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Chemical Bonding I: Basic Concepts

Chemical Bonding I: Basic Concepts. 6.1 The Octet Rule Lewis Structures Multiple Bonds 6.2 Electronegativity and Polarity Electronegativity Dipole Moment, Partial Charges and Percent Ionic Character 6.3 Drawing Lewis Structures 6.4 Lewis Structures and Formal Charge

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Chemical Bonding I: Basic Concepts

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  1. Chemical Bonding I: Basic Concepts 6.1 The Octet Rule Lewis Structures Multiple Bonds 6.2 Electronegativity and Polarity Electronegativity Dipole Moment, Partial Charges and Percent Ionic Character 6.3 Drawing Lewis Structures 6.4 Lewis Structures and Formal Charge 6.5 Resonance 6.6 Exceptions to the Octet Rule Incomplete Octets Odd Numbers of Electrons Expanded Octets 6

  2. 3- N F Valence Electrons – e- in last period (or shell) Lewis dot symbols of the main group elements represent the valence electrons in the atom. N 1s2 2s2 2p3 B 1s22s2 2p1 C 1s2 2s2 2p2 N 1s2 2s2 2p3 O 1s2 2s2 2p4 F 1s2 2s2 2p5 Ne 1s2 2s2 2p6 Octet rule: s2 p6 is very stable Explains rules for determining the formulas of ionic cpds

  3. Lewis Dot Symbols Lewis dot symbols of the main group elements represent the valence electrons in the atom. •• • • N • • 1s22s22p1 1s22s22p2 1s22s22p3 •• • • • • • • • • 5 valence electrons; first pair formed in the Lewis dot symbol B C O •• For nonmetals in the second period, the number of unpaired dots is the number of bonds the atom can form.

  4. Lewis dot structures and covalent bond tendencies Lone pair Radical (not stable) only needs 1 e- to achieve He structure H C N O F Ne Covalent Bonds Covalent Bond A pair of electrons ‘shared’ between two atoms This is an alternate method (to ion formation) for atoms to achieve a ‘stable octet’. 4 3 2 1 0 = # electrons needed to share to form ‘octet’

  5. Covalent Bond Elements that exist as diatomic gaseous molecules include …. H2 O2 N2 F2 Cl2 etc. Represent each diatomic molecule as a Lewis Dot structure. Each element should share electron pairs to obtain octet. (H only 2 e-.) O O O = O

  6. H F F H HF should be considered an … a) ionic b) covalent ….. Compound? Electronegativity & Bond polarity DEN = (4.0 – 2.1) = 1.9 electron rich region electron poor region e- poor e- rich d+ d-

  7. Electronegativity There is no sharp distinction between nonpolar covalent and polar covalent or between polar covalent and ionic. The following rules help distinguish among them: • A bond between atoms whose electronegativites differ by less than 0.5 is general considered purely covalent or nonpolar. • A bond between atoms who’s electronegativies differ by the range of 0.5 to 2.0 is generally considered polar covalent. • A bond between atoms whose electronegativities differ by 2.0 or more is generally considered ionic.

  8. Bond Type DEN examples Ionic ≥ 2.0 Na-Cl, Mg-O Polar covalent 0.5 – 1.9 N←H, C→O, O ← H, Co → O Nonpolar≤ 0.4 C — H, Si — H, Br — Cl Electronegativityof Elements What type of bond is found in Ru2O3? a. ionic b. polar c. nonpolar What type of bond is found in CS2? a. ionic b. polar c. nonpolar Electronegativity & Bond polarity

  9. Electronegativity and Polarity Electron density maps show the distributions of charge. • Electrons spend a lot of time in red and very little time in blue. Electrons are shared equally nonpolar covalent Electrons are not shared equally and are more likely to be associated with F polar covalent Electrons are not shared but rather transferred from Na to F ionic

