1. Chapter 5: Soap

2. Introductory Activity Fill a test tube with an inch of water Add a squirt of cooking oil to the test tube. Observe Stopper, shake & observe Add a few drops of soap. Observe Stopper, shake & observe With another test tube, add water & soap only. Observe. Compare the two test tubes. Make particle visualizations describing each test tube.

3. Introductory Activity What ideas do you have about how soap works? What kinds of things do advertising and marketing tell you? What do the soap companies want you to know about how soap works?

4. 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

5. Soap Inter-molecular forces Works based on Molecular Geometry Bonding types & Structures Determined by

6. Section 5.1—Types of Bonds

7. Why atoms bond Atoms are most stable when they’re outer shell of electrons is full Atoms bonds 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.

8. 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 Some atoms give electrons away to reveal a full level underneath. Some atoms gain electrons to fill their current valence shell. -

9. 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.

10. Ionic Bonding—Metal + Non-metal Metals have fewer valence electrons and much lower ionization energies (energy needed to remove an electron) 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) The cation & anion are attracted because of their charges—forming an ionic bond

11. Bonding between non-metals When two non-metals bond, neither one loses or gains electrons much more easily than the other one. Therefore, they share electrons Non-metals that share electrons evenly form non-polar covalent bonds Non-metals that share electrons un-evenly form polar covalent bonds

12. Metals 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 aren’t bonded to specific other atoms, but rather to the network as a whole

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

14. Melting/Boiling Points Ionic bonds tend to have very high melting/boiling points as it’s hard to pull apart those electrostatic attractions  They’re found as solids under normal conditions Polar covalent bonds have the next highest melting/boiling points  Most are solids or liquids under normal conditions Non-polar covalent bonds have lower melting/boiling points  Most are found as liquids or gases

15. Solubility in Water Ionic & polar covalent compounds tend to be soluble in water Non-polar & metallic compounds tend to be insoluble

16. 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

17. Section 5.2—Drawing Molecules

18. Drawing Molecules on Paper Lewis Structures (or Dot Structures) are one way we draw molecules on paper Since paper is 2-D and molecules aren’t, it’s not a perfect way to represent how molecules bond…but it’s a good way to begin to visualize molecules

19. Drawing Ionic Compounds

20. 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

21. 2: Placing electrons around an atom 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

22. 3: 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

23. 4: 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

24. 4: 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

25. A note about Ionic Dot Structures The atoms are not sharing the electrons— make sure you clearly draw the atoms separate!

26. Drawing Covalent Compounds

27. Tips for arranging atoms Hydrogen & Halogens (F, Cl, Br, I) can only bond with one other atom—they can’t go in the middle of a molecules  Always put them around the outside In general, write out the atoms in the same order as they appear in the chemical formula

28. Repeat first two steps from before 1.Use the periodic table to decide how many electrons are around each atom 2.Write the electrons around each atom Example: Draw the Lewis Structure for CH 4

29. H H Repeat first two steps from before 1.Use the periodic table to decide how many electrons are around each atom 2.Write the electrons around each atom 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 Each hydrogen has 1

30. H H 3: 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!

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

32. What if they’re not all full after that? Sometimes, the first 3 steps don’t leave you with full valence shells for all atoms Example: Draw the Lewis Structure for CH 2 O

33. 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

34. 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

35. Properties of multiple bonds Single Bond Double Bond Triple Bond Shorter bonds (atoms closer together) Stronger bonds (takes more energy to break)

36. Polyatomic Ions

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

38. Polyatomic Ions They are a group of atoms bonded together that have an overall charge 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

39. 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

40. Isomers

41. More than one possibility Often, there’s more than one way to correctly draw a Dot Structure HCCHC H H HCCHC H H Chemical Formula: C 3 H 4 Contains 2 sets of double bonds between carbons Contains 1 triple bond and 1 single bond between carbons Both structures have full valence shells!

42. Both are “correct” The chemical formula alone does not give you enough information to differentiate between the two structures HCCHC H H HCCHC H H Chemical Formula: C 3 H 4 You’ll learn in Chapter 11 how to differentiate between these two structures with chemical names

43. Isomers Isomers: Structures with the same chemical formula but different chemical structure Atoms must be bonded differently (multiple versus single bonds) or in a different order) but have the same overall chemical formula to be isomeric structures

44. Section 5.3—Molecules in 3D

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

46. 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 (that bonds repel each other because they’re like charges) attempts to explain why molecules form the shapes they form

47. What shapes do molecules form? Linear 2 bonds, no lone pairs Trigonal planar 3 bonds, no lone pairs Indicates a bond coming out at you Indicates a bond going away from you

48. What shapes do molecules form? Tetrahedron 4 bonds, no lone pairs Trigonal bipyramidal 5 bonds, no lone pairs

49. What shapes do molecules form? Octahedron 6 bonds, no lone pairs

50. 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!

