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_Presentations/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 transition metals 3 4 3 2 1 0

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 e-C: (1 atom)(4 e-) = 4 e-

H: (3 atoms)(1 e-) = 3 e-

I: (1 atom)(7 e-) = 7 e-

14 e-

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

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 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

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

• N2O4

– 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