2. Chapter 8 -Covalent Bonding A steroid alkaloid derived from skin secretions of the Phyllobates and Dendrobates genera of South American poison- arrow frogs. It is one of the most potent venoms known. Batrachotoxin

3. Bonds Forces that hold groups of atoms  Forces that hold groups of atoms together and make them function together and make them function as a unit. as a unit. Ionic bonds – transfer of electrons  Ionic bonds – transfer of electrons  Covalent bonds – sharing of electrons

4. Molecule – neutral group of atoms that are held together by covalent bonds. Diatomic molecule – molecule containing only two atoms H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2, Molecular compound – a chemical compound whose whole simplest unit is a molecule.

5. Molecular formula – shows the types and number of atoms combined in a single molecule of a molecular compound. Example: NH 3 Does not show arrangement or space. These models do! ammonia Space filling Perspective drawing Ball & stick Structural formula

6. Electron Dot Notation

7. The Octet Rule Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level. Diatomic Hydrogen H : H shared pair H- H

8. Comments About the Octet Rule 2nd row elements C, N, O, F observe the octet rule. 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals. When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

9. Lewis Structures (Structural formula) Shows how valence electrons are arranged among Shows how valence electrons are arranged among atoms in a molecule. atoms in a molecule. Covalent bond represented by the line Covalent bond represented by the line Unshared pair is the non- bonding pairs surrounding the peripheral atoms.

10. C H H H H.. Completing a Lewis Structure – CH 4 Add up available valence electrons: C = 4, H = (4)(1) Total = 8 Join peripheral atoms to the central atom with electron pairs. Complete octets on atoms other than hydrogen with remaining electrons Total bonds = total # of electrons – total valence electrons ÷ 2 Make carbon the central atom

11. Multiple Covalent Bonds: Double bonds Two pairs of shared electrons CH 4 ( )

12. Multiple Covalent Bonds: Triple bonds Three pairs of shared electrons C2H2C2H2

13. Coordinate Covalent Bonds One atom contributes both electrons in the bond. ·C· + : O :  : C :: O 1s 2s 2p C O : C :: O  : C ::: O : · ·· ···

14. Polyatomic Ions Tightly bound group of atoms that contain both covalent and coordinate covalent bonds. Lets draw: NH 4 +,

15. Exceptions to the Octet Rule Lets draw! BF 3, PCl 5, NO 2

16. Resonance Occurs when more than one valid Lewis structure can be written for a particular molecule. These are resonance structures. The actual structure is an average of the resonance structures.

17. Resonance in Ozone Neither structure is correct.

18. Resonance in a carbonate ion: CO 3 - Resonance in an acetate ion: Resonance in Polyatomic Ions

19. Bonding Forces  Electron – electron repulsive forces repulsive forces  Proton – proton repulsive forces  Electron – proton attractive forces Bond Length- average distance between two bonded atoms (minimal potential energy)

20. Bond Length Diagram

21. Bond Dissociation Energy (BDE) It is the energy required to break the bond between two covalently bonded atoms. It gives us information about the strength of a bonding interaction. (the larger the energy the stronger the bond. The greater the BDE the less reactive.

22. VSEPR Model The structure around a given atom is determined principally by minimizing electron pair repulsions. (Valence Shell Electron Pair Repulsion)

23. Predicting a VSEPR Structure Draw Lewis structure. Put pairs as far apart as possible. Put pairs as far apart as possible. Determine positions of atoms from the way electron pairs are shared. Determine positions of atoms from the way electron pairs are shared. Determine the name of molecular structure from positions of the atoms. Determine the name of molecular structure from positions of the atoms.

24. AB 2 Linear BeF 2 180° AB 3 Trigonal-planar BF 3 120° AB 4 Tetrahedral CH 4 109.5°

26. Table – VSEPR Structures

27. VSEPR and Phosphorus hexachloride

29. VSEPR and Unshared Pairs Example NH 3 1 unshared pair Lone pairs occupy space like the shared pairs, but shape refers to position of atoms only. AB 3 E Trigonal Pyramidal

30. VSEPR and the ammonia molecule

32. H 2 0 2 unshared pairs AB 2 E 2 Bent (angular) Oxygen is in the center w/2 corners occupied by the unshared and two corners occupied by the hydrogen.