51. What shapes do molecules form? Bent 2 bonds, 1 lone pair Trigonal pyramidal 3 bonds, 1 lone pair

52. What shapes do molecules form? Bent 2 bonds, 2 lone pairs

53. 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.

54. 109.5° C 105° O Example of angle distortion

55. Ionic Compound structures 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

56. Section 5.4—Polarity of Molecules

57. Electronegativity The pull an atom has for the electrons it shares with another atom in a bond. Electronegativity is a periodic trend  As atomic radius increases and number of electron shells increases, the nucleus of an atom has less of a pull on its outermost electrons

58. Periodic Table with Electronegativies increases decreases

59. Polar Bond A polar covalent bond is when there is a partial separation of charge One atom pulls the electrons closer to itself and has a partial negative charge. The atom that has the electrons farther away has a partial positive charge

60. Two atoms sharing equally N N Each nitrogen atom has an electronegativity of 3.0 They pull evenly on the shared electrons The electrons are not closer to one or the other of the atoms This is a non-polar covalent bond

61. Atoms sharing almost equally Electronegativities: H = 2.1 C = 2.5 The carbon pulls on the electrons slightly more, pulling them slightly towards the carbon Put the difference isn’t enough to create a polar bond This is a non-polar covalent bond CHH H H

62. Sharing unevenly Electronegativities: H = 2.1 C = 2.5 O = 3.5 The carbon-hydrogen difference isn’t great enough to create partial charges But the oxygen atoms pulls significantly harder on the electrons than the carbon does. This does create a polar covalent bond This is a polar covalent bond COH H

63. Showing Partial Charges There are two ways to show the partial separation of charges  Use of “  ” for “partial”  Use of an arrow pointing towards the partial negative atom with a “plus” tail at the partial positive atom COH H ++ -- COH H

64. Ionic Bonds Ionic bonds occur when the electronegativies of two atoms are so different that they can’t even share unevenly…one atom just takes them from the other

65. How to determine bond type Find the electronegativies of the two atoms in the bond Find the absolute value of the difference of their values  If the difference is 0.4 or less, it’s a non-polar covalent bond  If the difference is greater than 0.4 but less than 1.4, it’s a polar covalent bond  If the difference is greater than 1.4, it’s an ionic bond

66. Let’s Practice Example: If the bond is polar, draw the polarity arrow C – H O—Cl F—F C—Cl

67. Polar Bonds versus Polar Molecules Not every molecule with a polar bond is polar itself  If the polar bonds cancel out then the molecule is overall non-polar. The polar bonds cancel out. No net dipole The polar bonds do not cancel out. Net dipole

68. The Importance of VSEPR You must think about a molecule in 3-D (according to VSEPR theory) to determine if it is polar or not! Water drawn this way shows all the polar bonds canceling out. But water drawn in the correct VSEPR structure, bent, shows the polar bonds don’t cancel out! Net dipole H O H O H H

69. Let’s Practice Example: Is NH 3 a polar molecule?

70. Section 5.5—Intermolecular Forces

71. Intra- versus Inter-molecular Forces So far this chapter has been discussing intramolecular forces  Intramolecular forces = forces within the molecule (chemical bonds) Now let’s talk about intermolecular forces  Intermolecular forces = forces between separate molecules

72. Breaking Intramolecular forces Breaking of intramolecular forces (within the molecule) is a chemical change  2 H 2 + O 2  2 H 2 O  Bonds are broken within the molecules and new bonds are formed to form new molecules

73. Breaking Intermolecular forces Breaking of intermolecular forces (between separate molecules) is a physical change  Breaking glass is breaking the intermolecular connections between the glass molecules to separate it into multiple pieces.  Boiling water is breaking the intermolecular forces in liquid water to allow the molecules to separate and be individual gas molecules.

74. London Dispersion Forces All molecules have electrons. Electrons move around the nuclei. They could momentarily all “gang up” on one side This lop-sidedness of electrons creates a partial negative charge in one area and a partial positive charge in another. + Positively charged nucleus - Negatively charged electron + - - - - Electrons are fairly evenly dispersed. + - - - - As electrons move, they “gang up” on one side. ++ --

75. London Dispersion Forces Once the electrons have “ganged up” and created a partial separation of charges, the molecule is now temporarily polar. The positive area of one temporarily polar molecule can be attracted to the negative area of another molecule. ++ -- ++ --

76. Strength of London Dispersion Forces Electrons can gang-up and cause a non- polar molecule to be temporarily polar The electrons will move again, returning the molecule back to non-polar The polarity was temporary, therefore the molecule cannot always form LDF. London Dispersion Forces are the weakest of the intermolecular forces because molecules can’t form it all the time.

