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Chapter 1 Introduction and Review

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Title: Chapter 1 Introduction and Review


1
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

2
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

3
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

4
First Electron Shell
The 1s orbital holds two electrons.
5
Second Electron Shell
6
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

?
??
??
7
Electronic Configurations
gt
8
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.

9
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.

    gt

10
Bond Formation
  • Ionic bonding electrons are transferred.
  • Covalent bonding electron pair is shared.

11
Bonding Region
  • Electrons are close to both nuclei.

12
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.
    gt

13
H2 s-s overlap
gt
14
Cl2 p-p overlap
Constructive overlap along the same axis forms a
sigma bond.
gt
15
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.

    gt

16
sp Hybrid Orbitals
  • 2 VSEPR pairs
  • Linear electron pair geometry
  • 180 bond angle

17
sp2 Hybrid Orbitals
  • 3 VSEPR pairs
  • Trigonal planar e- pair geometry
  • 120 bond angle

18
sp3 Hybrid Orbitals
  • 4 VSEPR pairs
  • Tetrahedral e- pair geometry
  • 109.5 bond angle

19
Pi Bonding
  • Pi bonds form after sigma bonds.
  • Sideways overlap of parallel p orbitals.

20
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.

21
Lewis Structures
  • Bonding electrons
  • Nonbonding electrons or lone pairs

Satisfy the octet rule! gt
22
Chemical Formulas
  • CH3COOH
  • C2H4O2
  • CH2O gt
  • Full structural formula (no lone pairs shown)
  • Line-angle formula
  • Condensed structural formula
  • Molecular formula
  • Empirical formula

23
Multiple Bonding
gt
24
Rotation around Bonds
  • Single bonds freely rotate.
  • Double bonds cannot rotate unless the bond is
    broken.

25
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

26
Structural Isomers
27
Stereoisomers
Cis-trans isomers are also called geometric
isomers. There must be two different groups on
the sp2 carbon.
28
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.

29
Electronegativity and Bond Polarity
  • Greater ?EN means greater polarity

gt
30
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

31
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
32
Resonance Example
  • The real structure is a resonance hybrid.
  • All the bond lengths are the same.
  • Each oxygen has a -1/3 electrical charge.

    gt

33
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
34
Resonance Hybrid
35
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.

36
BrØnsted-Lowry Acids and Bases
  • Acids can donate a proton.
  • Bases can accept a proton.
  • Conjugate acid-base pairs.

37
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.

38
Structural Effects on Acidity
  • Electronegativity
  • Size
  • Resonance stabilization of conjugate base


    gt

39
Electronegativity
  • As the bond to H becomes more polarized, H
    becomes more positive and the bond is easier to
    break.

gt
40
Size
  • As size increases, the H is more loosely held and
    the bond is easier to break.
  • A larger size also stabilizes the anion.

41
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.

42
Lewis Acids and Bases
  • Acids accept electron pairs electrophile
  • Bases donate electron pairs nucleophile

43
End of Chapter 1
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