Chapter 1' Structure and Bonding Dr' Ralph Mead Fall 07 CHM 211 Organic Chemistry I

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Chapter 1. Structure and BondingDr. Ralph
MeadFall 07 CHM 211Organic Chemistry I
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Organic Chemistry

• Q What is Organic Chemistry?
• Organic compounds are those based on carbon
  structures and organic chemistry studies their
  structures and reactions
• Q Why do we care?
• Carbon forms strong bonds to other carbon atoms
  and elements. Chains and rings of carbon atoms
  can form biological molecules, drugs, solvents
  and dyes.

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Topics to Discuss

• Atomic Structure
• Orbitals
• Electron Configurations
• Bonding Theory
• Chemical Bonds
• Hybridization
• Molecular Orbital Theory

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Atomic Structure

• Structure of an atom
• Positively charged nucleus contains protons and
  neutrons.
• Electrons orbit the nucleus in orbitals.
• Electrons are important!
• Electrons form bonds and determine structure of
  molecules!

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Atomic Number and Atomic Mass

• The atomic number (Z) protons
• The mass number (A) protons neutrons
• All the atoms of a given element have the same
  atomic number
• The atomic weight of an element is the weighted
  average mass in atomic mass units (amu) of an
  elements naturally occurring isotopes

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Atomic Structure Orbitals

• Q How do electrons move?
• Electrons are so small and light that they show
  properties of particles and waves.
• Q How do we know the exact location of the
  electrons?
• We can never know EXACTLY where an electron is,
  but we can determine the probability of finding
  it in a particular location, using Quantum
  Mechanics.

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Quantum Mechanics

• Q What is quantum mechanics?
• Quantum mechanics describes electron energies
  and locations by a wave equation
• Wave function is a solution of wave equation
• Each Wave function is an orbital,?
• A plot of ? 2 describes where electron most
  likely to be. It is a mathematical description
  of the shape of a wave as it vibrates.
• Electron cloud has no specific boundary so we
  show most probable area

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Types of Orbitals

• Q How many different orbitals are there?
• Four different kinds of orbitals
• s, p, d, and f
• Q Which orbitals are important?
• in organic chemistry, s and p orbitals are the
  most important!

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Shapes of Orbitals

• s orbitals spherical
• p orbitals dumbbell-shaped

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Orbitals and Shells

• Q What is a shell?
• a group of orbitals of increasing size and
  energy
• Different shells contain different numbers and
  kinds of orbitals
• Q How many electrons are in each shell?
• Each orbital can be occupied by two electrons

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p-Orbitals

• In each shell there are three perpendicular p
  orbitals of equal energy px, py, and pz
• Lobes of a p orbital are separated by region of
  zero electron density, called a node

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Atomic Structure Electron Configurations

• Ground-state electron configuration of an atom
  lists orbitals occupied by its electrons.
• Rules
• 1. (Aufbau principle)
• Lowest-energy orbitals fill first 1s ? 2s ? 2p ?
  3s ? 3p ? 4s ? 3d
• 2. (Pauli exclusion principle)
• Electron spin can have only two orientations, up
  ? and down ?
• Only two electrons can occupy an orbital, and
  they must be of opposite spin to have unique wave
  equations
• 3. (Hund's rule)
• If two or more empty orbitals of equal energy are
  available, electrons occupy each with spins
  parallel until all orbitals have one electron

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Electron Configurations

• Q Write the electron configuration for the
  following
• A. Na
• B. C

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Development of Chemical Bonding Theory

• Chemists discovered 2 important things about
  carbon atoms
• Carbon always has FOUR bonds
• Carbons four bonds have DIFFERENT SPATIAL
  DIRECTIONS
• Atoms surround carbon as corners of a tetrahedron

Note that a dashed line indicates a bond is
behind the page
Note that a wedge indicates a bond is coming
forward
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Valence Shells and Bonding

• Q What is a valence shell?
• Outermost shell of an atom.
• Q What are valence electrons?
• electrons in the valence shell.
• Q Why are valence electrons so important?
• They are the electrons that are exposed to other
  atoms and molecules.
• They are the electrons that are involved in
  bonding and chemical reactivity.

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Valence Shells and Bonding

• Q What is an ionic bond?
• bonding that occurs by the attraction of
  oppositely charged ions.
• EXAMPLE sodium chloride (NaCl-)
• Q What is a covalent bond?
• bonding that occurs by the sharing of electrons
  in the region between two nuclei.
• EXAMPLE methane (CH4)

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Valence Shells and Bonding

• Q How many valence electrons do the following
  atoms have?
• A. Na
• C

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The Nature of the Chemical Bond

• Q Why do atoms form bonds?
• Because the compound that results is more stable
  than the separate atoms
• Q How do we know how many bonds an atom will
  form?
• Atoms with one, two, or three valence electrons
  form one, two, or three bonds
• Atoms with four or more valence electrons form as
  many bonds as they need electrons to fill the s
  and p levels of their valence shells to reach a
  stable octet

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Non-bonding electrons

• Q What are non-bonding electrons?
• Valence electrons not used in bonding
• also called lone-pair electrons