33. VSEPR and the water molecule

34. Double and triple bonds are treated in the same way as single bonds. Along with polyatomic ions.

35. VSEPR and a molecule of I 3 Which structure is the correct one? # 3

36. VSEPR and Xenon tetrafluoride Which one will it be???

37. Hybridization The Blending of Orbitals

38. We have studied electron configuration notation and the sharing of electrons in the formation of covalent bonds. Methane is a simple natural gas. Its molecule has a carbon atom at the center with four hydrogen atoms covalently bonded around it. Lets look at a molecule of methane, CH 4.

39. What is the expected orbital notation of carbon in its ground state? (Hint: How many unpaired electrons does this carbon atom have available for bonding?) Can you see a problem with this? Carbon ground state configuration

40. You should conclude that carbon only has TWO electrons available for bonding. That is not not enough! How does carbon overcome this problem so that it may form four bonds? Carbon’s Bonding Problem

41. The first thought that chemists had was that carbon promotes one of its 2s electrons… …to the empty 2p orbital. Carbon’s Empty Orbital

42. However, they quickly recognized a problem with such an arrangement… Three of the carbon-hydrogen bonds would involve an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom. A Problem Arises

43. This would mean that three of the bonds in a methane molecule would be identical, because they would involve electron pairs of equal energy. But what about the fourth bond…? Unequal bond energy

44. The fourth bond is between a 2s electron from the carbon and the lone 1s hydrogen electron. Such a bond would have slightly less energy than the other bonds in a methane molecule. Unequal bond energy #2

45. This bond would be slightly different in character than the other three bonds in methane. This difference would be measurable to a chemist by determining the bond length and bond energy. But is this what they observe? Unequal bond energy #3

46. The simple answer is, “No”. Chemists have proposed an explanation – they call it Hybridization. Hybridization is the combining of two or more orbitals of nearly equal energy within the same atom into orbitals of equal energy. Measurements show that all four bonds in methane are equal. Thus, we need a new explanation for the bonding in methane. Enter Hybridization

47. In the case of methane, they call the hybridization sp 3, meaning that an s orbital is combined with three p orbitals to create four equal hybrid orbitals. These new orbitals have slightly MORE energy than the 2s orbital… … and slightly LESS energy than the 2p orbitals. sp 3 Hybrid Orbitals

48. Here is another way to look at the sp 3 hybridization and energy profile… sp 3 Hybrid Orbitals

49. While sp 3 is the hybridization observed in methane, there are other types of hybridization that atoms undergo. These include sp hybridization, in which one s orbital combines with a single p orbital. Notice that this produces two hybrid orbitals, while leaving two normal p orbitals sp Hybrid Orbitals

50. Another hybrid is the sp 2, which combines two orbitals from a p sublevel with one orbital from an s sublevel. Notice that one p orbital remains unchanged. sp 2 Hybrid Orbitals

51. What shape corresponds with what hybridization? How many hybridized orbitals are made? sp 3 sp 2 sp ? ? ?

52. 1.Draw the Lewis Structure. 2.Count the “things” stuck to it. 3.Count the lone pairs. 4.Follow the flow chart.

53. Molecular Orbitals 1.Click the link below to watch the video 2.Print hybridization worksheet under handout section of my web page and complete.

54. Polus (Latin) for Pole A pole is a part of a system that has opposite electric or magnetic positions. –Earth (magnetic north and south) –Batteries (positive and negative poles) –Bonds (polar covalent and non-polar covalent)

55. Polar-Covalent bonds Nonpolar-Covalent bonds Covalent Bonds: electron sharing  Electrons are unequally shared  Electronegativity difference between.3 and 1.7  5- 50 % ionic character  Electrons are equally shared  Electronegativity difference of 0 to 0.3  0-5 % ionic character Ionic  Electrons are transferred  Electronegativity difference is greater than 1.7

56. Electronegativity The ability of an atom in a molecule to attract shared electrons to itself. The ability of an atom in a molecule to attract shared electrons to itself. Linus Pauling 1901 - 1994

