Sulfur Dioxide (SO2) Lewis Structure

[alingy1]

Hi,

I drew the lewis structure of SO2. The one with 0 formal charges gives me two double bonds. But, I see on the web that one single bond and another double bond is an acceptable resonance structure. What do they mean? Sulfur can expand its octet. Having 0 formal charge is much more advantageous!

 

[Qube]

Sulfur might be able to expand its octet but that doesn't mean other resonance forms can't exist. As long as there is no grouping of like charges - i.e. having positive formal charges on adjacent atoms - your Lewis structure should be fine.

The Lewis structure for sulfur dioxide (SO2) can be drawn as follows:

1. Start with the chemical formula SO2, which tells you that the molecule consists of one sulfur (S) atom and two oxygen (O) atoms.
2. Determine the total number of valence electrons for the atoms in the molecule. Sulfur has 6 valence electrons, and each oxygen atom has 6 valence electrons. So, the total number of valence electrons in SO2 is 6 (for sulfur) + 6 (for the first oxygen) + 6 (for the second oxygen) = 18 electrons.
3. Next, arrange the atoms in the molecule. Place the sulfur atom in the center since it is less electronegative than oxygen.
4. Connect the sulfur atom to each of the oxygen atoms using single bonds (one line) to represent the sharing of two electrons between each pair of atoms.
5. After forming the single bonds, you've used up 4 of the 18 valence electrons. Now, place lone pairs of electrons around the oxygen atoms to satisfy the octet rule (oxygen typically forms double bonds with two lone pairs).

 

[AGNuke]

The catch is, you can show Sulphur having double bonds with Oxygens or having a covalent bond with one AND dative bond with other. Both are equally acceptable.

 

[Qube]

The catch is, you can show Sulphur having double bonds with Oxygens or having a covalent bond with one AND dative bond with other. Both are equally acceptable.

Perhaps not to the more calculation-inclined chemists!

 

[DrDu]

Sulfur can expand its octet.

Nope. This has been disproven about 50 years ago.

 

[alingy1]

I don't understand. When we draw Lewis structures (at undergrad level), we must satisfy formal charges for sure if it can expand its octet. So, why is the other one with a single bond accepted!?

 

[Qube]

I don't understand. When we draw Lewis structures (at undergrad level), we must satisfy formal charges for sure if it can expand its octet. So, why is the other one with a single bond accepted!?

It's because the d-orbitals aren't generally accessed by sulfur. The d-orbitals are too high in energy. So sulfur just sticks with structures that are best described using charge-separated Lewis structures.

 

[alingy1]

Hmm. How can I know if an atom can expand octet?

 

[Qube]

Hmm. How can I know if an atom can expand octet?

Atoms generally don't expand their octets. Unless you're dealing with d-block elements.

 

[abitslow]

Qube is simply wrong (although this is possibly just a difference in semantics). Chlorine pentafluoride comes to mind. PCl5, SF6, the list is endless...
O=S=O ?? So, if I understand you, the sulfur atom, which you surely KNOW is less electronegative than oxygen, is what? - surrounded by 10 electrons? (that is :S(=O)₂ ). comparing the difference in electronegativity, is a zero formal charge reasonable? Ozone with the same (formal) bonding has a dipole moment of 0.53D compared to SO₂ 's 1.62 D. So, if that isn't telling you that there is a LOT of charge separation, I don't know what could. Why will having a zero formal charge result in the lowest energy state ("is advantageous"). News to me. Wikipedia says:" a formal charge is the charge assigned to an atom in a molecule, assuming that electrons in a chemical bond are shared equally between atoms, regardless of relative electronegativity." IOW, "regardless of energy", since electronegativity is a proxy for the extra energy of bonding 'due to' quantum mechanics and was an early (fairly successful) attempt at including canonical ionic contributions to bond energy (ie at 'explaining' why bond energy of A-B is usually greater than the average of the A-A and B-B bond energies). Formal charge is an accounting 'gimmick', and should be viewed as having only very narrow utility. (The same can be said for electronegativity, but formal charge is more often used in organic chemistry, while electronegativity is more often used in inorganic, I think.) (Both can be (carefully) applied, but the reality is that neither are as good as Density Functional, Molecular Orbital, or Valence Bond quantum mechanical treatments of what is inherently a QM problem.) Up till about the 1990's, chemists relied on a lot of rules of thumbs like these, DFT has allowed them to instead rely on (more accurate) calculations and models (but at the expense of intuitiveness). Instead of relying on our crude mental models of the geometry and bonding of chemicals, we can instead look at the results of DFT theory to determine where electron density is. Its an attempt to fit more classical electrostatics to what is properly only solved with the use of QM. As a student, it makes sense to learn the easy stuff first, and then move on to better (but more complex) stuff. The real foundation of chemistry IS quantum mechanics, but that is beyond the mathematical competence of most undergraduates. At the same time, there is a boat load of chemistry which can be (more or less) put into some sort of coherent body of knowledge without knowing the QM. Learning about bonding begins the process of connecting the two.

 

[Qube]

Qube is simply wrong (although this is possibly just a difference in semantics). Chlorine pentafluoride comes to mind. PCl5, SF6, the list is endless...
O=S=O ??

The concept itself has been criticized by Ronald Gillespie who, based on an analysis of electron localization functions, wrote in 2002 that "as there is no fundamental difference between the bonds in hypervalent and non-hypervalent (Lewis octet) molecules there is no reason to continue to use the term hypervalent."[7]

For hypercoordinated molecules with electronegative ligands such as PF5 it has been demonstrated that the ligands can pull away enough electron density from the central atom so that its net content is again 8 electrons or fewer. Consistent with this alternative view is the finding that hypercoordinated molecules based on fluorine ligands, for example PF5 do not have hydride counterparts e.g. phosphorane PH5 which is an unstable molecule.

http://en.wikipedia.org/wiki/Hypervalent_molecule

 

[DrDu]

Qube is simply wrong (although this is possibly just a difference in semantics). Chlorine pentafluoride comes to mind. PCl5, SF6, the list is endless...
.

Wrong in which respect?
Main group atoms not extending the octet?
Well, in the compounds you mention, bonds are highly polar, so that the important resonance structures don't contain more than 4 covalent bonds plus several ionic bonds. But as you suspect, this involves a lot of semantics:
The following paper "Chemical bonding in hypervalent molecules: Is the octet rule relevant?" sums up the discussion
http://alpha.chem.umb.edu/chemistry/Seminar/06-09 WQE/InorgI.Carter.pdf
see also
http://www.chem.umn.edu/groups/tonks/Organometallics/VSEPR/Gillespie_CCR_2002_233_53_Hypervalence.pdf
and also contains some examples where the valence shell of the central atom may contain more than 8 electrons.