Chemistry Formulas Successive Ionization Energies Adding electrons gets increasingly difficult as each added electron reduces the elecetron affinity Cations stay under 3 charges Anions stay above 3 charges Polyatomic Ions Charged groups of covalently bound atoms Forms ionic bonds with oppositely charged ions nonmetals held together by covalent bonds Naming Oxoanions Central atom is in it s highest oxidation state Group 7A elements form four so prefix suffix must be used Protonated Oxoanions Oxoanions with 2 or 3 charges can be portonated H ion added and still retain negative charge Hydrogen in front of the anion name Common unit of bond lengths Most bond lengths fall between 1 and 3 Angstroms Bond Dipole Bonds in atoms with different electronegativities is said to have a bond dipole Represented by the vector pointing from the center of positive chrage toward the negative charge Percent Ionic Character Bond types change continuously from covalent to ionic 50 ionic ionic Between 5 50 ionic polar covalent Below 5 ionic covalent Lewis Symbols For the first three periods of the periodic table They represent an atom that is about to bond not an isolated atom Lewis Structures Atoms bond seeking to fill their octets Electrons not involved in bonding are called lone pairs How calculating Lewis Structures works Bond Energies and Lengths Shows how the bond strength of a particular bond increases and it s bond length decreases as it s bond order increases Strength decreases Length Increased Bond Order Increased Bond Strength Bond order number of shared pairs present Formal Charge FC Valence e lone pairs bonds Formal Charges of Zero An atom has a zero formal charge in a molecule if the of bonds it s in equals the of unpaired electrons in it s Lewis Symbol Zero formal charge is preferred Oxidation States Way to count electrons Aj 1 atom more electronegative than j Aj 0 atom less electronegative than j Aj 1 2 atoms are identical VSEPR Valence Shell Electron Pair Repulsion Model The regions of negative charge around an atom adopt positions that minimize the repulsions between them Bond Angles Angles formed by the intersection of the bonds Valence Bond TheoryEach bond results form the overlap of two or more atomic orbitals on two adjacent atoms Shows the overlap of the atomic orbitals in valence bond theory Shows bonding in O2 Bond order is the sum of and bonds All bonds only have one bond Close up of the internuclear axis Always a change of energy here Mixing functions leads to blue constructive regions and red destructive regions Mixing s and p orbitals gives sp hybrid orbitals Occupied by bonds and lps Sp orbitals 1s 1p 180 leaves 2 p orbitals availible to form bonds One triple or two double bonds can form 2 Sp orbitals 1s 2p 120 only 1 p orbital left to form bonds Double bond forms Sp3 orbitals 1s 3p 109 no p orbitals left for bonds only single bonds form of hybrid orbitals of electron groups Molecular Orbital TheoryPresents a truer picture of delocalization and explains the electronic structure within molecules Shows orbitals on different atoms mixing to produce delocalized bonds Bonding when the interacting lobes of the Aos have the same phase Antibonding when the interacting lobes of the AOs are opposite phases When opposite phase lobes interact electron density is annihilated Lobes of the same phase interect to increase electron density H2 is shown with a single bond as a stable molecule He2 is shown with the and orbitals filled which means there is no bond and it isn t a stable molecule The bond order of H H is 1 The bond order of He He is 0 which indicates it is nonbonding O2 molecules are paramagnetic deflected in a magnetic field which means that they contain unpaired electrons MO region is in the yellow AO is shown off to the side The 2p is the lowest energy MO and the 2p is the highest energy MO Remaining 2p orbitals interect to form a pair of 2p and 2p orbitals These pi orbitals hace the same energy