Number Of Electrons



  1. Number Of Electrons In Chlorine
  2. Number Of Electrons In Nitrogen
  3. Number Of Electrons In Potassium
  4. Number Of Electrons In Aluminum

Quantum Numbers,
Atomic Orbitals, and
Electron Configurations

Contents:
Quantum Numbers and Atomic Orbitals
1. Principal Quantum Number (n)
2.Angular Momentum (Secondary, Azimunthal) Quantum Number (l)
3.Magnetic Quantum Number (ml)
4.Spin Quantum Number (ms)
Table of Allowed Quantum Numbers
Writing Electron Configurations
Properties of Monatomic Ions
References

Quantum Numbers and Atomic Orbitals

Before we can use these orbitals we need to know the number of electrons that can occupy an orbital and how they can be distinguished from one another. Experimental evidence suggests that an orbital can hold no more than two electrons. To distinguish between the two electrons in an orbital, we need a fourth quantum number. This number divides the subshell into individual orbitals which hold the electrons; there are 2l+1 orbitals in each subshell. Thus the s subshell has only one orbital, the p subshell has three orbitals, and so on. Spin Quantum Number (m s): m s = +½ or -½. Specifies the orientation of the spin axis of an electron. Here's what I would suggest. If a complete set of quantum numbers is given to you, then you can say for a fact that you're only dealing with one electron. As you know, quantum numbers are used to describe the location and spin of an electron that surrounds an atom's nucleus. More specifically, the four quantum numbers will tell you the energy level on which the electron resides - given by n. Generally, the number of electrons—alongside with protons and neutrons—in an atom can be determined from a set of simple rules. The number of protons in the nucleus of the atom is equal to the atomic number (Z). To sciencetutorial4u. I am showing how to find proton, neutron and electron numbers of atom and ions in this video.

By solving the Schrödinger equation (Hy = Ey), we obtain a set of mathematical equations, called wave functions (y), which describe the probability of finding electrons at certain energy levels within an atom.

A wave function for an electron in an atom is called an atomic orbital; this atomic orbital describes a region of space in which there is a high probability of finding the electron. Energy changes within an atom are the result of an electron changing from a wave pattern with one energy to a wave pattern with a different energy (usually accompanied by the absorption or emission of a photon of light).

Each electron in an atom is described by four different quantum numbers. The first three (n, l, ml) specify the particular orbital of interest, and the fourth (ms) specifies how many electrons can occupy that orbital.

  1. Principal Quantum Number (n): n = 1, 2, 3, …, ∞
    Specifies the energy of an electron and the size of the orbital (the distance from the nucleus of the peak in a radial probability distribution plot). All orbitals that have the same value of n are said to be in the same shell (level). For a hydrogen atom with n=1, the electron is in its ground state; if the electron is in the n=2 orbital, it is in an excited state. The total number of orbitals for a given n value is n2.
  1. Angular Momentum (Secondary, Azimunthal) Quantum Number (l): l = 0, ..., n-1.
    Specifies the shape of an orbital with a particular principal quantum number. The secondary quantum number divides the shells into smaller groups of orbitals called subshells (sublevels). Usually, a letter code is used to identify l to avoid confusion with n:
l012345...
Letterspdfgh...

The subshell with n=2 and l=1 is the 2p subshell; if n=3 and l=0, it is the 3s subshell, and so on. The value of l also has a slight effect on the energy of the subshell; the energy of the subshell increases with l (s < p < d < f).

  1. Magnetic Quantum Number (ml): ml = -l, ..., 0, ..., +l.
    Specifies the orientation in space of an orbital of a given energy (n) and shape (l). This number divides the subshell into individual orbitals which hold the electrons; there are 2l+1 orbitals in each subshell. Thus the s subshell has only one orbital, the p subshell has three orbitals, and so on.
Oxygen
  1. Spin Quantum Number (ms): ms = +½ or -½.
    Specifies the orientation of the spin axis of an electron. An electron can spin in only one of two directions (sometimes called up and down).
    The Pauli exclusion principle (Wolfgang Pauli, Nobel Prize 1945) states that no two electrons in the same atom can have identical values for all four of their quantum numbers. What this means is that no more than two electrons can occupy the same orbital, and that two electrons in the same orbital must have opposite spins.
    Because an electron spins, it creates a magnetic field, which can be oriented in one of two directions. For two electrons in the same orbital, the spins must be opposite to each other; the spins are said to be paired. These substances are not attracted to magnets and are said to be diamagnetic. Atoms with more electrons that spin in one direction than another contain unpaired electrons. These substances are weakly attracted to magnets and are said to be paramagnetic.

