Which lewis structure correctly represents kcl

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Problem : Choose the best Lewis structure for SF4. Frequently Asked Questions What scientific concept do you need to know in order to solve this problem? What is the difficulty of this problem? Our modeling should reflect this. Let's look at SO 4 2 - as an example. Lets try out several different models and eliminate all but the correct one using our rules as we know them so far.

In this first attempt we've met the normal valence of all the oxygens so they all have a formal charge of 0. The S looks like one of the configurations of noble gas valence of 8 with no lone pairs. Since the noble gases are two columns to the right, that gives a formal charge of - 2 on the S.

The sum of the formal charges adds up to the overall charge but there's a high charge on S which would not be stable. In this second model, we can see that there is only 1 bond between the S and each O.

The sum of the formal charges adds up to the overall charge so we've put the correct number of electrons but there's a high charge on the S.

If we move to satisfy the normal valence of two of the oxygens and the S by making double bonds to two of the O atoms we get the structure shown on the right. This time, two of the oxygens and the S are all in a normal or hyper- valence state so have a formal charge of zero.

Two of the oxygens still look like F and so have a formal charge of - 1. The sum of the formal charges adds up to - 2 which is the overall charge.

Thus, this structure works and it has the lowest formal charges of the three models we've tried here. A good final check is to count the electrons in your structure and see if they add up to the sum of the valence electrons and the charge. Sulfate has 32 valence electrons that are paired up in 16 pairs. Count the bars in our diagram representing pairs of electrons either bonding or non-bonding and we see 16 electron pairs are indeed drawn in on the diagram.

We have done a good Lewis-Dot structure. The one problem left is that this model seems to make two of the oxygens different from the other two. Experiment doesn't agree with this prediction. We'll see later how we modify the model to accommodate this. So far, we've been exposed to the rules of creating Lewis structures one at a time. Let's summarize here the rules we've learned so far.

With practice, these rules become simply intuitive and you will see that eventually you will be able to quickly draw Lewis structures without reference to them any further. Check here for the answers. For a numeric, method to find a Lewis structure, click here. As one last example, let's look at Carbon monoxide and work its Lewis Structure out using the rules.

Since we have only two atoms, they must be bonded to each other and with filled octets they must be configured the same. If we try to make the atoms look like C we would come up with the incorrect structure shown here. Here, the C looks OK, it has 4 bonds and no lone pairs, just like the table of valences describes so it has a formal charge of zero.

This time, The O is neutral but the C has a - 2 charge on it. This model has 2 too many electrons on it, since carbon monoxide is a neutral molecule. We can write off this picture too. So, We were unable to come up with a model where the normal valence states of either atoms is satisfied.

Perhaps, if we "split the difference" This is the only Lewis structure that has a sum of formal charges that add up to zero, the charge on the molecule. This one looks strange but is the only correct Lewis dot structure. However, the Lewis-dot model does not properly predict atomic charges here. In this final model, we see that both the C and the O have a charge and that, in fact, the charges seem opposite of what we might have expected considering that O is very electronegative, compared to C.

In all likelihood, the O will pull the three bonding pairs closer to it, making an unevenly shared triple bond with most of the negative charge closer to the O. That will bring the actual atomic charge closer to zero for each atom.

Remember, the Formal charge is not necessarily the real charge. It's just the results of a very simple model for keeping track of electrons. There are some few exceptions to the rules which may cause us trouble. These are the radicals, In radicals, there is an odd number of electrons so no matter how many models we try where all the electrons are paired up, we'll never be able to make the model match the real molecule. One such examples of this is NO 2.

If we add up the electrons on this species, there are 5 from the N and 6 each from the two Os for a total of 17 electrons. An odd number. Now, pretend we didn't do that. We don't know that we cannot use the normal rules to make a Lewis structure that will work. Let's see what trouble we might get into. Which of these is the correct one. You can't tell. This is the point where our rules for Lewis structures break down and we realize we need something better.

The latter of these has the lower formal charges so we might suspect that it is the favoured one. However, EPR experimental data puts the unpaired electron almost exclusively on the N as shown in the former of the two. Many molecules have electrons that are localized in simple orbitals bonds, lone pairs, etc. Such species are often well represented using a single Lewis dot structure. However, many molecules or ions have electrons that either are not localized in one position or perhaps where higher energy levels are easily reachable at room temperature.