  10. 4 + 3 + 7 = 14 Valence e- EN H 1s1 2.1 C 2s2 2p2 2.5 N 2s2 2p3 3.0 O 2s2 2p4 3.5 Cl 3s2 3p5 3.0 Rules for Lewis Dot structures: CH3Cl 1. Count the total # of valence e- 2. Arrange atoms into possible structure (trial & error). Atom with lowest EN placed in center. H and F will always occupy terminal positions. 3. Place single bond between central atom & neighbors. complete the octets of ‘neighbors’ (show all lone pairs). Add extra electrons as lone pairs on central atom. Check that total # e-is correct. 4. Complete central atom octet with = or  bonds using lone pairs from neighbors. H2O How many lone pairs are on N? a. 0 b. 1 c. 2 d. 3 CH4 NH3

  11. 4 + 3 + 7 = 14 Valence e- EN H 1s1 2.1 C 2s2 2p2 2.5 N 2s2 2p3 3.0 O 2s2 2p4 3.5 Cl 3s2 3p5 3.0 Rules for structures: CH3Cl 1. Count the total # of valence e- 2. Arrange atoms into possible structure (trial & error). Atom with lowest EN placed in center. H and F will always occupy terminal positions. 3. Place single bond between central atom & neighbors. complete the octets of ‘neighbors’ (show all lone pairs). Add extra electrons as lone pairs on central atom. Check that total # e-is correct. 4. Complete central atom octet with = or  bonds using lone pairs from neighbors. CO2 Draw the Lewis dot structure for CO CN- HCN

  12. Friday’s Quiz % composition – FW – empirical formulas – molecular formulas Bond polarity and ionic, polar covalent, and nonpolar covalent bonds Lewis dot structures A compound contains: 13.3% H 60.0% C 26.7% O What is its empirical formula? 13.3g H • 1 mol/1.0 g = 13.3 mol ÷ 1.67 = 8.0 60.0g C • 1 mol/12.0 g = 5 mol ÷ 1.67 = 3.0 26.7g O • 1 mol/16.0 g = 1.67 mol÷ 1.67 = 1.0 C3H8O

  13. Lewis Dot Structures Homonuclear diatomic molecules H - H Rules for structures: H – O – H 1. Count the total # of valence e- 2. Arrange atoms into possible structure (trial & error). Atom with lowest EN placed in center. H and F will always occupy terminal positions. H – N – H | H 3. Place single bond between central atom & neighbors. complete the octets of ‘neighbors’ (show all lone pairs). Add extra electrons as lone pairs on central atom. Check that total # e-is correct. 4. Complete central atom octet with = or  bonds using lone pairs from neighbors. Draw the Lewis dot structure for CO C ≡ O N ≡ N F - F O = O Neither C nor O has the typical # of bonds – This is partly why CO is not stable

  14. Bond length is defined as the distance between the nuclei of two covalently bonded atoms. Which has the greater bond length ….. a) an N – N bond b) an N ≡ N bond N−N N=N N≡N 147 pm 124 pm 110 pm Multiple bonds are shorter than single bonds.

  15. Bond Energy We quantify bond strength by measuring the quantity of energy required to break it. H2(g)+ 436.4 kJ/mol → H(g) + H(g) Bond energy = 436.4 kJ/mol Which has the greater bond energy? a) an N – N bond b) an N ≡ N bond Bond energy kJ/mol Bond length pm Multiple bonds are shorter and stronger

  16. Bond Type DEN examples Ionic ≥ 2.0 Na-Cl, Mg-O Polar covalent 0.5 – 1.9 N←H, C→O, Nonpolar ≤ 0.4 C — H, Si — H Given element electronegativity determine the bond type Electronegativity & Bond polarity

  17. Electronegativity and Polarity Electron density maps show the distributions of charge. • Electrons spend a lot of time in red and very little time in blue. Electrons are shared equally nonpolar covalent Electrons are not shared equally and are more likely to be associated with F polar covalent Electrons are not shared but rather transferred from Na to F ionic

  18. H F e- poor e- rich Electronegativity & Bond polarity +0.41 -0.41 d+ d- A quantitative measure of the polarity of a bond is its dipole moment (μ). • Q is the charge. • ris the distance between the charges. • μis always positive and expressed in debye units (D). 1 D = 3.336×10−30C∙m F H− •• •• •• μ = Q x r

