All oxygen atoms, however, are equivalent, and the Lewis structure could be drawn with the double bond between carbon and any one of the three oxygen atoms. One oxygen atom must have a double bond to carbon to complete the octet on the central atom. Thus, the electronic structure of the NO 2 − ion is shown as:
The actual electronic structure of the molecule (the weighted average of the resonance forms) is represented by a resonance hybrid of the individual resonance forms.
Since the bonds are equivalent, they are equally weighted, so each connection between nitrogen and oxygen has an N-O bond order of 1.5. The actual distribution of electrons in each of the NO bonds in NO 2 − is the weighted average of a double bond and a single bond. Both NO bonds have the same atoms involved, the same length, and require the same energy added to break them. The nitrite ion is an example of equivalent resonance, which means that the two bonds are identical. It communicates that we’re talking about resonance structures and not a chemical reaction (which is signified by a single-headed arrow, →) or equilibrium (which is signified by two half-arrows, ⇌). The two headed arrow that connects two (or more) valid Lewis structures is important. Therefore, two valid Lewis structures must be drawn to represent the bonding in the nitrite ion, NO 2 –. Instead, we use the concept of resonance: if two or more Lewis structures with the same arrangement of atoms can be written for a molecule or ion, the actual distribution of electrons is a weighted average of the valid Lewis structures. It is not possible to write a single Lewis structure for NO 2 − which accurately represents the electronic structure. The two N-O bonds and the two oxygen atoms in nitrite are equivalent in all chemical and physical properties. Experimental evidence, however, establishes that nitrite is symmetric and that both N–O bonds in NO 2 − have the same strength and length. As we will see from the discussion of formal charge below, we would also expect that the electrons would be distributed such that the negative charge would be present on one oxygen atom. A double bond between two atoms is shorter (and stronger) than a single bond between the same two atoms. If this representation is the only correct resonance structure, we would expect the molecule to be asymmetric, meaning the bond lengths between the central nitrogen and the oxygen atoms would be different. Resonance Structuresįollowing the five steps for drawing a Lewis structure we can determine a valid Lewis structure for NO 2 – : comparison of carbon-carbon single, double, and triple bonds. We can apply these same concepts to carbon-carbon single, double, and triple bonds to compare how bond order, bond length and bond strength are related ( Figure 1). The C=O double bond is shorter than the C-O single bond. Higher bond orders generally correlate with shorter bond distances. Bond length is the equilibrium distance between two nuclei. This means that a C=O double bond is stronger than a C-O single bond, and the C=O double bond requires more energy to break than a C-O single bond. The stronger the bond, the more energy that is required to break the bond. Bond strength is a measured quantity: the energy (in kJ/mol) required to break a chemical bond, tabulated in Appendix G. The carbon is singly bonded to each hydrogen atom, which would give each CH bond orders of one.īond order is an index of bond strength: the higher the bond order, the stronger the bond. Looking at the structure of formaldehyde we can see that there is a double bond between the central carbon atom and the oxygen atom giving a CO bond order of two. Textbook content produced by OpenStax is licensed under a Creative Commons Attribution License. We recommend using aĪuthors: Paul Flowers, Klaus Theopold, Richard Langley, William R.
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