1. READING: Chapter 9 sections 1 – 3 READING: Chapter 9 sections 1 – 3 HOMEWORK – DUE TUESDAY 11/10/15 HOMEWORK – DUE TUESDAY 11/10/15 HW-BW 9.1 (Bookwork) CH 9 #'s 5, 6, 10, 51, 53, 54, 57 – 67 (odd), 86, 88, 89 HW-BW 9.1 (Bookwork) CH 9 #'s 5, 6, 10, 51, 53, 54, 57 – 67 (odd), 86, 88, 89 HW-WS 16 (Worksheet) (from course website) HW-WS 16 (Worksheet) (from course website) HOMEWORK – DUE THURSDAY 11/12/15 HOMEWORK – DUE THURSDAY 11/12/15 HW-BW 9.2 (Bookwork) CH 9 #’s 27, 28, 29, 31, 33, 37, 38, 40, 47, 48, 49 HW-BW 9.2 (Bookwork) CH 9 #’s 27, 28, 29, 31, 33, 37, 38, 40, 47, 48, 49 HW-WS 17 (Worksheet) (from course website) HW-WS 17 (Worksheet) (from course website) Lab Lab Wednesday/Thursday – EXP 12 Wednesday/Thursday – EXP 12 No prelab No prelab Next Monday/Tuesday – EXP 13 Next Monday/Tuesday – EXP 13 Prelab Prelab

2. Types of Bonding : Ionic Compounds Ionic bonding involves the complete TRANSFER of electrons from one atom to another. Usually observed when a metal bonds to a nonmetal.

3. Types of Bonding : Ionic Compounds Ionic bonding involves the complete TRANSFER of electrons from one atom to another. Usually observed when a metal bonds to a nonmetal. Metals have low ionization energy, making it relatively easy to remove electrons from them Nonmetals have high electron affinities, making it advantageous to add electrons to these atoms The oppositely charged ions are then attracted to each other, resulting in an ionic bond

4. Ionic compounds tend to be hard, rigid, and brittle, with high melting points. Ionic compounds tend to be hard, rigid, and brittle, with high melting points. Types of Bonding: Ionic Compounds

5. Ionic compounds tend to be hard, rigid, and brittle, with high melting points. Ionic compounds tend to be hard, rigid, and brittle, with high melting points. Ionic compounds do not conduct electricity in the solid state. Ionic compounds do not conduct electricity in the solid state. In the solid state, the ions are fixed in place in the lattice and do not move. In the solid state, the ions are fixed in place in the lattice and do not move. Types of Bonding: Ionic Compounds

6. Ionic compounds tend to be hard, rigid, and brittle, with high melting points. Ionic compounds tend to be hard, rigid, and brittle, with high melting points. Ionic compounds do not conduct electricity in the solid state. Ionic compounds do not conduct electricity in the solid state. In the solid state, the ions are fixed in place in the lattice and do not move. In the solid state, the ions are fixed in place in the lattice and do not move. Ionic compounds conduct electricity when melted or dissolved. Ionic compounds conduct electricity when melted or dissolved. In the liquid state or in solution, the ions are free to move and carry a current. In the liquid state or in solution, the ions are free to move and carry a current. Types of Bonding: Ionic Compounds

7. Covalent bonding involves the SHARING of electrons Usually observed when a nonmetal bonds to a nonmetal. Types of Bonding: Covalent Compounds

8. Covalent bonding involves the SHARING of electrons Usually observed when a nonmetal bonds to a nonmetal. Nonmetal atoms have relatively high ionization energies, so it is difficult to remove electrons from them When nonmetals bond together, it is better in terms of potential energy for the atoms to share valence electrons Potential energy lowest when the electrons are between the nuclei, holding the atoms together by attracting nuclei of both atoms Types of Bonding: Covalent Compounds

9. Metallic bonding involves electron POOLING Occurs when a metal bonds to another metal. Types of Bonding: Metals The relatively low ionization energy of metals allows them to lose electrons easily Metallic bonding involves the metal atoms releasing their valence electrons to be shared as a pool by all the atoms/ions in the metal Organized metal cations islands in a sea of electrons Electrons delocalized throughout the metal structure Bonding results from attraction of cation for the delocalized electrons Explains many of the properties of metals

