative charge Represented by an arrow pointing toward the partial negative Ex. Since Cl is more electronegative (3.0) than H (2.1), electrons spend more time around Cl, giving it a partial negative charge OR Represented by the Greek letter delta d Ex. NEVER use both arrows and deltas!

17. Dipole Example NH3 (N―H bond) N-----H 3.0 - 2.1 = 0.9 Polar Covalent N--------H SO2 (S―O bond) S-----O 3.5 – 2.5 = 1.0 Polar Covalent S--------O

18. Lewis Dot Diagram Lewis dot diagram- shows how the valence electrons (those involved in bonding) are arranged among the atoms in a molecule • Uses valence electrons • One dot = one valence electron • One dash = a covalent bond = two electrons

19. Lewis Dot Structures http://www.roymech.co.uk/images14/lewis_elements.gif

20. Key Idea! The key to remember when drawing Lewis structures is that each atom should have a noble gas configuration. The following rules should guide you: Duet rule- elements are stable with 2 valence electrons (H only) Octet rule- most elements are stable with 8 valence electrons NOTE: There are some exceptions to these rules. For example, _B_ is stable with _6_ valence electrons, _Be_ is stable with _4_ valence electrons and _P , S, Xe_ is stable with more than 8 valence electrons

21. Lewis Structures of Molecules • Determine the number of valance electrons it wants • Determine the number of valance electrons it has • Subtract them to find how many are shared and then divide that by two because there are two electrons per bond. • Subtract the number of it has by how many it is sharing to find how many are left over. • Draw the Lewis Dot Diagram for each element and find the one with the most unpaired electrons and make that the center.

22. Multiple Bonds Single, Double, Triple, etc bonds exist. The bond length decreases and the bond strength increases with more bonds in the same bonding area. HINTS: nitrogen will often have 1 unshared pair of electrons oxygen will often have 2unshared pair of electrons carbon almost never has an unshared pair of electrons

23. Example Ex. Methane C H W 1(8) + 4(2) = 16 H 1(4) + 4(1) = - 8 Shared 8/2 = 4 bonds 8-8 = 0 lone electrons

24. Practice Examples Ex. ammonia

25. Practice Examples Ex. Carbon Dioxide

26. Try This One Ex. Nitrogen

27. Try This One Nitrate ion

28. Resonance Resonance is a method of describing the delocalized electrons in some molecules where the bonding cannot be explicitly expressed by a single Lewis structure. Each individual Lewis structure is called a contributing structure of the target molecule or ion. Contributing structures are not isomers of the target molecule or ion since they only differ by the position of delocalized electrons.

29. Determining Bond Order

30. Try It • Determine the bond order and resonance structures for the nitrite ion. Bond order is 1½

31. Molecular Structure Molecular structure / geometric shape- the 3-D arrangement of atoms in a molecule VSEPR theory- Valence Shell Electron Pair Repulsion theory; • Useful for predicting molecular shapes • Bonding and nonbonding electron pairs around the centralatom are positioned as far apart as possible (because they will repel each other) • Lone pairs can spread out a little more since not held in place by two nuclei

32. NONPOLAR SHAPE POLAR SHAPE

33. POLAR SHAPE NONPOLAR SHAPE

34. NONPOLAR SHAPE

35. NONPOLAR SHAPE NONPOLAR SHAPE

36. Hybridization of Orbitals • Hybrid orbitals: are orbitals of equal energy produced by the combination of two of more orbitals on the same atom.

37. Polarity of Polyatomic Molecules Some polyatomic molecules are polar. Polarity depends upon: • Symmetry and bond type • Shape of the molecule

38. Molecule Polarity Ex. Water (H2O) Shape of water molecule is bent or V-shaped Since oxygen is more electronegative than hydrogen, it attracts the shared electrons to itself giving it a partial negative charge.

39. Now: • Looking at some of the previously drawn Lewis Structures: NO3- CCl4 NH3

40. Apply VSEPR Theory • Draw the 3D Shape for the previous Lewis Structures NO3- CCl4 NH3

41. Continued • Now lets look at the bond types and polarity: NO3- N ---- O Polar Covalent Bond EN = 3.0 EN = 3.5 DEN= 0.5 CCl4 C ---- ClPolar Covalent Bond EN = 2.5 EN = 3.0 DEN= 0.5 NH3 N ---- H Polar Covalent Bond EN = 3.0 EN = 2.1 DEN= 0.9

42. Apply to the VSEPR 3D models • Draw the 3D Shape for the previous Lewis Structures NO3- CCl4 NH3 NP NP

43. Multiple Center Atoms • Remember That Hydrogen Is Never Central • Halogens (Cl, Br, I, F) Are Rarely Ever Central • Carbon is Always Central If Present and Will Make Chains • Otherwise Look For How Many Unpaired Electrons are Present In Each Atom

44. C2H6O Start with the math 2(8) + 6(2) + 1(8) = 36 2(4) + 6(1) + 1(6) = 20 36 - 20 = 16 Shared electrons 16/2 = 8 Bonds 20 – 16 = 4 lone electrons

45. H C2H6O - Isomers H .. .. C C O H H H H Without indication of the structure or name it is impossible to know which was intended. H H .. .. H C H C O H H

46. Lets try another one: Draw all of the isomers for C4H10

47. Inter vsIntramolecular Intramolecular forces- bonding forces that hold the atoms of a molecule together • Ex. covalent and ionic Intermolecular forces- forces among molecules that cause them to aggregate to form a solid or a liquid • These are the forces that are broken during phase changes

48. Dipole-dipole attraction • Dipole-dipole attraction- polar molecules have dipole moments; molecules line up so that the partial positive end of one molecule is next to the partial negative end of another molecule • 1% as strong as a covalent or ionic bond. • They are only effective when molecules are close together, such as in solids and liquids. • They are not a factor in gases because the molecules are too far apart.

49. Hydrogen bonding • Hydrogen bonding- special dipole-dipole attraction that occurs among molecules in which hydrogen is bonded directly to a highly electronegative element (ex. N, O, F) • the great polarity of the bond and the small size of the hydrogen atom makes this an unusually strong force. • Although called Hydrogen “bonding”- they are not a true bond like ionic or covalent!

50. London dispersion forces London dispersion forces- instantaneous induced dipole occurring in all atoms and molecules. • An instantaneous dipole in one atom can induce a similar dipole in a neighboring atom. • The molecules are then attracted to each other. These attractions are very weak and short-lived.