Drawing the Lewis Structure for ClO_{2}^{}Viewing Notes:
Transcript: Let's do the ClO2 Lewis structure. Chlorine has 7 valence electrons. Oxygen has 6, we have two Oxygens; plus, we have a valence electron up here we need to add in for a total of 20 valence electrons. Chlorine's the least electronegative, it goes in the center. We'll put an Oxygen on either side. Next, two valence electrons between atoms to form chemical bonds; we've used 4, 6, 8, 10, 12, 14, 16, now back to the center to complete the octet, 18 and 20. So this looks pretty good. Each of the atoms has an octet. We've used all 20 valence electrons. But Chlorine can be found in the third period of the periodic table. That means it can have more than eight valence electrons, so we need to check our formal charges. When we calculate the formal charges, we can see the Chlorine has a plus one formal charge, and the Oxygen has a minus one. And they're both symmetrical, so they both have minus one. That makes sense: minus one plus one, minus one, that gives us this negative charge here. But we want our formal charges to be as close to zero as possible and still have this negative. So let's see what we can do. If we move a pair of electrons from the outside here to form a double bond, that should get rid of this +1 charge here for Chlorine. So let's recalculate our formal charges and see how that worked. So when we recalculate, after moving these valence electrons here into the center to form a double bond, the Chlorine now has a formal charge of zero. The blue Oxygen right here didn't change, it's still 1. And then the Oxygen in black right here now has a formal charge of zero. This is a better structure because the formal charges are closer to zero while still retaining that negative one right there. One last thing: we do need to put brackets around this to show that it is a negative ion. And that's it. That's the best structure for ClO2. This is Dr. B., and thanks for watching. 
