In addition, BF 3 will react with ammonia NH 3 , to form a stable compound, NH 3 BF 3 , for which a Lewis structure can be drawn that shows boron with a complete octet. Boron trifluoride-ammonia complex : This covalent compound NH 3 BF 3 shows that boron can have an octet of electrons in its valence level. Compounds of aluminum follow similar trends. Aluminum trichloride AlCl 3 , aluminum hydride AlH 3 , and aluminum hydroxide Al OH 3 indicate a valence of three for aluminum, with six valence electrons in the bonded molecule.
However, the stability of aluminum hydride ions AlH 4 — indicates that Al can also support an octet of valence shell electrons. Although the octet rule can still be of some utility in understanding the chemistry of boron and aluminum, the compounds of these elements are harder to predict than for other elements.
Some elements, most notably nitrogen, can form compounds that do not obey the octet rule. One class of such compounds are those that have an odd number of electrons. As the octet rule requires eight electrons around each atom, a molecule with an odd number of electrons must disobey the octet rule. Recall that the Lewis structure of a molecule must depict the total number of valence electrons from all the atoms which are bonded together. Nitric oxide has the formula NO. Therefore, no matter how electrons are shared between the nitrogen and oxygen atoms, there is no way for nitrogen to have an octet.
It will have seven electrons, assuming that the oxygen atom does satisfy the octet. Nitric oxide : Nitric oxide NO is an example of a stable free radical. It does not obey the octet rule on the nitrogen atom. Each line around the atoms represents a pair of electrons.
Nitric oxide is a by-product of combustion reactions that occur in engines, like those in automobile engines and fossil fuel power plants. It is also produced naturally during the electrical discharge of lightning during thunderstorms. Nitrogen dioxide is the chemical compound with the formula NO 2.
Again, nitrogen dioxide does not follow the octet rule for one of its atoms, namely nitrogen. There is persistent radical character on nitrogen because it has an unpaired electron.
The two oxygen atoms in this molecule follow the octet rule. Nitrogen dioxide : Nitrogen dioxide is another stable molecule that disobeys the octet rule. Note the seven electrons around nitrogen. Nitrogen dioxide is an intermediate in the industrial synthesis of nitric acid, millions of tons of which is produced each year.
This reddish-brown toxic gas has a characteristic sharp, biting odor and is a prominent air pollutant. Main group elements in the third period and below form compounds that deviate from the octet rule by having more than 8 valence electrons.
A hypervalent molecule is a molecule that contains one or more main group elements that bear more than eight electrons in their valence levels as a result of bonding. As a result, the second period elements more specifically, the nonmetals C, N, O, F obey the octet rule without exceptions.
If any of this seems unfamiliar, I encourage you to watch the video on introduction to drawing Lewis diagrams. But what we'd wanna do is first think about our valence electrons. So xenon right over here, it's actually a noble gas. It already has a full octet in its outer shell, so it has eight valence electrons. So xenon has eight valence electrons.
And then fluorine, we've seen this multiple times, has one, two, three, four, five, six, seven valence electrons, but there's five of them. So five times seven. I'm gonna be drawing a lot of electrons in this. So this gives us a total of eight plus 35, which is 43 valence electrons. But we have to be careful. This is a cation. It is a positively charged molecule.
It has a positive one charge. So we have to take one electron away because of that. So let's take away one valence electron to get that cation. And so we are left with So the next step is to try to draw its structure with some basic single covalent bonds. And xenon would be our preferred central atom because fluorine is more electronegative. It's actually the most electronegative element. So let's put xenon in the middle, and then let's put some fluorines around it, five of them to be specific.
So one, two, three, four. I'm having trouble writing an F. Four and then five fluorines. And now let me make five covalent bonds. One, two, three, four, five. So just like that, I have accounted for 10 valence electrons because you have two valence electrons in each of these covalent bonds, two, four, six, eight, So let me subtract 10 valence electrons. And then we are left with 32 valence electrons.
