1. Chemical Bonding Chapter 8 Concepts of Chemical Bonding

2. Chemical Bonding 8.1 Chemical Bonds Three basic types of bonds:  Ionic Electrostatic attraction between ions  Covalent Sharing of electrons  Metallic Metal atoms bonded to several other atoms. Electrons are free to move around the structure.

3. Chemical Bonding Lewis Symbols Electrons involved in chemical bonding are the valence electrons. G.N. Lewis (1875-1946) suggested a simple way of showing the valence electrons in an atom. Lewis electron-dot structures: consist of the chemical symbol for the element plus a dot for each valence electron.

4. Chemical Bonding Octet Rule Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons. An octet of electrons consists of full s and p subshells in an atom. There are many exceptions to the octet rule, but it provides a useful framework for many important concepts of bonding.

5. Chemical Bonding 8.2 Ionic Bonding Use Lewis Symbols to represent the reaction that occurs between magnesium and bromine:

6. Chemical Bonding Example: Write electron configurations for the following ions, and determine which have noble gas configurations:  Sr +2

7. Chemical Bonding Ti +2 Se -2 Ni +2 Br -1

8. Chemical Bonding 8.3 Covalent Bonding In these bonds atoms share electrons. There are several electrostatic interactions in these bonds:  Attractions between electrons and nuclei  Repulsions between electrons  Repulsions between nuclei

9. Chemical Bonding Covalent Bonds The attractions between nuclei and the electrons cause electron density to concentrate between the nuclei. As a result, the overall electrostatic interactions are attractive. A shared pair of electrons in any covalent bond acts as a king of “glue” to bind atoms together.

10. Chemical Bonding Lewis Structures The formation of covalent bonds can be represented using Lewis symbols  Examples: Draw: H 2, Cl 2, NH 3, CH 4

11. Chemical Bonding Multiple Bonds: Doubles and Triples Examples: Draw CO 2 and N 2 As a rule, the distance between bonded atoms decreases as the number of shared electron pairs increases.

12. Chemical Bonding 8.4 Polar Covalent Bonds Although atoms often form compounds by sharing electrons, the electrons are not always shared equally. Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does. Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

13. Chemical Bonding Polarity Nonpolar covalent bond is one in which the electrons are shared equally between two atoms, as in the Cl 2 and N 2 examples we just drew. Polar covalent bond, one of the atoms exerts a greater attraction for the bonding electrons than the other. If the difference in relative ability to attract electrons is large enough, an ionic bond is formed.

14. Chemical Bonding Electronegativity: The ability of atoms in a molecule to attract electrons to itself. On the periodic chart, electronegativity increases as you go…  …from left to right across a row.  …from the bottom to the top of a column.

15. Chemical Bonding Vocabulary Electronegativity – the ability of an atom IN A MOLECULE to attract electrons to itself. Ionization energy – how strongly an atom holds on to its electrons. Electron affinity – the measure of how strongly an atom attracts additional electrons.

16. Chemical Bonding Polar Covalent Bonds and Electronegativity The greater the difference in electronegativity, the more polar is the bond.

17. Chemical Bonding F 2 4.0 – 4.0 =0 Nonpolar HF 4.0 – 2.1 = 1.9 Polar Covalent LiF 4.0 – 1.0 = 3.0 Ionic

18. Chemical Bonding Example Which bond is more polar? Indicate in each case which atom has the partial negative charge.  B-Cl or C-Cl

19. Chemical Bonding Solution Use Figure 8.6  The difference in the electronegativities of chlorine and boron is 3.0 – 2.0 = 1.0  The difference between chlorine and carbon is 3.0 – 2.5 = 0.5  Therefore B-Cl is more polar. The chlorine atom carries the partial negative charge because it has a higher electronegativity.

20. Chemical Bonding 8.5 Lewis Structures Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

21. Chemical Bonding Writing Lewis Structures 1.Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.  If it is an anion, add one electron for each negative charge.  If it is a cation, subtract one electron for each positive charge. PCl 3 5 + 3(7) = 26

22. Chemical Bonding Writing Lewis Structures 2.The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds. Keep track of the electrons: 26  6 = 20

23. Chemical Bonding Writing Lewis Structures 3.Fill the octets of the outer atoms. Keep track of the electrons: 26  6 = 20  18 = 2

24. Chemical Bonding Writing Lewis Structures 4.Fill the octet of the central atom. Keep track of the electrons: 26  6 = 20  18 = 2  2 = 0

25. Chemical Bonding Writing Lewis Structures 5.If you run out of electrons before the central atom has an octet… …form multiple bonds until it does.

