Chem 261

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Chem 261
2012-02-04 15:20:05

Ch 1-3
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  1. Anti conformation
    The geowmetric arrangement around a carbon-carbon single bond, in which the two largest substituents are 180 apart as viewed in a Newman projection
  2. Antibonding orbital
    A molecular orbital that is higher in enegry than the atomic orbitals from which it is formed
  3. Aufbau principle
    A guide for determing the ground-state electronic configuration to the rough plane of the ring (See Equatorial bond)
  4. Boat cyclohexane
    A three-dimensional conformation of cyclohexane that bears a slight resemblance to a boat. Boat cyclohexane has no angle starin but has a large number of eclipsing interactions that make it less stable than chair cyclohexane.
  5. Bond angle
    The angle formed between two adjacent bonds.
  6. Bond Length
    The equilibrium distance between the nuclei of two atoms that are bonded to each other.
  7. Bonding orbital
    A molecular orbital that is lower in energy than the atomic orbitals from which it is formed.
  8. Bronsted acid
    A substance that donates a hydrogen ion (proton) to a base.
  9. Chair cyclohexane
    A three-dimensional conformation of cyclohexane that resembles the rough shape of a chair. The chair form of cyclohexane, which has neither angle strain nor eclipsing strain, represents the lowest-energy conforamtion of the molecule
  10. Cis-trans isomers
    Certain kinds of steriosomers that differ in their stereochemistry about a double bond or on a ring. Cis-trans isomers are also called geometric isomers.
  11. Conformation
    The exact three-dimensional shape of a molecule at any given instant, assuming that rotation around single bonds is frozen.
  12. Conformational analysis
    A means of assessing the minimum energy conformation of a substituted cycloalkane by totalling the steric interactions present in the molecule. Confirmational analysis is particularly useful in assessing the relative stabilities of different conformations of substituted cyclohexane rings.
  13. Conjugate acid
    The product that results from protonation of a Bronsted acid.
  14. Conjugate base
    The anion that results from dissociation of a Bronsted base.
  15. Constitutional isomers
    Isomers taht have their atoms connected in a different order. For example, butane and 2-methylpropane are constitutional isomers.
  16. Covalent bond
    A bond formed by sharing electrons that have the same energy level.
  17. Degenerate orbitals
    Two or more orbitals that have the same energy level.
  18. Delocalization
    A spreading our of electron density over a conjugated pi electron system. For example, allylic cations and allylic anions are delocalized because their charges are spread out by resonance stabilization over the entire pi electron system.
  19. Dipole moment
    A measure of the polarity of a molecule. A dipole moment arises when the centers of gravity of positive and negative charges within a molecule do not coincide.
  20. Eclipsed conformation
    The geometric arrangement around a carbon-carbon single bond in which the bonds to substituents on one carbon are parellel to the bonds to substituents on the neighboring carbon as viewed in a Newman projection. For example, the eclipsed conformation of ethan has the C-H bonds on one carbon lined up with the C-H bonds on the neighboring carbon.
  21. Electron Affinity
    The measure of an atom`s tendency to gain an electron and form an anion. Elements on the right side of the periodic table such as the halogens have higher electron affinities than do elements on the left side.
  22. Electronegativity
    The ability of an atom to attract electrons and thereby polarize the bond. As a general rule, electronegativity increases in going across the periodic table from right to left and in going from bottom to top.
  23. Emperical formula
    A formula that gives the relative proportions of elements in a compound in smallest whole numbers (see molecular formula)
  24. Equatorial bond
    A bond to cyclohexane that lies along the rough equator of the ring (see axial bond).
  25. Exothermic
    A term used to describe reactions that release heat and that therefore have negative enthalpy changes. On reaction diagrams, the products of exothermic reactions have energy levels lower than those of starting materials.
  26. Formal charge
    The difference in the number of electrons owned by an atom in a molecule and by the same atom in it`s elemental state. The formal charge on an atom is given by the formula:

