2. Chapter 6 -Chemical 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. 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

5. Table of Electronegativities

6. Polar-Covalent bonds Nonpolar-Covalent bonds Covalent Bonds  Electrons are unequally shared  Electronegativity difference between 0.5 and 1.9  Electrons are equally shared  Electronegativity difference of 0 to 0.4

7. 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.

8. Bonding Forces  Electron – electron repulsive forces repulsive forces  Proton – proton repulsive forces  Electron – proton attractive forces

9. Electron Dot Notation

10. 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 Fluorine

11. Hydrogen Chloride by the Octet Rule

12. Formation of Water by the Octet Rule

13. 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.

14. Shows how valence electrons are arranged among atoms in a molecule. Reflects central idea that stability of a compound relates to noble gas electron configuration. Lewis Structures

15. C H H H Cl.. Completing a Lewis Structure - CH 3 Cl Add up available valence electrons: C = 4, H = (3)(1), Cl = 7 Total = 14 Join peripheral atoms to the central atom with electron pairs. Complete octets on atoms other than hydrogen with remaining electrons Make carbon the central atom..

16. Multiple bonds Add up available valence electrons: Add up optimum valence electrons for each atom Subtract available from optimum Divide by 2 HCN H (1)+ C(4) + N(5) = 10 H (2) + C(8) + N(8) = 18 18-10 / 2 = 4 bonds H – C N

17. Multiple Covalent Bonds: Double bonds Two pairs of shared electrons

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

19. 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.

20. Resonance in Ozone Neither structure is correct.

21. Resonance in a carbonate ion: Resonance in an acetate ion: Resonance in Polyatomic Ions

22. 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

23. Ionic Bonds  Electrons are transferred  Electronegativity differences are generally greater than 1.9  The formation of ionic bonds is always exothermic!

24. Sodium Chloride Crystal Lattice Ionic compounds form solids at ordinary temperatures. Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions.

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

26. Hybridization The Blending of Orbitals

27. 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.

28. 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

29. 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

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

31. 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

32. 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

33. 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

34. 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

35. 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

36. 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

37. 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

38. 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

39. Hydrogen Bonding in Water

40. Hydrogen Bonding between Ammonia and Water

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

42. 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.

43. London Forces in Hydrocarbons