The document discusses various concepts of chemical bonding, emphasizing the role of valence electrons in ionic and covalent bonds. It explains Lewis symbols, structures, formal charges, oxidation numbers, and the importance of electron configurations for stability. Additionally, it covers different types of bonds, bond lengths, and exceptions to the octet rule.

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

3. In the period 1916-1919, two Americans, G.N. Lewis and Irving Langmuir, and a German, Walther Kossel, advanced an important proposal about chemical bonding: Something unique in the electron configurations of noble gas atoms accounts for their inertness, and atoms of other elements combine with one another to acquire electron configurations like noble gas atoms. Some fundamental ideas of Lewis theory Electrons, especially those of the outermost (valence) electronic shell, play a fundamental role in chemical bonding. In some cased electrons are transferred from one atom to another. Positive and negative ions are formed and attract each other through electrostatic forces called ionic bonds . In other cases one or more pairs of electrons are shared between atoms, this sharing of electrons is called covalent bond . Electrons are transferred, or shared, in such a way that each atom acquires an especially stable electron configuration. Usually this is a noble gas configuration, one with eight outer shell electrons, or an octet .

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 .

18. Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. Count total number of valence e - . Add 1 for each negative charge. Subtract 1 for each positive charge. Complete an octet for all atoms except hydrogen If structure contains too many electrons, form double and triple bonds on central atom as needed. Writing Lewis Structures 9.6

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

20. Step 1 – C is less electronegative than O, put C in center Step 2 – Count valence electrons C - 4 (2s 2 2p 2 ) and O - 6 (2s 2 2p 4 ) -2 charge – 2e - 4 + (3 x 6) + 2 = 24 valence electrons Step 3 – Draw single bonds between C and O atoms and complete octet on C and O 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 Step 5 - Too many electrons, form double bond and re-check # of e - Write the Lewis structure of the carbonate ion (CO 3 2- ). O C O O 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 Total = 24

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 -

24. Formal Charge and Lewis Structures 9.7 For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. Lewis structures with large formal charges are less plausible than those with small formal charges. Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms. Which is the most likely Lewis structure for CH 2 O? H C O H -1 +1 H C O H 0 0

25. Oxidation number The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. Free elements (uncombined state) have an oxidation number of zero. Na, Be, K, Pb, H 2 , O 2 , P 4 = 0 In monatomic ions, the oxidation number is equal to the charge on the ion. Li + , Li = +1 ; Fe 3+ , Fe = +3 ; O 2- , O = -2 The oxidation number of oxygen is usually –2 . In H 2 O 2 and O 2 2- it is –1 . 4.4

26. The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1 . 6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion. Group IA metals are +1 , IIA metals are +2 and fluorine is always –1 . HCO 3 - O = -2 H = +1 3x( -2) + 1 + ? = -1 C = +4 4.4 Oxidation numbers of all the elements in HCO 3 - ?

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 )