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Lewis Structures

• Q Why do we need Lewis structures?
• They are used to represent organic compounds
  whenever discussing valence electrons and their
  role in bonding.
• shown valence electrons of an atom as dots
• Hydrogen has one dot, representing its 1s
  electron
• Carbon has four dots (2s2 2p2)
• Q How do we know if we have the correct Lewis
  structure?
• Stable molecule results in a completed shell,
  with the octet rule satisfied
• octet (eight dots) for main-group atoms (two for
  hydrogen)

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Valences of Carbon

• Carbon has four valence electrons
• (2s2 2p2), forms four bonds (CH4)

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Valences of Oxygen

• Oxygen has six valence electrons
• (2s2 2p4), forms two bonds (H2O)

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Valences of Nitrogen

• Nitrogen has five valence electrons
• (2s2 2p3), forms only three bonds (NH3)

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Drawing Lewis Structures

• Q Draw a Lewis Structure for ammonia, NH3.
• STRATEGY
• 1. Determine the number of valence electrons in
  each atom.
• 2. Determine the arrangement of the atoms.
• 3. Arrange the remaining electrons so that each
  atom has a complete outer shell.
• 4. Show bonding pairs as a single line,
  nonbonding electrons as dots.

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Drawing Lewis Structures

• Q Draw a Lewis Structure for ammonia, NH3.
• 1. Determine the number of valence electrons in
  each atom.
• nitrogen has 5 valence electrons, hydrogen has
  one.
• 2. Determine the arrangement of the atoms.
• since hydrogen can only form a single bond to
  other atoms, the only sensible way to arrange the
  atoms, is with the nitrogen in the center of the
  hydrogens.

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Drawing Lewis Structures

• 3. Arrange the remaining electrons so that each
  atom has a complete outer shell.
• 4. Show bonding pairs as a single line,
  nonbonding electrons as dots.

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Drawing Line-Bond Structures

• Q What are line-bond drawings?
• They show the carbon skeleton with any
  functional groups that are attached, such as OH
  or Br.
• Q Why are they so useful?
• it saves a lot of time when drawing molecules,
  b/c you dont have to draw every carbon and
  hydrogen!
• Lines are drawn in a zigzag format, so that the
  end of every line represents a carbon atom
• Hydrogen atoms are NOT shown.

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Drawing Line-Bond Structures

• Q How do we draw them?
• using this molecule as an example, focus on the
  carbon skeleton, and be sure to draw any atoms
  other than C and H. ALL ATOMS OTHER THAN C AND H
  MUST BE DRAWN.
• EXAMPLE
• Notice that the carbon atoms in the straight
  chain are drawn in a zigzag format.

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Drawing Line-Bond Structures

• MISTAKES TO AVOID
• 1. NEVER draw a carbon atom with more than 4
  bonds.
• 2. When drawing a molecule, you should either
  show all of the Hs and Cs, or draw a line-bond
  drawing without ANY Hs and Cs. If you draw the
  carbons, you MUST draw the hydrogens.
• 3. When drawing the carbon atoms in zigzag
  format, try to draw all atoms of the bonds as far
  apart as possible.
• 4. In line-bond drawings, we do draw any Hs
  that are connected to atoms other than carbon.

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Drawing Line-Bond Structures



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Valence Bond Theory

• How does a covalent bond form?
• When two atoms get so close that an orbital,
  occupied with 1e-, overlaps
• Head-on overlap forms a single bond, called a s
  bond
• s bonds are the most common type of bond in
  organic compounds
• In a s bond, the e- in the overlapping orbitals
  are attracted to both nuclei
• An example of a s bond is the H2 bond shown in
  picture

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Bond Energy

• Q What is Bond Energy?
• The amount of energy released when a bond forms,
  or the amount of energy required to break one.
• AKA Bond Strength

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Bond Length

• Q What is bond length?
• The optimum distance between nuclei that leads to
  maximum stability
• Q Why is a bond length fixed?
• If too close, they repel because both are
  positively charged
• If too far apart, bonding is weak

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Hybridization

• Q What is hybridization?
• There are two simple orbitals s and p.
• s orbitals are spherical and p orbitals are
  dumbbell shaped with two lobes, one in the front
  and one in the back.
• Hybridization results from the mixing of
  orbitals on the same atom.
• 2nd row elements (C, N, O, and F) have one s
  and 3 p orbitals in the valence shell. These
  orbitals are usually mixed together to give
  hybridized orbitals sp, sp2, and sp3.

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Hybridization

• Q Why does hybridization occur?
• It occurs to separate electron pairs more widely
  in space.
• It places more electron density in the bonding
  region between the nuclei.

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Hybridization sp3 Orbitals and the Structure of
Methane

• Carbon has 4 valence electrons (2s2 2p2)
• In CH4, all CH bonds are identical (tetrahedral)
  
• sp3 hybrid orbitals s orbital and three p
  orbitals combine to form four equivalent,
  unsymmetrical, tetrahedral orbitals (sppp sp3),
  Pauling (1931)

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Tetrahedral Structure of Methane

• sp3 orbitals on C overlap with 1s orbitals on 4 H
  atom to form four identical C-H bonds
• Each CH bond has a strength of 438 kJ/mol and
  length of 110 pm
• Bond angle each HCH is 109.5, the tetrahedral
  angle.

