1. Chapter 3 Chemical Bonding

2. Chemical Bonding How do atoms combine to form compounds? How do atoms react in chemical reactions? These questions are fundamental to chemistry since chemical changes are essentially alterations of chemical bonds. CHEMICAL BOND – the force that holds atoms together to from ionic or molecular compound. G.N. LEWIS (1916) “ Atoms interact to gain stability by changing the outermost (valence) configuration so as to attain the electronic configuration of a noble gas.”

4. Lewis Dot Symbol – shows valence electrons; symbols of the elements surrounded by dot to represent the valence electrons. *for main group elements  group # = # valence electrons

5. Octet Rule – the tendency to achieve an electronic configuration with eight valence electrons. An octet of electrons consists of a full s and p subshells of an atom. Lewis Structures – a combination of Lewis symbols representing the transfer or sharing of electrons in a chemical bond.

6. Types of Bonding 1. Ionic Bonding - involves complete transfer of electrons - usually metallic and nonmetallic elements are involved - metals lose electrons (cations) and nonmetals gain electrons (anions) Ex. Na + Cl  NaCl

7. 2. Covalent Bonding - sharing of electrons - between nonmetals or between metalloids - leads to formation of molecules (vs ionic compounds)

8. How about if you are asked to write the Lewis Structure of a polyatomic compound?

9. Rules for Writing Lewis Structures 1. Count the total number of valence electrons present – from atoms – from charge if polyatomic ion (add or subtract e–s for charge) 2. Write the skeleton structure of the compound – The atom with the least number will be the central atom – Put least EN atom as central atom and surround with other atoms – Note: H and F atoms will always be outer atoms 3. Connect all atoms by drawing single bonds between all atoms , then distribute the remaining valence electrons as lone pairs around outer atoms so each has 8 electrons octet rule: atoms form bonds such that all atoms get eight electrons

10. 4. If there are not enough electrons for each atom to have an octet, make double and/or triple bonds between central atom and surrounding atoms – BUT fluorine can only form a single bond – Note that double bonds are shorter than single bonds, and triple bonds are shorter than double bonds 5. For polyatomic ions, square brackets are drawn around the Lewis structure , and the charge is put in the upper right-hand corner

11. Figure 10.1 The steps in converting a molecular formula into a Lewis structure. Molecular formula Atom placement Sum of valence e - Remaining valence e - Lewis structure Place atom with lowest EN in center Add A-group numbers Draw single bonds. Subtract 2e - for each bond. Give each atom 8e - (2e - for H) Step 1 Step 2 Step 3 Step 4

12. Examples: 1.CCl 4 2.NF 3 3.NH 4 +

13. Molecular formula Atom placement Sum of valence e - Remaining valence e - Lewis structure For NF 3 N F F F N 5e - F 7e - X 3 = 21e - Total 26e - : : : : : : : : : :

14. SAMPLE PROBLEM 10.1 Writing Lewis Structures for Molecules with One Central Atom SOLUTION: Step 1: Carbon has the lowest EN and is the central atom. The other atoms are placed around it. C Steps 2-4: C has 4 valence e - , Cl and F each have 7. The sum is 4 + 4(7) = 32 valence e - . Cl Cl F F Make bonds and fill in remaining valence electrons placing 8e - around each atom. PROBLEM: Write a Lewis structure for CCl 2 F 2 , one of the compounds responsible for the depletion of stratospheric ozone. PLAN: Follow the steps outlined in Figure 10.1 . C Cl Cl F F : : : : : : : : : : : :

15. SAMPLE PROBLEM 10.2 Writing Lewis Structure for Molecules with More than One Central Atom SOLUTION: Hydrogen can have only one bond so C and O must be next to each other with H filling in the bonds. There are 4(1) + 4 + 6 = 14 valence e - . C has 4 bonds and O has 2. O has 2 pair of nonbonding e - . C O H H H H : : PROBLEM: Write the Lewis structure for methanol (molecular formula CH 4 O), an important industrial alcohol that is being used as a gasoline alternative in car engines.

16. SAMPLE PROBLEM 10.3 Writing Lewis Structures for Molecules with Multiple Bonds. PLAN: SOLUTION: PROBLEM: Write Lewis structures for the following: (a) Ethylene (C 2 H 4 ), the most important reactant in the manufacture of polymers (b) Nitrogen (N 2 ), the most abundant atmospheric gas For molecules with multiple bonds, there is a Step 5 which follows the other steps in Lewis structure construction. If a central atom does not have 8e - , an octet, then two e - (either single or nonbonded pair)can be moved in to form a multiple bond. (a) There are 2(4) + 4(1) = 12 valence e - . H can have only one bond per atom. : (b) N 2 has 2(5) = 10 valence e - . Therefore a triple bond is required to make the octet around each N. C C H H H H C C H H H H N : N : . . . . N : N : . . N : N :

17. Exceptions to the Octet Rule 1. s2 group 2. molecules with odd number of electrons; e.g. NO, NO2 3.molecules with less than an octet: atoms that participate in bonding but do not have enough valence e-s to form an octet. e.g. BF3 4. molecules with more than an octet: atoms in periods 3 to 7 with expanded d valence shells. e.g. PCl5, SiF62-, POl3, SF6 *Since the second-period elements only have 2s and 2p valence orbitals available for bonding, and these orbitals can hold a maximum of eight electrons, the second-period elements can never have more than an octet of electrons. *Elements in the 3rd to 7th period have unfilled nd orbitals that can be used in bonding Illustration: Draw the orbital diagram for period 3 elements.