    FC=# of outer shell e in free atom - #of outer shell e in bonded atom.
  27. Functional group
    An atom or group of atoms that is part of a larger molecule and has a characteristic chemical reactivity. Functional groups display the same chemistry in all molecules of which they are part.
  28. Gauche conformation
    The CH3 conformation of butane in which the two CH3 methyl groups lie 60o apart as viewed in a Gauche conformation Newman projection. This conformation has 0.9 k/mol steric strain.
  29. Geometric isomers
    An old term for cis-trans isomers.
  30. Ground state
    The most stable, lowest-energy electronic configuration of a molecule.
  31. Hybrid orbital
    An orbital that is mathematically dervided from a combination of ground-state (s, p, d) atomic orbitals. Hybrid orbitals, such as the sp3, sp2, and sp hybrids of carbons, are strongly directed and form stronger bonds than groud-state atomic orbitals do.
  32. Hydrogen bond
    A weak (5kcal/mol) attraction between a hydrogen atom bonded to an electronegative element and an electron lone pair on another atom. Hydrogen bonding plays an important role in determing the secondary structure of proteins and in stablizing the DNA double helix.
  33. Inductive effect
    The electron-attracting or electron-withdrawing effect that is transmitted through sigma bonds as teh result of a nearby dipole. Electronegative elements have an electron-withdrawing inductive effect, whereas electropositive elements ave an electron-donating inductive effect.
  34. Ionic bond
    A bond between two ions due to the electrical attraction of unlike charges. Ionic bonds are formed between strongly electronegative elements (such as the holgens) and strongly electropositive elements (such as the alkali metals).
  35. Ionization energy
    The amount of energy required to remove an electron from an atom. Elements on the far right of the periodic talbe have high ionization energies, and elements on the far left of the periodic table have low ionization energies.
  36. Isomers
    Compounds that have the same molecular formula but have different structures.
  37. Lewis acid
    A substance having a vacant low-energy orbital that can accept an electron pair from a base. All electropiles are Lewis acids, but transition metal salts such as AlCl3 and ZnCl2 are particularly good ones. (see Lewis base)
  38. Lewis base
    A substance that donates an electron pair to an acid. All nucleophiles are Lewis bases. (see Lewis acid)
  39. Line-bound structure
    A representation of a molecule showing covalent bonds as lines between atoms (see Kekule structure).
  40. Lone-pair electrons
    Nonbonding electron pairs that occupy valence orbitals. It is the lone-pair electrons that are used by nucleophiles in their reaction with electrophiles.
  41. Mechanism
    A complete description of how a reaction occurs. A mechanism must account for all starting materials and all products, and must describe the details of each individual step in the overall reaction process.
  42. Molecular formula
    An expression of the total numbers of each kind of atom present in a molecule. The molecular formula must be a whole-number multiple of the empirical formula.
  43. Molecular orbital
    An orbital that is the property of the entire molecule rather than of an inidividual atom. Molecular orbitals result from overlap of two or more atomic orbitals when bonds are formed and may be either bonding, nonbonding, or antibonding. Bondind molecular orbitals are lower in energy than the starting atomic orbitals, nonbonding MO's are equal in energy to the starting orbitals, and antibonding orbitals are higher in energy.
  44. Newman projection
    A means of indicating stereochemical relationships between substituent groups on neighboring carbons. The carbon-carbon bond is viewed end-on, and the carbons are indicated by a circle. Bonds radiating from the center of the circle are attached to the front carbon, and bonds radiating from the edge of the circle are attached to the rear carbon.
  45. Node
    The surface of zero electron density between lobes of orbitals. For example, a p orbital has a nodal plane passing through the center of the nuclues, perpendicular to the line of the orbital.
  46. Normal alkane
    A straight-chain alkane, as opposed to a branched alkane. Normal alkanes are denoted by the suffix n, as in n-C4 H10 (n-butane).
  47. Polarity
    The unsymmetrical distribution of electrons in molecules that results when one atom attracts electrons more strongly than another.
  48. Polarizability
    The measure of the change in a molecule's electron distribution in response to changing electric interaction with solvents or ionic reagents.
  49. Primary, secondary, tertiary, quaternary
    Terms used to describe the substitution pattern at a specific site. A primary site has one organic substituent attached to it, a secondary site has two organic substituents, a teriary site has three, and a quaternary site has four.
  50. Principle of maximum overlap
    The stronest bonds are formed when overlap between orbitals is greatest.
  51. Ring-flip
    The molecular motion that converts one chair conformations of cyclohexane into another chair conformation. The effect of a ring-flip is to convert an axial substituent into an equatorial substituent.
  52. Saturated
    A saturated molecule is one that has only single bonds and thus can't undergo addition reactions. Alkanes, for example, are saturated, but alkenes are unsaturated.
  53. Sawhorse structure
    A stereochemical manner of representation that portrays a molecule using a stick drawing and gives a perspective view of the conformation around single bonds.
  54. Sigma bond
    A covalent bond formed by head-on overlap at atomic orbitals.
  55. sp orbital
    A hybrid orbital mathematically derived from the combination of an s and a p atomic orbital. The two sp orbitals that result from hybridization are oriented at an angle of 180o to each other.
  56. sp2 orbital
    A hybrid orbital mathematically derived by combination of an s atomic orbital with two p atomic orbitals. The three sp2 hybrid orbitals that result lie in a plane at angles of 120o to each other.
  57. sp3 orbital
    A hybrid orbital mathematically derived by combination of an s atomic orbital with three p atomic orbitals. The four sp3 hybrid orbitals that result are directed toward the corners of a tetrahedron at angles of 109o to each other.
  58. Staggered conformation
    The three-dimensional arrangement of atoms around a carbon-carbon single bond in which the bonds on the one carbon exaclty bisect the bond angles on the second carbon as viewed end-on. (see eclipsed conformation)
  59. Stereoisomers
    Isomers that have their atoms connected in the same order but have different three-dimensional arrangements. The term does not include constitutional isomers.
  60. Steric strain
    The strain imposed on a molecule when two groups are too close together and try to occupy the same space. Steric strain in responsible both for the greater stability of trans versus cis alkenes and for the greater stability of equatorially substituted versus axially substituted cyclohexanes.
  61. Torsional strain
    The strain in a molecule caused by electron repulsion between eclipsed bonds. Torsional strain plays a major role in destabilizing boat cyclohexane relative to chair cyclohexane. (see eclipsing strain)
  62. Unsaturated
    An unsaturated molecule is one that has multiple bonds and can undergo addition reactions. Alkenes and alkynes, for example, are unsaturated. (see saturated)
  63. Van der Waals forces
    The attractive forces between molecules that are caused by dipole-dipole interactions. Van der Waals forces are one of the primary forces responsible for holding molecules together in liquid and solid states.
  64. Wave function
    The mathematical expression that defines the behavior of an electron. The square of the wave function is the probability function that defines the shapes of orbitals.