10. Types of Bonding: Metals

11. 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

12. Lewis Dot Symbols: Elements and Ions Lewis dot structures of elements Lewis dot structures of elements Use the symbol of element to represent nucleus and inner electrons Use the symbol of element to represent nucleus and inner electrons Use a dot to represent each valence electron in the atom Use a dot to represent each valence electron in the atom NaCaIn NOFH Sn Xe

13. Atoms bond because it results in a more stable electron configuration. Atoms bond because it results in a more stable electron configuration. more stable = lower potential energy more stable = lower potential energy Atoms bond together by either transferring or sharing electrons Atoms bond together by either transferring or sharing electrons Usually this results in all atoms obtaining an outer shell with eight electrons Usually this results in all atoms obtaining an outer shell with eight electrons octet rule octet rule there are some exceptions to this “rule”!! there are some exceptions to this “rule”!! Lewis Dot Symbols: Octet Rule

14. When atoms bond, they tend to gain, lose, or share electrons to result in eight valence electrons When atoms bond, they tend to gain, lose, or share electrons to result in eight valence electrons noble gas configuration - ns 2 np 6 noble gas configuration - ns 2 np 6 Many exceptions Many exceptions H, Li, Be, B attain an electron configuration like He H, Li, Be, B attain an electron configuration like He Helium = two valence electrons, a duet Helium = two valence electrons, a duet Lithium loses its one valence electron Lithium loses its one valence electron Hydrogen shares or gains one electron Hydrogen shares or gains one electron o commonly loses its one electron to become H + Beryllium loses two electrons to become Be 2+ Beryllium loses two electrons to become Be 2+ o commonly shares its 2 electrons in covalent bonds, resulting in 4 valence electrons Boron loses three electrons to become B 3+ Boron loses three electrons to become B 3+ o commonly shares its 3 electrons in covalent bonds, resulting in 6 valence electrons expanded octets for elements in Period 3 or below expanded octets for elements in Period 3 or below using empty valence d orbitals using empty valence d orbitals Basically, only C, N, O, F, and Ne MSUT follow the octet rule Basically, only C, N, O, F, and Ne MSUT follow the octet rule Lewis Dot Symbols: Octet Rule

15. Cations have Lewis symbols without valence electrons Cations have Lewis symbols without valence electrons lost in the cation formation lost in the cation formation Anions have Lewis symbols with eight valence electrons Anions have Lewis symbols with eight valence electrons electrons gained in the formation of the anion electrons gained in the formation of the anion Lewis Dot Symbols: Elements and Ions Li Li + F [ ] F –1

16. Orbital diagrams Lewis electron-dot symbols Electron configurations Li 1s 2 2s 1 + F 1s 2 2s 2 2p 5 Li↑↓ 1s1s2p2p ↑ 2s2s 1s1s2p2p 2s2s ↑F + 1s1s2p2p2s2s Li + ↑↓ 1s1s2p2p 2s2s ↑↓↑↓F-F- Li F Li + + F – Lewis Dot Symbols and Other Electron Stuff → Li + 1s 2 + F – 1s 2 2s 2 2p 6

17. Covalent bonding involves the SHARING of electrons Usually observed when a nonmetal bonds to a nonmetal. Nonmetal atoms have relatively high ionization energies, so it is difficult to remove electrons from them When nonmetals bond together, it is better in terms of potential energy for the atoms to share valence electrons Potential energy lowest when the electrons are between the nuclei, holding the atoms together by attracting nuclei of both atoms Types of Bonding: Covalent Compounds

18. Covalent Bonds

19. Atoms share electrons to achieve a full outer level of electrons. The shared electrons are called a shared pair or bonding pair. H HorH–HH–H The shared pair is represented as a pair of dots or a line: An outer-level electron pair that is not involved in bonding is called a lone pair, nonbonding pair, or unshared pair. F F F–FF–F or

20. Covalent Bonds The bond order is the number of electron pairs being shared by a given pair of atoms. A single bond consists of one bonding pair and has a bond order of 1 A double bond consists of two bonding pair and has a bond order of 2 A triple bond consists of three bonding pair and has a bond order of 3 The bond energy (BE) is the energy needed to overcome the attraction between the nuclei and the shared electrons. The stronger the bond the higher the bond energy. The bond length is the distance between the nuclei of the bonded atoms.

