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

7. 9.1

8. 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 -

9. 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

10. + + 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

11. 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

12. 9.4

13. 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

14. 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)

15. 9.5

16. 9.5

17. 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

18. 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 .

20. 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

22. 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

23. 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 -

24. 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 -

25. 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 -

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

30. 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 ?

31. 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

32. 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.

33. 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?

34. 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

35. 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

36. 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

37. 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

38. 9.10 H 2 ( g ) + Cl 2 ( g ) 2HCl ( g ) 2H 2 ( g ) + O 2 ( g ) 2H 2 O ( g )