# 1.5: Atomic Orbitals

- Page ID
- 13808

After completing this section, you should be able to

- describe the physical significance of an orbital.
- list the atomic orbitals from 1
*s*to 3*d*in order of increasing energy. - sketch the shapes of
*s*and*p*orbitals.

Make certain that you can define, and use in context, the key terms below.

- nodal plane
- node
- orbital
- quantum mechanics
- wave function

For a refresher on quantum numbers view this text

## Atomic Orbitals

An orbital is the quantum mechanical refinement of Bohr’s orbit. In contrast to his concept of a simple circular orbit with a fixed radius, orbitals are mathematically derived regions of space with different *probabilities* of having an electron.

One way of representing electron probability distributions was illustrated in Figure 6.5.2 for the 1*s* orbital of hydrogen. Because Ψ^{2} gives the probability of finding an electron in a given volume of space (such as a cubic picometer), a plot of Ψ^{2} versus distance from the nucleus (*r*) is a plot of the *probability density*. The 1*s* orbital is spherically symmetrical, so the probability of finding a 1*s* electron at any given point depends *only* on its distance from the nucleus. The probability density is greatest at *r* = 0 (at the nucleus) and decreases steadily with increasing distance. At very large values of *r*, the electron probability density is very small but *not* zero.

In contrast, we can calculate the *radial probability* (the probability of finding a 1*s* electron at a distance *r* from the nucleus) by adding together the probabilities of an electron being at all points on a series of *x* spherical shells of radius *r*_{1}, *r*_{2}, *r*_{3},…, *r*_{x} _{− 1}, *r*_{x}. In effect, we are dividing the atom into very thin concentric shells, much like the layers of an onion (part (a) in Figure \(\PageIndex{1}\)), and calculating the probability of finding an electron on each spherical shell. Recall that the electron probability density is greatest at *r* = 0 (part (b) in Figure \(\PageIndex{1}\)), so the density of dots is greatest for the smallest spherical shells in part (a) in Figure \(\PageIndex{1}\). In contrast, the surface area of each spherical shell is equal to 4π*r*^{2}, which increases very rapidly with increasing *r* (part (c) in Figure \(\PageIndex{1}\)). Because the surface area of the spherical shells increases more rapidly with increasing *r* than the electron probability density decreases, the plot of radial probability has a maximum at a particular distance (part (d) in Figure \(\PageIndex{1}\)). Most important, when *r* is very small, the surface area of a spherical shell is so small that the *total* probability of finding an electron close to the nucleus is very low; at the nucleus, the electron probability vanishes (part (d) in Figure \(\PageIndex{1}\)).

For the hydrogen atom, the peak in the radial probability plot occurs at *r* = 0.529 Å (52.9 pm), which is exactly the radius calculated by Bohr for the *n* = 1 orbit. Thus the *most probable radius* obtained from quantum mechanics is identical to the radius calculated by classical mechanics. In Bohr’s model, however, the electron was assumed to be at this distance 100% of the time, whereas in the quantum mechanical Schrödinger model, it is at this distance only some of the time. The difference between the two models is attributable to the wavelike behavior of the electron and the Heisenberg uncertainty principle.

Figure \(\PageIndex{2}\) compares the electron probability densities for the hydrogen 1*s*, 2*s*, and 3*s* orbitals. Note that all three are spherically symmetrical. For the 2*s* and 3*s* orbitals, however (and for all other *s* orbitals as well), the electron probability density does not fall off smoothly with increasing *r*. Instead, a series of minima and maxima are observed in the radial probability plots (part (c) in Figure \(\PageIndex{2}\)). The minima correspond to spherical nodes (regions of zero electron probability), which alternate with spherical regions of nonzero electron probability.

## s Orbitals

Three things happen to *s* orbitals as *n* increases (Figure \(\PageIndex{2}\)):

- They become larger, extending farther from the nucleus.
- They contain more nodes. This is similar to a standing wave that has regions of significant amplitude separated by nodes, points with zero amplitude.
- For a given atom, the
*s*orbitals also become higher in energy as*n*increases because of their increased distance from the nucleus.

Orbitals are generally drawn as three-dimensional surfaces that enclose 90% of the electron density, as was shown for the hydrogen 1*s*, 2*s*, and 3*s* orbitals in part (b) in Figure \(\PageIndex{2}\). Although such drawings show the relative sizes of the orbitals, they do not normally show the spherical nodes in the 2*s* and 3*s* orbitals because the spherical nodes lie inside the 90% surface. Fortunately, the positions of the spherical nodes are not important for chemical bonding.

## p Orbitals

Only *s* orbitals are spherically symmetrical. As the value of *l* increases, the number of orbitals in a given subshell increases, and the shapes of the orbitals become more complex. Because the 2*p* subshell has *l* = 1, with three values of *m*_{l} (−1, 0, and +1), there are three 2*p* orbitals.

The electron probability distribution for one of the hydrogen 2*p* orbitals is shown in Figure \(\PageIndex{3}\). Because this orbital has two lobes of electron density arranged along the *z* axis, with an electron density of zero in the *xy* plane (i.e., the *xy* plane is a nodal plane), it is a 2*p*_{z} orbital. As shown in Figure \(\PageIndex{4}\), the other two 2*p* orbitals have identical shapes, but they lie along the *x* axis (2*p*_{x}) and *y* axis (2*p*_{y}), respectively. Note that each *p* orbital has just one nodal plane. In each case, the phase of the wave function for each of the 2*p* orbitals is positive for the lobe that points along the positive axis and negative for the lobe that points along the negative axis. It is important to emphasize that these signs correspond to the *phase* of the wave that describes the electron motion, *not* to positive or negative charges.

The surfaces shown enclose 90% of the total electron probability for the 2*p*_{x}, 2*p*_{y}, and 2*p*_{z} orbitals. Each orbital is oriented along the axis indicated by the subscript and a nodal plane that is perpendicular to that axis bisects each 2*p* orbital. The phase of the wave function is positive (orange) in the region of space where *x*, *y*, or *z* is positive and negative (blue) where *x*, *y*, or *z* is negative.

Just as with the *s* orbitals, the size and complexity of the *p* orbitals for any atom increase as the principal quantum number *n* increases. The shapes of the 90% probability surfaces of the 3*p*, 4*p*, and higher-energy *p* orbitals are, however, essentially the same as those shown in Figure \(\PageIndex{4}\).

The electron configuration of an atom is the representation of the arrangement of electrons distributed among the orbital shells and subshells. Commonly, the electron configuration is used to describe the orbitals of an atom in its ground state, but it can also be used to represent an atom that has ionized into a cation or anion by compensating with the loss of or gain of electrons in their subsequent orbitals. Many of the physical and chemical properties of elements can be correlated to their unique electron configurations. The valence electrons, electrons in the outermost shell, are the determining factor for the unique chemistry of the element.