The sum or mixture of all possible Lewis structures for it. Resonance is indicated by a double-headed arrow.
Summary
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- 1 Resonance
- 2 Lewis structure
- 3 The oxidation number
- 4 Hypervalence
- 5 Traditional explanation of hypervalence
- 6 Sources
Resonance
Resonance (also called Mesomeria ) in chemistry is a tool used (predominantly in organic chemistry) to represent certain types of molecular structures . The resonance consists of the linear combination of structures of a molecule (resonant structures) that do not coincide with the real structure, but that by means of their combination brings us closer to its real structure.
The effect is used in a qualitative way, and describes the electron attraction or release properties of the substituents, based on relevant resonant structures, and is symbolized by the letter R or M (sometimes also by the letter K). The resonant or mesomeric effect is negative (-R / -M) when the substituent is an electron- attracting group , and the effect is positive (+ R / + M) when, after resonance, the substituent is a group that releases electrons.
The molecular resonance is a key component in the theory of the covalent bond . For its existence, the presence of double or triple bonds in the molecule is essential. The net flow of electrons to or from the substituent is also determined by the inductive effect . The mesomeric effect as a result of the p orbital overlap (resonance) has no effect on this inductive effect, since the inductive effect is exclusively related to the electronegativity of the atoms, and their structural chemistry (which atoms are connected to which).
Lewis structure
A single Lewis structure sometimes does not adequately describe a molecule. Thus, for example, for the ozone molecule two Lewis structures can be drawn (Figure 1). However, neither one, separately, correctly represents the geometry of said molecule . Each has a single and a double link, when the two distances O central-O terminal must be identical (1.28 Å). This value appears between that of a single bond (OO, 1.48 Å) and that of a double bond (O = O, 1.21 Å).
Figure 1. Lewis structures for ozone.
To explain this apparent deficiency in Lewis’s theory, the concept of resonance must be introduced. According to this, the structure of a molecule can be represented by the sum or mixture of all possible Lewis structures for it. Resonance is indicated by a double-headed arrow:
The resonant structure should be considered as a mixture of the different structures and not as a balance or rapid exchange between them. In terms of quantum mechanics , the electronic distribution of each of the structures is represented by a wave function, with the real wave function of the molecule Y being a linear combination of the wave functions corresponding to each of the structures. resonant or canonical forms:
Resonance is all the more important when there are several contributing structures with the same energy, as described for the O3 molecule.
In these cases, all the resonant structures contribute equally to the hybrid in resonance. But if the different resonant structures have different energies, the contribution to the hybrid in resonance will be less important the higher the energy of the structure. That is, those lower energy resonant forms more closely resemble the actual shape of the molecule. Two non-equivalent Lewis structures, I and II, can be written for the BF3 molecule. The wave function comes equally from? = cI? I + cII? II, but in this case cI? cII. Which of them contributes the most to the hybrid in resonance?
The decision to be made about which Lewis structure is the one with the lowest energy and, consequently, contributes predominantly to the hybrid in resonance depends markedly on the distribution of formal charges on each atom. Although the total charge of a molecule is globally distributed throughout the structure, it is often useful to assign a formal charge to each atom of it. The formal charge of an atom in a Lewis structure can be calculated using the following expression:
Formal Load = Nv – Nps – ½ Npc
Where Nv is the number of valence electrons of the atom, Nps is the number of electrons in solitary (or unshared) pairs that this atom presents in the Lewis structure and ½ Npc is the number of electrons in shared (or bonding) pairs. ). That is, the formal charge is the difference between the number of electrons that the free atom has and the number of electrons that it has when it is part of the molecule. In an ideal sense, formal charge indicates the number of electrons that an atom gains or loses when it is involved in a covalent bond with another atom. The sum of the formal charges for a Lewis structure is always equal to the total charge of the chemical species.
The resonant structure with the least energy is one in which: 1) the formal charges on the individual atoms are the lowest and 2) the most electronegative atoms bear negative charges and the least electronegative charges the positive. Examples: NO3¯ and BF3. The resonant structures for chemical species NO3¯ are: The calculation of the formal charges of N and O gives the following results. In structure I, each oxygen has a formal charge of -1 and nitrogen of +2. In structures II, III and IV, the oxygen bound to the nitrogen by double bond present a zero formal charge and those linked by a single bond, formal charge -1. the nitrogenit has a formal charge of +1 in the three structures. According to the criteria mentioned above, these three structures are more stable than the first, and contribute more to the resonance hybrid.
