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30 Cards in this Set
- Front
- Back
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Valence electrons
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are the outer shell electrons of an
atom. The valence electrons are the electrons that participate in chemical bonding. group number= #of valence e- |
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Electron configurations of Ions of the Representative Elements
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These are derived from the electron configuration of elements with the required number of electrons added or removed from the most accessible orbital.
Electron configurations can predict stable ion formation |
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covalent bond
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is a chemical bond in which two or more electrons are shared by two atoms.
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lengths of covalent bonds
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Triple bond < Double Bond < Single Bond; go in order of atomic radii size (dec right, inc down)
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polar-covalent
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In many cases the electrons are shared, but not equally.
These bonds are called polar-covalent and are considered to be partially positively charged and partially negatively charged. EX: HCl |
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Electronegativity
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the relative ability of a bonded atom to
attract the shared electrons in a chemical bond; inc right, dec down like bond energy; -will most readily GAIN an e- -most negative e- affinity -greatest attraction for an e- in a covalent bond -most polar; highest diploe |
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formal charge
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the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure; formula: FC= (#valence e-)-(#nonbonding e-) - 1/2(#bonding e-); it is the charge an atom would have if the bonding electrons were shared equally; The formal charges must add up to equal the actual charge on
the species; the lowest one is correct although zeros overide negatives; IF THE SING ON THE MOLECULE IS NEGATIVE AND 2 ANS YIELD THE SAME RESULTS, YOU LOOK FOR THE MOST ELECTRONEG OF THE TWO. IF SIGN OF MOLECULE IS POS, YOU LOOK FOR LEAST ELECTRONEG OF THE TWO |
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Resonance
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the use of two or more Lewis
structures to represent a particular molecule or ion; connected by a double-headed arrow. The true structure is a blend of the resonance structures; Resonance indicates that electron delocalization occurs in the molecule; one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. |
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Are covalent bonds weaker than ionic bonds?
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In covalent bonding, electron sharing leads to strong, localized bonds. But most compounds with covalent bonds are gases, liquids, or low-melting solids. Ionic compound are high-melting solids; Within a molecule there are strong intramolecular bonds between two atoms; Between molecules there are weak intermolecular forces;
The latter are what determine melting and boiling points. |
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bond energy
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The enthalpy change required to break a particular bond in one mole of gaseous molecules; Single bond < Double bond < Triple bond; same as electronegativity (inc right, dec down)
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the relationship between bond length, bond strength, and bond energy is :
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The shorter the bond length,
the greater the bond strength, the greater the bond energy! VS The longer the bond length, the lower the bond strength, the lower the bond energy! |
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Valence Shell Electron-Pair Repulsion
(VSEPR) Theory |
Each group of valence electrons around a
central atom is located as far away from the others as possible in order to minimize repulsion. An electron group may be a single, double, or triple bond or a lone pair. |
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molecular shape
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is defined by the relative positions of the atomic nuclei. If there are no lone pairs of electrons the molecular shape is the same as the electron-group arrangement.
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Valence Bond Theory
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a covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the nuclei. Usually this means that each bonding orbital should contain one electron. The electrons must have opposite spins; The bond strength depends on the attraction of nuclei for the shared electrons, so the greater the orbital overlap, the stronger the bond. |
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Hybridization
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is the mixing of atomic orbitals in an atom to generate a set of new orbitals called hybrid orbitals; The total number of orbitals doesn’t change.
Covalent bonds are formed by: Overlap of hybrid orbitals with atomic orbitals or Overlap of hybrid orbitals with other hybrid orbitals |
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sigle bond
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sigma bond
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dbl bond
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sigma + pi
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triple bond
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sigma + 2pi
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Delocalized electrons
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are not confined between two adjacent bonding atoms, but actually extend over three or more atoms; Structures with delocalized electrons always have greater stability than similar structures without delocalized electrons; resonance indicates that this occurs.
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Molecular orbital theory
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bonds are formed from interaction of atomic orbitals to form molecular orbitals.
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bonding molecular orbital
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has lower energy and greater stability than the atomic orbitals from which it was formed
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antibonding molecular orbital
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has higher energy and lower stability than the atomic orbitals from which it was formed
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axial
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distance between elements in the central atom plane. for trigonal bipyramidal, its 120. for octahedral, its 90
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equatorial
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distance from atom in central atom plane, to a point above or below that plane. in trigonal bipyramidal, its 90. in octahedral, its 90.
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stability
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the higher the bond order, the more stable the species. formula: 1/2(bonding e- - antibonding e-); energy inc as you go up. the top one (antibonding one w/ the star) has more energy and less stability than the lower one.
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exceptions to octet rule
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S, P, B, Be, Xe, N
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when you dont know what the central atom is...
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make it carbon
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atomic radiius
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(determines lenghts of covalent bonds)
-most readily lose and e- -least affinity for e- |
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intramolecular bonds
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Within a molecule (between 2 atoms) there are strong intra bonds.
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intermolecular
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Between molecules there are weak intermolecular forces.
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