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A Bit of Review

A Bit of Review. Valence electrons Only electrons in the s and p sublevels count! Most atoms will gain or lose electrons in order to get a full valence electron shell The number of valence electrons in a neutral atom can be found from the periodic table!!

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A Bit of Review

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  1. A Bit of Review • Valence electrons • Only electrons in the s and p sublevels count! • Most atoms will gain or lose electrons in order to get a full valence electron shell • The number of valence electrons in a neutral atom can be found from the periodic table!! \ • Octet “rule”: Most elements need 8 e- to have a full valence electron shell • Some major exceptions: • H and He (2 valence e-)

  2. Let’s bond Bond: • Ionic • Ionic: Pertaining to or occurring in the form of ions • Bond: something that binds, fastens, confines, or holds together • Must be both a cation and an anion present • Total charge is zero

  3. Let’s bond Co valent together or shared valence electrons Nonmetals will SHARE electrons in order to fool themselves into thinking they have a full valence electron shell nonmetal and nonmetal

  4. Writing Lewis Dot Structures for Molecular Compounds A pair of shared e- between atoms = a covalent bond 1 pair of shared e- = single bond

  5. single bond single bond

  6. Writing Lewis Dot Structures for Molecular Compounds H H C H H bonding pair electrons are written as lines single bond

  7. Writing Lewis Dot Structures for Molecular Compounds A pair of shared e- between atoms = a covalent bond 1 pair of shared e- = single bond 2 pair of shared e- = double bond

  8. Writing Lewis Dot Structures for Molecular Compounds O C H H bonding pair are written as lines double bond

  9. Writing Lewis Dot Structures for Molecular Compounds A pair of shared e- between atoms = a covalent bond 1 pair of shared e- = single bond 2 pair of shared e- = double bond 3 pair of shared e- = triple bond

  10. Writing Lewis Dot Structures for Molecular Compounds N C H bonding pair are written as lines triple bond

  11. Writing Lewis Dot Structures for Molecular Compounds lone pair (2 e-each) O N bonding pair (2 e-each) C C H H H

  12. Writing Lewis Dot Structures for Molecular Compounds • General Rules: • “Central” atoms will be surrounded other atoms • The element with the least complete valence shell tends to be the central atom HUH?!?!?!

  13. Writing Lewis Dot Structures for Molecular Compounds • General Rules: • “Central” atoms will be surrounded other atoms • The element with the least complete valence shell tends to be the central atom • The least electronegative element tends to be the central atom HUH?!?!?!

  14. ELECTRONEGATIVITY

  15. Writing Lewis Dot Structures for Molecular Compounds Which element in each of the following groups is most likely to be the central atom in a molecular compound? a) N, H, O b) S, N, C c) Cl, P, O, H d) N, O, P e) O, S, F

  16. a) N, H, O b) S, N, P c) Cl, C, O, H

  17. a) N, H, O b) S, N, P c) Cl, C, O, H

  18. a) N, H, O b) S, N, C c) Cl, C, O, H

  19. a) N, H, O b) S, N, C c) Cl, C, O, H

  20. a) N, H, O b) S, N, P c) Cl, P, O, H

  21. a) N, H, O b) S, N, P c) Cl, P, O, H

  22. d) N, O, P

  23. d) N, O, P

  24. e) O, S, F

  25. e) O, S, F

  26. Writing Lewis Dot Structures for Molecular Compounds • General Rules: • “Central” atoms will be surrounded other atoms • The element with the least complete valence shell or least electronegative tends to be the central atom • Must be at least one bond to each atom • # of valence e- short = # of bonds formed nitrogen has 5 valence e- so needs 3 more to be full chlorine has 7 valence e- so needs 1 more to be full 3 1 nitrogen tends to form 3 bonds chlorine tends to form 1 bonds 3 1 THIS IS ONLY A GUIDELINE!!!!!!

  27. Writing Lewis Dot Structures for Molecular Compounds How many bonds are each of the following likely to form in a molecular compound? a) C b) N c) Ar d) I e) O

  28. a) C b) N c) Ar d) I e) O 4 3 0 1 2

  29. Writing Lewis Dot Structures for Molecular Compounds

  30. Formal Charge Formal charge on an atom = # of valence e- - # of bonds - # of lone electrons From periodic table Individual e-, NOT lone pair Formal charge on C = 4 Formal charge on C = 4 - 4 Formal charge on C = 4 - 4 - 0 Formal charge on C = 4 - 4 - 0= 0 Formal charge on C Formal charge on O = 6 - 2 Formal charge on O = 6 - 2 - 4 Formal charge on O = 6 - 2 - 4 = 0 Formal charge on O = 6 Formal charge on O

