Chapter 9 : Chemical Bonding I : Lewis Theory

1. Chapter 9 : Chemical Bonding I : Lewis Theory • Outline • Types of Chemical Bonds • Representing Valence Electrons • Ionic Bonding : Lattice Energies • Covalent Bonding : Lewis Structures • Electronegativity and Bond Polarity • Resonance and Formal Charge

2. Types of Chemical Bonds Chemical bonds form because they lower the potential energy of the charged particles (protons and electrons) that compose atoms. Potential Energy between two charged particles (q1 and q2) If q1 and q2 have the same sign the charges repel each other. If q1 and q2 have opposite signs the charges attract each other. Chemical Bonds

3. e- p+ p+ e- Repulsive Attractive Chemical Bonds

4. Types of Chemical Bonds Chemical Bonds

5. Representing Valence Electrons It is the valence electrons that play a fundamental role in bonding. Bonding theories should therefore focus on valence electrons. 1s2 2s2 2p4 In a Lewis structure, dots are placed around the atom to represent the valence electrons. Valence Electrons

6. There is something special about the electronic configuration of the noble gases (ns2np6) that make them chemically inert. When atoms of other elements combine with one another (bonding) they acquire noble gas electronic configurations. Helium (1s2)is a bit of an exception as it contains two electrons (a duet) instead of an octet. Valence Electrons

7. Lewis Symbols and Structures In a Lewis symbol, the nucleus and core electrons are represented by the atomic symbol. The valence electrons are represented with dots. C : [He] 2s2 2p2 N : [He] 2s2 2p3 Lewis Structures

8. Examples Ionic bonding (NaCl) Lewis Structures for Ionic Compounds A Lewis structure, is a combination of Lewis symbols that represents either the transfer or sharing of electrons in a chemical bond. Lewis Structures

9. Na ● Cl Cl ● ● [Na]+ [ ]- ● ● ● ● ● ● ● ● ● ● ● ● ● Practice Problems for Ionic Lewis Structures Write Lewis Structures for the following ionic compounds; a) LiBr b) BaCl2 c) Mg3N2 Lewis Structures

10. Energetics of Ionic Bond Formation • the ionization energy of the metal is endothermic • the electron affinity of the nonmetal is exothermic • generally, the ionization energy of the metal is larger than the electron affinity of the nonmetal, therefore the formation of the ionic compound should be endothermic • but the heat of formation of most ionic compounds is exothermic and generally large; Why?

11. Ionic Bonds • electrostatic attraction is nondirectional!! • no ionic molecule • chemical formula is an empirical formula, simply giving the ratio of ions based on charge balance • ions arranged in a pattern called a crystal lattice

12. Lattice Energy • the lattice energy is the energy released when the solid crystal forms from separate ions in the gas state • hard to measure directly, but can be calculated from knowledge of other processes • lattice energy depends directly on size of charges and inversely on distance between ions

13. Trends in Lattice Energy Ion Size • the force of attraction between charged particles is inversely proportional to the distance between them • larger ions mean the centre of positive charge ( ) is farther away from negative charge

14. Lattice Energy vs. Ion Size

15. Lattice Energy = Lattice Energy = Trends in Lattice Energy Ion Charge • the force of attraction between oppositely charged particles is directly proportional to the product of the charges • larger charge means the ions are more strongly attracted • larger charge

16. Cl Cl ● ● ● Cl ● ● ● ● Cl ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● Lewis Structures for Covalent Compounds Covalent bonding (Cl2) Notice than an electron pair is now shared between the two atoms. Both now have an octet. Covalent Bonding

17. H H ● ● ● Cl Cl ● ● ● ● ● ● ● ● ● ● ● ● ● H Atoms in Covalent Bonds Covalent bonding (HCl) Covalent Bonding

18. Cl Cl ● ● Cl Cl ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● Bond Pairs and Lone Pairs Pairs of electrons that are shared between bonded atoms are called bonding pairs. Non-bonding pairs of electrons are called lone pairs. ● ● Bonding pairs are often represented with a dash. Covalent Bonding

19. Examples • Write Lewis structures for the following; • Ammonia • CCl4 Covalent Bonding

20. + H - H Cl H ● + H H N ● ● ● ● ● ● ● H N ● H ● ● ● ● ● ● ● ● ● ● ● Cl H ● ● ● ● ● ● ● ● ● ● ● ● Coordinate Covalent (Dative) Bonds Sometimes each bonding atom does not contribute one electron when forming a bonding pair. In a coordinate covalent bond, one atom donates both electrons to the bonding pair. Covalent Bonding

21. H H O O ● ● H H C C ● ● ● H H O ● H C C O H ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● Multiple Covalent Bonds Let’s look at the Lewis structure of formaldehyde, H2CO Double bond between C and O Covalent Bonding

22. Example Determine the Lewis structure of molecular nitrogen, N2 Covalent Bonding

23. O O ● ● O O ● ● O O ● ● ● ● ● ● ● ● ● ● ● ● O ● ● ● ● ● ● ● ● O ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● Lewis Structure of O2 Like the halogens, oxygen is more stable as a molecular species, O2, rather than as an atom O. Let’s look at its Lewis structure. Covalent Bonding

24. ● ● O ● ● O ● ● ● ● Covalent Bonding

25. Sometimes the bonding pair electrons are shared equally H ● ● Cl Cl ● ● Cl ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● Polar Covalent Bonds A covalent bond involves pairs of electrons shared between two atoms The electrons in the bonding pair are more attracted to the chlorine atom because it has a higher affinity for electrons Electronegativity and Bond Polarity

26. Polar Bond Non-polar Bonds Electronegativity and Bond Polarity

27. δ- δ+ H ● ● Cl ● ● ● ● ● ● In a Lewis structure a polar bond is sometimes shown with partial charges indicated. Electronegativity and Bond Polarity

