Octet rule ending

This violates the octet rule—the carbon atom would end up with live bonds. So we cannot push the arrows that way. There is no way to turn the lone pair into a pi bond in this example. 

Termination occurs when the active sites of two growing chains meet, as shown in Fig. 2.3 d). The unpaired electrons form a bond that couples the ends of the chains. Alternatively, disproportionation may occur. This happens when one chain transfers a hydrogen atom to the other and the electrons on both species rearrange themselves to satisfy the octet rule. 

Because they do not obey the octet rule, hypervalent molecules have often been thought to involve some type of bonding that is not found in period 2 molecules. Ideas concerning the nature of this bonding have developed along a somewhat tortuous path that it is interesting and instructive to follow. We will in the end conclude that the nature of the bonding in these molecules is not different in type from that in related period 2 molecules and that there is therefore little justification for the continued use of this concept. 

The concept of valence has been subject to revision over the years. Initially, valence was regarded as the combining power of an element and was derived from the composition of compounds. At the end of the period before the age of quantum chemistry, valence was generally formulated in relation to the octet rule [1—3), a simple relation which still finds useful application in modem chemistry. 

What does this mean for first- and second-row elements A first-row element like hydrogen can accommodate two electrons around it. This would make it like the noble gas helium at the end of the same row. A second-row element is most stable with eight valence electrons around it like neon. Elements that behave in this manner are said to follow the octet rule. 

Let us illustrate the concept of formal charge using the ozone molecule (O3). Proceeding by steps, as we did in Examples 9.3 and 9.4, we draw the skeletal structure of O3 and then add bonds and electrons to satisfy the octet rule for the two end atoms ... 

You can see that although all available electrons are used, the octet rule is not satisfied for the central atom. To remedy this, we convert a lone pair on one of the end atoms to a second bond between that end atom and the central atom, as follows ... 

Our drawing of the Lewis structure for ozone (O3) satisfied the octet rule for the central atom because we placed a double bond between it and one of the two end O atoms. In fact, we can put the double bond at either end of the molecule, as shown by these two equivalent Lewis structures ... 

Ionisation of atomised metal may seem an extreme case, so perhaps it is not important if students think of the Na+ ion as more stable than the atom. However students have a strong tendency to see any species with an octet of electrons as stable, and research shows that by the end of secondary education, students will commonly rate a whole range of dubious ions as more stable than atoms because they have foil shells or octets of electrons. So not only do students tend to think Na is a stable ion, they make the same judgement about the chemically quite ridiculous species Na shown in Figure 3.7. It is important, therefore, that teachers make sure that students do not over-generalise the octet rule from a very useful rule of thumb for identifying the most likely formulae for molecules and ions, and adopt it as an absolute principle to judge stability and explain why reactions occur. 

Step 4 We see that this structure satisfies the octet rule for aU the O atoms but not for the N atom. The N atom has only six electrons. Therefore, we move a lone pair from one of the end O atoms to form another bond with N. Now the octet rule... 

When you draw curved arrows to indicate the creation of a new contributing structure, the arrows always start on either a double (or triple) bond or a lone pair of electrons, as shown in tire examples above. Further, the arrows should end at an atom that can accept a bond or should create a lone pair of electrons on an atom. Often when a new bond to an atom is created, one of the existing bonds to that atom must break so as not to exceed the octet rule. In the two examples given above, the central N of nitrite and C of acetate both acquire one bond and break one bond when you redistribute valence electrons between the contributing structures. 

In the first example the head of the curved arrow is giving a fifth bond to the carbon atom. This violates the octet rule. The second example does not violate the octet rule, because the carbon atom is gaining one bond but losing another. In the end, the carbon atom never has more than four bonds. 

The elements in groups 1, 2 and 13 have only 1, 2 or 3 electrons in their outer shell. These elements at the beginning of a period lose electrons to form positive ions (cations). The resulting simple ions obey the octet rule (eight electrons in the outer shell) and have an electron arrangement like the noble gas at the end of the previous period. 

As depicted in Figure 2.1 and outlined in the Summary of Chapter 2, viable bonding theories started to emerge from the quantum-mechanical model of the atom in the 1920s. G. N. Lewis proposed his now familiar electron-dot diagrams and octet rule for simple compounds in the early 1920s, and by the end of the decade, Nevil Sidgwick applied these ideas to coordination compounds. It was he who first proposed the idea of the coordinate-covalent bond referred to in earlier chapters. 

Beginning with the element carbon, each nonmetal atom in a molecule will be represented as being surrounded by eight valence electrons. This is called the octet rule. The reason for the octet rule is that having eight valence electrons makes a nonmetal atom isoelectronic with the noble gas at the end of its period, which gives that atom considerable stabiUty. 

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