Quantum NumbersGeneral rule of Electron configuration

The electron configuration of one atomic species (neutral or ionic) allows us to understand the shape and also energy of its electrons. Plenty of general rules are taken into consideration when assigning the "location" the the electron to its prospective energy state, but these assignments room arbitrary and it is constantly uncertain as to which electron is gift described. Knowing the electron construction of a varieties gives us a far better understanding of its bonding ability, magnetism and also other yellowcomic.comical properties.

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The electron configuration is the conventional notation provided to define the digital structure of one atom. Under the orbital approximation, us let every electron occupy an orbital, which deserve to be resolved by a solitary wavefunction. In doing so, we acquire three quantum numbers (n,l,ml), which room the same as the ones obtained from resolving the Schrodinger"s equation because that Bohr"s hydrogen atom. Hence, plenty of of the rules that we usage to describe the electron"s address in the hydrogen atom can additionally be supplied in systems entailing multiple electrons. As soon as assigning electron to orbitals, we must follow a set of 3 rules: the Aufbau Principle, the Pauli-Exclusion Principle, and Hund"s Rule.

The wavefunction is the solution to the Schrödinger equation. By fixing the Schrödinger equation because that the hydrogen atom, we achieve three quantum numbers, namely the major quantum number (n), the orbital angular inert quantum number (l), and the magnetic quantum number (ml). Over there is a fourth quantum number, called the rotate magnetic quantum number (ms), which is not derived from solving the Schrödinger equation. Together, these 4 quantum numbers can be supplied to describe the location of one electron in Bohr"s hydrogen atom. This numbers have the right to be assumed of as an electron"s "address" in the atom.


To aid describe the appropriate notation because that electron configuration, the is ideal to perform so with example. Because that this example, we will use the iodine atom. There are two means in i m sorry electron configuration can be written:

I: 1s22s22p63s23p64s23d104p65s24d105p5


I: 5s24d105p5

In both that these species of notations, the order of the power levels have to be created by enhanced energy, reflecting the variety of electrons in every subshell together an exponent. In the brief notation, you location brackets approximately the preceding noble gas facet followed by the valence shell electron configuration. The periodic table reflects that kyrpton (Kr) is the previous noble gas provided before iodine. The noble gas construction encompases the power states lower than the valence shell electrons. Therefore, in this instance =1s22s22p63s23p64s23d104p6.

Principal Quantum Number (n)

The primary quantum number n indicates the covering or power level in i beg your pardon the electron is found. The value of n have the right to be set between 1 to n, whereby n is the worth of the outermost covering containing an electron. This quantum number deserve to only be positive, non-zero, and also integer values. The is, n=1,2,3,4,..

For example, an Iodine atom has its outmost electrons in the 5p orbital. Therefore, the principle quantum number for Iodine is 5.

Orbital Angular inert Quantum Number (l)

The orbit angular inert quantum number, l, shows the subshell the the electron. You can additionally tell the form of the atom orbital v this quantum number. One s subshell synchronizes to l=0, a p subshell = 1, a d subshell = 2, a f subshell = 3, and so forth. This quantum number can only be positive and also integer values, although it can take ~ above a zero value. In general, for every worth of n, there room n values of l. Furthermore, the worth of l arrays from 0 come n-1. For example, if n=3, l=0,1,2.

So in regards come the example used above, the l values that Iodine because that n = 5 are l = 0, 1, 2, 3, 4.

Magnetic Quantum Number (ml)

The magnetic quantum number, ml, represents the orbitals the a given subshell. Because that a provided l, ml can range from -l to +l. A ns subshell (l=1), because that instance, can have 3 orbitals matching to ml = -1, 0, +1. In other words, it defines the px, py and also pzorbitals of the ns subshell. (However, the ml numbers don"t necessarily correspond to a given orbital. The truth that there room three orbitals merely is indicative the the 3 orbitals of a p subshell.) In general, because that a provided l, there space 2l+1 possible values for ml; and in a n principal shell, there are n2 orbitals discovered in that energy level.

