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CHAPTER 4 : CHEMICAL BONDING 4.1 LEWIS STRUCTURE

CHAPTER 4 : CHEMICAL BONDING 4.1 LEWIS STRUCTURE. LECTURE 1. Learning outcomes:. (a) Write the Lewis symbol for an atom. (b) State the octet rule and describe how atoms obtain the octet configuration. Group. e - configuration. # of valence e -. ns 1. 1. 1. 2.

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CHAPTER 4 : CHEMICAL BONDING 4.1 LEWIS STRUCTURE

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  1. CHAPTER 4 : CHEMICAL BONDING 4.1 LEWIS STRUCTURE

  2. LECTURE 1 Learning outcomes: (a) Write the Lewis symbol for an atom. (b) State the octet rule and describe how atoms obtain the octet configuration.

  3. Group e- configuration # of valence e- ns1 1 1 2 ns2 2 13 ns2np1 3 14 ns2np2 4 15 ns2np3 5 16 ns2np4 6 17 ns2np5 7 Valence electrons are the outer shell electrons of an atom.The valence electrons are the electrons that participate in chemical bonding.

  4. Lewis symbol • Consist of the symbol of an element • One dot or one cross for each valence electron in an atom of the element • Element in the same group have similar outer electron configuration • Similar Lewis symbol • Example ; • Li and X Na

  5. Lewis Symbols

  6. OCTET RULE • Atoms combine in order to achieve a more stable electronic configuration. • Maximum stability results when an atom is isoelectronic with a noble gas. • - Atom can achieve noble gas configuration through : • ( a) transfer of electron (gaining or losing) • ( b) sharing of electron

  7. THE ELECTRON CONFIGURATION OF IONS Ions have stable electronic configuration of either: • The noble gas configuration • Pseudonoble gas configuration • Half-filled orbitals • The inert pair effect

  8. (a) Noble gas configuration Group 1, 2 and 13 elements transfer valence e- to form cation with noble gas configuration. Example Na : 1s22s22p63s1 Na+ : 1s22s22p6( isoelectronic with neon ) Mg : 1s22s22p63s2 Mg2+: 1s22s22p6 ( isoelectronic with neon )

  9. Group 15, 16 and 17elements accept e to form anion with noble gas configuration. Example O : 1s22s22p4 O2- : 1s22s22p6 ( isoelectronic with neon ) F : 1s22s22p5 F- : 1s22s22p6 ( isoelectronic with neon )

  10. (b)Pseudo noble gas configuration • For d block elements, e from 4s orbital will be transferred first before the 3d electrons. • d block elements donate electrons to achieve pseudonable gas configuration • Example : • Zn : 1s22s22p63s23p64s23d10 • Zn2+ : 1s22s22p63s23p63d10 • ( pseudonoble gas configuration )

  11. Pseudo means false • So, pseudonoble gas means that the atom has similar electron configuration with noble gases • The valence electron configuration needed is either ns2np6nd10or ns2np6nd10nf14

  12. (c) Half-filled orbital • d block element can also donate or receive electron to achieve half-filled orbitals which is stable • Example : • Mn : 1s22s22p63s23p64s23d5 • Mn2+ : 1s22s22p63s23p63d5 • ( stability of the half-filled orbital) • Fe : 1s22s22p63s23p64s23d6 • Fe3+ : 1s22s22p63s23p63d5 • ( stability of the half-filled orbital)

  13. (d) The inert pair effect - elements Group 13 and 14 in Period 5 or 6form cations with configuration of, ns2np6nd10(n+1)s2 - This is named the inert pair effect. Example : Sn : [ Kr ] 4d105s25p2 Sn2+ : [ Kr ] 4d105s2 ( the inert pair effect )

  14. LECTURE 2 Learning outcomes: (c) Describe the formation of the following bonds using Lewis symbol. i. Ionic or electrovalent bond ii. Covalent bond iii. Dative or coordinate bond (d) Draw Lewis structure of covalent species with single, double and triple bonds. (e) Compare the bond length between single, double and triple bonds.

  15. Ionic or electrovalent bonding • Electrostatic attraction force between positiveion (cation) with negative ion (anion) • eg . Formation of NaCl . Na – e- Na+ (metal) (cation) Cl + e- Cl- (non-metal)(anion) Na+Cl- Ionic bond

  16. The formation of electrovalent bonds, • 1. An ionic bond is formed by the electrostatic forces. • 2. Ionic bond formed between 2 ions with different • charge and through electron transfer. • 3. Metal elements will donate electron while nonmetal • elements receive electron to achieve stability. • 4. This happen because metals are more electropositive • while non metals more electronegative.

  17. Na Na+ + e- 1s22s22p63s1 1s22s22p6 (obey the octet rule) Cl + e- Cl- 1s22s22p63s23p5 1s22s22p63s23p6 (obey the octet rule) Example :

  18. xx xx Na• Cl Na Cl x xx • x xx xx xx Formation can be described by Lewis structure – valence e- represented as dot or cross. - + (formation of electrovalent bond ) NOTE: Remember to show arrow from donater to accepter!!!

  19. Properties of electrovalent compound • Solid at room temperature • High melting/ boiling point • Soluble in water • Molten ionic compounds conduct electricity because they contain mobile ions ( cations & anions)

  20. Excercise (ionic bond): • Draw Lewis structure for the formation of following • ionic compounds. • a) KF. .. • K. +  F  [ K ]+ [  F  ]- • .. .. • -When K and F atoms come in contact with each other, • the outer 2s1 valence electron of K is transferred • to F atom.

