1. Chapter 8 Bonding: General Concepts

2. Chapter 8 Questions to Consider  What is meant by the term “chemical bond”?  Why do atoms bond with each other to form compounds?  How do atoms bond with each other to form compounds? Copyright © Cengage Learning. All rights reserved 2

3. Chapter 8 Bonding Theories  Explain how and why atoms attach together to form molecules  Explain why some combinations of atoms are stable and others are not  why is water H 2 O, not HO or H 3 O  Can be used to predict the shapes of molecules  Can be used to predict the chemical and physical properties of compounds

4. Chapter 8 Types of Bonds Types of AtomsType of Bond Bond Characteristic metals to nonmetals Ionic electrons transferred nonmetals to nonmetals Covalent electrons shared metals to metals Metallic electrons pooled  We can classify bonds based on the kinds of atoms that are bonded together

5. Chapter 8 Lewis Bonding Theory  One of the simplest bonding theories is called Lewis Theory  Lewis Theory emphasizes valence electrons to explain bonding  Using Lewis theory, we can draw models – called Lewis structures  aka Electron Dot Structures  Lewis structures allow us to predict many properties of molecules  such as molecular stability, shape, size, polarity G.N. Lewis (1875-1946)

6. Chapter 8 Lewis Structures of Atoms  We use the symbol of element to represent nucleus and inner electrons  And we use dots around the symbol to represent valence electrons  pair first two dots for the s orbital electrons  put one dot on each open side for first three p electrons  then pair rest of dots for the remaining p electrons

7. Chapter 8 Lewis Structures of Atoms

8. Section 8.10 Lewis Structures Lewis Theory and Ionic Bonding  Lewis symbols can be used to represent the transfer of electrons from metal atom to nonmetal atom, resulting in ions that are attracted to each other and therefore bond + Atoms bond because it results in a more stable electron configuration i.e like a noble gas.

9. Section 8.10 Lewis Structures Example  Using Lewis theory to predict chemical formulas of ionic compounds Predict the formula of the compound that forms between calcium and chlorine Transfer all the valence electrons from the metal to the nonmetal, adding more of each atom as you go, until all electrons are lost from the metal atoms and all nonmetal atoms have eight electrons. CaCl 2

10. Section 8.10 Lewis Structures Tro: Chemistry: A Molecular Approach, 2/e Electrons that are shared by atoms are called bonding pairs. Electrons that are not shared by atoms but belong to a particular atom are called lone pairs Another way atoms can achieve an octet of valence electrons is to share their valence electrons with other atoms. The shared electrons would then count toward each atom’s octet. The sharing of valence electrons is called covalent bonding Lewis Theory of Covalent Bonding

11. Section 8.10 Lewis Structures Octet Rule  Elements form stable molecules when surrounded by eight electrons. Copyright © Cengage Learning. All rights reserved

12. Section 8.10 Lewis Structures Single Covalent Bond  A covalent bond in which two atoms share one pair of electrons. H–H Copyright © Cengage Learning. All rights reserved 12

13. Section 8.10 Lewis Structures Double Covalent Bond  A covalent bond in which two atoms share two pairs of electrons. O=C=O Copyright © Cengage Learning. All rights reserved 13

14. Section 8.10 Lewis Structures Triple Covalent Bond  A covalent bond in which two atoms share three pairs of electrons. Copyright © Cengage Learning. All rights reserved 14

15. Section 8.10 Lewis Structures Steps for Writing Lewis Structures 1.Write the skeletal structure: (Make least electronegative atom the central) 2.Sum the valence electrons from all the atoms: (addone for each – charge and subtract one for each + charge) 3.Use a pair of electrons to form a bond between each pair of bound atoms. 4.Atoms usually have noble gas configurations. Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen). Copyright © Cengage Learning. All rights reserved 15

16. Section 8.10 Lewis Structures Steps for Writing Lewis Structures 1.Sum the valence electrons from all the atoms. (Use the periodic table.) Example: H 2 O 2 (1 e – ) + 6 e – = 8 e – total Copyright © Cengage Learning. All rights reserved 16

17. Section 8.10 Lewis Structures Steps for Writing Lewis Structures 2.Use a pair of electrons to form a bond between each pair of bound atoms. Example: H 2 O Copyright © Cengage Learning. All rights reserved 17

18. Section 8.10 Lewis Structures Steps for Writing Lewis Structures 3.Atoms usually have noble gas configurations. Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen). Examples: H 2 O, PBr 3, and HCN Copyright © Cengage Learning. All rights reserved 18

19. Section 8.10 Lewis Structures Draw a Lewis structure for each of the following molecules: NH 3 CO 2 CCl 4 Copyright © Cengage Learning. All rights reserved 19 CONCEPT CHECK!

