Chapter 1 Introduction and Review

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Chapter 1Introduction and Review
Organic Chemistry, 6th EditionL. G. Wade, Jr.

• Jo Blackburn
• Richland College, Dallas, TX
• Dallas County Community College District
• ã 2006, Prentice Hall

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Atomic Structure

• Atoms protons, neutrons, and electrons.
• The number of protons determines the identity of
  the element.
• Some atoms of the same element have a different
  number of neutrons. These are called isotopes.
• Example 12C, 13C, and 14C.
  gt

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Electronic Structure

• Electrons outside the nucleus, in orbitals.
• Electrons have wave properties.
• Electron density is the probability of finding
  the electron in a particular part of an orbital.
• Orbitals are grouped into shells, at different
  distances from the nucleus.
  gt

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First Electron Shell
The 1s orbital holds two electrons.
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Second Electron Shell
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Electronic Configurations

• Aufbau Principle Place electrons in lowest
  energy orbital first.
• Hunds Rule Equal energy orbitals are
  half-filled, then filled.
• 6C 1s2 2s2 2p2

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Electronic Configurations
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Wave Properties of Electrons

• Standing wave vibrates in fixed location.
• Wave function, ?, mathematical description of
  size, shape, orientation.
• Amplitude may be positive or negative.
• Node amplitude is zero.

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Wave Interactions

• Linear combination of atomic orbitals
• on different atoms produce molecular orbitals
• on the same atom give hybrid orbitals.
• Conservation of orbitals.
• Waves that are in phase add together.Amplitude
  increases.
• Waves that are out of phase cancel out.
  
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Bond Formation

• Ionic bonding electrons are transferred.
• Covalent bonding electron pair is shared.

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Bonding Region

• Electrons are close to both nuclei.

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Sigma Bonding

• Electron density lies between the nuclei.
• A bond may be formed by s-s, p-p, s-p, or
  hybridized orbital overlaps.
• The bonding MO is lower in energy than the
  original atomic orbitals.
• The antibonding MO is higher in energy than the
  atomic orbitals.
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H2 s-s overlap
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Cl2 p-p overlap
Constructive overlap along the same axis forms a
sigma bond.
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Molecular Shapes

• Bond angles cannot be explained with simple s and
  p orbitals. Use VSEPR theory.
• Hybridized orbitals are lower in energy because
  electron pairs are farther apart.
• Hybridization is LCAO within one atom, just prior
  to bonding.
  
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sp Hybrid Orbitals

• 2 VSEPR pairs
• Linear electron pair geometry
• 180 bond angle

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sp2 Hybrid Orbitals

• 3 VSEPR pairs
• Trigonal planar e- pair geometry
• 120 bond angle

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sp3 Hybrid Orbitals

• 4 VSEPR pairs
• Tetrahedral e- pair geometry
• 109.5 bond angle

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Pi Bonding

• Pi bonds form after sigma bonds.
• Sideways overlap of parallel p orbitals.

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Multiple Bonds

• A double bond (2 pairs of shared electrons)
  consists of a sigma bond and a pi bond.
• A triple bond (3 pairs of shared electrons)
  consists of a sigma bond and two pi bonds.

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Lewis Structures

• Bonding electrons
• Nonbonding electrons or lone pairs

Satisfy the octet rule! gt
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Chemical Formulas

• CH3COOH
• C2H4O2
• CH2O gt

• Full structural formula (no lone pairs shown)
• Line-angle formula
• Condensed structural formula
• Molecular formula
• Empirical formula

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Multiple Bonding
gt
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Rotation around Bonds

• Single bonds freely rotate.
• Double bonds cannot rotate unless the bond is
  broken.

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Isomerism

• Same molecular formula, but different arrangement
  of atoms isomers.
• Constitutional (or structural) isomers differ in
  their bonding sequence.
• Stereoisomers differ only in the arrangement of
  the atoms in space. gt

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Structural Isomers
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Stereoisomers
Cis-trans isomers are also called geometric
isomers. There must be two different groups on
the sp2 carbon.
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Dipole Moment

• Amount of electrical charge x bond length.
• Charge separation shown by electrostatic
  potential map (EPM).
• Red indicates a partially negative region and
  blue indicates a partially positive region.

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Electronegativity and Bond Polarity

• Greater ?EN means greater polarity

gt
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Calculating Formal Charge

• For each atom in a valid Lewis structure
• Count the number of valence electrons
• Subtract all its nonbonding electrons
• Subtract half of its bonding electrons

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Resonance

• Only electrons can be moved (usually lone pairs
  or pi electrons).
• Nuclei positions and bond angles remain the same.
• The number of unpaired electrons remains the
  same.
• Resonance causes a delocalization of electrical
  charge.

Examplegt
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Resonance Example

• The real structure is a resonance hybrid.
• All the bond lengths are the same.
• Each oxygen has a -1/3 electrical charge.
  
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Major Resonance Form

• Has as many octets as possible.
• Has as many bonds as possible.
• Has the negative charge on the most
  electronegative atom.
• Has as little charge separation as possible.

Examplegt
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Resonance Hybrid
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Arrhenius Acids and Bases

• Acids dissociate in water to give H3O ions.
• Bases dissociate in water to give OH- ions.
• Kw H3O OH- 1.0 x 10-14 at 24C
• pH -log H3O
• Strong acids and bases are 100 dissociated.

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BrØnsted-Lowry Acids and Bases

• Acids can donate a proton.
• Bases can accept a proton.
• Conjugate acid-base pairs.

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Acid and Base Strength

• Acid dissociation constant, Ka
• Base dissociation constant, Kb
• For conjugate pairs, (Ka)(Kb) Kw
• Spontaneous acid-base reactions proceed from
  stronger to weaker.

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Structural Effects on Acidity

• Electronegativity
• Size
• Resonance stabilization of conjugate base
  
  
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Electronegativity

• As the bond to H becomes more polarized, H
  becomes more positive and the bond is easier to
  break.

gt
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Size

• As size increases, the H is more loosely held and
  the bond is easier to break.
• A larger size also stabilizes the anion.

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Resonance

• Delocalization of the negative charge on the
  conjugate base will stabilize the anion, so the
  substance is a stronger acid.
• More resonance structures usually mean greater
  stabilization.

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Lewis Acids and Bases

• Acids accept electron pairs electrophile
• Bases donate electron pairs nucleophile

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End of Chapter 1