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Hybridization sp3 Orbitals and the Structure of
Ethane

• Two Cs bond to each other by s overlap of an sp3
  orbital from each
• Three sp3 orbitals on each C overlap with H 1s
  orbitals to form six CH bonds
• CH bond strength in ethane 420 kJ/mol
• CC bond is 154 pm long and strength is 376
  kJ/mol
• All bond angles of ethane are tetrahedral

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Hybridization sp2 Orbitals and the Structure of
Ethylene

• sp2 hybrid orbitals 2s orbital combines with two
  2p orbitals, giving 3 orbitals (spp sp2)
• sp2 orbitals are in a plane with120 angles
• Remaining p orbital is perpendicular to the plane
• Another way.

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Bonds From sp2 Hybrid Orbitals

• Two sp2-hybridized orbitals overlap to form a s
  bond
• p orbitals overlap side-to-side to formation a pi
  (?) bond
• sp2sp2 s bond and 2p2p ? bond result in sharing
  four electrons and formation of C-C double bond
• Electrons in the s bond are centered between
  nuclei
• Electrons in the ? bond occupy regions are on
  either side of a line between nuclei

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Structure of Ethylene

• H atoms form s bonds with four sp2 orbitals
• HCH and HCC bond angles of about 120
• CC double bond in ethylene shorter and stronger
  than single bond in ethane
• Ethylene CC bond length 133 pm (CC 154 pm)

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Hybridization sp Orbitals and the Structure of
Acetylene

• C-C a triple bond sharing six electrons
• Carbon 2s orbital hybridizes with a single p
  orbital giving two sp hybrids
• two p orbitals remain unchanged
• sp orbitals are linear, 180 apart on x-axis
• Two p orbitals are perpendicular on the y-axis
  and the z-axis
• Another way

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Orbitals of Acetylene

• Two sp hybrid orbitals from each C form spsp s
  bond
• pz orbitals from each C form a pzpz ? bond by
  sideways overlap and py orbitals overlap similarly

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Bonding in Acetylene

• Sharing of six electrons forms C ºC
• Two sp orbitals form s bonds with hydrogens

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C-C, CC, and C?C

• Bond length
• C-C gt CC gt CC
• Bond Strength
• C-C lt CC lt CC
• Bond Angles
• C-C (109.5o), CC (120o), CC (180o)
• Hybridizations
• C-C (sp3, all s)
• CC (sp2, one s and one p)
• CC (sp, one s, and two p)

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HybridizationNitrogen in Ammonia

• Q Can elements other than C have hybridized
  orbitals?
• HNH bond angle in ammonia (NH3) 107.3
• Ns orbitals (sppp) hybridize to form four sp3
  orbitals
• One sp3 orbital is occupied by two nonbonding
  electrons, and three sp3 orbitals have one
  electron each, forming bonds to H

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HybridizationOxygen in Water

• The oxygen atom is sp3-hybridized
• Oxygen has six valence-shell electrons but forms
  only two covalent bonds, leaving two lone pairs
• The HOH bond angle is 104.5

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Molecular Orbital Theory

• A molecular orbital (MO) where electrons are
  most likely to be found (specific energy and
  general shape) in a molecule
• A bonding MO is lower in energy than the two
  orbitals from which it forms.
• An antibonding MO is the oppositeit is higher
  energy.

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Molecular Orbital vs. Valence Bond Theory

• Valence bond and molecular orbital theories are
  complimentary, simply used for different
  purposes.
• VALENCE BOND THEORY
• provides a good qualitative description of
  bonding in molecules for use in routine
  situations, and should be used to think about the
  molecules and reactions described this semester.
• MOLECULAR ORBITAL THEORY
• reserved for detailed computer calculations when
  more quantitative results are required, or when
  describing pi bonding, such as with aromaticity.

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Summary

• Atomic Structure
• Orbitals
• Electron Configurations
• Bonding Theory
• Chemical Bonds
• Hybridization
• Molecular Orbital Theory

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Hybridization

• Q What are the 3 kinds of hybridizations that
  carbon exhibits in compounds?
• sp
• sp2
• sp3
• Q What are the shapes and bond angles of the 3
  hybrids?
• sp linear 180o bond angles acetylene
  (H-CC-H)
• sp2 planar 120o bond angles ethylene
  (H2CCH2)
• sp3 tetragonal 109.5o bond angles ethane
  (H3C-CH3)

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Hybridization

• Q What is a pi (p) bond?
• a bond formed by the side-to-side overlap of p
  orbitals.
• Remember, a sigma (s) bond is formed by head on
  overlap of bonds.
• Q Is a p bond stronger or weaker than a s bond?
• a p bond is shorter and stronger than a s bond
• in a CC bond, there is one s, one p bond.