Now, the next step is to try to allocate some more of these valence electrons to the terminal atom so that they get to a full octet. So let me do that to the fluorines. Each of these fluorines already are participating in a covalent bond, so they already have two valence electrons hanging out with them, so let's give 'em each six more. So let's give that fluorine six, and that fluorine gets six, and that fluorine gets six valence electrons, and that fluorine gets six valence electrons, and then last but not least, this fluorine gets six valence electrons.
So I have just given away six valence electrons to each of five fluorine atoms. One of the situations where expanded octet structures are treated as more favorable than Lewis structures that follow the octet rule is when the formal charges in the expanded octet structure are smaller than in a structure that adheres to the octet rule, or when there are less formal charges in the expanded octet than in the structure a structure that adheres to the octet rule.
Such is the case for the sulfate ion, SO 4 A strict adherence to the octet rule forms the following Lewis structure:. If we look at the formal charges on this molecule, we can see that all of the oxygen atoms have seven electrons around them six from the three lone pairs and one from the bond with sulfur. This is one more electron than the number of valence electrons then they would have normally, and as such each of the oxygens in this structure has a formal charge of If instead we made a structure for the sulfate ion with an expanded octet, it would look like this:.
Looking at the formal charges for this structure, the sulfur ion has six electrons around it one from each of its bonds.
This is the same amount as the number of valence electrons it would have naturally. This leaves sulfur with a formal charge of zero. The two oxygens that have double bonds to sulfur have six electrons each around them four from the two lone pairs and one each from the two bonds with sulfur. This is the same amount of electrons as the number of valence electrons that oxygen atoms have on their own, and as such both of these oxygen atoms have a formal charge of zero.
The two oxygens with the single bonds to sulfur have seven electrons around them in this structure six from the three lone pairs and one from the bond to sulfur. That is one electron more than the number of valence electrons that oxygen would have on its own, and as such those two oxygens carry a formal charge of Remember that with formal charges, the goal is to keep the formal charges or the difference between the formal charges of each atom as small as possible. The ICl 4 - ion thus has 12 valence electrons around the central Iodine in the 5 d orbitals.
Expanded Lewis structures are also plausible depictions of molecules when experimentally determined bond lengths suggest partial double bond characters even when single bonds would already fully fill the octet of the central atom. Despite the cases for expanded octets, as mentioned for incomplete octets, it is important to keep in mind that, in general, the octet rule applies. There are three exceptions: 1 When there are an odd number of valence electrons, 2 When there are too few valence electrons, and 3 when there are too many valence electrons.
Mike Blaber Florida State University. Learning Objectives To assign a Lewis dot symbol to elements not having an octet of electrons in their compounds. Exception 1: Species with Odd Numbers of Electrons The first exception to the Octet Rule is when there are an odd number of valence electrons. Exception 2: Incomplete Octets The second exception to the Octet Rule is when there are too few valence electrons that results in an incomplete Octet. Solution 1.
Draw connectivities: 3. Add octets to outer atoms: 4. Does central electron have octet? It has 6 electrons Add a multiple bond double bond to see if central atom can achieve an octet: 6. The central Boron now has an octet there would be three resonance Lewis structures However In this structure with a double bond the fluorine atom is sharing extra electrons with the boron.
Thus, the structure of BF 3 , with single bonds, and 6 valence electrons around the central boron is the most likely structure BF 3 reacts strongly with compounds which have an unshared pair of electrons which can be used to form a bond with the boron:. Exception 3: Expanded Valence Shells More common than incomplete octets are expanded octets where the central atom in a Lewis structure has more than eight electrons in its valence shell.
Draw the connectivities: 3. Add octet of electrons to outer atoms: 4. The ICl 4 - ion thus has 12 valence electrons around the central Iodine in the 5 d orbitals Expanded Lewis structures are also plausible depictions of molecules when experimentally determined bond lengths suggest partial double bond characters even when single bonds would already fully fill the octet of the central atom. Summary Following the Octet Rule for Lewis Dot Structures leads to the most accurate depictions of stable molecular and atomic structures and because of this we always want to use the octet rule when drawing Lewis Dot Structures.
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