Chem 1210 Exam 2

The flashcards below were created by user Anonymous on FreezingBlue Flashcards.

  1. Principal Quantum Number
    • (n) - size and energy of the orbital
    • N= 1, 2, 3, 4,
  2. Angular momentum quantum number
    (l) - shape of atomic orbitals (sometimes called a subshell) l=0,1,2,3...(highest number it can be= n-1)
  3. Magnetic quantum number
    (ml) - orientation of the orbital in space relative to the other orbitals in the atom. ex. l=0 then ml=0... l=1 then ml can = -1,0,1, etc.  max # of ml= 2l+1
  4. nodes
    • # of nodes= n-1 or l
    • Shell shaped area where electron can't be found
  5. S orbitals
    always spherical
  6. P orbitals
    dumbbell shaped (can have different orientation on x y and  axis)
  7. D orbitals
    clover shaped: can have different orientation on x y and z axis
  8. Pauli exclusion principle
    cannot have more than on electron with the same set of 4 numbers, therefore each shell can only hold two electrons with opposite spin.
  9. Polyelectronic Atoms
    atoms with more than one electron
  10. Electron correlation problem
    since the electron pathways are unknown, the electron repulsions cannot be calculated exactly
  11. Penetration effect
    • A 2s electron penetrates to the nucleus more than on in the 2p orbital
    • This causes an electron in a 2s orbital to be attracted to the nucleus more strongly than an electron in a 2p orbital.
    • Thus, the 2s orbital is lower in energy than the 2p orbitals in a polyelectronic atom.
  12. Aufbau Principle
    As protons are added to the nucleus to build up the elements, electrons are similarly added to hydrogen-like orbitals.
  13. Hund's Rule
    The lowest energy configuration for an atom is the one having the max number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate (same energy) orbitals.
  14. Ionization Energy
    • Energy required to remove an electron from a gaseous atom or ion.  
    • Ionization energy increases left to right and as you go up the periodic table.
    • (why? the electrons being removed are farther from the nucleus)
  15. Electron Affinity
    • Energy change associated with the addition of an electron to a gaseous atom.
    • Electron Affinity increases from left to right and going up the periodic table.
  16. Atomic Radius
    Atomic radius decreases going left to right and increases going down a group.
  17. Group 1A name
    Alkali Metals
  18. Group 2A name
    Alkaline earth metals
  19. Middle elements name
    transition elements
  20. Group 7A name
  21. Group 8A name
    Noble gases
  22. Alkali Metals
    Most chemically reactive of the metals
  23. amu
    • atomic mass unit
    • mass of an atom compared to 12C
  24. Avogadro's number
    a mole= 6.022x10^23
  25. Covalent Bonds
    bonds formed between atoms by sharing electrons
  26. Ionic Bonds
    Bonds formed due to force of attraction between oppositely charged ions.
  27. Ion
    atom or group of atoms that has a net positive or negative charge.
  28. Cation
    positive ion; lost electron(s)
  29. Anion
    negative ion; gained electron(s)
  30. Binary Compounds
    • Composed of two elements
    • ionic and covalent compounds included
    • 1. cation is named first and anion second
    • 2. anion is named by taking the root of the element name and adding -ide.
    • 3. type (II)- metals form more than one type of positive charge; charge on metal ion must be specified. (transition metal cations usually require a Roman numeral.)
  31. Polyatomic Ions
    multiple elements... just need to be memorized.
  32. Hg2^2+
  33. NH4(+)
  34. NO2(-)
  35. No3(-)
  36. SO3(2-)
  37. SO4(2-)
  38. HSO4(-)
    Hydrogen sulfate (or bisulfate)
  39. OH(-)
  40. CN(-)
  41. PO4(3-)
  42. HPO4(2-)
    Hydrogen phosphate
  43. H2PO4(-)
    Dihydrogen phosphate
  44. NCS(-) or SCN(-)
  45. CO3(2-)
  46. HCO3(-)
    Hydrogen Carbonate (or bicarbonate)
  47. ClO(-) or OCl(-)
  48. ClO2(-)
  49. ClO3(-)
  50. ClO4(-)
  51. C2H3O2(-)
  52. MnO4(-)
  53. Cr2O7(2-)
  54. CrO4(2-)
  55. O2(2-)
  56. C2O4(2-)
  57. S2O3(2-)
  58. Binary Covalent Compounds
    • Formed between two nonmetals
    • 1. first element is named first
    • 2. second element is named as an anion
    • 3. prefixes are used to denote the numbers of atoms present.
    • 4. prefis mono- never used for naming first element.
  59. Acids
    • Hydrogen appears first in the formula
    • Molecule with one or more H+ ions attached to an anion
  60. Naming Acids without oxygen
    acid is name with the prefix hydro- and the suffix ic
  61. Naming Acids with oxygen
    • suffix -ic is added to the root name if the anion name ends in -ate
    • suffix -ous is added to the root name if the anion name ends in -ite
  62. HF
    Hydrofluoric acid
  63. HCl
    hydrochloric acid
  64. HBr
    hydrobromic acid
  65. HI
    Hydroiodic acid
  66. HCN
    Hydrocyanic acid
  67. H2S
    hydrosulfuric acid
  68. HNO3
    Nitric acid
  69. HNO2
    nitrous acid
  70. H2SO4
    Sulfuric acid
  71. H2SO3
