An atomic orbital is the description of the behavior of an electron in an atom according to quantum mechanics. It describes the probability of the electron being in any location, and its energy (see Electron orbital for more background details). The orbitals of the hydrogen atom are used to describe the electron distribution of all other (many-electron) atoms. This does not fully include the interaction between the different electrons, so it is an approximation, although a good and very useful one.

A given (hydrogen-like) atomic orbital is identified by unique values of three quantum numbers: n, l, and m_l. The rules restricting the values of the quantum numbers, and their energies (see below), explain the electron configuration of the atoms and the periodic table.

Contents

1 The various types of orbitals 2 The shapes of orbitals 3 Orbital energy 4 Electron placement and the periodic table 5 Related topics 6 External links

The various types of orbitals

An atomic orbital is uniquely identified by the values of the three quantum numbers, and each set of the three quantum numbers corresponds to exactly one orbital, but the quantum numbers only occur in certain combinations of values. The rules governing the possible values of the quantum numbers are as follows:

The principal quantum number n is always a positive integer. In fact, it can be any positive integer, but for reasons discussed below, large numbers are seldom encountered. Each atom has, in general, many orbitals associated with each value of n; these orbitals together are sometimes called a shell.

The orbital angular momentum quantum number <math>\ell<math> is also a positive integer. Within a shell where n is some integer n₀, <math>\ell<math> ranges across all (integer) values satisfying the relation <math>0 \le \ell \le n_0-1<math>. For instance, the n = 1 shell has only orbitals with <math>\ell=0<math>, and the n = 2 shell has only orbitals with <math>\ell=0<math>, and <math>\ell=1<math>. The set of orbitals associated with a particular value of <math>\ell<math> are sometimes collectively called a subshell.

The magnetic quantum number <math>m_\ell<math> is also always an integer. Within a subshell where <math>\ell<math> is some integer <math>\ell_0<math>, <math>m_\ell<math> ranges thus: <math>-\ell_0 \le m_\ell \le \ell_0<math>.

The above results may be summarized in the following table. Each cell represents a subshell, and lists the values of <math>m_\ell<math> available in that subshell. Empty cells represent subshells that do not exist.

	<math>l=0<math>	1	2	3	4	...
<math>n=1<math>	<math>m_l=0<math>
2	0	-1, 0, 1
3	0	-1, 0, 1	-2, -1, 0, 1, 2
4	0	-1, 0, 1	-2, -1, 0, 1, 2	-3, -2, -1, 0, 1, 2, 3
5	0	-1, 0, 1	-2, -1, 0, 1, 2	-3, -2, -1, 0, 1, 2, 3	-4, -3, -2 -1, 0, 1, 2, 3, 4
...	...	...	...	...	...	...

Subshells are usually identified by their <math>n<math>- and <math>\ell<math>-values. <math>n<math> is represented by its numerical value, but <math>\ell<math> is represented by a letter as follows: 0 is represented by 's', 1 by 'p', 2 by 'd', 3 by 'f', and 4 by 'g'. For instance, one may speak of the subshell with <math>n=2<math> and <math>\ell=0<math> as a '2s subshell'.

(Historical note: The names 's', 'p', 'd', and 'f' originate from a now-discredited system of categorizing spectral lines as "strong", "principal", "diffuse", or "fundamental". When the first four types of orbitals were described, they were associated with these spectral line types, but there were no other names. The designations 'g' and 'h' were derived by following alphabetical order.)

The shapes of orbitals

Any discussion of the shapes of electron orbitals is necessarily imprecise, because a given electron, regardless of which orbital it occupies, can at any moment be found at any distance from the nucleus and in any direction.

However, the electron is much more likely to be found in certain regions of the atom than in others. Given this, a boundary surface can be drawn so that the electron has a high probability to be found anywhere within the surface, and all regions outside the surface have low values. The precise placement of the surface is arbitrary, but any reasonably compact determination must follow a pattern specified by the behavior of <math>\psi^2<math>, the square of the wavefunction. This boundary surface is what is meant when the "shape" of an orbital is mentioned.

Generally speaking, the number <math>n<math> determines the size and energy of the orbital: as <math>n<math> increases, the size of the orbital increases.

