Chem301 Ch1 USD

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Mattyj1388
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169279
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Chem301 Ch1 USD
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2012-10-02 21:26:40
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Ch1 Organic Chem USD chem301
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general terms and definitions text book "Organic Chemistry" University of San Diego/ Klein: copyright 2012 John Wiley & Sons, Inc. sigma bonds, pi bonds, ionic bonding, covalent bonding, basics, histories
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  1. Reactions occur
    as a result of the motion of electrons.
  2. Vitalism
    • "vital force": stipulated that it should be impossible to convert inorganic compounds into organic compounds without the introduction of an outside vital force.
    • Destroyed by German chemist Friedrich Wohler in 1828 by the conversion of ammonium cyanate (a known inorganic salt) into urea, a known organic compound found in urine:
    • NH4OCN --Heat-->
  3. Constitutional isomers
    • have different physical properties and different names but have the same elements and the same number of each element (same moleculare formula) but attached differently.
    • EX: Ammonium cyanate and Urea
    • or Dimethyl ether and Ethanol
  4. Valence
    • describes the number of bonds usually formed by each elemnt.
    • EX: Carbon makes four bonds and is called tetravalent.
  5. Tetravalent
    • Makes four bonds.
    • EX: Carbon "C"
  6. Trivalent
    • Makes three bonds.
    • EX: Nitrogen "N"
  7. Divalent
    • Makes two bonds.
    • EX: Oxygen "O"
  8. Monovalent
    • Makes one bond.
    • EX: Hydrogen "H"
  9. Existence of electron
    • Was first introduced in 1874 by George Johnstone Stoney.
    • In 1897 J.J. Thomson demonstrated this and is credited to discovering the electon.
  10. Covalent Bond
    • Two atoms sharing a pair of electorns.
    • Note: Gilbert Lewis defined covalent bond in 1916.
    • Electro negativity value under (1.8??)
  11. Valence electrons
    The electrons in the outer most shell of an atom.
  12. Lone pair
    nonbonding pair of electrons.
  13. Formal charge
    is associated with any atom that does not exhibit the appropriate number of valence electrons.
  14. Electronegativity
    • The measure of the ability of an atom to attract electrons.
    • Note: typically Florine is the strongest and the electonegitivity decresses as you move away from "F".
  15. Nonpolar covalent bonding
    • will have an electronegativity of less than 0.5.
    • This means that the two atoms sharing the electons are doing so rather if not completly evenly.
  16. Polar covalent bonds
    • Will have an electronegativity value between 0.5 and 1.7.
    • This meas that although the two atoms are sharing a pair of electrons, the pull (or attraction) is not even and will cause each atom to be slightly charged, one positive and the other negative.
  17. Ionic binding
    • Will have an electronegativity value above 1.7.
    • Is when one atom "steals" an electron from another.
  18. Octet rule
    Second-row elements generally obey this rule, bonding to achive noble gas electron configuration.
  19. Quantum mechanics
    • 1924 french physicist Louis de Broglie suggested that electrons, heretofore considered as particles, also exhibited wavelike properties.
    • 1926 Erwin Schrodinger, Warner Heisenberg, and Paul Dirac independently proposed a mathmatical description of the electron that incorperated its wavelike properties. 
  20. Wave equation
    An equation is constructed to describe the total energy of a hydrogen atom (i.e. one proton plus one electron). This equation takes into account the wavelike behavior of an electron that is in the electric field of a proton.
  21. Wavefunctions
    is given when the wave equation is solved. The greek symbol psi () is used to denote each wavefunction with a subscript number(2). Each of these wavefunctions corresponds to an allowed enegy level for the electron. This result is incredably important because it suggests that an electron, when contained in an atom, can onlyexist at discrete energy levels. In other words, the energy of the electron is quantized.
  22. Electron density
    The probability of finding an electron in a particular region of space.
  23. "shape"
    The shape of an orbital refers to the region of space that contains 90 - 95% of the electron density. Beyond this region, the remaining 5-10% of the electron density tapers off but never ends. In fact, if we want to consider the region of space that contains 100% of the electron density, we must consider the entire universe.
  24. Atomic orbital
    • An occupied orbital must be treated as a cloud of electron density. Because it is a region of space defined with respect to the nucleus of a single atom.
    • EX: s, p, d, f
  25. Node
    Locations where a wave hits a zero value. or when =0.
  26. Degenerate orbitals
    Orbitals with the same energy level.
  27. Aufbau principle
    The lowest-energy orbital is filled first. 
  28. Pauli exclusion principle
    Each orbital can accomodate a maximum of 2 e- that have opposite spin. To understand what "spin" means, we can imagine an e- spinning in space (although this is an oversimplification explanation of the term "spin"). For reasons that are beyond the scope of this coarse, electrons only have two positive spin states (designated by  or ). In order for the orbital to accomodate 2 e-, the e- must have oppisite spin states. 
  29. Hund's rule
    When dealing with degenerate orbitals, such as p orbitals, 1 e- is placed in each degenerate orbital first, before electrons are paired up.
  30. Constructive interference
    produces a wave with larger amplitude.
  31. Destructive interference
    results in waves canceling each other, which produces a node.
