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 • Region where an e- pair is most likely to exist – Formed by overlapping atomic orbitals Bond Length • Avg. dist. b/w 2 bonded atoms – Occur at min. PE Bond E • E required to break a bond b/w 2 atoms & separate them • Stronger bonds are shorter – Single = long = weak – Triple = short = strong Electronegativity • Tendency of an atom to attract bonding e- to itself • Inc. across a period, dec. down a group Electron Density • The more EN atom, has a higher electron density than the less EN atom – Pulls more e- to it 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_Pres entations/Bonding_part_III_polar.ppt Dipole • Molecule in which 1 end has a partial + charge & the other end has a partial charge Dipole Moment (EN dif.) • Determines polarity of a bond & molecule • Larger d.m. higher polarity stronger bond (6-2) Valence Electrons • e- in the outer-most E level of an atom, where it can participate in bonding 1 0 1 2 3 4 3 2 1 0 1 2 3 4 3 2 1 0 1 2 3 4 3 2 1 0 transition metals Lewis Structure • Lewis structure: represents the valence e- in a molecule Lewis Dot Structure • Place 1 e- on each side of atom before pairing any e- Unshared Pair • (Lone pair): pair of valence e- not involved in bonding Rules for Drawing Lewis Structures • H & halogens bond to only 1 other atom • Atom w/ the lowest EN is often the central atom Lewis Structure Practice Draw CH3I 1. Count valence eC: (1 atom)(4 e-) = 4 eH: (3 atoms)(1 e-) = 3 eI: (1 atom)(7 e-) = 7 e14 e- Lewis Structure Practice 2. Arrange atoms & form single bonds H H:C:I H 3. Complete the octets & verify # of eH H:C:I: H Multiple Bonds • C, N, & O commonly form double bonds • N & C can form triple bonds Lewis Structure Practice Draw SO3 1. Count valence e• (1 x 6 e-) + (3 x 6 e-) = 24 val. e- 2. Arrange atoms & form single bonds Lewis Structure Practice 3. Complete octets • Already used 24, no remaining pairs for the central atom Lewis Structure Practice 5. Try double bonds, then triple bonds if necessary Resonance Structure • Multiple Lewis structures possible for 1 molecule • Intermediate structure • Ex: O3 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 Polyatomic Ion Practice Draw NO31. Count valence e• (1 x 5 e- ) + (3 x 6 e-) + 1 = 24 e- 2. Connect atoms 3. Add octet to atoms bonded to central atom Polyatomic Ion Practice 4. Place leftover e- on central atom • Already used 24 5. If no octet, try double bond 6. Check for resonance structures 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 Metal Practice (Ionic Cmpds) Draw the Lewis structure for BaBr2 (1 x 2 e-) + (2 x 7 e-) = 16 e- : Br : Ba : Br : 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 Covalent Naming Practice • SCl4 – Sulfur tetrachloride • P4O6 – Tetraphosphorus hexoxide • N 2O 4 – Dinitrogen tetroxide – Drop vowel on prefix if root begins w/ vowel (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 Molecular Shapes • Linear: • Bent: • Trigonal planar: • Tetrahedral: • Trigonal pyramidal: Shape Affects Properties • Generally, greater polarity higher bp – Harder to break • Molecular dipole: – Ex: H2O – Ex: CO2 (6-4) Intermolecular Forces • Attraction b/w molecules • W/out these forces all covalent substances would be gases • Weaker than ionic forces 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) London Forces • (Dispersion forces): attraction b/w atoms & molecules caused by formation of instantaneous dipoles • Weakest forces