Covalent Lewis Structures

1. Covalent Lewis Structures • Add up the valence electrons. • Put the element with the fewest number of atoms in the middle and surround with the remaining elements. • Draw a line from the central atom to the adjacent atoms • Subtract two electrons for each line drawn. • Make octets as needed giving the most electronegative atoms electrons first. B only needs 6, H &He 2 e’s. P, S, Cl, Argon and everything below can violate the octet rule and have more than 8 e’s if you must. • If needed form double and triple bonds. • Assign formal charges to atoms. Formal Charge = # of valence electrons – (# of dots+ # of lines) • Assign formal charges to molecules.

2. VSEPR • Beryllium chloride • Aluminum chloride • Methane (CH4) • Phosphine (PH3) • Water • Niobium (V) Bromide • Sulfur hexafluoride

3. Notes Water CO2 CO

13. Steps for Drawing Lewis Structures • Add up the total number of valence electrons available (EA). Be sure to take into account the total number of atoms present. • Exception: if the compound is negatively charged, add 1 electron for each negative value • Exception: if the compound is positively charged, remove 1 electron for each negative value • Example: CH4 C= 4 VE H= 1VE = 4 + 4(1) = 8 EA • Example: CO2 • Example: SO42- • Draw out a skeletal arrangement of lewis structure. Place the least electronegative atom in the center, surrounded by the remaining atoms. Draw lines to represent the bonds between the central atom and each surrounding atom. • Exception: Hydrogen can never be a central atom because it is found in the first energy level and can only have 2 electrons at most. • Draw the skeletal arrangement in the following examples • CO2 • H2SO4 (Just underline the central atom) • NCl3 • PF5 • Count and draw in the number of electrons necessary (EN) for each to have a full octet. These are the electrons needed. Be sure to include the bonded electrons. • Example: CO2 • Subtract the number of valence electrons available (from step one) from the number of electrons needed (EN –EA). This is the number of valence electrons that you still need. We make up for these by adding double or triple bonds. If you get a negative number, you have extra electrons. These may become part of your expanded octet. • Example: CO2 • To add double or triple bonds: divide the number you get from step 4 by 2 (why? Because we are adding more bonds and each bond represents 2 electrons). The number you get is the number of bonds you need to add in the form of either double or triple bonds. • Example: CO2 • Replace necessary electron pairs with double or triple bonds, in accordance with your answer from step 5. You may need to redraw your skeleton with any necessary double bonds and repeat steps 3-5. • Example: Add in bonds for CO2 • But where do they go? Double? Triple? Considering the following: • How many bonds does this atom want to make? • Octet rule • Exceptions: 3 reasons: • H: 2 • Be: 4 • B: 6 • P: 10 • S: 12 • Expanded octets can be in the form of bonds or lone pairs of electrons • Avoid charges on individual atoms unless ion. The formal charge of the atom is the number of electrons before bonding – number of electrons surrounding the atom after bonding (bonds count as 1 electron). • Calculate the formal charge of each atom as written in lewis structure: CO2 • Assign formal charge to entire molecule: sum of individual atoms’ formal charges. This number should add up to the original charge of the molecule. • Example: calculate formal charge of CO2 • Lewis Structure Practice • Directions: • Name each molecule or ion. • Draw each Lewis structure. • Draw any resonance structures that are possible. • If there is an exception to the octet rule, circle the example and state what the exception is. • NH3 • PO43- • SO42- • NO2- • O3 • NH4+ • BH3 • XeF4 • ClO2 • CN- • C2H4 • C2H2

14. Steps for Drawing Lewis StructuresAdd up the total number of valence electrons available (EA). Be sure to take into account the total number of atoms present. Exception: if the compound is negatively charged, add 1 electron for each negative valueException: if the compound is positively charged, remove 1 electron for each negative valueExample: CH4 C= 4 VE H= 1VE = 4 + 4(1) = 8 EAExample: CO2Example: SO42-Draw out a skeletal arrangement of lewis structure. Place the least electronegative atom in the center, surrounded by the remaining atoms. Draw lines to represent the bonds between the central atom and each surrounding atom. Exception: Hydrogen can never be a central atom because it is found in the first energy level and can only have 2 electrons at most. Draw the skeletal arrangement in the following examplesCO2 H2SO4 (Just underline the central atom)NCl3 PF5 Count and draw in the number of electrons necessary (EN) for each to have a full octet. These are the electrons needed. Be sure to include the bonded electrons. Example: CO2Subtract the number of valence electrons available (from step one) from the number of electrons needed (EN –EA). This is the number of valence electrons that you still need. We make up for these by adding double or triple bonds. If you get a negative number, you have extra electrons. These may become part of your expanded octet.Example: CO2To add double or triple bonds: divide the number you get from step 4 by 2 (why? Because we are adding more bonds and each bond represents 2 electrons). The number you get is the number of bonds you need to add in the form of either double or triple bonds. Example: CO2