For Chem 1: Significanceof the ELectron in Bonding The Octet Rule Lewis Symbol/Structures Formal Charge Polyatomic Ions Types of Bonds (Ionic, Covalent, Coordinate Covalent, Metallic Bonds, Multiple Bonds) Exceptions to the Octet Rules Oxidation Number is not included in the class discussion and exam. ;D

1. Chemical Bonding I: Basic Concepts

2. 9.1 Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons that particpate in chemical bonding. 1A 1 ns 1 2A 2 ns 2 3A 3 ns 2 np 1 4A 4 ns 2 np 2 5A 5 ns 2 np 3 6A 6 ns 2 np 4 7A 7 ns 2 np 5 Group # of valence e - e - configuration

4. Lewis symbol – consists of a chemical symbol to represent the nucleus and core (inner shell) electrons of an atom, together with dots placed around the symbol to represent the valence (outer shell) electrons. Write the Lewis symbol of the following elements 1. Si 2. N 3. P 4. As 5. Sb 6. Bi 7. Al 8. I 9. Se 10. Ar Write the Lewis symbol of the following 1. Sn 2. Br – 3. Na + 4. S 2-

5. Lewis structure – is a combination of Lewis symbols that represents either the transfer or sharing of electrons in a chemical bond. Write Lewis structures for the following compounds (a) BaO (b) MgCl 2 (c) Al 2 O 3 (d) Na 2 S (e) Mg 3 N 2 (f) calcium iodide (g)barium sulfide (h) lithium oxide

6. 9.1

7. 9.2 The Ionic Bond 1s 2 2s 1 1s 2 2s 2 2p 5 1s 2 1s 2 2s 2 2p 6 [He] [Ne] Li + F Li + F - Li Li + + e - e - + F F - F - Li + + Li + F -

8. A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Lewis structure of F 2 9.4 Why should two atoms share electrons? F F + 7e - 7e - F F 8e - 8e - F F F F lone pairs lone pairs lone pairs lone pairs single covalent bond single covalent bond

9. + + Lewis structure of water Double bond – two atoms share two pairs of electrons or Triple bond – two atoms share three pairs of electrons or 9.4 8e - H H O O H H O H H or 2e - 2e - single covalent bonds O C O O C O 8e - 8e - 8e - double bonds double bonds N N 8e - 8e - N N triple bond triple bond

10. Lengths of Covalent Bonds Bond Lengths Triple bond < Double Bond < Single Bond 9.4 Bond Type Bond Length (pm) C - C 154 C  C 133 C  C 120 C - N 143 C  N 138 C  N 116

11. 9.4

12. Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms electron rich region electron poor region e - rich e - poor  +  - 9.5 H F F H

13. Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. Electron Affinity - measurable , Cl is highest Electronegativity - relative , F is highest 9.5 X ( g) + e - X - ( g)

14. 9.5

15. 9.5

16. Classification of bonds by difference in electronegativity Difference Bond Type 0 Covalent  2 Ionic 0 < and <2 Polar Covalent 9.5 Covalent share e - Polar Covalent partial transfer of e - Ionic transfer e - Increasing difference in electronegativity

17. Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent N – 3.0 N – 3.0 3.0 – 3.0 = 0 Covalent 9.5 Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H 2 S; and the NN bond in H 2 NNH 2 .

19. Step 1 – N is less electronegative than F, put N in center Step 2 – Count valence electrons N - 5 (2s 2 2p 3 ) and F - 7 (2s 2 2p 5 ) 5 + (3 x 7) = 26 valence electrons Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms. Step 4 - Check, are # of e - in structure equal to number of valence e - ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons 9.6 Write the Lewis structure of nitrogen trifluoride (NF 3 ). F N F F

21. 9.7 Two possible skeletal structures of formaldehyde (CH 2 O) An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion. H C O H H C O H formal charge on an atom in a Lewis structure = 1 2 total number of bonding electrons ( ) total number of valence electrons in the free atom - total number of nonbonding electrons -

22. formal charge on C = 4 - 2 - ½ x 6 = -1 formal charge on O = 6 - 2 - ½ x 6 = +1 -1 +1 9.7 H C O H C – 4 e - O – 6 e - 2H – 2x1 e - 12 e - 2 single bonds (2x2) = 4 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 formal charge on an atom in a Lewis structure = 1 2 total number of bonding electrons ( ) total number of valence electrons in the free atom - total number of nonbonding electrons -

23. formal charge on C = 4 - 0 - ½ x 8 = 0 formal charge on O = 6 - 4 - ½ x 4 = 0 0 0 9.7 C – 4 e - O – 6 e - 2H – 2x1 e - 12 e - 2 single bonds (2x2) = 4 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 H C O H formal charge on an atom in a Lewis structure = 1 2 total number of bonding electrons ( ) total number of valence electrons in the free atom - total number of nonbonding electrons -

27. Figure 4.10 The oxidation numbers of elements in their compounds 4.4

28. NaIO 3 Na = +1 O = -2 3x( -2 ) + 1 + ? = 0 I = +5 IF 7 F = -1 7x( -1 ) + ? = 0 I = +7 K 2 Cr 2 O 7 O = -2 K = +1 7x( -2 ) + 2x( +1 ) + 2x( ?) = 0 Cr = +6 4.4 Oxidation numbers of all the elements in the following ?

29. Polyatomic Ions – consists of two or more atoms, and the forces holding atoms together within such ions are covalent bonds Coordinate covalent bond -a covalent bond in which one atom contributes both electrons -once such bond is formed we cannot tell it from regular covalent bond Example: HCl + NH 3 forms NH 4 Cl

30. Metallic Bonds -Electron Sea Model -pictures a solid metal as a network of positive ions immersed in “sea of electrons” -electrons in the sea are free (not attached to any particular ion) and they are mobile.

31. A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. 9.8 O O O + - O O O + - O C O O - - O C O O - - O C O O - - What are the resonance structures of the carbonate (CO 3 2 -) ion?

32. Exceptions to the Octet Rule The Incomplete Octet BeH 2 BF 3 9.9 H H Be Be – 2e - 2H – 2x1e - 4e - B – 3e - 3F – 3x7e - 24e - F B F F 3 single bonds (3x2) = 6 9 lone pairs (9x2) = 18 Total = 24

33. Exceptions to the Octet Rule Odd-Electron Molecules NO The Expanded Octet (central atom with principal quantum number n > 2) SF 6 9.9 N – 5e - O – 6e - 11e - N O S – 6e - 6F – 42e - 48e - S F F F F F F 6 single bonds (6x2) = 12 18 lone pairs (18x2) = 36 Total = 48

34. The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy . Bond Energy 9.10 H 2 ( g ) H ( g ) + H ( g )  H 0 = 436.4 kJ Cl 2 ( g ) Cl ( g ) + Cl ( g )  H 0 = 242.7 kJ HCl ( g ) H ( g ) + Cl ( g )  H 0 = 431.9 kJ O 2 ( g ) O ( g ) + O ( g )  H 0 = 498.7 kJ O O N 2 ( g ) N ( g ) + N ( g )  H 0 = 941.4 kJ N N Bond Energies Single bond < Double bond < Triple bond

35. Average bond energy in polyatomic molecules 9.10 H 2 O ( g ) H ( g ) + OH ( g )  H 0 = 502 kJ OH ( g ) H ( g ) + O ( g )  H 0 = 427 kJ Average OH bond energy = 502 + 427 2 = 464 kJ

36. 9.10 H 2 ( g ) + Cl 2 ( g ) 2HCl ( g ) 2H 2 ( g ) + O 2 ( g ) 2H 2 O ( g )