Covalent Bonds

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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_Presentat
  ions/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_Presentat
  ions/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

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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
• 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

• Arrange atoms form single bonds
• H
• H C I
• H
• 3. 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
• Count valence e-
• (1 x 6 e-) (3 x 6 e-) 24 val. e-
• Arrange atoms form single bonds

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Lewis Structure Practice

• 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-
• Count valence e-
• (1 x 5 e- ) (3 x 6 e-) 1 24 e-
• Connect atoms
• Add octet to atoms bonded to central atom

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Polyatomic Ion Practice

• Place leftover e- on central atom
• Already used 24
• 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