covalent bonds & molecular forces
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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. - PowerPoint PPT PresentationTRANSCRIPT
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Covalent Bonds & Molecular Forces
Ch.6
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(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
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Molecular Orbital
• Region where an e- pair is most likely to exist– Formed by overlapping atomic orbitals
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Bond Length
• Avg. dist. b/w 2 bonded atoms– Occur at min. PE
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Bond E
• E required to break a bond b/w 2 atoms & separate them
• Stronger bonds are shorter– Single = long = weak– Triple = short = strong
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Electronegativity
• Tendency of an atom to attract bonding e- to itself
• Inc. across a period, dec. down a group
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Electron Density
• The more EN atom, has a higher electron density than the less EN atom– Pulls more e- to it
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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
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Dipole
• Molecule in which 1 end has a partial + charge & the other end has a partial - charge
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Dipole Moment (EN dif.)
• Determines polarity of a bond & molecule
• Larger d.m. higher polarity stronger bond
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(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 transition metals 3 4 3 2 1 0
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Lewis Structure
• Lewis structure: represents the valence e- in a molecule
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Lewis Dot Structure
• Place 1 e- on each side of atom before pairing any e-
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Unshared Pair
• (Lone pair): pair of valence e- not involved in bonding
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Rules for Drawing Lewis Structures
• H & halogens bond to only 1 other atom
• Atom w/ the lowest EN is often the central atom
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Lewis Structure Practice
Draw CH3I
1. 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-
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Lewis Structure Practice
2. Arrange atoms & form single bonds H
H : C : I H3. Complete the octets & verify # of e-
H H : C : I :
H
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Multiple Bonds
• C, N, & O commonly form double bonds
• N & C can form triple bonds
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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
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Lewis Structure Practice
3. Complete octets
• Already used 24, no remaining pairs for the central atom
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Lewis Structure Practice
5. Try double bonds, then triple bonds if necessary
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Resonance Structure
• Multiple Lewis structures possible for 1 molecule
• Intermediate structure
• Ex: O3
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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
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Polyatomic Ion Practice
Draw NO3-
1. 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
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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
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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
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Metal Practice (Ionic Cmpds)
Draw the Lewis structure for BaBr2
(1 x 2 e-) + (2 x 7 e-) = 16 e-
: Br : Ba : Br :
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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
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Covalent Naming Practice
• SCl4– Sulfur tetrachloride
• P4O6
– Tetraphosphorus hexoxide
• N2O4
– Dinitrogen tetroxide – Drop vowel on prefix if root begins w/
vowel
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(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
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Molecular Shapes
• Linear: • Bent:
• Trigonal planar:
• Tetrahedral:
• Trigonal pyramidal:
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Shape Affects Properties
• Generally, greater polarity higher bp– Harder to break
• Molecular dipole:– Ex: H2O
– Ex: CO2
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(6-4) Intermolecular Forces
• Attraction b/w molecules
• W/out these forces all covalent substances would be gases
• Weaker than ionic forces
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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)
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London Forces
• (Dispersion forces): attraction b/w atoms & molecules caused by formation of instantaneous dipoles
• Weakest forces