Quantum NumbersGeneral rule of Electron construction

The electron configuration of an atomic types (neutral or ionic) allows us to understand the shape and energy of its electrons. Many general rules space taken into factor to consider when assigning the "location" of the electron come its prospective power state, yet these assignments are arbitrary and it is constantly uncertain as to which electron is gift described. Learning the electron construction of a varieties gives us a much better understanding that its bonding ability, magnetism and other bsci-ch.orgical properties.

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The electron configuration is the standard notation provided to explain the electronic structure of an atom. Under the orbital approximation, us let each 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 space the exact same as the ones acquired from resolving the Schrodinger"s equation for Bohr"s hydrogen atom. Hence, many of the rules that we use to explain the electron"s deal with in the hydrogen atom can also be supplied in systems including multiple electrons. As soon as assigning electrons to orbitals, we should 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 solving the Schrödinger equation because that the hydrogen atom, we acquire three quantum numbers, namely the major quantum number (n), the orbital angular momentum quantum number (l), and the magnetic quantum number (ml). Over there is a fourth quantum number, referred to as the turn magnetic quantum number (ms), i m sorry is not acquired from resolving the Schrödinger equation. Together, these 4 quantum numbers deserve to be used to explain the ar of an electron in Bohr"s hydrogen atom. This numbers have the right to be assumed of as an electron"s "address" in the atom.


To assist describe the proper notation because that electron configuration, that is ideal to do so through example. For this example, us will use the iodine atom. There room two methods in i beg your pardon electron configuration deserve to be written:

I: 1s22s22p63s23p64s23d104p65s24d105p5


I: 5s24d105p5

In both the these types of notations, the bespeak of the energy levels must be created by increased energy, mirroring the variety of electrons in each subshell as an exponent. In the brief notation, you place brackets approximately the preceding noble gas aspect followed by the valence shell electron configuration. The regular table reflects that kyrpton (Kr) is the vault noble gas listed before iodine. The noble gas configuration encompases the energy states lower than the valence covering electrons. Therefore, in this situation =1s22s22p63s23p64s23d104p6.

Principal Quantum Number (n)

The primary quantum number n indicates the shell or energy level in which the electron is found. The value of n deserve to be set between 1 to n, wherein n is the value of the outermost shell containing an electron. This quantum number have the right to only be positive, non-zero, and integer values. That is, n=1,2,3,4,..

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

Orbital Angular inert Quantum Number (l)

The orbital angular momentum quantum number, l, shows the subshell of the electron. Friend can also tell the shape of the atom orbital v this quantum number. One s subshell coincides to l=0, a p subshell = 1, a d subshell = 2, a f subshell = 3, and also so forth. This quantum number have the right to only it is in positive and integer values, return it have the right to take top top a zero value. In general, for every value of n, there room n worths of l. Furthermore, the worth of l varieties from 0 come n-1. For example, if n=3, l=0,1,2.

So in regards come the instance 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, to represent the orbitals the a given subshell. For a provided l, ml can range from -l to +l. A p subshell (l=1), for instance, deserve to have three orbitals corresponding to ml = -1, 0, +1. In various other words, it specifies the px, py and pzorbitals the the ns subshell. (However, the ml numbers don"t necessarily exchange mail to a offered orbital. The truth that there room three orbitals merely is indicative the the three orbitals that a p subshell.) In general, because that a offered l, there room 2l+1 feasible values because that ml; and in a n primary shell, there are n2 orbitals discovered in that power level.

Continuing on native out instance from above, the ml worths of Iodine room 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 also 4dyz orbitals.

Spin Magnetic Quantum Number (ms)

The turn magnetic quantum number can only have actually a worth of either +1/2 or -1/2. The value of 1/2 is the turn quantum number, s, which defines the electron"s spin. Due to the spinning of the electron, it generates a magnetic field. In general, one electron v a ms=+1/2 is called ­­an alpha electron, and also one v a ­ms=-1/2 is dubbed a beta electron. No two paired electrons have the right to have the very same spin value.

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

When managing multi-electron systems, us must consider the electron-electron interactions. Hence, the previously explained postulate breaks down in that the energy of the electron is now identified by both the primary quantum number, n, and the orbit angular inert quantum number, l. Return the Schrodinger equation for many-electron atoms is extremely daunting to settle mathematically, we deserve to still explain their electronic structures via electron configurations.