18. SAMPLE PROBLEM 10.5 Writing Lewis Structures for Octet Rule Exceptions PLAN: SOLUTION: Draw the Lewis structures for the molecule and determine if there is an element which can be an exception to the octet rule. BFCl 2 will have only 1 Lewis structure. PROBLEM: Write Lewis structures for BFCl 2 .

19. Formal Charge: Selecting the Best Resonance Structure An atom “owns” all of its nonbonding electrons and half of its bonding electrons. Formal charge of atom = # valence e - - # of covalent bonds - # of nonbonding electrons The smaller the formal charge of the atoms, the more stable the structure For O A # valence e - = 6 # nonbonding e - = 4 # bonding e - = 4 X 1/2 = 2 Formal charge = 0 For O B # valence e - = 6 # nonbonding e - = 2 # bonding e - = 6 X 1/2 = 3 Formal charge = +1 For O C # valence e - = 6 # nonbonding e - = 6 # bonding e - = 2 X 1/2 = 1 Formal charge = -1

20. Theories of Covalent Bonding

21. Valence Bond Theory Hybridization Theory Valence Shell Electron Pair Repulsion Theory Molecular Orbital Theory

22. The Central Themes of VB Theory Basic Principle A covalent bond forms when the orbitals of two atoms overlap and the overlap region, which is between the nuclei, is occupied by a pair of electrons. The two wave functions are in phase so the amplitude increases between the nuclei.

23. The Central Themes of VB Theory Themes A set of overlapping orbitals has a maximum of two electrons that must have opposite spins. The greater the orbital overlap, the stronger (more stable) the bond. The valence atomic orbitals in a molecule are different from those in isolated atoms. There is a hybridization of atomic orbitals to form molecular orbitals.

24. Figure 11.1 Orbital overlap and spin pairing in three diatomic molecules. Hydrogen, H 2 Hydrogen fluoride, HF Fluorine, F 2

25. Hybridization Theory Basic Principle the process of mixing of atomic orbitals of the same atom to form degenerate orbitals called hybrid orbitals - the number of hybrid orbitals formed is equal to the number of pure atomic orbitals that combine.

26. Hybrid Orbitals The numbe r of hybrid orbitals obtained equals the number of atomic orbitals mixed. The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed. Key Points sp sp 2 sp 3 sp 3 d sp 3 d 2 Types of Hybrid Orbitals

27. Figure 11.2 The sp hybrid orbitals in gaseous BeCl 2 . atomic orbitals hybrid orbitals orbital box diagrams

28. Figure 11.2 The sp hybrid orbitals in gaseous BeCl 2 (continued). orbital box diagrams with orbital contours

29. Figure 11.3 The sp 2 hybrid orbitals in BF 3 .

30. Figure 11.4 The sp 3 hybrid orbitals in CH 4 .

31. Figure 11.5 The sp 3 hybrid orbitals in NH 3 .

32. Figure 11.5 continued The sp 3 hybrid orbitals in H 2 O.

33. Figure 11.6 The sp 3 d hybrid orbitals in PCl 5 .

34. Figure 11.7 The sp 3 d 2 hybrid orbitals in SF 6 .

35. Figure 11.8 Shortcut in getting the hybridization of the central atom: Molecular formula Lewis structure Count no. of atoms attached and lone pairs hybridization Figure 10.1 Step 1 Figure 10.12 Step 2 Step 3 Table 11.1

36. Number of atoms attached + Lone Pairs Hybrid Orbitals 2 sp 3 sp 2 4 sp 3 5 sp 3 d 6 sp 3 d 2

37. sp 3 hybridized Hybridization of the central atom: No. of lone pairs = 1 No. of atoms attached = 3 Lewis structure 1.NF 3 N F F F : : : : : : : : : :

38. sp 2 hybridized Hybridization of the central atom: No. of lone pairs = 1 No. of atoms attached = 2 Lewis structure 2. O 3

39. sp 3 d hybridized Hybridization of the central atom: No. of lone pairs = 1 No. of atoms attached = 4 Lewis structure 3. SF 4

40. What are the hybrid orbitals for the central atom in each of the following molecules? a. PCl 3 d. IBr 4 –: b. NH 2 – e.SCN– c. SeF 4 f. SnCl 5 –