  19. 4 + 3 + 7 = 14 Valence e- EN H 1s1 2.1 C 2s2 2p2 2.5 N 2s2 2p3 3.0 O 2s2 2p4 3.5 Cl 3s2 3p5 3.0 Rules for structures: CH3Cl 18 1. Count the total # of valence e- 2. Arrange atoms into possible structure (trial & error). Atom with lowest EN placed in center. H and F will always occupy terminal positions. O OO .. .. .. :O – O – O: ̈̈ ̈ 3. Place single bond between central atom & neighbors. complete the octets of ‘neighbors’ (show all lone pairs). Add extra electrons as lone pairs on central atom. Check that total # e-is correct. .. .. .. :O = O – O: or …. ̈ 4. Complete central atom octet with = or  bonds using lone pairs from neighbors. .. .. .. :O – O = O: ̈ Draw the Lewis dot structure for Ozone, O3. Resonance – A term used to describe the electronic state for molecules in which multiple, identical Lewis dot structures can be drawn. Indicates delocalization of electrons. In reality, each bond is identical and has partial double bond character. This adds stability to the molecule.

  20. Draw the Lewis dot structure for CO2 Formal Charges (FC) A means to determine the best Lewis dot structure when more than one structure meets requirements. Not the same as resonance, where all structures are equivalent FC = # valence e-– (#bonds + unshared e-) CO2 0 0 0 -1 0 +1 .. .. .. :O = C = O: vs. :O – C ≡ O: ̈ • Best formulas have lowest FC • (-) FC more likely on more electronegative atom • Adjacent atoms should not have same sign FC

  21. Rules for structures: Count e- → arrange atoms → single bonds → lone pairs → central atom CO32- How many e-? a. 30 b. 32 c. 24 d. 22 Best frame? A B C O C O O C O O O O C O O Does it matter which O atom gets the C = O? a. yes b. no Resonance – A term used to describe the electronic state for molecules in which multiple, identical Lewis dot structures can be drawn. Indicates delocalization of electrons.

  22.   BeH2 & BF3 Incomplete octets expanded octet odd # of e- PF5, SF6, SF4 nonmetals ≥ period 3 NO & NO2 Octet Exceptions

  23. Octet Exceptions Incomplete octets Common only to Be and B chemistry . ∙Be∙∙B∙ Be 2s2 B 2s2 2p1 BeH2BF3 .. :F: .. | .. H – Be – H :F – B – F: ̈̈̈ ̈̈̈

  24. Octet Exceptions odd # of e- 5 + 6 = 11 5 + 12 = 17 NO NO2 Nitric oxide or nitrogen dioxide Nitrogen oxide . .. :N = O: .. . .. :O - N = O: ̈ has resonance Molecules with radicals (unpaired electrons) are very reactive.

  25. Octet Exceptions expanded octet The central atom must be a nonmetals ≥ period 3 PF5SF6SF4 In expanded octet molecules the d orbitals become available for extra electrons. s2 p6 d2 or s2p6d4 We will introduce the concept of hybridization of orbitals next chapter.

  26. Rules for structures: Count e- → arrange atoms → single bonds → lone pairs → central atom FC = # valence e-– (#bonds + unshared e-) • Best formulas have lowest FC • (-) FC more likely on more electronegative atom • Adjacent atoms should not have same sign FC [SO4]2- H2SO4 When an expanded octet is not required, but can lead to a better formal charge either structure is acceptable.

  27. Chapter Summary: Key Points 6

  28. Drawing Lewis Structures Follow these steps when drawing Lewis structure for molecules and polyatomic ions. Draw the skeletal structure of the compound. The least electronegative atom is usually the central atom. Draw a single covalent bond between the central atom and each of the surrounding atoms. Count the total number of valence electrons present; add electrons for negative charges and subtract electrons for positive charges. For each bond in the skeletal structure, subtract two electrons from the total valence electrons. Use the remaining electrons to complete octets of the terminal atoms by placing pairs of electrons on each atom. Complete the octets of the most electronegative atom first. Place any remaining electrons in pairs on the central atom. If the central atom has fewer than eight electrons, move one or more pairs from the terminal atoms to form multiple bonds between the central atom and terminal atoms. 6.3

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