1. Chapter 5: Soap

2. Soap This chapter will introduce the chemistry needed to understand how soap works  Section 5.1: Types of bonds  Section 5.2: Drawing Molecules  Section 5.3: Compounds in 3D  Section 5.4: Polarity of Molecules  Section 5.5: Intermolecular Forces  Section 5.6: Intermolecular Forces and Properties

3. Section 5.1—Types of Bonds

4. Why atoms bond Atoms are most stable when they’re outer shell of electrons is full Atoms bond to fill this outer shell For most atoms, this means having 8 electrons in their valence shell  Called the Octet Rule Common exceptions are Hydrogen and Helium which can only hold 2 electrons  Called the Duet Rule

5. Remember Valence Electrons Valence electrons: the electrons found in the highest energy level Short Cut Rule: the group # next to the letter A of the Representative elements represents the valence number

6. 1: How many valence electrons are in an atom? The main groups of the periodic table each have 1 more valence electron than the group before it. 12345678

7. Remember, Lewis Dot diagrams show their valence electrons When atoms bond, they have 4 orbitals available (1 “s” and 3 “p”s). There are 4 places to put electrons Put one in each spot before doubling up! Example: Draw the Lewis Structure for an oxygen atom

8. LEWIS DOT DIAGRAMS

9. What is a Bond & Why does it form? Like glue holding atoms together It’s really forces of attraction between valence electrons holding atoms together They form because it lowers the potential energy of the atoms and creates stability

10. Three Types of Bonding

11. Ionic Bonding—Metal + Non-metal Metals have fewer valence electrons and much lower ionization energies than non-metals Therefore, metals tend to lose their electrons and non-metals gain electrons Metals become cations (positively charged) Non-metals become anions (negatively charged) There is a transfer of electrons The cation & anion are electrostatically attracted because of their charges—forming an ionic bond

13. One way valence shells become full Na - - -- - - - - - - Cl - - -- - - - - - - - - - - - - - Sodium has 1 electron in it’s valence shell Chlorine has 7 electrons in it’s valence shell METALS: Some atoms give electrons away to reveal a full level underneath. NONMETALS: Some atoms gain electrons to fill their current valence shell. -

14. One way valence shells become full Na - - -- - - - - - - Cl - - -- - - - - - - - - - - - - - - + - The sodium now is a cation (positive charge) and the chlorine is now an anion (negative charge). These opposite charges are now attracted, which is an ionic bond.

15. IONIC BONDING Sodium + Chlorine  Sodium Chloride

16. Transfer electrons in ionic bonding Transfer electrons from metal atoms to non-metal atoms, keeping track of their new charge Example: Draw the Lewis Structure for KCl

17. Transfer electrons in ionic bonding Transfer electrons from metal atoms to non-metal atoms, keeping track of their new charge ClK Potassium has 1 electron Chlorine has 7 electrons Example: Draw the Lewis Structure for KCl

18. +1 Transfer electrons in ionic bonding Transfer electrons from metal atoms to non-metal atoms, keeping track of their new charge ClK Potassium has 1 electron Chlorine has 7 electrons Example: Draw the Lewis Structure for KCl

19. Add more atoms if needed If the transfer from one atom to another doesn’t result in full outer shells, add more atoms Example: Draw the Lewis Structure the ionic compound of Barium fluoride

20. Add more atoms if needed If the transfer from one atom to another doesn’t result in full outer shells, add more atoms FBa Barium has 2 electron Fluorine has 7 electrons Example: Draw the Lewis Structure the ionic compound of Barium fluoride The fluorine is full, but the Barium isn’t!

21. Add more atoms if needed If the transfer from one atom to another doesn’t result in full outer shells, add more atoms FBa Barium has 2 electron Fluorine has 7 electrons Example: Draw the Lewis Structure the ionic compound of Barium fluoride F Add another fluorine atom

22. Add more atoms if needed If the transfer from one atom to another doesn’t result in full outer shells, add more atoms FBa Barium has 2 electron Fluorine has 7 electrons Example: Draw the Lewis Structure the ionic compound of Barium fluoride F +2 Now all have full valence shells and the charges are balanced, just as when you learned to write in Chpt 2—BaF 2 !