77. Strength of London Dispersion Forces Larger molecules have more electrons The more electrons that gang-up, the larger the partial negative charge. The larger the molecule, the stronger the London Dispersion Forces Larger molecules have stronger London Dispersion Forces than smaller molecules. All molecules have electrons…all molecules can have London Dispersion Forces

78. Dipole Forces Polar molecules have permanent partial separation of charge. The positive area of one polar molecule can be attracted to the negative area of another molecule. ++ -- ++ --

79. Strength of Dipole Forces Polar molecules always have a partial separation of charge. Polar molecules always have the ability to form attractions with opposite charges Dipole forces are stronger than London Dispersion Forces

80. Hydrogen Bonding Hydrogen has 1 proton and 1 electron.  There are no “inner” electrons. It bonds with the only one it has. When that electron is shared unevenly (a polar bond) with another atom, the electron is farther from the hydrogen proton than usual.  This happens when Hydrogen bonds with Nitrogen, Oxygen or Fluorine This creates a very strong dipole (separation of charges) since there’s no other electrons around the hydrogen proton to counter-act the proton’s positive charge.

81. Strength of Hydrogen Bond Hydrogen has no inner electrons to counter-act the proton’s charge It’s an extreme example of polar bonding with the hydrogen having a large positive charge. This very positively- charged hydrogen is highly attracted to a lone pair of electrons on another atom. This is the strongest of all the intermolecular forces.

82. Hydrogen Bond N H H N H H Hydrogen bond

83. Section 5.6—Intermolecular Forces & Properties

84. IMF’s and Properties IMF’s are Intermolecular Forces  London Dispersion Forces  Dipole interactions  Hydrogen bonding The number and strength of the intermolecular forces affect the properties of the substance. It takes energy to break IMF’s Energy is released when new IMF’s are formed

85. IMF’s and Changes in State Some IMF’s are broken to go from solid  liquid. All the rest are broken to go from liquid  gas. Breaking IMF’s requires energy. The stronger the IMF’s, the more energy is required to melt, evaporate or boil. The stronger the IMF’s are, the higher the melting and boiling point

86. Water Water is a very small molecule In general small molecules have low melting and boiling points Based on it’s size, water should be a gas under normal conditions However, because water is polar and can form dipole interactions and hydrogen bonding, it’s melting point is much higher This is very important because we need liquid water to exist!

87. IMF’s and Viscosity Viscosity is the resistance to flow  Molasses is much more viscous than water Larger molecules and molecules with high IMF’s become inter-twined and “stick” together more The more the molecules “stick” together, the higher the viscosity

88. Solubility In order from something to be dissolved, the solute and solvent must break the IMF’s they form within itself They must then form new IMF’s with each other

89. Solubility -+-+ -+-+ -+-+ -+-+ -+-+ Solvent, water (polar) ++ -- -+-+ Solute, sugar (polar) Water particles break some intermolecular forces with other water molecules (to allow them to spread out) and begin to form new ones with the sugar molecules.

90. Solubility Solvent, water (polar) ++ -- -+-+ Solute, sugar (polar) As new IMF’s are formed, the solvent “carries off” the solute—this is “dissolving” -+-+ -+-+ -+-+ -+-+ -+-+

91. Solubility If the energy needed to break old IMF’s is much greater than the energy released when the new ones are formed, the process won’t occur  An exception to this is if more energy is added somehow (such as heating)

92. Oil & Water Water has London Dispersion, Dipole and hydrogen bonding. That takes a lot of energy to break Water can only form London Dispersion with the oil. That doesn’t release much energy Much more energy is required to break apart the water than is released when water and oil combine. Water is polar and can hydrogen bond, Oil is non-polar. Therefore, oil and water don’t mix!

93. Surface Tension Surface tension is the resistance of a liquid to spread out.  This is seen with water on a freshly waxed car The higher the IMF’s in the liquid, the more the molecules “stick” together. The more the molecules “stick” together, the less they want to spread out. The higher the IMF’s, the higher the surface tension.

94. Soap & Water Soap has a polar head with a non-polar tail The polar portion can interact with water (polar) and the non-polar portion can interact with the dirt and grease (non- polar). Polar head Non-polar tail Soap

95. Soap & Water The soap surrounds the “dirt” and the outside of the this Micelle can interact with the water. The water now doesn’t “see” the non-polar dirt. Dirt

96. Soap & Surface Tension The soap disturbs the water molecules’ ability to form IMF’s and “stick” together. This means that the surface tension of water is lower when soap is added. The lower surface tension allows the water to spread over the dirty dishes.

97. What did you learn about soap?

98. Soap Inter-molecular forces Works based on Molecular Geometry Bonding types & Structures Determined by