Steps for Lewis Structures 1 Count the total number of valence electrons 2 Select a central atom a The atom that is isolated Carbon the least electronegative 3 Attach each outside atom to the central atom with a single bond 4 Give every outside atom enough lone pairs for an octet up to the valence total from 1 5 Count the number of electrons used in the structure a Any excess electrons apply to the central atom 6 If the central atom does not have an octet 7 8 If still no octet apply a triple bond and check for resonance a Apply a double bond and check for resonance b c Be sure to remove lone pairs from the atom that donated the double triple bond If resonance draw all structures and choose the structure that minimizes formal charge If the structure is an ion a Positive charges mean subtract electrons from the valence total b Negative charges mean add electron to the valence total c Be sure you place brackets around the structure and notate its charge outside the bracket in the upper right hand corner 9 Common Exceptions a NO2 has a single electron instead of a lone pair and is a resonance structure However there cannot be 2 double bonds because N does not have enough space in his orbitals to hold more than 8 electrons b H atoms only have space to hold 2 electrons so they can only be bonded to one atom with a single bond and have no lone pairs c BF3 is a structure that could have a double bond and resonance but doesn t because of formal charge The octet for B is 6 electrons 10 Common Resonance Structures a Sulfur Trioxide b Nitrate ion c Carbonate ion d Nitrogen dioxide e Ozone Calculating Formal Charge Formal Charge Valence e of element 1 2 Bonding e 2 of lone pairs In other words 1 Determine the number of valence electrons that the atom should have normally 2 Calculate the number of electrons that are actually around it in the structure a Each lone pair is 2 e and each line of a bond is 1 b Think of it as how many e is the atom contributing to the structure If it s a S O double bond then O has 4 electrons that it owns and it shares 2 of them with the S in the double bond so it s formal charge is 6 4 2 0 Condition Bonding sites lone pairs Electron Domain Molecular Geometry Hybridization 2 0 3 0 3 1 4 0 4 1 4 2 5 0 5 1 5 2 6 0 6 1 6 2 Linear Linear Trigonal Planar Trigonal Planar Trigonal Planar Angular Bent Tetrahedral Tetrahedral Tetrahedral Trigonal Pyramidal Tetrahedral Bent Trigonal Bipyramidal Trigonal Bypyramidal Trigonal Bipyramidal Trigonal Bipyramidal Octahedral Octahedral Octahedral See Saw T Shaped Octahedral Square Pyramid Square Planar sp sp2 sp2 sp3 sp3 sp3 sp3d sp3d sp3d sp3d2 sp3d2 sp3d2 Notes Mol Geo 180 120 120 109 5 107 105 90 120 180 90 120 180 90 180 90 180 90 180 90 180 In the notes section of this table is the most likely common bond angles found in the molecular geometry You should notice that as the number of lone pairs increases the bond angle decreases This is the point of the VSEPR theory and model It is because electrons repel each other and a greater number of lone pairs results in more repulsion which decreases the angle between the bonded atoms Also you should notice that for the electron domain the only thing that is relevant is the number of bonding sites defined as a single double or triple bond lone pair or free radical and not the number of lone pairs In other words the electron domain will always be the same if the number of bonding sites is the same The molecular geometry however changes based on the number of lone pairs in relation to the number of bonding sites When reading the conditions section of this table you should read 4 1 as the Lewis structure has 4 bonding sites total one of which is a lone pair so the geometry would be trigonal pyramidal
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