Continuing on from out example from above, the ml worths of Iodine space ml = -4, -3, -2, -1, 0 1, 2, 3, 4. These arbitrarily correspond to the 5s, 5px, 5py, 5pz, 4dx2-y2, 4dz2, 4dxy, 4dxz, and 4dyz orbitals.

Spin Magnetic Quantum Number (ms)

The turn magnetic quantum number can only have actually a value of one of two people +1/2 or -1/2. The worth of 1/2 is the turn quantum number, s, which explains the electron"s spin. As result of the turn of the electron, it generates a magnetic field. In general, an electron through a ms=+1/2 is called ­­an alpha electron, and also one with a ­ms=-1/2 is dubbed a beta electron. No two paired electrons can have the exact same spin value.

Out that these four quantum numbers, however, Bohr postulated that only the major quantum number, n, determines the energy of the electron. Therefore, the 3s orbit (l=0) has actually the same power as the 3p (l=1) and also 3d (l=2) orbitals, nevertheless of a difference in l values. This postulate, however, stop true only for Bohr"s hydrogen atom or other hydrogen-like atoms.

When handling multi-electron systems, us must take into consideration the electron-electron interactions. Hence, the previously explained postulate breaks under in that the power of the electron is now established by both the major quantum number, n, and also the orbital angular momentum quantum number, l. Back the Schrodinger equation for many-electron atoms is extremely an overwhelming to settle mathematically, we can still describe their electronic structures via electron configurations.

General rule of Electron Configuration

There space a collection of basic rules the are supplied to number out the electron construction of an atomic species: Aufbau Principle, Hund"s Rule and also the Pauli-Exclusion Principle. Prior to continuing, it"s essential to understand that each orbital can be populated by two electron of opposite spin (which will be further questioned later). The complying with table mirrors the possible variety of electrons that can occupy every orbital in a provided subshell.

subshell number of orbitals total variety of possible electrons in each orbital
s 1 2
p 3 (px, py, pz) 6
d 5 (dx2-y2, dz2, dxy, dxz, dyz) 10
f 7 (fz3, fxz2, fxyz, fx(x2-3y2), fyz2, fz(x2-y2), fy(3x2-y2)


Using our example, iodine, again, we view on the regular table that its atom number is 53 (meaning it has 53 electron in that neutral state). Its complete electron configuration is 1s22s22p63s23p64s23d104p65s24d105p5. If you counting up all of these electrons, you will check out that that adds up to 53 electrons. An alert that each subshell deserve to only save the max amount of electrons as shown in the table above.

Aufbau Principle

The indigenous "Aufbau" is German for "building up". The Aufbau Principle, also called the building-up principle, states that electron"s accounting orbitals in stimulate of enhancing energy. The stimulate of occupation is as follows:



Hund"s ascendancy states that as soon as electrons occupy degenerate orbitals (i.e. Very same n and l quantum numbers), they must very first occupy the north orbitals before dual occupying them. Furthermore, the many stable construction results once the spins are parallel (i.e. All alpha electrons or every beta electrons). Nitrogen, for example, has 3 electron occupying the 2p orbital. According to Hund"s Rule, lock must an initial occupy every of the three degenerate ns orbitals, specific the 2px orbital, 2py orbital, and the 2pz orbital, and with parallel spins (Figure 2). The configuration listed below is incorrect since the 3rd electron rectal does no occupy the north 2pz orbital. Instead, it occupies the half-filled 2px orbital. This, therefore, is a violation that Hund"s ascendancy (Figure 2).

Figure 2. A visual depiction of the Aufbau Principle and Hund"s Rule. Keep in mind that the filling of electrons in each orbital(px, py and also pz) is arbitrarily as long as the electrons space singly filled prior to having 2 electrons occupy the exact same orbital.(a)This diagram to represent the correct filling the electrons for the nitrogen atom. (b) This diagramrepresents the incorrectfilling that the electrons because that the nitrogen atom.

