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76 Cards in this Set
- Front
- Back
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orbital |
probability distribution map showing where an electron is likely to be found |
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n |
principal quantum number overall size/energy n= 1,2,3,... |
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l |
angular momentum number shape l= n-1 |
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ml |
magnetic quantum number orientation of orbital -l to l |
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ms |
spin quantum number orientation of spin + or - 1/2 |
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values of l |
0- s 1- p 2- d 3- f |
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each combination of n, l, and ml is |
one atomic orbital (two electrons) |
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same n |
in same principal shell |
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same n and l |
in same sub-shell |
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energy of an orbital |
-2.18 x 10^-18 J |
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change in energy |
-2.18 x 10^-18 J (1/n^2 f - 1/n^2 i) |
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orbitals on periodic table |
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periodic properties |
based on elements location |
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modern periodic table |
mendeleev |
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periodic law |
when the elements are arranged in order of increasing mass, certain sets of properties recur periodically |
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main group elements |
properties largely predictable |
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transition elements |
properties less predictable |
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family/group |
column |
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period |
row |
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electron configuration |
particular orbitals that electrons occupy for that atom |
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ground state |
lowest energy state 1s1 |
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orbital diagram |
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pauli exclusion act |
no two electrons in an atom can have the same four quantum numbers same orbital, opposite spins |
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degenerate |
same energy ex: 3s,3p,3d |
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coulomb's law |
potential energy of two charged particles depends on their charges and on their separation |
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shielding |
innermost orbital shields outer electrons from force of nucleus |
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effective nuclear charge |
net charge experienced by electrons atomic number - number of shielding electrons |
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bc of penetration |
1) sublevels aren't degenerate 2) 4s is lower in energy than 3d |
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aufbau principle |
electrons fill the lowest available energy levels before filling higher levels
ex: 1s before 2s |
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hund's rule |
when filing degenerate orbitals, electrons fill them singly first |
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valence electrons |
last s sub-shell and any f and d sub-shells that aren't filled |
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s sub-shell |
2 electrons |
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p sub-shell |
6 electrons |
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d sub-shell |
10 electrons |
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f sub-shell |
14 electrons |
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group number |
number of valence electrons |
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row number |
highest principle quantum number |
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metals |
lower left, lose electrons |
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non-metals |
upper right, gain electrons |
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group 1A |
alkali metals, 1+ |
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group 2A |
alkali earth metals, 2+ |
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group 7A |
halogens, 1- |
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group 8A |
noble gases, stable |
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paramagnetic |
contains unpaired electrons |
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diamagnetic |
contains only paired electrons |
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covalent bonding |
nonmetals (molecular) |
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metallic bonding |
metals |
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ionic bonding |
metal and nonmetal (cation and an anion) |
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size of atoms |
cations < neutral atoms < anions |
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ionization energy (IE) |
energy required to remove an electron from the atom or ion in the gaseous state |
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size exceptions |
boron is smaller than beryllium aluminum is smaller than gallium |
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electron affinity (EA) |
normally negative bc releases energy when gaining an electron |
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metallic character increases |
moving down the periodic table |
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empirical formula |
relative number of atoms (HO) |
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molecular formula |
actual number of atoms (H2O2) |
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duet rule |
stable lewis structure with two dots ex: He, H |
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octet rule |
most stable electron configurations contain eight electrons |
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element (III) |
chromium, iron, cobalt, copper, tin, mercury, lead -ous < -ic |
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hydrates |
contain specific number of water molecules |
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hemi |
1/2 |
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mono |
1 |
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di |
2 |
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tri |
3 |
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tetra |
4 |
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penta |
5 |
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hexa |
6 |
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hepta |
7 |
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octa |
8 |
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nona |
9 |
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deca |
10 |
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single bond |
1 electron pair 2 electrons |
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double bond |
2 electron pairs 4 electrons |
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triple bond |
3 electron pairs 6 electrons |
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formula mass |
average mass of a molecule =(# of atoms in 1st element x atomic mass of 1st element) + (# of atoms in 2nd element x atomic mass of 2nd element) |
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composition of compounds |
mass % of element = mass of element in 1 mol/ mass of 1 mol of compound x100% |
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periodic table trends
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