57. Table of Electronegativities

58. What is the electronegativity difference between the two atoms, determine if it is polar or non-polar covalent. N= 3.0 H = 2.1 F = 4.0 Cl = 3.0 Al= 1.5 Ca = 1.0 a. N and H b. F and F c. Ca and Cl d. Al and Cl 0.9, moderately polar 0, non polar 2.0, ionic 1.5, very polar

59. Polarity A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment. A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment. What is the difference in elecronegativity? H= 2.1 and F= 4.0

60. Sample Problem 1 Classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs; and chlorine, Cl. Draw the dipole moment? S= 2.5 H= 2.1 Cs= 0.8 Cl= 3.0

61. Table of dipole moments

62. If the bond is polar…..is the molecule also polar? CO 2 O=C=O Label the dipoles: H 2 O O H Which is a polar molecule?

63. Homework for January 16, 2013 1. Click on the link below to watch the video. 2. Print and complete the accompanying worksheet “Intermolecular vs. Intra molecular forces” Intra and Intermolecular forces

64. 1. What is an intramolecular bond? The electrostatic force between two atoms in a molecule. 2. Which type of intramolecular bond has the lowest melting and boiling point and why? Covalent bonds- less of a difference in electronegativity; weak bond Think about the structure of ionic vs. covalent molecules!! 3. Name three intra molecular bond types and compare their relative strength? Covalent bond < Metallic bond < Ionic Bond 4. What is an intermolecular bond? Force of attraction between neighboring molecules. 5. Name the three types of intermolecular bonds? Place them in order of increasing strength? London dispersion forces< dipole-dipole forces < hydrogen bonding

65. So…..why do I need to know all of this stuff??? 1.How a bullet proof vest work? 2.How does soap clean your dishes? 3.Why does wax bead up on your car?

66. Intra-molecular bonds: the force of attraction between the atoms in a molecule. Covalent bonds Inter-molecular bonds: the force of attraction between two or more molecules. Hydrogen bonds dipole-dipole forces London dispersion forces

67. Relative magnitudes of forces The types of bonding forces vary in their strength as measured by average bond energy. Covalent bonds (400 kcal) Hydrogen bonding (12-16 kcal ) Dipole-dipole interactions (2-0.5 kcal) London forces (less than 1 kcal) Strongest Weakest

68. Hydrogen Bonding Hydrogen bonding in Kevlar, a strong polymer used in bullet- proof vests. Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen Cohesive forces (hydrogen bonding) pull the water dropplet into a spherical shape.

69. Hydrogen Bonding in Water

70. Hydrogen Bonding between Ammonia and Water

71. Dipole-Dipole Attractions Attraction between oppositely charged regions of neighboring molecules.

72. The water dipole

73. The ammonia dipole

74. London Dispersion Forces The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors. Fritz London 1900-1954 London forces increase with the size of the molecules.

75. London Forces in Hydrocarbons

77.  Cat’s Meow Lab

78.  Problem: What will happen if you mix milk, soap and food coloring?  Hypothesis:  If _____________ then ___________

79. Hypothesis:  If ….Then statement  Do you think these items will react together?  What type of reaction could occur?  Why?

80. Procedures:  1. Pour milk into the petri dish  2. Place four drops of food coloring into the milk.  Use four different colors  Place in the center  3. Touch the surface of the milk with the q-tip.  4. Dip the q-tip in detergent and then in the milk.  5. Record Observations.

81. Observations  Data Table 1: Observations Milk + food coloringMilk, food coloring & soap

82. Conclusion:  1 Paragraph (3 to 5 sentences) describing:  What do you think happened?  Was your hypothesis correct?  HINTS!!  Milk is a colloid; made of water, fat, sugar, and protein.  Soap has a hydrophilic and a hydrophobic end.

83. Videos of Alternate Experiments  Heavy Cream Heavy Cream  Almond Milk Almond Milk  Half and Half Half and Half

84.  2 % milk

85. Covalent Network Compounds Some covalently bonded substances DO NOT form discrete molecules. Diamond, a network of covalently bonded carbon atoms Graphite, a network of covalently bonded carbon atoms

86. Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world. Models can be physical as with this DNA model Models can be mathematical Models can be theoretical or philosophical

87. Fundamental Properties of Models A model does not equal reality. Models are oversimplifications, and are therefore often wrong. Models become more complicated as they age. We must understand the underlying assumptions in a model so that we don’t misuse it.