Table of Allowed Quantum Numbers

nlmlNumber of
orbitals		Orbital
Name		Number of
electrons
10011s2
20012s2
1-1, 0, +132p6
30013s2
1-1, 0, +133p6
2-2, -1, 0, +1, +253d10
40014s2
1-1, 0, +134p6
2-2, -1, 0, +1, +254d10
3-3, -2, -1, 0, +1, +2, +374f14

Writing Electron Configurations

The distribution of electrons among the orbitals of an atom is called the electron configuration. The electrons are filled in according to a scheme known as the Aufbau principle ('building-up'), which corresponds (for the most part) to increasing energy of the subshells:

Electrons

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f

It is not necessary to memorize this listing, because the order in which the electrons are filled in can be read from the periodic table in the following fashion:

Or, to summarize:

In electron configurations, write in the orbitals that are occupied by electrons, followed by a superscript to indicate how many electrons are in the set of orbitals (e.g., H 1s1)

Another way to indicate the placement of electrons is an orbital diagram, in which each orbital is represented by a square (or circle), and the electrons as arrows pointing up or down (indicating the electron spin). When electrons are placed in a set of orbitals of equal energy, they are spread out as much as possible to give as few paired electrons as possible (Hund's rule).

examples will be added at a later date

In a ground state configuration, all of the electrons are in as low an energy level as it is possible for them to be. When an electron absorbs energy, it occupies a higher energy orbital, and is said to be in an excited state.

Properties of Monatomic Ions

The electrons in the outermost shell (the ones with the highest value of n) are the most energetic, and are the ones which are exposed to other atoms. This shell is known as the valence shell. The inner, core electrons (inner shell) do not usually play a role in chemical bonding.

Elements with similar properties generally have similar outer shell configurations. For instance, we already know that the alkali metals (Group I) always form ions with a +1 charge; the 'extra' s1 electron is the one that's lost:

IALi1s22s1Li+1s2
Na1s22s22p63s1Na+1s22s22p6
K1s22s22p63s23p64s1K+1s22s22p63s23p6

The next shell down is now the outermost shell, which is now full — meaning there is very little tendency to gain or lose more electrons. The ion's electron configuration is the same as the nearest noble gas — the ion is said to be isoelectronic with the nearest noble gas. Atoms 'prefer' to have a filled outermost shell because this is more electronically stable.

  • The Group IIA and IIIA metals also tend to lose all of their valence electrons to form cations.
IIABe1s22s2Be2+1s2
Mg1s22s22p63s2Mg2+1s22s22p6
IIIAAl1s22s22p63s23p1Al3+1s22s22p6
  • The Group IV and V metals can lose either the electrons from the p subshell, or from both the s and p subshells, thus attaining a pseudo-noble gas configuration.
IVASn[Kr]4d105s25p2Sn2+[Kr]4d105s2
Sn4+[Kr]4d10
Pb[Xe]4f145d106s26p2Pb2+[Xe]4f145d106s2
Pb4+[Xe]4f145d10
VABi[Xe]4f145d106s26p3Bi3+[Xe]4f145d106s2
Bi5+[Xe]4f145d10
  • The Group IV - VII non-metals gain electrons until their valence shells are full (8 electrons).
IVAC1s22s22p2C4-1s22s22p6
VAN1s22s22p3N3-1s22s22p6
VIAO1s22s22p4O2-1s22s22p6
VIIAF1s22s22p5F-1s22s22p6
Number Of Electrons
  • The Group VIII noble gases already possess a full outer shell, so they have no tendency to form ions.
  • Transition metals (B-group) usually form +2 charges from losing the valence s electrons, but can also lose electrons from the highest d level to form other charges.
B-groupFe1s22s22p63s23p63d64s2Fe2+1s22s22p63s23p63d6
Fe3+1s22s22p63s23p63d5

References

Martin S. Silberberg, Chemistry: The Molecular Nature of Matter and Change, 2nd ed. Boston: McGraw-Hill, 2000, p. 277-284, 293-307.

How would you find out the number of electrons from just the 4 quantum numbers?

1 Answer

Explanation:

If a complete set of quantum numbers is given to you, then you can say for a fact that you're only dealing with one electron.

As you know, quantum numbers are used to describe the location and spin of an electron that surrounds an atom's nucleus.

More specifically, the four quantum numbers will tell you

Number Of Electrons In Chlorine

  • the energy level on which the electron resides - given by #n#, the principal quantum number
  • the subshell in which the electron resides - given by #l#, the angular momentum quantum number
  • the exact orbital in which you can find the electron - given by, #m_l#, the magnetic quantum number
  • the spin of the electron - given by #m_s#, the spin quantum number

Now, each complete set of quantum numbers can only describe one electron. This means that if you're given something like this

#n=2, l=1, m_l = 0, m_s = +1/2#

then you know for a fact that only one electron can have that set of quantum numbers.

Now, let's assume that you are not given all four quantum numbers. These are some rules to help you figure out how many electrons can share an incomplete set of quantum numbers

  • You are given the value of#n#

If this is the case, then use the fact that the number of electrons you get per energy level is equal to

#color(blue)('no. of electrons' = 2n^2)#

Number Of Electrons In Nitrogen

  • You are given#n#and#l#

This time, you know the energy level and the subshell in which the electrons can be found. To determine how many electrons can share these quantum numbers, use #m_l#, which can take values from #-l# to #l#.

Number Of Electrons In Potassium

Once you know how many orbitals you have per subshell, multiply that value by #2# to get the number of electrons

#color(blue)('no. of electrons' = 2 xx 'no. of subshells')#

  • You are ginve#n#, #l#, and#m_l#

Now you know the energy level, subshell, and specific orbital in which the electrons can reside.

Number Of Electrons In Aluminum

Since the only possible values for #m_s# are #+-1/2#, it follows that every orbital can hold a maximum of two electrons.

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