Exam 2

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  1. electronegativity
    the ability of an atom to attract TOWARD itself the pair of electrons in a chemical bond
  2. definition of covalent, polar covalent and ionic in terms of electronegativities:
    ionic: 1.7 < d [difference of more than 1.7]

    polar covalent: .4 < d ≤ 1.7 [greater than .4, or less than or = to 1.7]

    covalent: 0 < d ≤ .4 [less than or = to .4; 0 = pure covalent]

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    • -if a complex absorbs a particular color, it will have the appearance of whatever color is across from it on the wheel
    • -ex. if you determine that a complex absorbs photons in the orange range, you can conclude that the complex will look blue
    • -the color that we see is the complementary color of the color absorbed
  4. quantum numbers
    -n (principle): related to the size of the orbital; shell/energy level

    • -l (azimuthal): related to the shape of the orbital; subshell
    • l=0 s orbital (spherical)
    • l=1 p orbital (2 lobes)
    • l=2 d orbital (4 lobes)
    • l=3 f orbital (6 lobes)
    • l=4 g orbital

    -ml (magnetic): related to spatial orientation of the orbital; = -l ... +l; indicated number of orbitals in that sublevel

    -ms (spin): + 1/2 or -1/2
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    • Energy = planck's constant * frequency
    • = p.c. (3E8 m/s) / wvlngth
  6. l (azimuthal)
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  7. energy of an electron
    -the more DIFFICULT it is to remove an electron, the more negative the energy of an electron is (low shielding)

    - if an outer electron is easier to remove, it's said to have a HIGER (less negative) energy... (high shielding)

    -energy increases with l, due to both distance & shielding
  8. atomic orbitals
    regions of space where the probability of finding an electron about an atom is highest
  9. Pauli Exclusion Principle
    no two electrons within an atom may have the same 4 quantum numbers
  10. Aufbau Principle
    electrons will occupy the lowest energy orbital that they can enter
  11. Hund's Rule
    electrons in a subshell containing more than one orbital will spread out among the orbitals with as many of their spins in the same direction as possible
  12. in order of 'stability'
    filled, empty, 1/2 filled, anything else

    -PS: the attractive force felt by an outer electron is always less than that which would be felt if the nucleus were not shielded
  13. the 3d subshell is higher in energy than the 4s subshell so appears in the 4th period
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  14. rules for electron configuration
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    • basicaly you would have to get all the way down to 6s before 4f came into play
  15. Atomic Radii
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    • increases down, and to the left (due to shielding/nucleus size)
  16. ionization energies
    • energy required to REMOVE an electron from (gaseous) atom in ground state; increases up and to the right (standard)
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  17. electron affinity
    • the energy change that occurs when an electron is added to a (gaseous) atom: X(g) + e- --> X-(g)
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    • + EA(endothermic): requires energy; such an atom doesn't readily accept an electron

    -EA (exothermic): acceptance of electrons is energetically favorable
  18. electronegativity
    • a number that describes the relativity ability of an atom, when bonded, to attract electrons
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  19. oxidation numbers
    • H: +1
    • O: -2
    • F: -1
    • -same thing for all elements in those colums, but everything else has to be determined
    • ex. Cr2O7 -2
    • O 7 -2 -14 -14 + 2? = -2
    • Cr 2 ? 2?
    • Oxd #s: Cr: +6, O: -2
  20. Formal Charges
    group # - number of e- surrounding element (if bonded, split bond in 2, so it gets 1 electron)
  21. types of covalent bonds
    Covalent bonds are formed when atoms share electrons.

    • - Nonpolar covalent bonds have a symmetrical charge distribution
    • - Polar covalent bonds have an asymmetrical charge distribution

    • • If the atoms share 2 electrons a single covalent bond is formed.
    • • If the atoms share 4 electrons a double covalent bond is formed.
    • • If the atoms share 6 electrons a triple covalent bond is formed
  22. lewis dot structures
    • C: likes 4 bonds NO lone pairs
    • N: likes to triple bond, 1 lone pair
    • O: likes to double bond, 2 lone pairs
    • F: likes to single bond, 3 lone pairs
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    • Ionic Compounds Without a Transition Metal:
    • -the metal ion's name does not change regardless of charge; the non-metal's name ends in ide
    • -ex. calcium phosphide = Ca3P2
    • Polyatomic Ions
    • -when metals are bonded to polyatomic ions (consist of two or more atoms with one overall charge), the same rules apply, but you have to learn the names and charges of common polyatomic ions
    • -ex. Aluminum sulfate; has Al+3 and SO4-2; = Al2(SO4)3

    • Ionic Compounds With a Transition Metal
    • -the only difference here is that we have to specify the charge of the transition metal ion by using a Roman numeral
    • -ex. Copper(I) oxide = Cu2O
  24. NAMING Covalent Compounds
    these are formed from non-metals that share electrons. Because there are many sharing possibilities between two non-metals, the formula cannot be guessed unless we have a naming system that reveals the number of atoms involved.

    -mono=1, di=2, tri=3, tetra=4, penta=5, hexa=6, hepta=7

    -the only time we drop a prefix is if the mono is to appear at the beginning of the name

    • -ex. N2O5 = dinitrogen pentoxide
    • PCl3 = phosphorus trichloride

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  25. dipole moments
    • think of them as vectors, arrows point to - elements, + sides to + charged
    • - if they don't cancel, there's asymmetrical charge distribution = dipole moment
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  26. hybridization
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  27. Common Negative Ions (Anions)
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Card Set:
Exam 2
2012-06-11 10:50:29
Gen Chem

Models of Chemical Bonding
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