As for BF3, whose Lewis structures appear above, it is I that contributes the most to the resonant hybrid, since II presents charge separation and also fluorine supports positive formal charge, being more electronegative than the boron atom. Therefore, it can be affirmed that since the resonant form I has less energy than II, it will contribute to a greater extent to the resonance hybrid, that is, cI> cII. Figure 2 shows, qualitatively, the energy order of the resonant forms and of the real structure for the molecules of ozone and boron trifluoride. In that one the two forms I and II contribute equally to the resonant hybrid, while in this one one of the forms is more stable in energy terms, and therefore contributes more to the resonant form.
Figure 2. Energy order of the resonant forms and resonance energies for O3 and BF3.
From the previous figure it can be extended how, in the case that for a molecule only one Lewis structure can be written, that is coincident with the real structure, and its energy is that of the real molecule, with no contribution from the resonant energy.
The oxidation number
The concept of formal charge is based on the assumption that each and every one of the bonds that an atom forms in a molecule is of the covalent type. On the contrary, the oxidation number is a concept that arises from the opposite situation, that is, from the ionic character of all these bonds. It is usually defined as the effective ionic charge that an atom would have if the pair of electrons in the bond belonged to the most electronegative atom. Consider again the Ion NO3-:
Each oxygen atom (element more electronegative than N) has one more pair of electrons than it has in its valence shell in the free state (6): therefore the oxidation number for this atom is -2. The N atom would not have any electron and its oxidation state is +5. When an element is assigned a certain oxidation number, that element is said to be in a specific oxidation state. Thus, when the nitrogen has the oxidation number +5 it is said to have the oxidation state +5.
In practice, oxidation numbers are assigned by applying the following rules:
- The sum of the oxidation numbers of all the atoms is equal to the total charge of the chemical species.
- Atoms in their elemental form have oxidation numbers equal to zero.
- The atoms of the elements in group 1 have the oxidation number +1. Those of group 2 the oxidation number +2. Those in group 13, except B, have oxidation numbers +3 and +1. Those in group 14, except for C and Si, have oxidation numbers +4 and +2.
- Hydrogen has oxidation number +1 in its combinations with nonmetals and -1 when combined with metals.
- Fluorine has oxidation number -1 in all its compounds.
- Oxygen has oxidation numbers -2 (as long as it is not combined with fluorine), -1 when it is as peroxide ion, O22-, -1/2 in superoxides, O2¯ and -1/3 in ozones, O3¯.
- Halogens have the oxidation number -1 in all their compounds as long as they do not combine with oxygen or another more electronegative halogen.
Hypervalence
The elements of period 2 fulfill the octet rule, but those of period 3 and following show deviations from it. For example, the existence of the PCl5 molecule can only be explained, in terms of Lewis’s theory, if the P atom has 10 electrons in its valence shell. Similarly, in the very stable compound SF6 molecule, the S atom must have 12 electrons in its valence shell if each fluorine atom is linked to the central S atom by an electron pair. Species of this type, for which the Lewis structures demand the existence of more than 8 electrons in the valence shell of one of their atoms, are called hypervalent. Those species whose resonant structures include expansions of the octet, but which do not necessarily have more than 8 valence electrons,
Traditional explanation of hypervalence
The traditional explanation of hypervalence and expansion of the octet ruler makes use of the empty, low-energy d orbitals of the central atom, which can accommodate the additional electrons. According to this explanation, a P atom can accommodate more than 8 electrons in its valence shell if it makes use of its empty 3d orbitals. In the PCl5 molecule , for example, P uses at least one of its 3d orbitals. The absence of hypervalence in the second period is due to the absence of d-orbitals at level 2. A more compelling argument to explain this absence could be the geometric difficulty in accommodating more than four atoms around a small central atom. Later, an alternative explanation will be offered that obviates the contest of the orbitals. D.