21. For a given pair of atoms, a higher bond order results in a shorter bond length and higher bond energy. Between any two atoms, more bonds = shorter bonds Between any two atoms, more bonds = larger bond energy Covalent Bonds

22. For a given pair of atoms, a higher bond order results in a shorter bond length and higher bond energy. Between any two atoms, more bonds = shorter bonds Between any two atoms, more bonds = larger bond energy Bond length increases down a group in the periodic table and decreases across the period. Bond energy shows the opposite trend. Covalent Bonds Internuclear distance (bond length) Covalent radius 133 pm Internuclear distance (bond length) Covalent radius 72 pm

23. Covalent Bonds:  H rxn The heat released or absorbed during a chemical change is due to differences between the bond energies of reactants and products.  º rxn =  Hº reactant bonds broken   Hº product bonds formed reactantsproducts always + always  Bond breaking is endothermic,  H (breaking) = + Bond making is exothermic,  H (making) = −

24. Covalent Bonds:  H rxn 4(C–H bond = 413 kJ) O=O bond = 498 kJ C–H bond = 413 kJ  º rxn =  Hº reactant bonds broken   Hº product bonds formed 2(O=O bond = 498 kJ)2(O=O bond = 498 kJ) = 996 kJ 4(C–H bond = 413 kJ) = 1652 kJ2(C=O bond = 799 kJ)C=O bond = 799 kJ2(C=O bond = 799 kJ) = 1598 kJ O–H bond = 467 kJ4(O–H bond = 467 kJ)4(O–H bond = 467 kJ) = 1868 kJ  º rxn = (1652 kJ + 996 kJ)  (1598 kJ + 1868 kJ)  º rxn =  818 kJ

25. Electronegativity and Polarity A covalent bond in which the shared electron pair is not shared equally, but remains closer to one atom than the other, is a polar covalent bond. If “X” and “Y” share bonding e - equally: If “X” and “Y” do NOT share bonding e - equally: Unequal sharing of bonding e - leads to polar covalent bonds

26. Electronegativity and Polarity A covalent bond in which the shared electron pair is not shared equally, but remains closer to one atom than the other, is a polar covalent bond. The ability of an atom in a covalent bond to attract the BONDING electrons towards itself is called its electronegativity.

27. Electronegativity and Polarity A covalent bond in which the shared electron pair is not shared equally, but remains closer to one atom than the other, is a polar covalent bond. Unequal sharing of electrons causes the more electronegative atom of the bond to be partially negative and the less electronegative atom to be partially positive. The ability of an atom in a covalent bond to attract the BONDING electrons towards itself is called its electronegativity.

28. Polar Covalent Bonds Polar: having poles NS One end opposite from the other

29. Polar Covalent Bonds Polar covalent bonds result from differences in electronegativity. HUH?!?!?! Electronegativity: The ability of an atom to pull BONDING electrons towards itself.

30. If EN difference is: EN < 0.5 the bond is considered to be pure covalent 2.5-2.5=0 2.8-2.8=02.5-2.1=0.4 2.0 > EN ≥ 0.5 the bond is considered to be polar covalent 3.5-2.5=14.0-2.1=1.93.0-2.5=0.53.5-1.8=1.7 EN ≥ 2.0 the bond is considered to be ionic Al:FCa:ONa:ClRb:N 1.5-4.0=2.51.0-3.5=2.50.9-3.0=2.10.8-3.0=2.2 Electronegativity and Polarity

31. The lowercase Greek letter delta, , is used to indicate a polar bond. The MORE EN element has extra e -, so it is negative and is indicated by the symbol  –. The LESS EN element is short of e -, so it is positive and is indicated by the symbol  +. H – Cl ++ –– 3.02.1 Electronegativity and Polarity

32. Give delta notation and polarity arrows for the following: ++ –– ++ –– ++ –– ++ –– 2.53.54.02.13.02.51.83.5 Electronegativity and Polarity

33. Lewis Structures 1) Determine the total number of valence electrons available in the chemical If ion, add 1 electron for each negative charge and subtract 1 electron for each positive 2) Draw the skeletal structure of the molecule using single bonds to connect the atoms Central atom(s) will be surrounded by other atoms Central atom(s) tend to be the element that is the least electronegative H and F always exterior atoms 3) Fill the octets for all atoms except hydrogen (2), beryllium (4) and, boron (6) Count total electrons drawn and subtract this from the number of valence electrons available found in step #1. If you have not drawn enough electrons, add the missing ones to the central atom If you have drawn too many electrons, remove lone pair(s) and add multiple bonds # of valence e- needed = # of bonds formed (guideline only)