General rules of Electron Configuration

There room a collection of basic rules the are provided to figure out the electron configuration of an atom species: Aufbau Principle, Hund"s Rule and the Pauli-Exclusion Principle. Prior to continuing, it"s important to know that every orbital deserve to be inhabited by two electrons of opposite turn (which will be further debated later). The adhering to table mirrors the possible variety of electrons that deserve to occupy every orbital in a given subshell.

subshell number the orbitals total number of possible electron in every 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 check out on the regular table the its atom number is 53 (meaning it contains 53 electron in the neutral state). Its complete electron configuration is 1s22s22p63s23p64s23d104p65s24d105p5. If you counting up all of these electrons, you will view that the adds up to 53 electrons. An alert that every subshell can only contain the max quantity of electrons as suggested in the table above.

Aufbau Principle

The word "Aufbau" is German because that "building up". The Aufbau Principle, additionally called the building-up principle, claims that electron"s occupy orbitals in order of raising energy. The bespeak of occupation is together follows:



Hund"s dominance states that when electrons accounting degenerate orbitals (i.e. Same n and also l quantum numbers), they must first occupy the north orbitals before dual occupying them. Furthermore, the many stable construction results once the spins room parallel (i.e. All alpha electron or every beta electrons). Nitrogen, because that example, has 3 electrons occupying the 2p orbital. According to Hund"s Rule, they must first occupy each 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 rectal the half-filled 2px orbital. This, therefore, is a violation the Hund"s dominance (Figure 2).

Figure 2. A visual depiction of the Aufbau Principle and Hund"s Rule. Note that the filling of electron in every orbital(px, py and also pz) is arbitrary as long as the electrons are singly filled before having two electrons accounting the exact same orbital.(a)This diagram to represent the correct filling that electrons because that the nitrogen atom. (b) This diagramrepresents the incorrectfilling that the electrons for the nitrogen atom.

Electronic construction of Cations and also Anions

The method we designate digital configurations because that cations and anions is essentially comparable to that for neutral atom in their ground state. The is, we follow the three vital rules: Aufbau Principle, Pauli-exclusion Principle, and Hund"s Rule. The electronic configuration of cations is assigned by removing electrons first in the outermost p orbital, adhered to by the s orbital and finally the d orbitals (if any much more electrons need to be removed). Because that instance, the soil state electronic configuration that calcium (Z=20) is 1s22s22p63s23p64s2. The calcium ion (Ca2+), however, has actually two electrons less. Hence, the electron construction for Ca2+ is 1s22s22p63s23p6. Due to the fact that we have to take away 2 electrons, we very first remove electrons from the outermost shell (n=4). In this case, every the 4p subshells space empty; hence, we begin by removing indigenous the s orbital, which is the 4s orbital. The electron construction for Ca2+ is the same as the for Argon, which has 18 electrons. Hence, we can say that both are isoelectronic.

The electronic configuration the anions is assigned by including electrons follow to Aufbau Principle. We include electrons to fill the outermost orbital the is occupied, and then add more electrons to the next greater orbital. The neutral atom chlorine (Z=17), because that instance has 17 electrons. Therefore, its floor state digital configuration have the right to be written as 1s22s22p63s23p5. The chloride ion (Cl-), on the various other hand, has secondary electron for a total of 18 electrons. Adhering to Aufbau Principle, the electron rectal the partly filled 3p subshell first, make the 3p orbital completely filled. The electronic configuration because that Cl- can, therefore, it is in designated as 1s22s22p63s23p6. Again, the electron configuration for the chloride ion is the exact same as the for Ca2+ and Argon. Hence, they room all isoelectronic to every other.


1. Which of the princples explained above tells us that electrons that space paired cannot have actually the very same spin value?

2. Uncover the worths of n, l, ml, and also ms for the following:

a. Mg

b. Ga

c. Co

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

4. I beg your pardon of the following cannot exist (there may be much more than one answer):

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

b. N = 3; l = 2; ml = 1; ms = 1

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

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

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

5. Create 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 can be either +1/2 or -1/2

b. N = 4; l = 0, 1, 2, 3; ml = -3, -2, -1, 0, 1, 2, 3; ms can be one of two people +1/2 or -1/2

c. N = 3; l = 0, 1, 2; ml = -2, -1, 0, 1, 2, 3; ms can be one of two people +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 can not be 4, since l varieties from (0 - n-1)

b. Ms can only it is in +1/2 or -1/2

c. Okay

d. Okay

e. The worth of n cannot be zero.

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5. A. 1s22s22p63s23p3

b. 1s22s22p63s23p6

c. 1s22s22p63s23p64s23d7


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