20. Section 8.11 Exceptions to the Octet Rule  Boron tends to form compounds in which the boron atom has fewer than eight electrons around it (it does not have a complete octet). BH 3 = 6e – Copyright © Cengage Learning. All rights reserved 20

21. Section 8.11 Exceptions to the Octet Rule  When it is necessary to exceed the octet rule for one of several third-row (or higher) elements, place the extra electrons on the central atom. SF 4 = 34e – AsBr 5 = 40e – Copyright © Cengage Learning. All rights reserved 21

22. Section 8.11 Exceptions to the Octet Rule Draw a Lewis structure for each of the following molecules: BF 3 PCl 5 SF 6 Copyright © Cengage Learning. All rights reserved 22 CONCEPT CHECK!

23. Section 8.11 Exceptions to the Octet Rule Let’s Review  C, N, O, and F should always be assumed to obey the octet rule.  B and Be often have fewer than 8 electrons around them in their compounds.  Second-row elements never exceed the octet rule.  Third-row and heavier elements often satisfy the octet rule but can exceed the octet rule by using their empty valence d orbitals. Copyright © Cengage Learning. All rights reserved 23

24. Section 8.11 Exceptions to the Octet Rule Let’s Review  When writing the Lewis structure for a molecule, satisfy the octet rule for the atoms first. If electrons remain after the octet rule has been satisfied, then place them on the elements having available d orbitals (elements in Period 3 or beyond). Copyright © Cengage Learning. All rights reserved 24

25. Section 8.11 Exceptions to the Octet Rule Rules Governing Formal Charge  To calculate the formal charge on an atom: 1.Take the sum of the lone pair electrons and one-half the shared electrons. 2.Subtract the number of assigned electrons from the number of valence electrons on the free, neutral atom. Copyright © Cengage Learning. All rights reserved 25

26. Section 8.11 Exceptions to the Octet Rule Rules Governing Formal Charge  The sum of the formal charges of all atoms in a given molecule or ion must equal the overall charge on that species. Copyright © Cengage Learning. All rights reserved 26

27. Section 8.11 Exceptions to the Octet Rule Formal Charge  Formal charge = (# valence e – on free neutral atom) – (# valence e – assigned to the atom in the molecule).  Assume:  Lone pair electrons belong entirely to the atom in question.  Shared electrons are divided equally between the two sharing atoms. Copyright © Cengage Learning. All rights reserved 27

28. Section 8.11 Exceptions to the Octet Rule Consider the Lewis structure for POCl 3. Assign the formal charge for each atom in the molecule. P: 5 – 4 = +1 O: 6 – 7 = –1 Cl: 7 – 7 = 0 Copyright © Cengage Learning. All rights reserved 28 CONCEPT CHECK!

29. Section 8.12 Resonance.... OSO.................. OSO.............. When there is more than one Lewis structure for a molecule that differ only in the position of the electrons, they are called resonance structures The actual molecule is a combination of the resonance forms – a resonance hybrid Look for multiple bonds or lone pairs

30. Section 8.12 Resonance Drawing Resonance Structures 1.Draw first Lewis structure that maximizes octets 2.Assign formal charges 3.Move electron pairs from atoms with (−) formal charge toward atoms with (+) formal charge 4.If (+) fc atom 2 nd row, only move in electrons if you can move out electron pairs from multiple bond 5.If (+) fc atom 3 rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet

31. Section 8.12 Resonance  More than one valid Lewis structure can be written for a particular molecule. NO 3 – = 24e – Copyright © Cengage Learning. All rights reserved 31

32. Section 8.12 Resonance  Actual structure is an average of the resonance structures.  Electrons are really delocalized – they can move around the entire molecule. Copyright © Cengage Learning. All rights reserved 32

33. Section 8.2 Electronegativity The Pauling Electronegativity Values Copyright © Cengage Learning. All rights reserved 33

34. Section 8.2 Electronegativity Copyright © Cengage Learning. All rights reserved

35. Section 8.2 Electronegativity Arrange the following bonds from most to least polar: a) N–FO–FC–F b)C–FN–OSi–F c)Cl–ClB–ClS–Cl a) C–F, N–F, O–F b) Si–F, C–F, N–O c) B–Cl, S–Cl, Cl–Cl Copyright © Cengage Learning. All rights reserved 35 EXERCISE!

36. Section 8.3 Bond Polarity and Dipole Moments Dipole Moment 36

37. Section 8.3 Bond Polarity and Dipole Moments No Net Dipole Moment (Dipoles Cancel) Copyright © Cengage Learning. All rights reserved 37

38. Section 8.8 Covalent Bond Energies and Chemical Reactions Predict Δ H for the following reaction: Given the following information: Bond Energy (kJ/mol) C–H 413 C–N 305 C–C 347 891  Δ H = –42 kJ Copyright © Cengage Learning. All rights reserved 38 CONCEPT CHECK!