Electronic construction of Cations and Anions

The means we designate digital configurations for cations and also anions is essentially comparable to the for neutral atoms in their ground state. That is, us follow the three vital rules: Aufbau Principle, Pauli-exclusion Principle, and Hund"s Rule. The digital configuration the cations is assigned by remove electrons an initial in the outermost ns orbital, followed by the s orbital and finally the d orbitals (if any more electrons must be removed). For instance, the ground state digital configuration the calcium (Z=20) is 1s22s22p63s23p64s2. The calcium ion (Ca2+), however, has two electrons less. Hence, the electron construction for Ca2+ is 1s22s22p63s23p6. Because we have to take away 2 electrons, we an initial remove electrons from the outermost covering (n=4). In this case, every the 4p subshells room empty; hence, we start by removing from the s orbital, i beg your pardon is the 4s orbital. The electron configuration for Ca2+ is the very same as that for Argon, which has actually 18 electrons. Hence, we can say the both space isoelectronic.

The electronic configuration of anions is assigned by including electrons according to Aufbau Principle. We add electrons to to fill the outermost orbital that is occupied, and then add an ext electrons to the next greater orbital. The neutral atom chlorine (Z=17), because that instance has actually 17 electrons. Therefore, its floor state electronic configuration deserve to be written as 1s22s22p63s23p5. The chloride ion (Cl-), on the other hand, has secondary electron for a total of 18 electrons. Complying with Aufbau Principle, the electron occupies the partially filled 3p subshell first, do the 3p orbital fully filled. The electronic configuration because that Cl- can, therefore, be designated as 1s22s22p63s23p6. Again, the electron configuration for the chloride ion is the very same as that for Ca2+ and Argon. Hence, they room all isoelectronic to each other.


1. I m sorry of the princples explained over tells united state that electron that room paired cannot have the same spin value?

2. Find the worths of n, l, ml, and also ms because that the following:

a. Mg

b. Ga

c. Co

3. What is a possible mix for the quantum number of the 5d orbital? Give an example of an facet which has actually the 5d orbital as it"s most outer orbital.

4. I m sorry of the complying with cannot exist (there may be more than one answer):

a. N = 4; l = 4; ml = -2; multiple sclerosis = +1/2

b. N = 3; l = 2; ml = 1; multiple sclerosis = 1

c. N = 4; l = 3; ml = 0; multiple sclerosis = +1/2

d. N = 1; l = 0; ml = 0; multiple sclerosis = +1/2

e. N = 0; l = 0; ml = 0; multiple sclerosis = +1/2

5. Compose electron configurations for the following:

a. P

b. S2-

c. Zn3+


1. Pauli-exclusion Principle

2. A. N = 3; l = 0, 1, 2; ml = -2, -1, 0, 1, 2; ms have the right to be one of two people +1/2 or -1/2

b. N = 4; l = 0, 1, 2, 3; ml = -3, -2, -1, 0, 1, 2, 3; ms have the right to be either +1/2 or -1/2

c. N = 3; l = 0, 1, 2; ml = -2, -1, 0, 1, 2, 3; ms have the right to be either +1/2 or -1/2

3. N = 5; l = 3; ml = 0; ms = +1/2. Osmium (Os) is an example.

4. A. The worth of l cannot be 4, because l ranges from (0 - n-1)

b. Ms have the right to only it is in +1/2 or -1/2

c. Okay

d. Okay

e. The value of n can not be zero.

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5. A. 1s22s22p63s23p3

b. 1s22s22p63s23p6

c. 1s22s22p63s23p64s23d7


Atkins, P. W., & De Paula, J. (2006). Physical yellowcomic.comistry for the Life Sciences. new York, NY: W. H. Freeman and Company. Petrucci, R. H., Harwood, W. S., & Herring, F. G. (2002). General yellowcomic.comistry: principles and modern-day Applications. top Saddle River, NJ: Prentice-Hall, Inc. Shagoury, Richard. yellowcomic.comistry 1A great Book. 4th Ed. Tradition Publishing. 2006. Print