26. Chemical Bonding Writing Lewis Structures Then assign formal charges.  For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms.  Subtract that from the number of valence electrons for that atom: The difference is its formal charge.

27. Chemical Bonding Writing Lewis Structures The best Lewis structure…  …is the one with the fewest charges.  …puts a negative charge on the most electronegative atom.

28. Chemical Bonding 8.6 Resonance This is the Lewis structure we would draw for ozone, O 3. - +

29. Chemical Bonding Resonance But this is at odds with the true, observed structure of ozone, in which…  …both O—O bonds are the same length.  …both outer oxygens have a charge of  1/2.

30. Chemical Bonding Resonance One Lewis structure cannot accurately depict a molecule such as ozone. We use multiple structures, resonance structures, to describe the molecule.

31. Chemical Bonding Resonance Just as green is a synthesis of blue and yellow… …ozone is a synthesis of these two resonance structures.

32. Chemical Bonding Resonance In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. They are not localized, but rather are delocalized.

33. Chemical Bonding Resonance The organic compound benzene, C 6 H 6, has two resonance structures. It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

34. Chemical Bonding 8.7 Exceptions to the Octet Rule There are three types of ions or molecules that do not follow the octet rule:  Ions or molecules with an odd number of electrons.  Ions or molecules with less than an octet.  Ions or molecules with more than eight valence electrons (an expanded octet).

35. Chemical Bonding Odd Number of Electrons Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons. Examples: ClO 2, NO, and NO 2 Complete pairing of electrons is impossible.

36. Chemical Bonding Fewer Than Eight Electrons Consider BF 3 :  Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine.  This would not be an accurate picture of the distribution of electrons in BF 3.

37. Chemical Bonding Fewer Than Eight Electrons Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.

38. Chemical Bonding Fewer Than Eight Electrons The lesson is: If filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom.

39. Chemical Bonding More Than Eight Electrons The only way PCl 5 can exist is if phosphorus has 10 electrons around it. It is allowed to expand the octet of atoms on the 3rd row or below.  Presumably d orbitals in these atoms participate in bonding.

40. Chemical Bonding More Than Eight Electrons Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.

41. Chemical Bonding More Than Eight Electrons This eliminates the charge on the phosphorus and the charge on one of the oxygens. The lesson is: When the central atom is on the 3rd row or below and expanding its octet eliminates some formal charges, do so.

42. Chemical Bonding 8.8 Covalent Bond Strength Most simply, the strength of a bond is measured by determining how much energy is required to break the bond. This is the bond enthalpy. Always a positive quantity. The bond enthalpy for a Cl—Cl bond, D(Cl—Cl), is measured to be 242 kJ/mol.

43. Chemical Bonding Average Bond Enthalpies Pg 330 This table lists the average bond enthalpies for many different types of bonds. Average bond enthalpies are positive, because bond breaking is an endothermic process.

44. Chemical Bonding Average Bond Enthalpies NOTE: These are average bond enthalpies, not absolute bond enthalpies; the C—H bonds in methane, CH 4, will be a bit different than the C—H bond in chloroform, CHCl 3.

45. Chemical Bonding Enthalpies of Reaction Yet another way to estimate  H for a reaction is to compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed. In other words,  H rxn =  (bond enthalpies of bonds broken)   (bond enthalpies of bonds formed)

46. Chemical Bonding Enthalpies of Reaction CH 4 (g) + Cl 2 (g)  CH 3 Cl (g) + HCl (g) In this example, one C—H bond and one Cl—Cl bond are broken; one C—Cl and one H—Cl bond are formed.

47. Chemical Bonding Enthalpies of Reaction So,  H rxn = [D(C—H) + D(Cl—Cl)]  [D(C—Cl) + D(H—Cl)] = [(413 kJ) + (242 kJ)]  [(328 kJ) + (431 kJ)] = (655 kJ)  (759 kJ) =  104 kJ

48. Chemical Bonding Bond Enthalpy and Bond Length We can also measure an average bond length for different bond types. As the number of bonds between two atoms increases, the bond length decreases.