Transcript Ch.6 Notes - Green Local Schools

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