These species may not be so well represented by a single Lewis dot structure because the species itself does not stay in one configuration. Lewis dot diagrams give us a static picture of what the molecule or ion might look like.

So how do we deal with situations were a static model does not tell us the entire story. We assume that we can represent each different configuration by a Lewis dot structure and in some cases, all these different diagrams add to our overall understanding of the real species. We will restrict our discussion here to situations where we can find more than one equivalent ground-state Lewis structure.

In these cases, we must consider whether the multiple Lewis dot structures represent chemically different species or different electronic configurations of the same species.

In our sulphate ion case, there are six identical structures and they an interchange one for the other as indicated below. We use a single double-headed arrow to represent movement of electrons only. We assume Atoms do not move because they are several thousand times more massive and hence several thousands times slower than electrons.

Since there are six structures which interchange on a rapid basis, we need to rethink our ideas of formal charges and of Bond order. Hence, for the top oxygen pretend we have nailed down the ion so it can't rotate. That way we can keep track of the motion of the electron within the ion. We can now draw a new type of structure using the average bond order and the average charges.

This average model better represents the reality of the situation in the real sulphate ion. Hence, the solid and the dotted lines joining the pair in the final "average" diagram. Some people will also argue that higher energy structures must be included in the list of possible resonance structures. We saw two such structures already for the sulphate ion. In the purest view, they are also possible resonance structures but they represent higher energy configurations of the sulphate ion.

Since the high-energy configurations do not last long first law of thermodynamics , they don't contribute as much to the overall picture as do configurations represented by the six ground-state Lewis dot diagrams above. So, we will not count the high-energy diagrams in the calculations of average bond order or formal charge. In the previous example as in most resonance structures you will encounter, the average bond order is greater than one for the bonds where the resonance is actually occurring.

That is because, typically, one bond remains in place throughout the resonance process and a second or third bond moves from place to place in the different resonant structures. In certain special situations, where there are insufficient electrons for normal bonding to occur, resonant structures can be formed between atoms of a species to form a bond order that is actually less than one. In these cases, there must be at least one of the series of resonance structures where the bond is completely missing.

Obviously, these compounds would not be as stable as those we normally think of as having resonance but they do exist. Take the molecule B 2 H 6. It does exist but if we try to draw a normal Lewis-dot structure, we cannot find a way to bond all the atoms together at one time. Using resonance structures we show how the bonds can move in such a way as to create a bond for at least part of the time. In this resonance model, we see that the B? An average picture that shows the entire structure as being connected might be.

Here, we see that the central hydrogens are bridging between two borons in a curved, electron-deficient 3-centre bond sometimes called a banana bond. This three-centre bond has only two electrons in it and hence the bond order of the individual B?

Molecules with this kind of bonding are stable mostly in boron compounds although there are examples of other three-centre electron deficient bonds. It is possible to create valid Lewis structure models of such bonding in any analogous compound made up of column 3 elements like Al, Ga, In, Tl.

We can draw the analogous structures but for other reasons, beyond the scope of this discussion, actual chemical analogues such as Al 2 H 6 or Ga 2 H 6 are not as stable and if created, some of these would quickly reform into other, more stable compounds.

We can find other examples of electron-deficient three-centre bonds that are stable and analogous to the B 2 H 4 molecule, for example, Al 2 Cl 4 exists, is stable and used electron-deficient three-centre bonds that are quite analogous to those shown above. We can now use the ideas developed here to describe a process where a bond is formed when a species with an extra electron pair forms a covalent bond with a species which is deficient in an electron pair.

This obvious acid-base reaction is exactly analogous to the reaction between ammonia and the hydrogen cation. In this reaction, still would refer to the ammonia as a base and the hydrogen cation as the acid. In this case, the F - ion has the extra electron pair and uses it to bond with the BF 3 in which the B is missing one electron pair to complete the octet.

The Fluoride ion donates an electron pair to the Boron. The Fluoride acts just like the ammonia or the hydroxide ion did in the previous two examples where they clearly the base. Thus, the electron-pair donor is a Lewis Base. One last type of bond can be explained now using the Lewis dot structures. In this case, however, a bond is not formed, merely a strong attraction.

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