  32. Valence bond theory
    a bond is simply the sharing of electron density between two atoms as a result of the constructive interference of their atomic orbitals.
  33. Sigma bond
    • The overlap of of an orbital between to atoms; the electon density of this bond is primarily located on the bond axis (the line that can be drawn between the 2 atoms).
    •  
  34. Molecular orbital (MO) theory
    describes a bond in terms of the constructive interference between 2 overlapping atomic orbitals (like bond theory). However it goes further by using mathmatics as a tool to explore the consequences of atomic orbital overlap. The mathmatical methode is called the linear combination of atomic orbitals (LCAO). According to this theory, atomic orbitals are mathmatically combined to produce new orbitalls, called molecular orbitals.
  35. Molecular orbitals
    the new orbitals formed from overlapping orbitals.
  36. Bonding MO
    the lower energy molecular orbital is the result of constructive interference of the original 2 atomic orbitals.
  37. Antibonding MO
    The higher the energy molecular orbital is the result of destructive interference. Notice that the antibonding MO has 1 node, which explains why it is higher in energy.
  38. Highest occupied molecular orbital
    (HOMO); the highest energy orbital from among the occupied orbitals.
  39. Lowest unoccupied molecule orbital
    (LUMO); the lowest energy orbital from among the unoccupied orbitals.
  40. hybridized atomic orbitals
    A hybrid orbital is an orbital formed by the combination of two or more atomic orbitals.
  41. sp3-hybridized orbitals
    • s orbital and 3 p orbitals combine to give a different structure and properties.
    • It also explains mathanes shape.
  42. pi bond
    • The p orbitals actually overlap with each other as well as in ethylene.
    • note: that there is only 1 pi bond above, NOT 2.
  43. Example of 1 sigma bond and 2 pi bonds
    • Acetylene:  
    • or rather
  44. What does VSEPR stand for
    Valence Shell Electron Pair Repulsion
  45. Steric number
    •    total sigma bonds
    • + total pi bonds
    • = Steric #
  46. Steric number 4 (sp3) could be what geometries? 
    • Tetrahedral
    • Trigonal pyramidal
    • Bent
  47. Steric number 3 (sp2) could be what geometries? 
    • Bent
    • Trigonal Planar
  48. Steric number 2 (sp) could be what geometries? 
    Linear
  49. Dipole moment
    • is used as an indicator of polarity, where () is used as an indicator of polarity, where () is defined as the amount of partial charge () on either end of the dipole multiplied by the distance of seporation (d):
    • =  + d
  50. Dipole-dipole interactions
    occur between two molecules that possess permanent dipole moments.
  51. Hygrogen bonding
    a special type of dipole-dipole interaction that occures when the lone pairs of an electonegative atom interact with an electron-poor hydrogen atom. Compounds that exhibit hydrogen bonding have higher boiling points than similare compounds that lack hygrogen bonding.
  52. debye
    1 debye = 10-18 esu  cm
  53. intermolecular forces
    The attractive forces between the individual molecules.
  54. hydrogen bond
    • a special type of dipole-dipole interaction that occurs  between an electronegative atom and a "H" atom that is connected to another electronegative atom.
    • Note: except when attached to F,O,N.
  55. Protic
    • any compound that has a proton connected to an electnegative atom.
    • EX: ethanol.
  56. Hydrocarbons
    Compounds that contain carbon and hydrogen and only have single bonds.
  57. London dispersion forces
    weak,
  58. Hydrophylic
    "loves water" 
  59. Hydrophobic
    "ware fearing"
  60. Soap
    Have a polar head and nonpolar tail (extnsive at times). compounds that will contain both hydrophylic and hydrophobic regions. The hydrophobic tails surrond nonpolar compounds, forming a water-soluble micelle.
  61. Micelle
    a ball of polar heads on the outside with nonpolar tails on the inside.
  62. Hydrogen deficiency index
    • [HDI] = [(2C + 2 + N - X) - H] / 2
    • Note: C = # of carbon, N = # of nitrogen, X = # of all halogens, H = # of hydrogens.
    • answers:
    • 0 = all single bonds
    • 1 = double bond or ring
    • 2 = Triple and/or double and/or enclosed (ring).
    • Note: there could be a double and a ring for an answer of 2 or any combination of the possabiliyies.
  63. What happens when you cross a node with the electon density?
    • The node itself means no density, after you cross it it will be the oppisite spin (charge).
  64. What are the electronegativity values for bonds?
    • Non-polar covalent:   < 0.5
    • polar covalent:           = .5 - 1.7
    • Ionic:                          = 1.7 and up
  65. Ionic bond
    Electrostatic attraction between ions of oppisite charges.
  66. Delta Electronegativity
    Delta EN is always a prediction, so we write WE PREDICT
  67. Steric #
    • the # of occupied sites around the atom.
    • # of atoms attached + # of lone pairs (or partial lone pairs).
    • Steric # shapes:
    • 2 = Linear
    • 3 = Bent, triaganol planer
    • 4 = Bent, triaganal paramidail, tetrahedral
  68. Excited states
    typiclly . . .1s' 2p3 . . . whenever the first half of an orbital is filled.

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