1. © Prentice Hall 2001Chapter 11 Bonding Outer shell = valence electrons Octet rule - An atom is most stable if it has an outer shell of eight electrons and no electrons of higher energy Elements at the left of the periodic table generally have one or two electrons in excess of a stable noble gas structure (octet of electrons) These electrons are easily removed - ionization energy is low Such elements are electropositive

2. © Prentice Hall 2001Chapter 12 Bonding Elements at the right of the periodic table generally are just one or two electrons short of a noble gas structure (octet of electrons) These elements easily add electrons

3. © Prentice Hall 2001Chapter 13 Ionic Bonds Opposite charges attract Attractions between ions hold a crystal together and are called ionic bonds

4. © Prentice Hall 2001Chapter 14 Covalent Bonds Instead of giving up or acquiring electrons, an atom can also achieve eight electrons by sharing covalent bond

5. © Prentice Hall 2001Chapter 15 Covalent Bonds Other covalent bonds

6. © Prentice Hall 2001Chapter 16 Polar Covalent Bonds In the fluorine-fluorine bond as well as in the hydrogen-hydrogen bond, electrons are shared equally In hydrogen fluoride and in water, electrons are attracted more toward one atom

7. © Prentice Hall 2001Chapter 17 Polar Covalent Bonds This tendency of atoms to attract electrons is known as electronegativity There is a continuum of bonding types ionic bond polar covalent bondnonpolar covalent bond

8. © Prentice Hall 2001Chapter 18 Electronegativity

9. © Prentice Hall 2001Chapter 19 Lewis Structures The chemical symbols we have been using in which valence electrons are shown as dots are called Lewis structures

10. © Prentice Hall 2001Chapter 110 Drawing Lewis Structures Write the symbols for the elements in the correct structural order Consider nitric acid, HNO 3

11. © Prentice Hall 2001Chapter 111 Drawing Lewis Structures Calculate the number of valence electrons for all atoms in the compound 1H@1 electron= 1 3O@6 electrons=18 1N@5 electrons= 5 Total=24

12. © Prentice Hall 2001Chapter 112 Drawing Lewis Structures Put a pair of electrons between each symbol - at least one bond needed between each atom and its neighbor

13. © Prentice Hall 2001Chapter 113 Drawing Lewis Structures Beginning with the outer atoms, place remaining electrons in pairs around atoms until each has eight (except for hydrogen)

14. © Prentice Hall 2001Chapter 114 Drawing Lewis Structures If you run out of electrons before each atom (other than hydrogen) has eight electrons, move unshared pairs to form multiple bonds

15. © Prentice Hall 2001Chapter 115 Formal Charges A bookkeeping system for electrons Used to show the approximate distribution of electron density in a molecule or polyatomic ion

16. © Prentice Hall 2001Chapter 116 Formal Charges Assign each atom half of the electrons in each pair it shares

17. © Prentice Hall 2001Chapter 117 Formal Charges Also give each atom all electrons from its unshared pairs 6 7 4 6 1

18. © Prentice Hall 2001Chapter 118 Formal Charges Subtract the number of assigned electrons from the number of valence electrons for an uncombined atom of the same element 6 - 6 = 0 6 - 7 = -1 5 - 4 = +16 - 6 = 0 1 - 1 = 0

19. © Prentice Hall 2001Chapter 119 Formal Charges The algebraic sum of all formal charges on a species (molecule or ion) must equal the actual charge on the species Zero for molecules Positive for cations Negative for anions

20. © Prentice Hall 2001Chapter 120 Kekulé Structures In most cases we show electron pairs between atoms as bonds and represent them with a dash Also, unless there is a particular need to show unpaired electrons, we generally do not show them becomes

21. © Prentice Hall 2001Chapter 121 Condensed Structural Formulas Kekulé formulas also are called structural formulas Often, structural formulas are condensed becomes