  21. b) BaO . .. Ba +  O  [ Ba ]2+ [  O  ]2- . .. When Ba and O atoms come in contact with each other, 2 e of 2s orbital of Ba are transferred to O. c) Na2O Na . . . . + O 2[ Na ]+ [  O  ]2- . . . Na .

  22. Covalent Bond Formed by sharing 1 or more pairs of valence electrons between nonmetal atoms (group 14 , 15 , 16 , 17 ,18)

  23. Why should two atoms share electrons? + 7e- 7e- 8e- 8e- F F F F F F F F lonepairs lonepairs single covalent bond single covalent bond lonepairs lonepairs Lewis structure of F2

  24. .. + :O: + H+ O H H H H H Coordinate covalent or dative bond Defination:A bond in which the pair of electrons is supplied by one of the two bonded atoms Eg: hydroxonium ion, H3O+

  25. Eg : F3BNH3 molecule BH3 + NH3 H3BNH3 F F .. + N B B H F F F F H H N H H H

  26. Single Bond • A covalent bond formed when 2 atoms share a pair of electrons • Represent by dash (-) between 2 atom • A single bond is made up of a sigma (σ) bond • Example: HCl and HF

  27. Double Bond • A covalent bond formed when 2 atoms share 2 pairs of electrons • Represent by double dash (=) between 2 atoms • A double bond is made up of sigma bond (σ) and pi bonds (π) • Example: O2

  28. Triple bond • A covalent bond formed when 2 atoms share 3 pairs of electrons • Represent by triple dash (Ξ) between 2 atoms • A triple bond is made up of 1 sigma bond (σ) and 2 pi bonds (π) • Example : N2

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

  30. Comparison of the bond length between single, double and triple bonds. • Multiple bonds are shorter than single covalent bonds. • Bond length is defined as the distance between the nuclei of two covalently bonded atoms in a molecule. Refer figure given.

  31. For a given pair of atoms such as carbon and nitrogen, triple bonds are shorter than double bond, which, in turn are shorter than single bond. Refer table given.

  32. 74pm 161pm H2 HI Figure : Bond length in H2 and HI

  33. Table : Comparison of the bond length

  34. LECTURE 3 Learning outcomes: f) Determine the formal charge and the most plausible Lewis structure. g) Explain the exception to the octet rule: incomplete octet, expanded octet and odd number electrons. h) Explain the concept of resonance using appropriate examples.

  35. .. .. .. .. C O Cl: formal charge C = 4-[4+½(4)] = -2 :Cl: formal charge O = 6-[0+½(8)] = 2 formal charge Cl = 7-[6+½(2)] = 0 Formal charge Cl = 7-[6+½(2)] = 0 .. structure A formal charge =valence e – (no. lone pair e + ½ no. bond-pair e)

  36. :O: C Cl: formal charge C = 4-[0+½(8)] = 0 :Cl: formal charge O = 6-[6+½(2)] = -1 formal charge Cl = 7-[4+½(4)] = +1 .. .. .. Structure B formal charge Cl = 7-[6+½(2)] = 0 formal charge =valence e – (no. lone pair e + ½ no. bond-pair e)

  37. eg : :O C Cl: formal charge C = 4-[ 0+½(8)] = 0 :Cl: formal charge O = 6-[4+½(4)] = 0 formal charge Cl = 7-[6+½(2)] = 0 .. .. .. Structure C formal charge Cl = 7-[6+½(2)] = 0 formal charge =valence e – (no. lone pair e + ½ no. bond-pair e) ..

  38. Conclusion … Refer to the obtained formal charge, the most plausible structure is… Figure C

  39. Exercise 1: Calculate the formal charge for each atoms of NOCl molecule Answer

  40. Exercise 2: Draw the right Lewis structure of CO2. Answer .

  41. Exception to the octet rule • In octet configuration , atom should have e configuration of noble gas • But…there’s an exception * incomplete octet * expanded octet *odd electron molecule

  42. Incomplete octet • Elements in groups 2 & 13 • Period 2 • Less metallic character • Do not donate e but share e • Central atom have less than 8 e (not achieve octet configuration)

  43. . : : : : .. .. .. .. Cl Cl Be Cl Cl Be + + Eg : BeCl2 .. .. .. .. . . . Be shared e with Cl (covalent bond) but in BeCl2 molecule Be only have 4 e in the outer shell (less than 8 e, does not achieve octet configuration)

  44. Expanded Octet • Involves period 3 and onwards (non metals) • Has d orbital that involves in bonding • Central atoms having 10 or even 12 valence e.

  45. F F F S F F F S shared e with F (covalent bond) but in SF6 molecule S have 12 e in the outer shell (more than 8 e, does not obeys octet rule) eg : SF6

  46. PF5 F F P F F F P shared e with F (covalent bond) but in PF5 molecule P have 10 e in the outer shell (more than 8 e, does not obeys octet rule)

  47. •• •• •• • •• + N=O O=N-O - •• •• •• •• • Odd electron molecule • Molecule with odd no. of e, complete pairing is impossible and an octet around each atom cannot be achieved • Mostly from an odd numbered group ( Nitrogen- Group 15, Chlorine-Group 17) • Example : nitric oxide (NO) and nitrogen dioxide (NO2) : • Since we need even no. of electrons for complete pairing (to reach 8) • The octet rule cannot be satisfied for all the atoms in any of these molecule

  48. : .. .. .. .. .. .. : .. .. .. .. S O S O O O RESONANCE Same atomic structure but different arrangement of e- Resonance structure : 2 or more Lewis structure for single molecule that cannot be represented with 1 accurate Lewis structureEg: Sulphur dioxide, SO2

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