Use formal charges to identify the better Lewis structure.

Video transcript

hey everyone in this example, we need to determine what is the most acceptable lewiS structure for the given molecules below. So we should recall our formula for formal charge. And that's due to the fact that for whatever molecule that is the most acceptable lewis structure, it's going to have the least amount of formal charges. We should recall our formula for formal charge is taking our valence electrons for a selected atom in the molecule, subtracting that from our number of lone pairs or non bonding electrons. And then subtracting that from then our number of bonding electrons divided by two. So we're going to go ahead and begin with part A. We have the structure for ethanol and we can go ahead and select any of the atoms in the molecule. So let's go ahead and label our carbon. Since we have multiple carbons, we'll label this 11 and this one, number two, beginning with carbon one, We want to go ahead and recall that carbon is in group four a of our periodic table. So we would recall that that corresponds to four valence electrons. Were then subtracting that from any lone pairs on this carbon. And if we look at this molecule we don't see any lone pairs. So we say zero lone pairs. And then as far as bonding electrons, we would recall that within a bond in a molecule we have two electrons. So we have a total of 24 bonding electrons on this carbon. So that would be four bonding electrons divided by two. So beginning with our division, we have four divided by two. Which is going to give us two, meaning we would have a difference of four minus two giving us a formal charge of plus two for this carbon here. So we can go ahead and write that in parentheses. Moving on to the second carbon in this molecule, We would recognize again because carbon is in group four a. It has four valence electrons. We do not see any lone pairs based on this carbon position in the molecule. So we can say zero lone pairs yet again. And then as far as bonding electrons, we have four bonds giving us a total of 2468 electrons. So we would say eight electrons divided by two. So we start with the division we have 2, 8 divided by two. That's going to give us four, giving us 4 -4. And we would then have a formal charge equal to zero. So this is a neutral carbon here. Moving on to our oxygen here, We would recognize that oxygen is in group six a of our periodic table corresponding to six valence electrons. We would then recognize that in our molecule this oxygen has one lone pair here. So we would say -1 lone pair. And then just to make more room, I'll put this further down. So then we're going to take the difference of bonding electrons. We would recognize that this oxygen has a total of four bonds around itself. So that's 2468 bonding electrons. So that's eight electrons divided by two. We know that eight divided by two is going to be four. So we would have six minus four which would give us a formal charge When we also subtract the one lone pair of plus one. So we can go ahead and write that next to that oxygen there and now let's move forward and go to choice be to compare how many formal charges it has in the molecule. So we have another structure of ethanol. And we should recall from our first structure that when we have a carbon with four charges or four bonds to itself and no loan pairs, it will have a neutral charge or neutral formal charge. So these two carbons here will have both formal charges of zero Because they both have four bonds around themselves. And in this case now we have an oxygen atom here. So let's go ahead and calculate the formal charge of this oxygen. So I'll just label it oxygen be For sake of our room here. So we should again recall that oxygen is located in group six a of our periodic tables which we recall would correspond to six valence electrons. We would then recognize that within this molecule this oxygen has a total of two lone pairs. So we would say -2 lone pairs. And then we want to subtract bonding electrons divided by two. And on this oxygen we would see we have a total of 24 bonding electrons Divided by two, which is going to give us two. So we would have 6 -2. That's going to leave us with 4 -2, which would leave us with a formal charge of plus two. So we can label this oxygen as having a formal charge of plus two. And you may be wondering why we haven't calculated the formal charges of the hydrogen? Well, we would recognize that hydrogen are located in group one A of our periodic table. We recognize that none of the hydrogen have any lone pairs, they just have one single bond. So that would give us just the difference between our bonding electrons for these hydrogen where they only have a total of two bonding electrons in their single bonds and two divided by two. We recall as one, so all of our hydrants and each of these molecules are going to simply have a formal charge of zero. So we can disregard them. Moving on to the third molecule in part C. We would recognize from part A are given molecule in part A. That when we have a carbon with just two bonds that leaves us with a formal charge of plus two. So we can label that here, we also recall that above when we have a carbon with four bonds around itself and no lone pairs, it will have a formal charge of zero because it's neutral. And then with oxygen we recall from the first molecule in part A that when we have an oxygen atom with four bonds and one lone pair. That leaves us with a formal charge of plus one. So we can rule out A and C. So far because they yielded the same amount of formal charges. And now let's move on to molecule D. So we would recognize that in this molecule we have an extra carbon here. So this is not even an ethanol molecule, but we can go ahead and calculate the formal charge regardless. For these two carbons here, we would recognize they have formal charges of zero and then a formal charge of plus two for this carbon just like above and the other molecules. And then for the oxygen here, we'll just label this oxygen D. We have again six valence electrons. We would subtract the total number of lone pairs. Where we would count a total of 12 and three lone pairs on this oxygen atom, and then subtracting the number of bonding electrons. We recognize that only has one bond corresponding to two bonding electrons divided by two which would give us one. So we would have six minus three which is three and then three minus one which is going to yield a formal charge of Plus two for this oxygen here. And so we can go ahead and to determine our final answer for this question because we recognize that molecule D. Has a formal charge of plus four total. We can rule it out here. It had to formal charges of plus two. However, when we looked at molecule B, it only had one formal charge of a plus two charge on the oxygen atom in the structure of ethanol. So we can confirm that to complete this example, choice B is the most correct choice because it yielded the least amount of formal charges. It only had a net formal charge of plus two. So this would be our final answer to be the most acceptable LeWIS structure out of all of the given molecules here. So I hope that everything I reviewed was clear. If you have any questions, please leave them down below and I will see everyone in the next practice video.