Lecture 1 - Introduction and Hydrocarbon structure

1. Review of Chemical concepts
   1. Basic Structure of atoms
      1. Nucleus - protons, neutrons
         1. Protons - positive, identify the atom
         2. neutrons - neutral, determine isotope
      2. Electrons
         1. Shells
            1. Shells can hold specific numbers of electrons
            2. Can be several orbitals in a shell
         2. Valence
            1. electrons that occupy the outer shell
   2. Covalent
      1. Lewis Structures
         1. Named after Gilbert Lewis
         2. Stable electron configuration is a octet
         3. A full octet is a closed shell configuration
         4. Lewis structure depicts the covalent structure of atoms by showing all valence electrons
         5. In covalent bonds electrons can be shared. Called bonding electrons
         6. Valence electrons not involved in bonding are Non-bonding electrons
         7. Bond depicted as two dots or a line
         8. In a atom in a molecule is most stable with 8 electrons. Except for H with is given 2
         9. Carbon is tetravalent and will always have 4 bonds
      2. Formal Charges
         1. Atoms that have an octet don't always have a neutral charge
         2. Atom "owns" one electron from a covalent bond
         3. Also count in non-bonded electrons
         4. non-bonded + 1/2 number of bonds
   3. Connectivity
      1. Saturated compounds
         1. Molecular formulas say nothing about bond connections
         2. One molecular formula can specify many different structures.
         3. Must be aware of how atoms are connected. Use idea of Octet Rule to promote this.
         4. Carbon is different from other elements in that it tends to form chains
         5. Chains can have branches
         6. Saturated compounds have a common formula C(n)H(2n+2)
      2. Unsaturated compounds
         1. Remember carbon can had multiple bonds.
         2. Multiple bonds still give can complete octet
         3. Idea of "unit of saturation"
   4. Structure
      1. Tetrahedron
         1. Unlike ionic bonds, covalent bonds, have a property of distance and direction
         2. A bonds between two atoms will have an average length, or bond length
         3. Three atoms, bonded continuously will have an average angle between them, or bond angle
         4. Singly bonded carbon form bond angles of about 109. Making a tetrahedron
         5. In 2D this is depicted with wedges and lines.
         6. Even in chains the geometry of the tetrahedron holds.
         7. Here the bond lengths will differ (C-C, C-hetro), but the bonds angles will be similar
         8. Other atoms can have "close" tetrahedron angles (triangular, bent ...)
      2. Conformational isomers
         1. Single bonds are not rigid, atoms and twist and turn using the bond as an axis.
         2. Thus a given molecule can have many different shapes over a period of time.
         3. Molecule can interchange between isomers
         4. No bonds are broken though, so even if it looks different it is the same molecule
         5. The bond connections do not change.
         6. Some forms of the molecule will be favored, due to there energy
      3. Planer molecules
         1. Double bonds are more rigid than single bonds
         2. The atoms can not rotate freely as they can in single bonds
         3. Result is that atoms in the double bond and those attached will be in a plane
      4. Linear molecules
         1. Have even less flexibility than single bonds.
         2. Atoms in the bond and those connect lye along the tripple bond axis.
      5. Geometric isomers.
         1. Since structure is more hindered, have potential for other structures
         2. Example butene: unlike conformational isomers the structure can not interchange
         3. To interchange would require breaking the double bond. Not easily done.