0

Why doesn't unstable odd electron compound $\ce{NO}$ dimerize to $\ce{N2O2}$? Why doesn't this structure of ozonide dimerize? But why then does $\ce{BH3}$ dimerizes to $\ce{B2H6}$, and $\ce{AlCl3}$ to $\ce{Al2Cl6}$, both with something like bridged bonds?

enter image description here

Mithoron
  • 4,546
  • 14
  • 40
  • 61
RE60K
  • 3,866
  • 3
  • 21
  • 50

1 Answers1

8

Odd electron species do dimerize.

$\ce{NO}$

Nitric oxide dimerizes, but only at low temperature (and probably high pressure}.

$\ce{NO2}$

Nitrogen dioxide does dimerize. In fact, this is a well known property of $\ce{NO2}$. $\ce{NO2}$ (orange-brown) is in equilibrium with $\ce{N2O4}$ (colorless).

$$\ce{2NO2 <=> N2O4}$$ The position of this equilibrium is dependent on temperature and pressure (like all gas equilibria), and so it makes a nice demonstration of gas properties.

$\ce{O3}$

Ozone does not dimerize. Dimerization would lead to $\ce{O6}$, which likely is a 6-membered ring of all oxygen atoms. While this structure violates no rules, the oxygen-oxygen single bond is not stable. This instability of the peroxide bond makes it both useful and dangerous.

So why does $\ce{NO2}$ dimerize readily, $\ce{NO}$ dimerize under duress, and $\ce{O3}$ not dimerize at all?

The extra electron in the nitrogen oxides is found in an antibonding orbital. Conversion to the dimer gets that electron into a bonding orbital. There are no antibonding electrons in $\ce{O3}$.

$\ce{BH3}$

Borane dimerizes because the boron atom is electron deficient - it has an empty p orbital. This species is so electron deficient that it will form bonding interactions with nearly any electron pairs, including those already in sigma bonds. (More properly the B-H-B bonds are three-center-two-electron bonds). Aluminum compounds dimerize the same way.

So why does $\ce{BH3}$ dimerize and $\ce{O3}$ does not?

Ozone has no electron deficient atoms. Count the electrons. All atoms have an octet, and while one has a positive formal charge, the compound as a whole is not electron deficient.

Jan
  • 67,989
  • 12
  • 201
  • 386
Ben Norris
  • 42,831
  • 8
  • 123
  • 181
  • Molecular oxygen is pretty stable despite its single bond. – Dissenter Nov 28 '14 at 22:22
  • 4
    Molecular oxygen in its ground state has a bond order of 2. – Ben Norris Nov 28 '14 at 22:28
  • http://chemwiki.ucdavis.edu/@api/deki/files/4228/image017.png?revision=1 – Dissenter Nov 28 '14 at 22:30
  • 4
    http://en.wikipedia.org/wiki/Molecular_orbital_diagram#Dioxygen - The Lewis dot structure of $\ce{O2}$ is insufficient to represent the molecule as a diradical with bond order of 2. https://www.youtube.com/watch?v=PFuYWFWC7-w – Ben Norris Nov 28 '14 at 22:42
  • will the edit help explain more, i edited the question too – RE60K Nov 29 '14 at 04:05
  • i'm not talking about ozone, it's ozonide! – RE60K Nov 30 '14 at 12:52
  • Apologies. Your structure for ozonide looks like it has a positive charge on the rightmost oxygen atom: $\ce{^{-}O-O-O+ <-> ^{-}O-O^{+}=O}$ and not a radical: $\ce{^{-}O-O-O.}$ – Ben Norris Nov 30 '14 at 21:11
  • You might include http://pipeline.corante.com/archives/2014/10/10/things_i_wont_work_with_peroxide_peroxides.php in your useful and dangerous list. – Dan D. Dec 12 '14 at 12:59