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What is the octet rule?
What are the exceptions? |
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Definition
The octet rule states than an atom tends to bond with others atoms so that it has 8 electrons in its outermost shell.
Incomplete octet-elements stable with fewer than 8 electrons: Hydrogen (2), Helium (2), lithium (2), beryllium (4), boron (6)
expanded octet-an element in period 3 and great can hold more than 8 electrons
odd numbers of electrons-any molecule with an odd number of valence electrons can't distribute them evenly to give 8 to each element |
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| electrons are transferred from one element, typically a metal (low ionization energy), to another element, typically nonmetal (high electron affinity). The resulting force holding them together is between the opposite charges. This bonding creates lattice structures. |
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| an electron pair is shared between two atoms. Nonpolar is equal sharing. Polar is unequal sharing. |
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| a covalent bond in which both electrons come from the same atom. |
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| atom that loses an electron and is positive |
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| atom that gains electrons and is negative |
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| characteristic physical properties of ionic compounds |
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-high melting point
-high BP
-aqueous state are good conductors
-ready dissolve in aqueous solvent
-form crystalline lattice |
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| distance between two nuclei; single bond is the longest/weakest and triple bond is the shortest/strongest |
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| energy needed to break bonds between components; the higher the energy the stronger the bond |
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| due to differences in electronegativity; forms a dipole |
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| vector quantity of a polar bond or polar molecule that represents the polarity of a dipole |
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| chemical symbol of an element surrounded by dots representing its s and p orbital electrons |
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| normal number of electrons in the atom's valence shell minus the number of nonbonding electrons minus half the number of bonding electrons. Formal charge underestimates teh effect of electronegativity in assuming perfect electron sharing. Oxidation numbers overestimate electronegativity in assuming 100% share to one atom |
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| Rules for assessing stability of resonance structures |
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Definition
-a lewis structure with small or no formal charges is preferred over a structure with large formal charges
-less separation between opposite charges is preferred over a Lewis structure with a large separation of opposite charges
-a lewis structure in which negative formal charges are placed on more electronegative atoms is more stable than one in which the negative formal charges are placed on less electronegative atoms |
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| Valence shell electron pair repulsion (VSEPR) theory |
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| theory that uses lewis dot structures to predict the molecular geometry of covalently bonded molecules |
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| describes the spatial arrangement of all pairs of electrons around the central atom, bonded or not; lone electron pairs take up more space than bonding electrons thus decreasing bond angles |
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| describes spatial arrangement of only the bonding pairs of electrons |
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| two bonded or nonbonded pairs of electrons (electronic geometry); 180 degree bond angle |
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| three bonded or nonbonded pairs of electrons (electronic geometry); 120 degree bond angle |
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| four bonded or nonbonded pairs of electrons (electronic geometry); 109.5 degree bond angle |
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| five bonded or nonbonded pairs of electrons (electronic geometry); 90, 120 and 180 degree bond angles |
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| six bonded or nonbonded pairs of electrons (electronic geometry); 90 and 180 degree bond angles |
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| consider vector addition of bond dipole moments when figuring overall molecular polarity; if all dipole moments cancel out molecule is nonpolar; if they don't cancel out there will be a net dipole moment equal to the sum of the vectors |
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| What are the quantum numbers |
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Definition
principal quantum number (n) n = 1,2,3... - main energy level or shell
angular momentum quantum number (l) l = 0...n-1 - referred to as subshells; l=0 s-orbital, l=1 p-orbital
magnetic quantum number (m(sub)l) ml = -l to +l - the orientation of the electron
spin quantum number (m(sub)s) ms = +1/2 or -1/2, up or down spin respectively |
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| london dispersion forces, dipole-dipole interactions, hydrogen bonds |
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| results from electron density being unequally distributed between two atoms causing tiny rapid polarization and depolarization |
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| dipole-dipole interactions |
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| interactions between polar and nonpolar parts of molecules |
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| when H bonds to N, O or F due to huge differences in electronegativity |
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