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MCAT Gen. Chem Ch. 3

Octet Rule Elements will be most stable with eight valence electrons.
Elements With An Incomplete Octet And Are Stable: H, He, Li, Be, and B
Elements With An Expanded Octet And Are Stable: Elements in Period 3 or greater
Compounds With An Odd Number Of Electrons Cannot Have: Eight electrons on each element
Ionic Bond Formed via the transfer of one or more electrons form an element with a relatively-low ionization energy to an element with a relatively high electron affinity
Ionic Bonds Occur Between Elements With: Large electronegativity (Delta EN > 1.7). This occurs usually between metals and nonmetals.
Cation Positively charged ion
Anion Negatively charged ion.
Crystalline Lattices Formed by ionic compounds. These are large, organized arrays of ions.
Ionic Compounds Have Unique: Physical and chemical properties.
Ionic Compounds Tend To: Dissociate in water and other polar solvents.
Ionic Solids Tend To: Have high melting points.
Covalent Bond Formed via the sharing of electrons between two elements of similar electronegativities.
Bond Order Refers to whether a covalent bond is a single bond, double bond, or triple bond.
As Bond Order Increases: Bond strength increases, bond energy increases, and bond length decreases.
Covalent Bonds Can Be Categorized As: Nonpolar or polar depending on the nature of the elements involved
Nonpolar Bonds Result in molecules in which both atoms have exactly the same electronegativity.
Some Bonds Are Considered Nonpolar When: There is a small difference in electronegativity between the atoms (Delta EN < 0.5)
Polar Bonds Form when there is a sig. difference in electronegativities (Delta EN = 0.5 to 1.7) but not enough to transfer electrons and form an ionic bond.
In A Polar Bond, The More Electronegative Element Takes On: A partial negative charge
In A Polar Bond, The Less Electronegative Element Takes On: A partial positive charge
Coordinate Covalent Bonds Result when a single atom provides both bonding electrons while the other atom does not contribute any
Coordinate Covalent Bonds Are Most Often Found In: Lewis acid-base chemistry.
Lewis Dot Symbols Chemical representation of an atom's valence electrons
Formal Charges Exist when an atom is surrounded by more or fewer valence electrons than it has in its neutral state
Resonant Structures Represent all of the possible configurations of electrons (stable and unstable) that contribute to the overall structure
Valence Shell Electron Pair Repulsion (VSEPR) Theory Predicts the 3-D molecular geometry of covalently bonded molecules
Nonbonding Electrons Exert More Repulsion Than Bonding Electrons Because: They reside closer to the nucleus.
Electronic Geometry Position of all electrons in a molecule, whether bonding or nonbonding.
Molecular Geometry Refers to the position of only the bonding pairs of electrons in a molecule.
Polarity Of Molecules Dependent on the dipole moment of each bond and the sum of the dipole moments in a molecular structure.
Nonpolar Molecules May Contain: Nonpolar bonds, or polar bonds with dipole moments that cancel each other
Sigma Bonds Result of head-to-head bond overlap.
Pi Bonds Result of the overlap of two parallel electron cloud densities.
Intermolecular Forces Electrostatic attractions between molecules. They are significantly weaker than covalent bonds.
London Dispersion Forces Weakest interactions but are present in all atoms.
As The Size Of An Atom Or Structure Increases: London dispersion force increases
Dipole-dipole Interactions Occur between the oppositely charged ends of polar molecules, and are stronger than London forces.
Dipole-dipole Interactions Are Evident In: Solid and liquid phases but are negligible in the gas phase due to the distance between particles
Hydrogen Bonds Specialized subset of dipole-dipole interactions involved in intra- and intermolecular attraction
Hydrogen Bonding Occurs When Hydrogen Is Bonded To: One of three very electronegative atoms, NOF: Nitrogen, Fluorine, Oxygen
Eq. 3.1: Dipole Moment p = qd. p = dipole moment. q = mag. of charge. d = displacement vector between two partial charges.
Eq. 3.2: Formal Charge Formal Charge = V - N nonbonding - 1/2 N bonding. V = normal number of electrons in the atom's valence shell. N = number of electrons.
Created by: SamB91
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