Also in general terms, <math>\ell<math> determines an orbital's shape, and <math>m_\ell<math> its orientation. However, since some orbitals are described by equations in complex numbers, the shape sometimes depends on <math>m_\ell<math> also. <math>s<math>-orbitals (<math>\ell=0<math>) are shaped like spheres. <math>p<math>-orbitals have the form of two ellipsoids with a point of tangency at the nucleus. The three <math>p<math>-orbitals in each shell are oriented at right angles to each other, as determined by their respective values of <math>m_\ell<math>.

Four of the five <math>d<math>-orbitals look similar, each with four pear-shaped balls, each ball tangent to two others, and the centers of all four lying in one plane, between a pair of axes. Three of these planes are the <math>xy<math>-, <math>xz<math>-, and <math>yz<math>-planes, and the fourth has the centres on the <math>x<math> and <math>y<math> axes. The fifth and final <math>d<math>-orbital consists of three regions of high probability density: a torus with two pear-shaped regions placed symmetrically on its <math>z<math> axis.

Orbital energy

In atoms with a single electron (essentially hydrogen), the energy of an orbital (and, consequently, of any electrons in the orbital) is determined exclusively by <math>n<math>. The <math>n=1<math> orbital has the lowest possible energy in the atom. Each successively higher value of <math>n<math> has a higher level of energy, but the difference decreases as <math>n<math> increases. At a certain point, the level of energy becomes high enough that the electron escapes from the atom.

In atoms with multiple electrons, the energy of an electron depends not only on the intrinsic properties of its orbital, but also on its interactions with other electrons. These interactions depend on the detail of its probability distribution, and so the energy levels of orbitals depend not only on <math>n<math> but also on <math>\ell<math>. Higher values of <math>\ell<math> are associated with higher values of energy; for instance, the 2p state is higher than the 2s state. When <math>\ell<math> = 3, the increase in energy of the orbital becomes so large as to push the energy of orbital above the energy of the s-orbital in the next higher shell; when <math>\ell<math> = 4 the energy is pushed into the shell two steps higher.

The energy order of the first 24 subshells is given in the following table. Each cell represents a subshell with <math>n<math> and <math>\ell<math> given by its row and column indices, respectively. The number in the cell is the subshell's position in the sequence. Empty cells represent subshells that either do not exist or stand farther down in the sequence.

	<math>s<math>	<math>p<math>	<math>d<math>	<math>f<math>	<math>g<math>
1	1
2	2	3
3	4	5	7
4	6	8	10	13
5	9	11	14	17	21
6	12	15	18	22
7	16	19	23
8	20	24

Electron placement and the periodic table

Several rules govern the placement of electrons in orbitals. The first dictates that no two electrons in an atom may have the same set of values of quantum numbers (this is the Pauli exclusion principle). These quantum numbers include the three that define orbitals (<math>n<math>, <math>\ell<math>, and <math>m_\ell<math>), as well as (the hitherto unmentioned) s. Thus, two electrons may occupy a single orbital, so long as they have different values of <math>s<math>.

Additionally, an electron always tries to occupy the lowest possible energy state. It is possible for it to occupy any orbital so long as it does not violate the Pauli exclusion principle, but if lower-energy orbitals are available, this condition is unstable. The electron will eventually lose energy (by releasing a photon) and drop into the lower orbital. Thus, electrons fill orbitals in the order speficied by the energy sequence given above.

This behavior is responsible for the structure of the periodic table. The table may be divided into several rows (called 'periods'), numbered starting with 1 at the top. The presently known elements occupy seven periods. If a certain period has number <math>i<math>, it consists of elements whose outermost electrons fall in the <math>i<math>th shell.

The periodic table may also be divided into several numbered rectangular 'blocks'. The elements belonging to a given block have this common feature: their highest-energy electrons all belong to the same <math>\ell<math>-state (but the <math>n<math> associated with that <math>\ell<math>-state depends upon the period). For instance, the leftmost two columns constitute the 's-block'. The outermost electrons of Li and Be respectively belong to the 2s subshell, and those of Na and Mg to the 3s subshell.

The number of electrons in a neutral atom increases with the atomic number. The electrons in the outermost shell, or valence electrons, tend to be responsible for an element's chemical behavior. Elements that contain the same number of valence electrons can be grouped together and display similar chemical properties.

Related topics

• Electron configuration
• Molecular orbital
• Atomic electron configuration table

External links

• A visualization of all common and uncommon atomic orbitals, from 1s to 7g (http://www.shef.ac.uk/chemistry/orbitron/)
• David Manthey's Orbital Viewer (http://www.orbitals.com/orb/index.html) renders orbitals with n ≤ 30

fr:Orbitale atomique