  31. Formal Charge Formal charge on an atom = # of valence e- - # of bonds - # of lone electrons From periodic table Individual e-, NOT lone pair Formal charge on N = 5 Formal charge on N = 5 - 4 Formal charge on N = 5 - 4 - 0 Formal charge on N = 5 - 4 - 0= +1 Formal charge on N Formal charge on O = 6 - 1 - 6 Formal charge on O = 6 - 1 - 6 = -1 Formal charge on O Formal charge on O = 6 Formal charge on O = 6 - 1 Formal charge on O = 6 - 1 - 6 = -1 Formal charge on O = 6 Formal charge on O = 6 - 2 Formal charge on O = 6 - 2 - 4 Formal charge on O = 6 - 2 - 4 = 0 Formal charge on O

  32. Shapes of Molecules • Electron pairs surrounding an atom repel each other. This is referred to as Valence Shell Electron Pair Repulsion (VSEPR) theory. • NOT based on how many electron pair there are, but where they are. • Where they are: electron domains (charge clouds) HUH?!?!?!

  33. Shapes of Molecules What counts as an electron domain? 1) A connection between the atom of interest and any other atom a) single, double, and triple bonds each count one electron domain 2) A lone pair on the atom of interest 2 electron domains 4 electron domains 4 electron domains 4 electron domains 4 electron domains 2 electron domains 3 electron domains 3 electron domains

  34. Shapes of Molecules • The molecular shape gives the arrangement of only the atoms around the central atom as a result of electron repulsion. (just the atoms) • The electron pair geometry gives the arrangement of atoms AND the lone pair electrons around the central atom.

  35. Linear Molecules • In carbon dioxide, CO2, the central C atom has two atoms attached (the 2 oxygen atoms) and has no lone pair. • 2 electron domains • The electron pair geometry is linear. • The molecular shape is linear. • Bond angle is 180o

  36. Trigonal Planar • In formaldehyde, H2CO, the central C atom has three atoms attached (2 hydrogen and 1 oxygen) and has no lone pair. • 3 electron domains • The electron pair geometry is trigonal planar. • The molecular shape is trigonal planar. • Bond angle is 120o

  37. Trigonal Planar • In sulfur dioxide, SO2, the central S atom has two atoms attached (2 oxygens) and has one lone pair. • 3 electron domains • The electron pair geometry is trigonal planar. • The molecular shape is bent. • Bond angle is 120o

  38. Tetrahedral • In water, H2O, the central O atom has two atoms attached (2 hydrogen atoms) and has two lone pair. • 4 electron domains • The electron pair geometry is tetrahedral. • The molecular shape is bent. • Bond angle is ~109.5o (~104.5o actually)

  39. Tetrahedral • In ammonia, NH3, the central N atom has three atoms atached (3 hydrogen atoms) and has one lone pair. • 4 electron domains • The electron pair geometry is tetrahedral. • The molecular shape is trigonal pyramidal. • Bond angle is ~109.5o (~107o actually)

  40. Tetrahedral • In methane, CH4, the central C atom has four atoms attached (4 hydrogen atoms) and has no lone pair. • 4 electron domains • The electron pair geometry is tetrahedral. • The molecular shape is tetrahedral. • Bond angle is 109.5o

  41. Let’s bond Co valent together or shared valence electrons Nonmetals will SHARE electrons in order to fool themselves into thinking they have a full valence electron shell If “X” and “Y” share bonding e- equally: If “X” and “Y” do NOT share bonding e- equally: Unequal sharing of bonding e- leads to polar covalent bonds

  42. N S Polar Covalent Bonds Polar: having poles One end opposite from the other

  43. Polar Covalent Bonds Polar covalent bonds reslut from differences in electronegativity. HUH?!?!?! Electronegativity: The ability of an atom to pull BONDING electrons towards itself.

  44. Polar Covalent Bonds

  45. Polar Covalent Bonds If EN difference is: < 0.5 the bond is considered to be nonpolar 2.5-2.5=0 2.5-2.5=0 2.8-2.8=0 2.5-2.1=0.4 ≥0.5 the bond is considered to be polar 3.5-2.5=1 4.0-2.1=1.9 3.0-2.5=0.5 3.5-1.8=1.7

  46. Delta (δ) Notation for Polar Bonds • The lowercase Greek letter delta, d, is used to indicate a polar bond. • The MORE EN element has extra e-, so it is negative and is indicated by the symbol d–. • The LESS EN element is short of e-, so it is positive and is indicated by the symbol d+. 2.1 3.0 H–Cl • d– • d+

  47. Polar Covalent Bonds Give delta notation and polarity arrows for the following: • d+ • d– • d+ • d– • d+ • d– • d+ • d– 3.5 1.8 3.0 2.5 2.5 3.5 4.0 2.1

  48. Nonpolar Molecules with Polar Bonds • CCl4 has polar bonds, but the overall molecule is nonpolar • Using VSEPR theory, the four chlorine atoms are at the four corners of a tetrahedron

  49. Nonpolar Molecules with Polar Bonds • The chlorines are each δ–, while the carbon is δ+. • The net effect of the polar bonds is zero, so the molecule is nonpolar.

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