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121TutorialWritingLewisStrucSp2006

121TutorialWritingLewisStrucSp2006 - Steps to Writing Lewis...

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Steps to Writing Lewis Structures Dr. Yau 1. Examine formula to determine whether substance is ionic or molecular . If it is ionic, work with each ion separately. 2. Calculate the total # of valence electrons from the Group #. Remember to adjust for the charge if it is an ion. Divide the total by 2 to get # of pairs of valence electrons . 3. Work out a reasonable skeleton : Usually the element shown on left in the formula is the central atom. Sometimes there can be more than one central atom.(e.g.For N 2 H 4 , both N are considered "central".) H N N H H H .. .. Avoid stringing out the atoms. SO 4 2– is not S-O-O-O-O, but ) Remember H can have only one bond, so H can never be in the middle. ( NEVER X H X ) It also helps to remember that the unpaired electrons in the e - dot symbol of each element show the preferred number of bonds. (Note: This is not the required , but the preferred number of bonds. Some structures may not be able to get the preferred number of bonds .) For example, •C• prefers 4 bonds, : N prefers 3 bonds, : S : prefers 2 bonds, : Cl : prefers 1 bond. Sometimes more than one arrangement of atoms is possible, giving rise to isomers . e.g. H C C O H H H H H .. .. H C O H H H H H C .. .. 4. Write only one bond between adjacent atoms, counting each as a pair of electrons. Continue putting in pairs of e - as “lone pairs” on the outside atoms (not central atom). Spread them out evenly, as pairs, amongst the outside atoms. Remember not to put “lone pairs” on H !! ( NEVER X–H: ) When the outside atoms are “full”, and if there are any e - left over, put them as “lone pairs” on the central atom. 5. If and only if you run out of e before all the atoms get 4 pairs (the octet), borrow a lone pair from a neighbor and share (forming double or triple bonds). Do not exceed the octet with multiple bonds in this class. 6. Check to see all atoms have octets (except H and noted exceptions.) Sometimes more than one arrangement of electrons is possible, giving rise to resonance structures . Consider what is the preferred number of bonds for each element to determine which resonance structures are reasonable or contribute most to the overall structure. : Cl=N–N–Cl : ←→ : Cl–N=N–Cl : ←→ : Cl–N–N=Cl : most reasonable resonance structure 7. If it is an ion, put square brackets around the structure and write in the net charge.
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