How do resonance structures affect molecular shape

The double-headed arrow indicates that the actual electronic structure is an average of those shown, not that the molecule oscillates between the two structures. When it is possible to write more than one equivalent resonance structure for a molecule or ion, the actual structure is the average of the resonance structures. The electrons appear to "shift" between different resonance structures and while not strictly correct as each resonance structure is just a limitation of using the Lewis structure perspective to describe these molecules.

A more accurate description of the electron structure of the molecule requires considering multiple resonance structures simultaneously. Most arrows in chemistry cannot be used interchangeably and care must be given to selecting the correct arrow for the job.

At this point, the carbon atom has only 6 valence electrons, so we must take one lone pair from an oxygen and use it to form a carbon—oxygen double bond. In this case, however, there are three possible choices:. As with ozone, none of these structures describes the bonding exactly. Each predicts one carbon—oxygen double bond and two carbon—oxygen single bonds, but experimentally all C—O bond lengths are identical.

We can write resonance structures in this case, three of them for the carbonate ion:. Like ozone, the electronic structure of the carbonate ion cannot be described by a single Lewis electron structure.

While each resonance structure contributes to the total electronic structure of the molecule, they may not contribute equally. Assigning Formal charges to atoms in the molecules is one mechanism to identify the viability of a resonance structure and determine its relative magnitude among other structures. The formal charge on an atom in a covalent species is the net charge the atom would bear if the electrons in all the bonds to the atom were equally shared.

Alternatively the formal charge on an atom in a covalent species is the net charge the atom would bear if all bonds to the atom were nonpolar covalent bonds. To determine the formal charge on a given atom in a covalent species, use the following formula:.

Resonance: All elements want an octet, and we can do that in multiple ways by moving the terminal atom's electrons around bonds too. Remember to determine the number of valence electron each atom has before assigning Formal Charges. The total of valence electrons is Find the most ideal resonance structure. Note: It is the one with the least formal charges that adds up to zero or to the molecule's overall charge.

The most electronegative atom usually has the negative formal charge, while the least electronegative atom usually has the positive formal charges.

Although, in this course, we shall draw benzene and its derivatives as a single resonance hybrid Structure 5. We conclude with a resonance hybrid which is an ion. Structure 5. Three resonance structures, all equivalent to Structure 5. As you work through this course you will need various resources to help you complete some of the activities. Making the decision to study can be a big step, which is why you'll want a trusted University. Take a look at all Open University courses. If you are new to University-level study, we offer two introductory routes to our qualifications.

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Free learning from The Open University. Featured content. Free courses. Electron delocalization frequently reveals itself through a distorted molecular geometry. The most frequently cited situation is a symmetric set of bond distances in a molecule that does not have a symmetric Lewis structure, but delocalization can also create other kinds of geometrical distortions. The first two sections in this essay examine distortions in bond distances and bond angles.

The last section looks at recent research on the "real" factors that make benzene and possibly many other molecules symmetric. Bond Distances Bond Angles Benzene. Ozone, O 3 , is an excellent example of a symmetrical resonance hybrid with distorted bond distances.

As it happens, the OO distances in ozone are identical, and at pm, they tends toward "short", but we could call them "in between. This kind of analysis is applied to all resonance hybrids. We always expect a hybrid's bond distances to reflect the hybrid's bond pattern. If the hybrid contains a partial double bond, the bond distance should be somewhat longer than a complete double bond, and somewhat shorter than a single bond. Analogous predictions would be made for partial single bonds and partial triple bonds.

VSEPR rules link bond angles to the number of electron pairs surrounding an atom. The counting of electron pairs, and the prediction of bond angles, becomes uncertain when a lone pair is delocalized. Therefore, a resonance hybrid with a delocalized lone pair may display distorted bond angles. Because of charge delocalization, each oxygen atom has two-thirds of a full negative charge. Charge delocalization helps to stabilize the whole species.

The stability a species gains from having charge delocalization through resonance contributors is called resonance stabilization effect. The greater the number of resonance contributors, the greater the resonance stabilization effect, and the more stable the species is. The actual structure of the carbonate anion is a combination of all the three equivalent resonance structures, that can be called a hybrid. What does the actual structure look like, and can we draw one structure on paper to show the actual structure?

The actual structure can not be shown with a conventional Lewis structure, because the regular Lewis structures do not include partial charges, and there is two-thirds of a full negative charge on each oxygen atom in CO 3 An attempt to show the hybrid structure can be by using dashed lines to show that the bond between carbon and oxygen is somewhere between a single and double bond, and each oxygen atom has partial charges.

The delocalized charges can also be represented by the calculated electrostatic potential map of the electron density in the CO 3 2- anion.