28. Dipole Moments Dipole moments allow molecules to be oriented by electric fields. Electronegativity and Bond Polarity

29. Calculating Dipole Moments Polar bonds have dipole moments. The bigger the dipole moment, the more polar the bond. Units of C·m or debye (D). 1D = 3.34x10-34 C·m. Calculate the dipole moment between an electron and proton separated by 130 pm. Electronegativity and Bond Polarity

30. Electronegativity Electronegativity (EN) describes an atom’s ability to compete for electrons with other atoms to which it is bonded. Related to ionization energy (IE) and electron affinity (EA) but not quite the same thing. IE and EA describe the affinity of an isolated atom for electrons. EN describes the affinity of a bonded atom for electrons. Electronegativity and Bond Polarity

31. The Pauling Scale EN values range from ~ 0.7 to 4.0. Higher values indicate larger electronegativity (more ownership of electrons). High EN – more nonmetallic (more likely to hog the shared electron) Electronegativity and Bond Polarity

32. A Continuum of Bond Types Electronegativity and Bond Polarity

33. Measured dipole moment of bond X 100% % ionic character = Dipole moment if e- was fully transferred Percent ionic character The percent ionic character is a measure of how ionic a bond is. Recall that a 100% ionic bond has complete transfer of an electron from atom A to atom B. Example : calculate the percent ionic character of a bond that has a bond length of 130 pm and a dipole moment of 3.5D. Electronegativity and Bond Polarity

34. Electronegativities allow one to determine the relative polarities of covalent bonds Electronegativity and Bond Polarity

35. Examples • Without using the EN tables place the following molecules in order of the most polar bond ? • CCl4 ; CBr4 ; CF4 ; CI4 Electronegativity and Bond Polarity

36. Which of the following bonds is the most polar and which is least polar ? H-Cl ; H-O ; H-C • What is the percent ionic character of each of these bonds ? Electronegativity and Bond Polarity

37. Writing Lewis Structures Basic Rules (summary of what we’ve discussed so far) 1) All the valence electrons of all the atoms in the molecular formula must appear in the Lewis Structure. 2) Usually, all the electrons are paired in the Lewis structure. 3) Usually, every atom in the Lewis structure acquires an octet of electrons. Biggest exception is H (two electrons). There are a few other exceptions. 4) Sometimes, multiple bonds are required to satisfy the octet rule. Multiple covalent bonds are most often formed by C,N, O, P and S atoms. Lewis Structures

38. Strategy for Writing Lewis Structures Lewis Structures

39. Example Write plausible Lewis structures for ; a) CO2 and b) CO c) H2O Lewis Structures

40. .. .. .. + + O=N=O .. O N O .. .. .. .. number of valence e- in free atom number of associated valence e- in bonded atom F.C. = - Formal Charges Write a plausible Lewis Structure for NO2+ Both obey the rules for Lewis Structures given so far. Are both correct ? Formal charges are apparent charges on atoms in a Lewis structure that arise when atoms have not contributed equal numbers of electrons to the covalent bonds joining them. Resonance and Formal Charge

41. number of valence e- in free atom number of associated valence e- in bonded atom F.C. = - Resonance and Formal Charge

42. .. + .. O N O .. .. Resonance and Formal Charge

43. .. .. .. + + O=N=O .. O N O .. .. .. .. Extra rules for determining Lewis structures • The sum of the formal charges in a Lewis structure must equal the charge on the molecule or ion. • Where formal charges are required, they should be as small as possible. • Negative formal charges usually appear on the more electronegative atoms. Positive formal charges appear on the least electronegative atoms. Resonance and Formal Charge

44. .. .. .. .. .. .. .. .. .. .. .. O=Cl-N .. O-Cl=N O=N-Cl .. .. .. .. .. .. .. .. .. O-N=Cl .. .. .. Example Four possible Lewis structures for nitrosyl chloride (NOCl) are given below. Which one is most plausible ? (a) (b) (c) (d) Resonance and Formal Charge

45. Bonus Question Provide reasons why is N-O-Cl unlikely to be the skeleton structure for this molecule. Resonance and Formal Charge

46. Resonance Ozone, (O3), is the stuff you smell after a thunderstorm and when you make photocopies. Possible Lewis structures for ozone are ; Both obey all the rules of Lewis Structures including formal charge rules. Do they agree with experiment ? O-O bond length is 145 pm O=O bond length is 121 pm Resonance and Formal Charge

47. The true Lewis structure of ozone is neither the two structures given above but rather a composite of both. (a) (b) The situation in which two or more plausible Lewis structures can be written but the “correct” structure cannot be written at all is called resonance. Example : write the Lewis structure for the carbonate ion (CO3-2) Resonance and Formal Charge

48. •• • O-H • O-H •• Exceptions to the Octet Rule 1) Odd-electron species What’s the Lewis structure of NO ? NO is a stable, odd-electron molecule  pretty rare Most odd-electron species are highly reactive and are called free radicals transitory species. Exceptions to the Octet Rule

49. 2) Incomplete octets What’s the Lewis structure of BH3 ? Species with incomplete octets are limited to some Be, B and Al compounds Exceptions to the Octet Rule

50. Cl Cl P Cl P Cl Cl Cl Cl Cl • • F F F •• •• •• •• •• •• •• •• •• •• •• •• •• •• •• •• •• •• •• •• •• •• •• •• •• •• •• •• S •• •• •• •• •• •• •• •• •• •• •• •• •• •• F F F 3) Expanded Valence Shells There are a few atoms that can accommodate 10 or even 12 valence electrons. This occurs for atoms in the third period or beyond that are bonded to very electronegative atoms. Exceptions to the Octet Rule