There are notable exceptions to the octet rule, however, such as the ability of sulfur to form 6 bonds not shown here. Fluorine symbol F is found in column 7 on the periodic table. It has 7 valence electrons. It needs to make 1 bond to get an octet. In the diagram above, we show fluorine making 1 bond. The bond could be drawn up, down, left, or right.
There are 6 more fluorine valence electrons appearing as 3 pairs of dots. The same goes for the other halogens. How many electrons does chlorine have? There are 7 valence electrons for chlorine, and it would have 1 bond and three pairs of dots. Same for bromine Br and iodine I.
The noble gasses are noble. They are already born with 8 electrons, generally, so they already obey the octet rule without bonding. The figure above shows neon Ne surrounded by 4 pairs of dots, without any bonds. The same would apply for Ar, Kr, and Xe. There are more exceptions to the octet rule for the noble gasses. Note that helium He is noble.
It is element number 2 with only 2 electrons, so it cannot possibly have an octet of 8 electrons. Still, helium is noble, and elements near helium on the periodic table are stabilized with 2 electrons not an octet of 8. Speaking of exceptions to the octet rule, apparently somebody made XeF 2 in a fancy lab. Yet it did, apparently twice. Hydrogen is very different than the other elements. Hydrogen is another of the exceptions to the octet rule.
It is born with 1 electron. It takes an electron to form a bond. So hydrogen just forms a bond. It is drawn as an H with one stick or line. No dots showing any hydrogen valence electrons. The hydrogen valence electrons act like helium, element number 2. Atoms with an expanded octet. To have an expanded octet more than 8 electrons you need more than 4 orbitals.
In the next chapter we will learn that molecular orbitals are different than atomic orbitals, but one way to describe molecular orbitals is by assuming they are the result of different atomic orbitals mixing with each other, and the number of molecular orbitals must equal the number of atomic orbitals that were mixed to create them.
This means that you need to start using d-orbitals, and those are only available for atoms in the third period or below So atoms in the third period or greater can have expanded octets.
Phosphorous often has 5 orbitals 10 electrons and sulfur often has 6 orbitals 12 electrons because they are in the third period, but nitrogen and oxygen can never have expanded octets because they are in the second period and there is not such thing as a 2d orbital. Why can't elements in the second row have an expanded octet? For the second row, the principle quantum number is 2, and so you have s and p orbitals in the valence shell. NOTE: in the next chapter we will study molecular orbitals, and the orbitals in molecules are not atomic orbitals s,p,d,f , but are different.
That said, the number of molecular orbitals is determined by the number of atomic orbitals in the valence shell, but they are different.
Exceptions to the Octet Rule The "octet rule" says that in many compounds the most stable correct electron configuration is when there are 8 electrons four filled orbitals. In the hydrogen molecule, for example, each hydrogen atom acquires some control over two electrons, thus achieving something resembling the helium structure.
Similarly the formation of a fluorine molecule from its atoms can be represented by. Again a pair of electrons is shared, enabling each atom to attain a neon structure with eight electrons i. Similar diagrams can be used to describe the other halogen molecules:. In each case a shared pair of electrons contributes to a noble-gas electron configuration on both atoms.
Since only the valence electrons are shown in these diagrams, the attainment of a noble-gas structure is easily recognized as the attainment of a full complement of eight electron dots an octet around each symbol.
In other words covalent as well as ionic compounds obey the octet rule. The octet rule is very useful, though by no means infallible, for predicting the formulas of many covalent compounds, and it enables us to explain the usual valence exhibited by many of the representative elements. In order to attain a noble-gas structure, therefore, they need only to participate in the sharing of one pair of electrons.
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