    Sulfurous acid
  72. H3PO4
    Phosphoric acid
  73. HC2H3O2
    Acetic acid
  74. Chemical Bond
    Forces that hold groups of atoms together and make them function as a unit.
  75. Polar covalent Bond
    • unequal sharing of electrons between atoms in a molecule. 
    • Results in a charge separation in the bond
  76. Electronegativity
    • ability of an atom in a molecule to attract shared electrons to itself.
    • electronegativity increases left to right and up the periodic table.
    • values range from 0.7 to 4.0
  77. Difference in electronegativity between atoms
    • 0-0.4 : Covalent bond
    • 0.4-2.0 : Polar Covalent
    • 2.0-larger: ionic
  78. Dipole Moment
    • Property of a molecule whose charge distribution can be represented by a center of positive charge and a center of negative charge.
    • Arrow represents dipole moment (points to the negative charge center)
  79. Electron configurations in Stable Compounds
    • two nonmetals react to form covalent bond and share electrons so that it completes the valence electron configuration of both atoms
    • nonmetal and representative-group metal react to form binary ionic compound, form so that valence electron config. of the nonmetal achieves the electron configuration of the next noble gas atom. valence orbitals of the metal are emptied.
  80. Isoelectronic series
    series of ions/atoms containing the same number of electrons.
  81. Bond Energies
    • To break bonds, energy must be added to the system (endothermic)
    • To form bonds, energy is released (exothermic)
  82. Lattice Energy
    • the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid.
    • Lattice Energy= k(Q1Q2/r)
    • Q1 and Q2= charges on the ions
    • R=shortest distance between the centers of the cations and anions
    • k=proportionality constant
  83. Ionic compound characteristic
    any compound that conducts and electric current when melted will be classified as ionic.
  84. Localized Electron Model
    • Electron pairs are assumed to be localized on a particular atom or in the space between two atoms:
    •     Lone pairs- electrons localized on an atom
    •     Bonding pairs- pairs of electrons found in the space between the atoms.
  85. Lewis Structure
    Shows how valence electrons are arranged among atoms in a molecule.
  86. Duet Rule
    Hydrogen forms stable molecules where it shares two electrons
  87. Octet Rule
    Elements form stable molecules when surrounded by eight electrons
  88. Single Covalent Bond
    covalent bond in which two atoms share one pair of electrons
  89. Double covalent bond
    covalent bond in which two atoms share two pairs of electrons
  90. Triple Covalent Bond
    covalent bond in which two atoms share three pairs of electrons
  91. Steps for writing Lewis Structures
    • 1. sum the valence electron from all atoms
    • 2. use pair of electrons to form a bond between each pair of bound atoms.
    • 3. Arrange the remaining electrons to satisfy the octet rule (or duet for hydrogen)
  92. Boron exception to octet rule
    tends to form compounds in which the boron atom has fewer than eight electrons around it.
  93. Exceeding octet rule
    sometimes necessary to exceed octet rule for several of the third-row(or higher) elements, place the extra electrons on the central atom.
  94. Rsonance
    More than one valid Lewis structure can be written for a particular molecule.
  95. Formal Charge
    • used to evaluate nonequivalent lewis structures
    • atoms try to achieve formal charges as close to zero as possible.
    • Formal Charge= (# valence e on free atom) - (# valence e assigned to the atom in the molecule)
    • lone pairs- count both of the electrons
    • bonds- count one of the electrons
  96. VSEPR Model
    valence Shell Electron-Pair Repulsion: the structure around a given atom is determined principally by minimizing electron pair repulsions.
  97. Electronic Geometry arrangements
    • # of electron pairs         Arrangement of pairs
    •        2                                    Linear
    •        3                                 Trigonal Planar
    •       4                                   Tetrahedral
    •       5                             Trigonal bipyramidal
    •       6                                    Octahedral
  98. Hybridization
    mixing of the native atomic orbitals to form special orbitals for bonding.  examples... sp, sp2, sp3, dsp3, etc
  99. Sigma bond
    electron pair is shared in an area centered on a line running between the atoms
  100. Pi Bond
    forms double and triple bonds by sharing electron pair(s) in the space above and below the sigma bond. Uses the unhybridized p orbitals.
  101. Using the localized Electron Model
    • 1. Draw the Lewis structure(s)
    • 2. Determine the arrangement of electron pairs using the VSEPR model.
    • 3. Specify the hybrid orbitals needed to accommodate the electron pairs.
Card Set:
Chem 1210 Exam 2
2014-10-20 02:00:42
Chem 1210

Exam 2
Show Answers: