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×08th Feb 2020 @ 14 min read

The atomic orbital is a complex mathematical function called a wave function, which decides the energy, angular momentum, and location of an electron. A better way to define the atomic orbital is the space around the nucleus which has a high probability of finding the electron.

The simplest orbital of all is 1s orbital, which is spherical in shape (see figure below). The nucleus is at the center of the orbital, and the electron revolves around the nucleus.

The region surrounding the nucleus is the probabilistic region, not discrete. The probability of finding the electron in this region is high (say 90 %), and the electron might exist beyond it. Thus, an orbital is an electronic probabilistic cloud around the nucleus.

In quantum mechanics, a wave function is a complex mathematical description of a quantum state. It provides the probability distribution of an electron and is denoted by Ψ. Every orbital has its own Ψ. |Ψ|^{2} represents the probability density function. The probability of an electron at the position *x* between a and b is given by the equation below.

We can normalize the equation.

Thus, the size of the orbital, when the probability of finding an electron is 100 %, is equal to the size of the universe, i.e. −∞ to ∞. Practically, we define the orbital based on the probability of 90 %, not 100 %.

The hydrogen atom has only one electron and only atom for which the Schrodinger equation can be solved exactly. The wave function for the hydrogen atom in the spherical coordinate system is as follows:

*R*(*r*) is the radial function, which decides the nuclear distance while *Y ^{m}_{l}*(

Here, *a*_{0} is the Bohr radius, , and *L*^{2l + 1}_{n − l − 1} is a generalized Laguerre polynomial.

As you can see from the above, the quantum numbers *n*, *l*, and *m* decide the equation of the wave function Ψ. For each set of these three quantum numbers, we get a new wave function. Hence, we can conclude these three quantum numbers represent an orbital.

The principal quantum number *n* is a set of positive integers; *n* = 1, 2, 3… The azimuthal quantum number *l* is a set of non-negative integers ranging from 0 to (*n* − 1); *l* = 0, 1, 2 … (*n* − 2), (*n* − 1). Finally, the magnetic quantum number *m* is a set of integers from −*l* to *l*; *m* = −*l*, −(*l* − 1) … −1, 0, 1 … (*l* − 1), *l*.

An atom can have a large number of orbitals. The three quantum numbers mentioned in the section above are used to identify an orbital in an atom.

In an atom, the electronic space is divided into shells. Each shell is represented by a value of the principal quantum number *n*. For example, *n* = 1 is the lowest energy shell, called K shell, *n* = 2 is the second shell, L shell, and so on.

Within every shell, there are subshells. The number of subshells in a shell is decided by the azimuthal quantum number *l*. For every value of *n*, the possible values of *l* are 0, 1, 2 … *n* − 2, *n* − 1. Thus for K shell (*n* = 1), *l* = 0, for L shell (*n* = 2), *l* = 0, 1; for M shell (*n* = 3), *l* = 0, 1, 2, and so forth. Each value of *l* corresponds to a subshell. *l* = 0 is an s subshell, *l* = 1 is a p subshell, *l* = 2 is a d subshell, *l* = 3 is an f subshell etc. The table below summarizes the same.

Principal quantum number n | Shell | Azimuthal quantum number l | Subshell | No. of subshells |
---|---|---|---|---|

1 | K | 0 | s | 1 |

2 | L | 0, 1 | s, p | 2 |

3 | M | 0, 1, 2 | s, p, d | 3 |

4 | N | 0, 1, 2, 3 | s, p, d, f | 4 |

5 | O | 0, 1, 2, 3, 4 | s, p, d, f, g | 5 |

6 | P | 0, 1, 2, 3, 4, 5 | s, p, d, f, g, h | 6 |

7 | Q | 0, 1, 2, 3, 4, 5, 6 | s, p, d, f, g, h, i | 7 |

8 | R | 0, 1, 2, 3, 4, 5, 6, 7 | s, p, d, f, g, h, i, k | 8 |

… | … | 0, 1, 2, 3, 4, 5, 6, 7… | s, p, d, f, g, h, i, k… | … |

Note: In the naming of subshells, the letter j is skipped. |

From the above table, the *n*-th shell has *n*'s subshells. Each of these subshells consists of a number of orbitals, which is determined by the magnetic quantum number *m*. For every *l*, the possible *m*'s are 0, ±1, ±2 … ±(*l* − 1), ±*l*. Each value of *m* corresponds to an orbital.

