Ammonia nitrogen orbitals

The processes of cosmochemistry constrain planetary materials to be one of three types -gas (hydrogen and helium - the most abundant elements of the Universe), ice (water, methane, ammonia, nitrogen - the next most abundant elements in nucleosynthesis), and rock - principally made up of Mg-Fe silicates (Stevenson, 2004). If the standard model of planetary formation is applied to the formation of the Earth and the terrestrial planets, then temperatures in the solar nebula in the vicinity of the Earth s orbit would have been between 500 and 800 K. At these temperatures rock - Fe-Mg silicates, and Fe-Ni metal would condense but not water ice. Micron-sized particles of these minerals would have grown in series of stages into the present configuration of planetary bodies as outlined below and summarized in Table 2.3. 

In ammonia, nitrogen resembles the carbon of methane. Nitrogen is sp -hybridized, but (Table 1.1) has only three unpaired electrons they occupy three of the sp orbitals. Overlap of each of these orbitals with the s orbital of a hydrogen atom results in ammonia (Fig. 1.12). The fourth sp orbital of nitrogen contains a pair of electrons. 

These are octahedral ions with only sigma interactions. The ammonia ligands have no 77 orbitals available for significant bonding with the metal ion. The donor orbital of NH3 is mostly nitrogen orbital in composition, and the other p orbitals are used in bonding to the hydrogens. ... 

Hybridized orbitals readily form chemical bonds because they tend to maximize overlap with other orbitals. However, if the central atom of a molecule contains lone pairs, hybrid orbitals can also accommodate them. For example, the nitrogen orbitals in ammonia are sp hybrids. Three of the hybrids are involved in bonding with three hydrogen atoms, but the fourth hybrid contains a lone pair. The presence of the lone pair lowers the tendency of nitrogen s orbitals to hybridize. (Remember that the tendency to hybridize increases with the number of bonds formed.) Therefore the bond angle in NH3 is 107°, a bit closer to the unhybridized p orbital bond angle of 90°. 

We consider first some experimental observations. In general, the initial heats of adsorption on metals tend to follow a common pattern, similar for such common adsorbates as hydrogen, nitrogen, ammonia, carbon monoxide, and ethylene. The usual order of decreasing Q values is Ta > W > Cr > Fe > Ni > Rh > Cu > Au a traditional illustration may be found in Refs. 81, 84, and 165. It appears, first, that transition metals are the most active ones in chemisorption and, second, that the activity correlates with the percent of d character in the metallic bond. What appears to be involved is the ability of a metal to use d orbitals in forming an adsorption bond. An old but still illustrative example is shown in Fig. XVIII-17, for the case of ethylene hydrogenation. 

Ammonia is a colourless gas at room temperature and atmospheric pressure with a characteristic pungent smell. It is easily liquefied either by cooling (b.p. 240 K) or under a pressure of 8-9 atmospheres at ordinary temperature. Some of its physical and many of its chemical properties are best understood in terms of its structure. Like the other group head elements, nitrogen has no d orbitals available for bond formation and it is limited to a maximum of four single bonds. Ammonia has a basic tetrahedral arrangement with a lone pair occupying one position ... 

Describe the bonding in ammonia assuming sp hybridization of nitrogen In what kind of orbital is the unshared pair What orbital overlaps are involved in the N—H bonds , ... 

Both the LUMO and LUMO + 1 energy levels of the nitronaphthyridines and the values of the coefficients at the carbon and ring nitrogen in the LUMO and LUMO -f 1 orbitals were determined. Using homo = -11.9 eV for ammonia, the values of the stabilization energy AE) for each position of the nitronaphthyridines were calculated. Tire results of the cal-... 

The valence-bond concept of orbital hybridization described in the previous four sections is not limited to carbon compounds. Covalent bonds formed by-other elements can also be described using hybrid orbitals. Look, for instance, at the nitrogen atom in methylamine, CH3NH2, an organic derivative of ammonia (NH3) and the substance responsible for the odor of rotting fish. 

Let s look at an example. In ammonia (NH3), the nitrogen atom is sp hybridized, so all four orbitals arrange in a tetrahedral structure, just as we would expect. But only three of the orbitals in this arrangement are responsible for bonds. So, if we look just at the atoms that are connected, we do not see a tetrahedron. Rather, we see a trigonal pyramidal arrangement ... 

Ammonia is a prime example of a Lewis base. In addition to its three N—H bonds, this molecule has a lone pair of electrons on its nitrogen atom, as Figure 21-1 shows. Although all of the valence orbitals of the nitrogen atom in NH3 are occupied, the nonbonding pair can form a fourth covalent bond with a bonding partner that has a vacant valence orbital available. 

The simplest type of Lewis acid-base reaction is the combination of a Lewis acid and a Lewis base to form a compound called an adduct. The reaction of ammonia and trimethyl boron is an example. A new bond forms between boron and nitrogen, with both electrons supplied by the lone pair of ammonia (see Figure 21-21. Forming an adduct with ammonia allows boron to use all of its valence orbitals to form covalent bonds. As this occurs, the geometry about the boron atom changes from trigonal planar to tetrahedral, and the hybrid description of the boron valence orbitals changes from s p lo s p ... 

The exact reverse of the above is seen with aniline (13), which is a very weak base (pKa = 4-62) compared with ammonia (pKa = 9-25) or cyclohexylamine (pKa = 10-68). In aniline the nitrogen atom is again bonded to an sp2 hybridised carbon atom but, more significantly, the unshared electron pair on nitrogen can interact with the delocalised 7r orbitals of the nucleus ... 

If nitrogen uses only its p orbitals in bond formation, the angle between N-H bonds would be 90°. However, compounds prefer formations in which electrons are as far apart as possible. For ammonia this is made possible by forming a tetrahedral structure in which the angle between the bonds (M-H) is 107°. This is only possible by undergoing sp3 hybridization. 

The simplest compounds to consider here are ammonia and water. It is apparent from the above electronic configurations that nitrogen will be able to bond to three hydrogen atoms, whereas oxygen can only bond to two. Both compounds share part of the tetrahedral shape we saw with 5/ -hybridized carbon. Those orbitals not involved in bonding already have their full complement of electrons, and these occupy the remaining part of the tetrahedral array (Figure 2.21). These electrons are not inert, but play a major role in chemical reactions we refer to them as lone pair electrons. 

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