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Level 52

Molecular Orbitals


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Linear
2 Bonds
Linear diatomic
CN=2 LP-1
trigonal planar
a molecule whose shape is triangular and in one plane; ex: BH3, BF3
Bentangle
CN=3 LP-1
tetrahedral
3 dimensional shape showing one central atom with 4 atoms coming off of it. Tetra means four!
trigonal pyramidal
4 electron groups with lone pair
Bent angle
CN=4 LP-2
trigonal bipyramidal
Bond Angle - 90°, 120° and 180°
Seesaw
4 bonds, 1 pair
T-shaped
Bond Angle - 90° and 180°
octahedral
Bond Angle - 90° and 180°
Square Pyramid
5 Bonds, 1 pair
Square Planar
4 bonds, 2 pairs
CN 2
Linear - sp
CN 3
Trigonal planar - sp²
CN 4
Tetrahedral - sp³
CN 5
Trigonal bipyramidal - sp³d
CN 6
Octahedral - sp³d²
hybrid orbitals
_______ are eqivalent because they have the same size, shape, and energy
Coordination number (CN)
Number of atoms bound to atom of interest and number of lone pairs of electrons on the atom of interest
CN alone
Electronic geometry and hybridization
Molecular shape
CN + LP
Formal charge
(Valence electrons) - (# of bonds + # of electrons)
Total FC's
Want an even distribution of charge, want (-) on most electronegative atoms, (+) on least electronegative atoms.
Obey octet rule
C, N, O, F, Ne
Metallic radius
One-half the shortest distance between nuclei of adjacent, individual atoms in a crystal of the element
Covalent radius
One-half the shortest distance between nuclei of bonded atoms
Atomic size decreases
As the effective nuclear charge (Zeff) increases, outer electrons are pulled closer to the nucleus
Atomic size increases
As the principal quantum number (n) increases, the probability that outer electrons spend most of their time farther from the nucleus increases as well
Ionization Energy
Energy needed to remove an electron from the valence shell of an atom - how easily an atom can become an ion (COMPARE REACTIVITY OF A METAL)
Electron Affinity
Energy emitted upon addition of an electron - tendency to gain an electron (COMPARE REACTIVITY OF A NONMETAL)
Isoelectronic
Atoms having the same electronic configuration
Amphoteric
A substance that can act as both an acid and a base
Ionic radius
A measure of the size of an ion and is obtained from the distance between the nuclei of adjacent ions in a crystalline ionic compound
Ionic bonding
Bonding between metals and nonmetals
Covalent bonding
Bonding between nonmetals
Metallic bonding
Bonding between metals with electron pooling. Electrons are delocalized, moving freely throughout the entire piece of metal
Lattic energy
The enthalpy change that accompanies the reverse of the previous equation- 1 mol of ionic solid separating into gaseous ions
Coulomb's Law
A mathematical formula whose consequence is that negatively and positively charged particles attract each other and similarly charged species repel each other.
Electrostatic energy
(cation charge*anion charge)/ (cation radius + anion radius)
Bond Order
# of bonding electons-# of antibonding electrons/2 *More electrons in antibonding orbitals leads to a less stable molecule, which means it is weaker and is less likely to exist.
Bond energy (BE)
The energy needed to overcome this attraction and is defined as the standard enthalpy change for breaking the bond in 1 mol of gaseous molecules
Electronegativity
a measure of the ability of an atom in a chemical compound to attract electrons
Paramagnetic
A substance that is attracted to a magnetic field
Valence-shell electron-pair repulsion theory
To minimize repulsions, each group of valence electrons around a central atom is located as far as possible from the others.
Dipole moment
A measure of molecular polarity
Valence bond theory
A covalent bond forms when orbitals of two atoms overlap and a pair of electrons occupy the overlap region
sigma bond
a bond formed when two atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting the two atomic nuclei
pi bond
A bond formed when parallel p orbitals overlap creating two regions of electron density, one above and one below the internuclear axis.