For an s subshell (*l* = 0), *m* = 0. Thus, it has only one orbital, called s orbital. For a p subshell (*l* = 1), *m* = 0, ±1. Therefore, a p subshell has three orbitals: p_{z} (*m* = 0), and p_{x} and p_{y} (*m* = ±1).

In a d subshell (*l* = 2), there are five orbitals: d_{z2} (*m* = 0), d_{xz} and d_{yz} (*m* = ±1), and d_{xy} and d_{x2 − y2} (*m* = ±2).

In an f subshell, (*l* = 3), there are seven orbitals: f_{z3} (*m* = 0), f_{xz2} and f_{yz2} (*m* = ±1), f_{xyz} and f_{z(x2 − y2)} (*m* = ±2), and f_{x(x2 − 3y2)} and f_{y(3x2 − y2)} (*m* = ±3).

Subshells | Azimuthal quantum number l | Magnetic quantum number m | Orbital |
---|---|---|---|

s | 0 | 0 | s |

p | 1 | 0 | p_{z} |

±1 | p_{x} and p_{y} | ||

d | 2 | 0 | d_{z2} |

±1 | d_{xz} and d_{yz} | ||

±2 | d_{xy} and d_{x2 − y2} | ||

f | 3 | 0 | f_{z3} |

±1 | f_{xz2} and f_{yz2} | ||

±2 | f_{xyz} and f_{z(x2 − y2)} | ||

±3 | f_{x(x2 − 3y2)} and f_{y(3x2 − y2)} |

Other higher orbitals, viz. g, h, i, and so on, are very complicated and hardly encountered.

If notice from the previous table, the number of orbitals increases oddly with the azimuthal quantum number. The number of orbitals of a given *l* is 2*l* + 1.

For a given value of the principal quantum number *n*, the number of orbitals is given by *n*^{2}.

The possible values of *l* for given *n* are 0, 1, 2 … *n* − 2, *n* − 1, and for every *l* there are 2*l* + 1 orbitals. The number of orbitals for a given *n*:

An orbital can hold the upmost two electrons, so the number of electrons for a given *n* is 2*n*^{2}.

Commonly, the orbital is named with a combination of numbers and alphabets. Consider 4d_{z2} orbital; the first number represents the principal quantum number (or shell). Here, it is 4 (N shell). The number is followed by an alphabet, here d, it stands for a subshell. The third is the subscript to the alphabet. It tells the orientation of an orbital. Thus, 4d_{z2} orbital is in the d subshell of the 4th shell oriented along *z*-axis.

Orbitals can also be represented simply by the three quantum numbers, for example, 4d_{z2} is 420 (*n* = 4, *l* = 2, *m* = 0).

Let take one more example: *n* = 5, *l* = 3, *m* = ±1 corresponds 5f_{xz2} and 5f_{yz2} orbitals. Both lie in the f subshell of the 5th shell. 5f_{xz2} rests in the *xz* plane while 5f_{yz2} rests in the *yz* plane.

Every orbital has a unique shape, and the shape becomes more complex and difficult to follow as we move toward higher orbitals.

The s orbital is spherical in shape; the nucleus resides at the center of the sphere. It does not orient itself in any direction. In other words, it is non-directional.

There are three dumbbell-shaped p orbitals. Each orbital has two lobes aligned in one of the three axes. The p_{z} orbital is aligned with the *z*-axis, p_{x} with the *x*-axis and finally p_{y} with the *y*-axis.

The d_{z2} orbital has two lobes in the *z*-axis and a doughnut in the *xy* plane. The remaining four d orbitals have two dumbbells to each. d_{xz} and d_{yz} lie in the *xz* and *yz* planes while the other two, d_{xy} and d_{x2 −y2}, are in the *xy* plane. d_{xz} and d_{yz} are identical to each other. They have the same radial component of the wave function and only differ by an angle of 90°. Similarly, d_{xy} and d_{x2 −y2} are identical to each other and differ by an angle of 45°.

