Unpaired electrons represent places where electrons can be gained in ionic compounds, or electrons that can be shared to form molecular compounds. The valence electrons of helium are better represented by two paired dots, since in all of the noble gases, the valence electrons are in filled shells, and are unavailable for bonding. Covalent bonds generally form when a nonmetal combines with another nonmetal. Both elements in the bond are attracted to the unpaired valence electrons so strongly that neither can take the electron away from the other unlike the case with ionic bonds , so the unpaired valence electrons are shared by the two atoms, forming a covalent bond :.
The shared electrons act like they belong to both atoms in the bond, and they bind the two atoms together into a molecule. The shared electrons are usually represented as a line — between the bonded atoms. In Lewis structures, a line represents two electrons.
Atoms tend to form covalent bonds in such a way as to satisfy the octet rule , with every atom surrounded by eight electrons. Hydrogen is an exception, since it is in row 1 of the periodic table, and only has the 1 s orbital available in the ground state, which can only hold two electrons.
The shared pairs of electrons are bonding pairs represented by lines in the drawings above. The unshared pairs of electrons are lone pairs or nonbonding pairs. All of the bonds shown so far have been single bonds , in which one pair of electrons is being shared. It is also possible to have double bonds , in which two pairs of electrons are shared, and triple bonds , in which three pairs of electrons are shared:.
Examples 1. This uses up all of the valence electrons. The octet rule is satisfied everywhere, and all of the atoms have formal charges of zero. This uses up six of the eight valence electrons. All of the valence electrons have now been used up, the octet rule is satisfied everywhere, and all of the atoms have formal charges of zero. This uses up four of the valence electrons. This uses up six of the valence electrons. The remaining two valence electrons must go on the oxygen:.
All of the valence electrons have been used up, and the octet rule is satisfied everywhere. The remaining six valence electrons start out on the N:. The octet rule can be satisfied if we move two pairs of electrons from the N in between the C and the N, making a triple bond:. This uses up the sixteen valence electrons The octet rule is not satisfied on the C, and there are lots of formal charges in the structure:.
The octet rule can be satisfied, and the formal charges diminished if we move a pair of electrons from each oxygen atom in between the carbon and oxygen atoms:. Now, all of the valence electrons have been used up, the octet rule is satisfied everywhere, and all of the atoms have formal charges of zero. Place the remaining valence electrons on the O and Cl atoms:. Making a carbon-chlorine double bond would satisfy the octet rule, but there would still be formal charges, and there would be a positive formal charge on the strongly electronegative Cl atom structure 2.
Making a carbon-oxygen double bond would also satisfy the octet rule, but all of the formal charges would be zero, and that would be the better Lewis structure structure 3 :. Examples continued from section B 9. We can satisfy the octet rule on the central O by making a double bond either between the left O and the central one 2 , or the right O and the center one 3 :. In this example, we can draw two Lewis structures that are energetically equivalent to each other — that is, they have the same types of bonds, and the same types of formal charges on all of the structures.
The actual molecule is an average of structures 2 and 3 , which are called resonance structures. Structure 1 is also a resonance structure of 2 and 3 , but since it has more formal charges, and does not satisfy the octet rule, it is a higher-energy resonance structure, and does not contribute as much to our overall picture of the molecule.
The real molecule does not alternate back and forth between these two structures; it is a hybrid of these two forms. The ozone molecule, then, is more correctly shown with both Lewis structures, with the two-headed resonance arrow between them:.
In contrast, the lone pairs on the oxygen in water are localized — i. Resonance delocalization stabilizes a molecule by spreading out charges, and often occurs when lone pairs or positive charges are located next to double bonds.
Resonance plays a large role in our understanding of structure and reactivity in organic chemistry. A more accurate picture of bonding in molecules like this is found in Molecular Orbital theory, but this theory is more advanced, and mathematically more complex topic, and will not be dealt with here.
Examples We can satisfy the octet rule and make the formal charges smaller by making a carbon-oxygen double bond. Once again, structure 1 is a resonance structure of 2 , 3 , and 4 , but it is a higher energy structure, and does not contribute as much to our picture of the molecule.
Multi-Center Molecules Molecules with more than one central atoms are drawn similarly to the ones above. The octet rule and formal charges can be used as a guideline in many cases to decide in which order to connect atoms. C 2 H 6 ethane C 2 H 4 ethylene The octet rule is not satisfied on the B, but the formal charges are all zero. In fact, trying to make a boron-fluorine double bond would put a positive formal charge on fluorine; since fluorine is highly electronegative, this is extremely unfavorable.
In this structure, the formal charges are all zero, but the octet rule is not satisfied on the N. Since there are an odd number of electrons, there is no way to satisfy the octet rule.
Nitric oxide is a free radical, and is an extremely reactive compound. In the body, nitric oxide is a vasodilator, and is involved in the mechanism of action of various neurotransmitters, as well as some heart and blood pressure medications such as nitroglycerin and amyl nitrite. Notice that the formal charge on the phosphorus atom is zero.
Notice that the formal charge on the sulfur atom is zero. Notice that the formal charge on the xenon atom is zero. Structures 1 and 2 are resonance structures of each other, but structure 2 is the lower energy structure, even though it violates the octet rule.
Sulfur can accommodate more than eight electrons, and the formal charges in structure 2 are all zero. Lone pairs go in the equatorial positions, since they take up more room than covalent bonds. The Lewis structures of the previous examples can be used to predict the shapes around their central atoms:.
Formula Lewis Structure Bonding Shape 1. CH 4 4 bonds. NH 3 3 bonds. The nitrogen and and one oxygen are bonded through a double bond which counts as "one electron pair".
Hence the molecule has three electron pairs and is trigonal planar for electron pair geometry. The one lone electron exerts a less repulsion than normal on the two bonding oxygen atoms so they are able to spread out more to a o bond angle from the ideal of o. All oxygen atoms have an octet of electrons. Nitrogen atom does not have an octet because the whole molecule is short an electron.
This electron deficient feature of the molecule make is very reactive because it will try to react with some other molecule to complete the octet. Elmhurst College. Lewis Diagrams. Trigonal Planar. Trigonal Pyrimid. Chemistry Department. Virtual ChemBook. Bent molecules occur most frequently in group 16 elements because they have 6 electrons in their outer shell, allowing for 2 bonds and 2 lone pairs.
First, the term bent molecule refers to the geometric shape that bonded atoms have taken. Second, each bent molecule must have a central atom with 2 bonded atoms. Third, every bent molecule must have 1 or 2 lone pairs of electrons. I hope that helps.
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