23. Covalent Bonding When two non-metals share electrons 2 Identical Non-metals that share electrons evenly form non-polar covalent bonds 2 Different Non-metals that share electrons un-evenly form polar covalent bonds

24. COVALENT BONDING NONPOLAR COVALENT BONDING Chlorine + Chlorine  Chlorine gas POLAR COVALENT BONDING Hydrogen + Flourine  Hydrogen Fluoride + 

25. Metallic Bonding Metals form a pool of electrons that they share together. The electrons are free to move throughout the structure—like a sea of electrons Atoms are bonded as a network

26. Metallic Bonding

27. Bonding in Never Purely Ionic or Covalent: Use EN to indicate primary type 0% 5% 50% 100% 0-----.3 --------------------1.7-------------------3.3 np polar ionic covalent If the electronegativity difference is: Greater than 1.7IONIC Between 0-.29NONPOLAR COVALENT Between.3 and 1.69POLAR COVALENT

28. Figure 12.4 The three possible types of bonds. nonpolar polar ionic Sharing electrons equally Sharing electrons unequally Transfer of electrons

29. Examples CH 4 F2F2 H2OH2O NaF 2.5 – 2.1 =.4 polar covalent 4.0 -4.0= 0 nonpolar covalent 3.5 – 2.1 = 1.4 polar covalent 4.0 –.9 = 3.1 ionic

30. Bond type affects properties There are always exceptions to these generalizations (especially for very small or very big molecules), but overall the pattern is correct

31. Melting/Boiling Points Of Compounds Ionic Compounds tend to have very high melting/boiling points as it’s hard to pull apart those electrostatic attractions of the ionic bond  These compounds are found as solids under normal conditions Metals have INTERMEDIATE melting/boiling points. Metallic bonds vary in strength. o All are found as solids under normal conditions except Hg

32. Melting/Boiling Points Of Compounds Covalent Compounds Polar Covalent molecules have the next highest melting points  These compounds are soft solids or liquids under normal conditions Non-polar covalent molecules have the lowest melting/boiling points  These compounds are found as liquids or gases

33. Solubility in Water Ionic & polar covalent compounds tend to be soluble in water Non-polar & metallic compounds tend to be insoluble in water Use this: Like Dissolves Like rule of thumb. Polar solutes dissolve in polar solvents Nonpolar solutes dissolve in nonpolar solvents

34. Conductivity of Electricity In order to conduct electricity, charge must be able to move or flow Metallic bonds have free-moving electrons— they can conduct electricity in solid and liquid state Ionic bonds have free-floating ions when dissolved in water or in liquid form that allow them conduct electricity Covalent bonds never have charges free to move and therefore cannot conduct electricity in any situation

35. Ionic Compounds ONLY CONDUCT ☺ X

36. Like Dissolves Like LIke

37. Structures: Ionic Compounds Ionic compounds are made of positive and negative ions. They pack together so that the like-charge repulsions are minimized while the opposite-charge attractions are enhanced. Na +1 Cl -1

38. Structures Metals: sea of electrons Covalent compounds:separate discrete molecules

39. Visual Representation of 3 Bonding Types

40. Section 5.2—Drawing Molecules of Covalent Compounds

41. Tips for arranging atoms In general, write out the atoms in the same order as they appear in the chemical formula Hydrogen & Halogens (F, Cl, Br, I) can only bond with one other atom  Always put them around the outside The least electronegative atom is usually in the middle; Carbon always goes in the middle

42. Steps to Drawing Lewis Structures 1. Decide how many valence electrons are around each atom 2. Arrange the atoms in a skeletal structure and connect them with a bonding pair of electrons. 3. Place remaining electrons around atoms so they each acquire 8 electrons. Exception is H.

43. H H Example: Carbon Tetrachloride Example: Draw the Lewis Structure for CH 4 Remember, “H” can’t go in the middle… put them around the Carbon! C H H Carbon has 4 electrons = 8 electrons Each hydrogen has 1

44. H H Count electrons around each atom Any electron that is being shared (between two atoms) gets to be counted by both atoms! All atoms are full with 8 valence electrons (except H—can only hold 2) Example: Draw the Lewis Structure for CH 4 C H H Carbon has 8 Each Hydrogen has 2 All have full valence shells—drawing is correct!

45. Bonding Pair Pair of electrons shared by two atoms…they form the “bond” H H C H H Bonding pair

46. Try These CH 3 IPCl 3

47. What if they’re not all full after that? Multiple Covalent Bonds Are needed when there is not enough electrons to complete an octet To satisfy: move lone pair in between atoms to satisfy the duet/octet rule

48. Example: Draw the Lewis Structure for CH 2 O H C O H Remember that hydrogen atoms can’t go in the middle!

49. Example: Draw the Lewis Structure for CH 2 O H C O H The two hydrogen atoms are full But the carbon and oxygen only have 7 each!

50. Example: Draw the Lewis Structure for CH 2 O H C O H But they each have a single, unshared electron. They could share those with each other!

51. Example: Draw the Lewis Structure for CH 2 O H C O H Now the carbon and oxygen both have a full valence!

52. Double Bonds & Lone Pairs Double bonds are when 2 pairs of electrons are shared between the same two atoms Lone pairs are a pair of electrons not shared—only one atom “counts” them H C O H Double Bond Lone pair

53. And when a double bond isn’t enough… Sometimes forming a double bond still isn’t enough to have all the valence shells full Example: Draw the Lewis Structure for C 2 H 2

54. … Example: Draw the Lewis Structure for C 2 H 2 H C C H Remember that hydrogen atoms can’t go in the middle!