The f_{z2} orbital has two lobes along the *z*-axis and two doughnuts in-between. f_{xz2} and f_{yz2} are similar to the d_{xz} and d_{yz} except they have two extra bean-shaped lobes, which are aligned in the *x*-axis in f_{xz2} and in the *y*-axis in f_{yz2}. f_{xyz} and f_{z(x2 − y2)} have eight lobes; four of them are below the *xy* plane and the rest four above. f_{x(x2 − 3y2)} and f_{y(3x2 − y2)} have six lobes each; all the lobes are placed in the *xy* plane. All the above three pairs are identical to each other and have the same radial component but a different angular component.

Except for the s orbital, all orbitals are directional—they orient themselves in a specific direction.

For s orbitals, *l* and *m* are always zero, the only *n* varies. When *n* = 1, it is 1s orbital. The wave function of 1s orbital for the hydrogen atom can be obtained by substituting *n*, *l*, and *m* as 1, 0, 0 in the generalized wave function mentioned earlier.

Ψ^{2} is the probability density function. The probability density function is the probability of finding an electron per unit volume. It is maximum at *r* = 0 and tends to zero with increasing *r* (see the first graph below). We can obtain the radial probability if we multiply the area of the sphere, 4π*r*^{2}, to the probability density. The third graph shows the radial probability with the radius. At *r* = 0 and *r* = ∞, the probability is zero. From the graph, the probability reaches a peak at *r* = 52.9 pm, which is also the Bohr's radius (*a*_{0}).

When *n* = 2 and *n* = 3, we get 2s and 3s orbitals. Their wave functions are as follows:

The graph below is the plot of the radial probability (Ψ^{2}*r*^{2}) versus the radius for 1s, 2s, and 3s orbitals.

As seen from the above graph, 1s orbital has only one peak while 2s and 3s have two and three peaks respectively. As a consequence, there is point of the minimum probability between every two peaks. Besides the extrema, the probability becomes zero at these points. These points, where the probability becomes zero, are called nodes. 1s orbital has zero nodes, 2s orbital has one node, and 3s has two nodes; in general, we can say *n*s orbital has (*n* − 1) nodes.

Note: The different colors in the diagrams indicate different phases of the orbital, just like the crest and trough in the wave.

The nodes in orbitals are similar to the nodes of the standing wave, where they are defined as the points of zero amplitudes. One can imagine the nodes as the gap between spheres that placed one within another.

One can also observe from the above diagram the size of the orbital increasing with the principal quantum number—i.e., 3s > 2s > 1s.

The p orbital corresponds to *l* = 1. The p orbital, unlike the s orbital, is not spherical in shape; it is dumbbell-shaped. As discussed earlier, there are three p orbitals: p_{z} (*l* = 1; *m* = 0), and p_{x} and p_{y} (*l* = 1; *m* = ±1).

For *n* = 2, the wave functions are as follows:

Ψ_{210} is the wave function for 2p_{z} orbital and Ψ_{21±1} is for 2p_{x} and 2p_{y}. The *θ* and *ϕ* part of Ψ decides the orientation of the orbital. In the graphs below, only the radial part of Ψ is considered, i.e., *θ* and *ϕ* is ignored.

Unlike the s orbital, the probability density of the p orbital is not maximum at *r* = 0. We can see from the above figure it is zero at *r* = 0. The value attained the peak in-between. The radial probability graph also shows a similar behavior. Since the probability density is zero at the nucleus, the electron will spend more time away from the nucleus in the p orbital in comparison to the s orbital, where the probability density is maximum at the nucleus.

Also, the probability is zero in the *xy* plane for 2p_{z} orbital. Similarly, the corresponding planes for 2p_{x} and 2p_{y} are *yz* and *xz*. Such planes, where the probability becomes zero, are called nodal planes. In p orbitals, there are three nodal planes as shown below.

