Nitric oxide molecular orbitals

The molecular orbital description of the bonding in NO is similar to that in N2 or CO (p. 927) but with an extra electron in one of the tt antibonding orbitals. This effectively reduces the bond order from 3 to 2.5 and accounts for the fact that the interatomic N 0 distance (115 pm) is intermediate between that in the triple-bonded NO+ (106 pm) and values typical of double-bonded NO species ( 120 pm). It also interprets the very low ionization energy of the molecule (9.25 eV, compared with 15.6 eV for N2, 14.0 eV for CO, and 12.1 eV for O2). Similarly, the notable reluctance of NO to dimerize can be related both to the geometrical distribution of the unpaired electron over the entire molecule and to the fact that dimerization to 0=N—N=0 leaves the total bond order unchanged (2 x 2.5 = 5). When NO condenses to a liquid, partial dimerization occurs, the cis-form being more stable than the trans-. The pure liquid is colourless, not blue as sometimes stated blue samples owe their colour to traces of the intensely coloured N2O3.6O ) Crystalline nitric oxide is also colourless (not blue) when pure, ° and X-ray diffraction data are best interpreted in terms of weak association into... 

The nitric oxide molecule has one unpaired electron residing in an antibonding -rr molecular orbital. When that electron is removed, the bond order increases from 2.5 to 3, so in coordinating to metals, NO usually behaves as though it donates three electrons. The result is formally the same as if one electron were lost to the metal,... 

The unusual properties of nitric oxide result from an unpaired electron in the highest occupied molecular orbital (HOMO). Nitrogen contains five valence electrons (electrons in the outermost shell with the greatest influence on bonding), while oxygen contains six valence electrons (Fig. 1). Therefore, nitric oxide contains a total of eleven valence electrons. Because orbitals can hold only two electrons with each electron possessing an opposite spin, there must be a single... 

Lewis dot diagrams of nitric oxide compared to the nitrosonium ion and molecular nitrogen. Each bond contains one electron from each atom. These simple diagrams fail to properly account for the effective bond order of 2.5 predicted by molecular orbital theory and must be only considered as illustrative. The dimer of two nitric oxide molecules has five bonds, which is the same as two individual molecules. Thus, nitric oxide remains dissociated at room temperatures. 

Most of biological chemistry can be understood in terms of simple ball and stick models. The chemistry of nitric oxide and related oxides is more intimidating because its patterns of bonding depend strongly on quantum mechanics and molecular orbital theory. But the basics can be grasped by comparison to other molecules and a simple consideration of where nitrogen sits in the periodic table. 

A second illustrative example of the utility of TRPES and TRCIS for studying complex molecular photodissociation dynamics that involve multiple electronic state is the case of the weakly bound cis-planar C2V nitric oxide dimer [174], The weak (Do = 710cm-1) 1 A ground-state covalent bond is formed by the pairing of two singly occupied ji orbitals, one from each NO(X2II) monomer. The very intense UV absorption spectrum of the NO dimer appears... 

An example of free radical formation is molecular oxygen, which can accept electrons from a variety of sources to produce reactive oxygen species (ROS) such as the superoxide radical, the hydroxyl radical, and the nitric oxide radical. The superoxide anion radical is formed when one electron is taken up by one of the 2p orbitals of molecular oxygen. Certain drugs and other xenobiotics have the capacity to undergo so-called redox cycles, whereby they provide electrons to molecular oxygen and form super oxide. 

A Lewis structure cannot be written for NO, for it has an odd number of electrons. Pauling interprets its structure as having a three-electron bond and estimates that the extra electron adds about half as much extra stability to the molecule as would an ordinary covalent bond. Alternatively, the molecular orbital picture describes this molecule as having one a bond and f ir bonds (p. 71). At any rate, nitric oxide is, as it should be, paramagnetic. 

General. Nitric oxide readily forms complexes with transition metals and is in many ways similar to carbon monoxide. It has a single electron in a tt antibonding electron, which is easily lost. When NO is not bridging, molecular orbital theory would predict a linear M-NO+ moiety and that M-NO would be bent. In principal, this would seem to allow a ready indication of the metal oxidation state from an X-ray... 

The molecular orbital energy-level diagram for nitric oxide (NO). The bond order is 2.5, or (8 - 3)/2. 

The case of nitric oxide is more straightforward and the formation of the molecular orbital may be represented as,... 

A comparison of this structure with the electronic configuration according to the molecular orbital treatment brings out an important similarity. The molecular orbitals of nitric oxide are... 

The various classes of reactions that are commonly treated can be rationalized by a consideration of the structure of nitric oxide which is represented in molecular orbital terminology as ... 

Reactions which involve addition of an electron to produce NO . This species is produced when nitric oxide is bubbled into a solution of potassium in liquid ammonia (15). Although on the basis of molecular orbital theory a para-... 

Resonance forms for the sulfite ion have not been included. There has been discussion concerning whether the compound should be considered a sulfonated hyponitrite (16) or a nitrosated hydroxylamine sulfonate (9). The compound can be prepared by nitrosating a sulfonated hydroxylamine (9). There is little to be gained by this dispute, for the structure is best described as in the above diagrams or in molecular orbital terminology. The reaction of sulfite ion and nitric oxide is best considered as a Lewis acid-base reaction in which nitric oxide behaves as the acid and sulfite ion as the donor. A reaction sequence can be formulated with the following equations ... 

Prepare a molecular orbital energy level diagram for nitric oxide (NO) and predict the bond order of this molecule. On the basis of the molecular orbital diagram, what do you predict for the bond orders of NO+ and NO Which of these diatomic species would you expect to have the shortest bond length Why ... 

Molecular orbital theory then explains why NO is stable with respect to N and O atoms, and predicts that NO+ will form fairly easily. We have not explained why nitric oxide gas consists of NO molecules rather than N202 molecules, or why NO is unstable with respect to oxidation by oxygen to N02. To do this, we would have to consider orbital energy-level diagrams for N202 or N02 and 02 as well. 

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