Orbitals angular location probability

Figure 117 The angular location probability pattern for a 2p,-orbital. The pattern indicates that the electron is most likely to be found on either side of the nucleus along the r -axis. There is no chance of finding the...

Figure 118 The angular location probability patterns for the d- bitals. The lobes of the d,-orbital are located between the x- and y-axes the lobes of the d, -orbital are between the y- and z-axes perpendicular to the x-axis and the lobes of the d -orbital ate between the x- and z-axes perpendicular to the y-axis.

Another measure of the size of an orbital is the most probable distance of the electron from the nucleus in that orbital. Figure 5.4c shows that the most probable location of the electron is progressively farther from the nucleus in ns orbitals for larger n. Nonetheless, there is a finite probability for finding the electron at the nucleus in both 2s and 3s orbitals. This happens because electrons in s orbitals have no angular momentum ( = 0), and thus can approach the nucleus along the radial direction. The ability of electrons in s orbitals to penetrate close to the nucleus has important consequences in the structure of many-electron atoms and molecules (see later). 

There are always 2n2 possible combinations of quantum numbers. We divide these into orbitals. Orbitals are maps of the probability of the electron being located at a certain region in space. They are designated by their angular momentum quantum numbers. The values of magnetic and spin quantum numbers define the electrons within an orbital. 

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