These nodal planes are known as angular nodes. Thus, p orbitals have three angular nodes. The s orbital does not have angular nodes; they only have radial nodes. p orbitals have both radial and as angular nodes.

The wave functions of 3p_{z}, 3p_{x} and 3p_{y} are as follows:

The plot below shows the radial node in 3p_{z} orbital. 2p_{z} does not have any radial nodes.

In 4p_{z} orbital, not shown in the above figure, has two radial nodes. In short, the *n*p orbital has (*n* − 2) radial nodes.

There are five d orbitals: d_{z2}, d_{xz}, d_{yz}, d_{xy}, and d_{x2 − y2}. Of these orbitals, d_{z2} is unique; it has two lobes in the *z*-axis and a doughnut-shaped lobe in the *xy* plane. d orbitals start from the 3rd shell and their wave functions are mentioned below.

In all the above wave functions, the angular parts are different while the radial parts are a multiple of one another.

Like p orbitals, d orbitals have angular nodes. Each d orbital has 2 angular nodes.

The radial nodes for the *n*d orbital are (*n* − 3) nodes. So 3d has no radial nodes. The radial nodes in the d orbital start from the 4d.

The f orbital is far more complex than the d orbital and observed in heavy elements. There are seven f orbitals: f_{z3}, f_{xz2}, f_{yz2}, f_{xyz}, f_{z(x2 − y2)}, _{x(x2 − 3y2)}, and f_{y(3x2 − y2)}.

Each of these orbitals has 3 angular nodes. The radial nodes start from 4f; *n*f orbital has (*n* − 4) radial nodes.

Nodes, as discussed above, are the region with zero probabilities of finding electrons. They can be divided into two types: angular and radial. The radial nodes are obtained from the radial component of the wave function while the angular nodes from the angular component, i.e., spherical harmonics.

orbital | n | l | Radial nodes | Angular nodes | Total nodes |
---|---|---|---|---|---|

1s | 1 | 0 | 0 | 0 | 0 |

2s | 2 | 0 | 1 | 0 | 1 |

3s | 3 | 0 | 2 | 0 | 2 |

2p | 2 | 1 | 0 | 1 | 1 |

3p | 3 | 1 | 1 | 1 | 2 |

4p | 4 | 1 | 2 | 1 | 3 |

3d | 3 | 2 | 0 | 2 | 2 |

4d | 4 | 2 | 1 | 2 | 3 |

5d | 5 | 2 | 2 | 2 | 4 |

4f | 4 | 3 | 0 | 3 | 3 |

From the above table, we generalized the formula for nodes. The angular nodes only depend on the value of *l* and is equal to *l*. On the other hand, the radial nodes depend on both *n* and *l* and is given by *n* − *l* − 1. Therefore, the total number of nodes is the sum of both and is *n* − 1.

Orbits and orbitals are often confused with each other, particularly by beginners. They sound the same but are completely different concepts and should not be interchangeably used. The table below explained the difference between both.

Orbits | Orbitals |
---|---|

Orbits are fundamental to the Bohr model, a predecessor of the quantum mechanics model; it describes the orbit as a circular path followed by an electron. Orbits are concentric circular paths of electrons. | An orbital is a space around that the nucleus of an atom that has a high probability of finding an electron. |

Orbits originated from the Bohr-Rutherford atomic model. | The notion of orbitals was derived from the quantum model of an atom. |

Orbits are always concentric circular paths. In the modified Bohr model, orbits have an elliptical shape. | Orbitals come in multiple shapes and their shapes become incomprehensible as we move toward higher orbitals. The simplest of them is the s orbital, which has a spherical shape. |

The orbit is a two-dimensional figure. | Orbitals are in three dimensions. |

They are non-directional. | They are directional except s orbitals. |

Because orbits have a definite path, we can predict its position and momentum. This contradicts the Heisenberg uncertainty principle. | We cannot estimate the exact position of an electron in an orbital. We can only find the probability of an electron; the uncertainty always remains. |

Orbits are not real; they do not exist, and electrons do not revolve around the nucleus in a definite path. | Orbitals are realistic. |

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