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Covalent Bonds & Molecular Forces

Covalent Bonds & Molecular Forces. Ch.6. (6-1) Covalent Bond. e - are shared b/w 2 atoms Single bond : 1 shared pair Double bond : 2 shared pairs Triple bond : 3 shared pairs http://facweb.eths.k12.il.us/weinerj/PPT_Presentations/covalent_bonding.ppt. Molecular Orbital.

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Covalent Bonds & Molecular Forces

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  1. Covalent Bonds & Molecular Forces Ch.6

  2. (6-1) Covalent Bond • e- are shared b/w 2 atoms • Single bond: 1 shared pair • Double bond: 2 shared pairs • Triple bond: 3 shared pairs • http://facweb.eths.k12.il.us/weinerj/PPT_Presentations/covalent_bonding.ppt

  3. Molecular Orbital • Region where an e- pair is most likely to exist • Formed by overlapping atomic orbitals

  4. Bond Length • Avg. dist. b/w 2 bonded atoms • Occur at min. PE

  5. Bond E • E required to break a bond b/w 2 atoms & separate them • Stronger bonds are shorter • Single = long = weak • Triple = short = strong

  6. Electronegativity • Tendency of an atom to attract bonding e- to itself • Inc. across a period, dec. down a group

  7. Electron Density • The more EN atom, has a higher electron density than the less EN atom • Pulls more e- to it

  8. Bonding • Nonpolar covalent: bonding e- shared equally • EN difference 0 to 0.5 • Polar covalent: bonding e- are localized on the more EN atom • EN dif. 0.6 to 2.1 • Ionic: e- transferred, not shared • EN dif. larger than 2.1 • http://facweb.eths.k12.il.us/weinerj/PPT_Presentations/Bonding_part_III_polar.ppt

  9. Dipole • Molecule in which 1 end has a partial + charge & the other end has a partial - charge

  10. Dipole Moment (EN dif.) • Determines polarity of a bond & molecule • Larger d.m.  higher polarity  stronger bond

  11. 1 0 1 2 3 4 3 2 1 0 1 2 3 4 3 2 1 0 1 2 transition metals 3 4 3 2 1 0 (6-2) Valence Electrons • e- in the outer-most E level of an atom, where it can participate in bonding

  12. Lewis Structure • Lewis structure: represents the valence e- in a molecule

  13. Lewis Dot Structure • Place 1 e- on each side of atom before pairing any e-

  14. Unshared Pair • (Lone pair): pair of valence e- not involved in bonding

  15. Rules for Drawing Lewis Structures • H & halogens bond to only 1 other atom • Atom w/ the lowest EN is often the central atom

  16. Lewis Structure Practice Draw CH3I • Count valence e- C: (1 atom)(4 e-) = 4 e- H: (3 atoms)(1 e-) = 3 e- I: (1 atom)(7 e-) = 7 e- 14 e-

  17. Lewis Structure Practice • Arrange atoms & form single bonds H H : C : I H 3. Complete the octets & verify # of e- H H : C : I : H

  18. Multiple Bonds • C, N, & O commonly form double bonds • N & C can form triple bonds

  19. Lewis Structure Practice Draw SO3 • Count valence e- • (1 x 6 e-) + (3 x 6 e-) = 24 val. e- • Arrange atoms & form single bonds

  20. Lewis Structure Practice • Complete octets • Already used 24, no remaining pairs for the central atom

  21. Lewis Structure Practice 5. Try double bonds, then triple bonds if necessary

  22. Resonance Structure • Multiple Lewis structures possible for 1 molecule • Intermediate structure • Ex: O3

  23. Polyatomic Ion Structure • Account for charge in the total # of val.e- • Negative = add e- • Positive = subtract e- • Put structure in brackets & write charge on the top right

  24. Polyatomic Ion Practice Draw NO3- • Count valence e- • (1 x 5 e- ) + (3 x 6 e-) + 1 = 24 e- • Connect atoms • Add octet to atoms bonded to central atom

  25. Polyatomic Ion Practice • Place leftover e- on central atom • Already used 24 • If no octet, try double bond 6. Check for resonance structures

  26. Octet Rule Exceptions • H never has more than 2 val. e- • B & Al may have 6 val. e- • Ionic bonds: only non-metals have octet

  27. Metal Practice (Ionic Cmpds) Draw the Lewis structure for BaBr2 (1 x 2 e-) + (2 x 7 e-) = 16 e- : Br : Ba : Br :

  28. Naming Covalent Cmpds • 1st element named is least EN • Add prefix if more than 1 atom • Table 6-5, p.212 • 2nd element is most EN • Add prefix & suffix -ide • Ex: CO2 = carbon dioxide

  29. Covalent Naming Practice • SCl4 • Sulfur tetrachloride • P4O6 • Tetraphosphorus hexoxide • N2O4 • Dinitrogen tetroxide • Drop vowel on prefix if root begins w/ vowel

  30. (6-3) VSEPR • Valence shell e- pair repulsion theory: predicts molecule shape based on the repulsion b/w e- clouds • e- pairs position themselves as far apart as possible

  31. Molecular Shapes • Linear: • Bent: • Trigonal planar: • Tetrahedral: • Trigonal pyramidal:

  32. Shape Affects Properties • Generally, greater polarity  higher bp • Harder to break • Molecular dipole: • Ex: H2O • Ex: CO2

  33. (6-4) Intermolecular Forces • Attraction b/w molecules • W/out these forces all covalent substances would be gases • Weaker than ionic forces

  34. Dipole Force • Force b/w + & - ends of polar molecules • Hydrogen bond: strong dipole attraction in which a H atom is bonded to a strongly EN atom • N, O, F (halogens)

  35. London Forces • (Dispersion forces): attraction b/w atoms & molecules caused by formation of instantaneous dipoles • Weakest forces

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