55. … Example: Draw the Lewis Structure for C 2 H 2 H C C H Each carbon atom only has 7 electrons…not full

56. Example: Draw the Lewis Structure for C 2 H 2 H C C H But they each have an un-paired electron left!

57. Example: Draw the Lewis Structure for C 2 H 2 H C C H Now they each have 8 electrons!

58. Triple Bonds A Triple Bond occurs when two atoms share 3 pairs of electrons Triple Bond H C C H

59. TRY These: HCN CO 2

60. Properties of multiple bonds Single Bond Double Bond Triple Bond Longer & Weaker bonds (Lower Bond Energy: takes less energy to break) Shorter (atoms closer together & Stronger bonds (Higher Bond Energy: takes more energy to break)

61. Bond Dissociation Energy

62. Polyatomic Ions They are a group of atoms bonded together that have an overall charge Example: Draw the Lewis Structure for CO 3 -2

63. Drawing Polyatomic Ions Example: Draw the Lewis Structure for CO 3 -2 C O When there’s a single atom of one element, put it in the middle O O

64. Drawing Polyatomic Ions Example: Draw the Lewis Structure for CO 3 -2 C O None of the atoms have full valence shells…they all have 7! O O The carbon can double bond with one of the oxygen atoms

65. Drawing polyatomic Ions Example: Draw the Lewis Structure for CO 3 -2 C O Now the Carbon and the one oxygen have 8…but the other two oxygen atoms still only have 7 O O This is a polyatomic ion with a charge of “-2”…that means we get to “add” 2 electrons!

66. Drawing Polyatomic Ions Example: Draw the Lewis Structure for CO 3 -2 C O Now the Carbon and the one oxygen have 8…but the other two oxygen atoms still only have 7 O O This is a polyatomic ion with a charge of “-2”…that means we get to “add” 2 electrons! -2

67. Covalent bond within…ionic bond between Polyatomic ions have a covalent bond within themselves… But can ionic bond with other ions Covalent bonds within C O O O Na

68. Covalent bond within…ionic bond between Polyatomic ions have a covalent bond within themselves… But an ionic bond with other ions Covalent bonds within Ionic bond with other ions C O O O -2 Na +1

69. Section 5.3—Molecules in 3D

70. Bonds repel each other Bonds are pairs electrons. Electrons are negatively charged Negative charges repel other negative charges SO bonds repel each other Molecules arrange themselves in 3-D so that the bonds are as far apart as possible

71. Valence Shell Electron Pair Repulsion Theory Valence Shell Electron Pair Repulsion Theory (VSEPR Theory) Outer shell of electrons involved in bonding Bonds are made of electron pairs Those electron pairs repel each other Attempts to explain behavior This theory attempts to explain the 3-D shape of molecules.

72. Coding for a Shape A code can help you connect the molecule to its shape. “A” stands for the central atom “B” stands for the number of bonding atoms off the central atom. “E” stands for the number of lone pair coming off the CENTRAL atom.

73. What shapes do molecules form? Linear 2 bonds, no lone pairs OR any 2 atom molecule AB 2 AB Linear BeCl 2 HCl

74. What shapes do molecules form? Trigonal planar 3 bonds, no lone pairs Indicates a bond coming out at you Indicates a bond going away from you AB 3 BF 3

75. What shapes do molecules form? Tetrahedral 4 bonds, no lone pairs AB 4 CH 4

76. What shapes do molecules form? Trigonal bipyramidal 5 bonds, no lone pairs AB 5 PCl 5

77. What shapes do molecules form? Octahedral 6 bonds, no lone pairs AB 6 SF 6

78. Lone Pairs Lone pairs are electrons, too…they must be taken into account when determining molecule shape since they repel the other bonds as well. But only take into account lone pairs around the CENTRAL atom, not the outside atoms!

79. What shapes do molecules form? Bent 2 bonds, 1 lone pair AB 2 E AB 2 E 2 SO 2 Bent 2 bonds, 2 lone pairs H2OH2O

80. What shapes do molecules form? Trigonal pyramidal 3 bonds, 1 lone pair AB 3 E NH 3

81. Lone Pairs take up more space Lone pairs aren’t “controlled” by a nucleus (positive charge) on both sides, but only on one side. This means they “spread out” more than a bonding pair. They distort the angle of the molecule’s bonds away from the lone pair.

82. 109.5° C 105° O Example of angle distortion

84. Rotating Molecular Shapes http://intro.chem.okstate.edu/1314F00/Lec ture/Chapter10/VSEPR.html http://intro.chem.okstate.edu/1314F00/Lec ture/Chapter10/VSEPR.html