Atoms in the periodic table follows periodicity and ground state electron configuration. In this article, we are going to investigate ground state electron configuration in detail.
Ground State Electron Configuration: Historical review
Madame Marie Curie, a Polish-born French chemist, and physicist with her husband, Pierre, announced the discovery of new atom radium in the year of 1898. They had separated these from a very radioactive mixture from pitchblende, an ore of uranium element. This radioactive mixture was primarily a compound of barium element.
When this radioactive mixture was heated in a flame, however, it gave a new atomic line in the atomic emission spectrum, in addition to the spectrum for the barium element. This new line in the atomic spectrum is based on their discovery of a new element, which they called radium element, on this experimental finding.
It took them a few more years to isolate a pure compound of radium element from the radioactive mixture. Finally, in the year of 1910, Marie Curie isolated the pure metallic element in the laboratory. Radium element, like uranium, is a very radioactive element.
But in most of radium’s chemical and physical properties, it is similar to the nonradioactive element barium in the same column. It was this Physico-chemical similarity that made the final separation of the new element so difficult to get in pure form.
Chemists had a long history known to show that groups of elements have similar properties in the table. In 1869 Dmitri Mendeleev found that the elements could be arranged in a particular way by columns in the table form, with elements in the same column displaying similar chemical and physical properties.
Thus, Mendeleev placed beryllium, calcium, strontium, and barium in one column in the table. Now, with the Curies’ discovery, radium was added to this column in the last. Mendeleev’s arrangement of the atomic elements, the periodic table, was originally based on the practically observed chemical and physical properties of the atomic elements and their compounds.
We now explain this arrangement in terms of the electronic structure of atoms in the periodic table. In this article, we will look at this electronic configuration structure and its relationship to the periodic table of elements.
What is Ground state Electron Configuration of Atoms
We found that an electron in an atom in the periodic table has four quantum numbers— n, l, m(l) , and m(s)—associated with it. “n” is the principal quantum number, “l” is the angular quantum number, “m(l)” is the magnetic quantum number, and “m(s)” is the spin quantum number.
The first three quantum numbers characterize the orbital that describes the region of space where an electron is most likely to be found; we say that the electron “occupies” this orbital in the electronic configuration.
The spin quantum number, m(s), describes the spin orientation of an electron in the orbital of the complete ground state electron configuration of an atom. In the first section of this article, we will look further at electron spin; then we will discuss how electrons are distributed among the possible orbitals of an atom.
Electron Spin and the Pauli Exclusion Principle
Otto Stern and Walther Gerlach first observed and experimented with electron spin magnetism in the year of 1921. They directed a beam of less dense silver atoms into the magnetic field of a specially designed magnet. The same experiment with less dense silver atoms can be done with hydrogen atoms.
The beam of less dense hydrogen atoms is split into two by the magnetic field; half of the atoms are bent in one direction and half in the other direction. The reason that the hydrogen atoms are affected by the laboratory magnet shows that they themselves act as tiny small magnets.
The beam of less dense hydrogen atoms is split into two because the electron in each hydrogen atom behaves as a tiny small magnet with only two possible orientations. In effect, the electron in the hydrogen atom acts as though it were a ball of spinning charge, and, like a circulating electric charge in the hydrogen atom, it would create a magnetic field around the tiny small hydrogen atom.
Electron spin in the hydrogen atom, however, is subject to a quantum restriction on the possible directions of the spin axis. The resulting directions of spin magnetism correspond to spin quantum numbers m(s)=+½ and m(s)=-½ .
Electron Configurations and Orbital Diagrams
We describe the electrons in an atom by the atom’s complete ground state electron configuration. An electron configuration structure of an atom is a particular distribution of electrons among the available orbital subshells. A subshell consists of a group of orbitals having the same n and l quantum numbers but different m(l) values.
Remember that we denote each subshell of orbitals by its principal quantum number, n, followed by a letter standing for its l quantum number (s, p, d, f, etc.). The notation for the complete ground state electron configuration of atom lists the subshell symbols one after the other, with a superscript giving the number of electrons in that subshell of orbitals.
For example,ground state electron configuration of the lithium atom (atomic number 3) with two electrons in the 1s subshell orbital and one electron in the 2s subshell orbital is written 1s2 2s1 . Whereas the ground state electron configuration of an atom gives the number of electrons in each subshell orbital, we use a diagram to show how the orbitals of a subshell orbital are occupied by electrons in it.
It is called an orbital diagram of an atom in the given complete ground state electron configuration. An orbital subshell is represented by a circle. Each group of orbitals in a subshell is labeled by its subshell orbital group notation.
An electron in an orbital subshell is shown by an arrow; the arrow points up when m(s)=+1/2 and down when m(s)=-1/2. The orbital diagram shows the electronic configuration structure of an atom in which there are two electrons in the 1s subshell orbital, or orbital (one electron with m(s)=+1/2, the other with m(s)=-1/2); two electrons in the 2s subshell (m(s)=+1/2, m(s)=-1/2); and one electron in the 2p subshell (m(s)=+1/2). The complete ground state electron configuration of the atom is 1s2 2s2 2p1 .
Pauli Exclusion Principle
Not all of the possible arrangements of electrons among the subshell orbitals of an atom are physically possible. The Pauli exclusion principle, which summarizes hydrogen split experimental observations, states that no two electrons in an atom can have the same four quantum numbers.
If one electron in the hydrogen atom has the quantum numbers n = 1, l = 0, m(l) = 0, and m(s)=+1/2, no other electron can have these same quantum numbers. In other words, you cannot place two electrons with the same value of spin quantum number m(s) in a 1s orbital.
The orbital diagram is not a possible arrangement of electrons in the complete ground state electron configuration of atom structure. Because there are only two possible values of spin quantum number m(s), an orbital subshell can hold no more than two electrons—and then only if the two electrons have different spin quantum numbers.
In an orbital diagram, an orbital subshell with two electrons must be written with arrows pointing in opposite directions. The two electrons are said to have opposite spins. We can restate the definition of Pauli exclusion principle as follows:
Pauli exclusion principle: An orbital subshell can hold at most two electrons, and then only if the electrons have opposite spin directions.
Each subshell orbital holds a maximum of twice as many electrons as the number of orbitals in the subshell. Thus, a 2p subshell orbital, which has three subshell orbitals (with m(l) = –1, 0, and +1), can hold a maximum of six electrons. The maximum number of electrons in various orbital subshells is given in the following table.
|Subshell Orbital||Number of subshell Orbitals||Maximum Number of Electrons in subshell orbital|
|s (l = 0)||1||2|
|p (l = 1)||3||6|
|d (l = 2)||5||10|
|f (l = 3)||7||14|
Building-Up Principle and the Periodic Table
Each and every atom in the periodic table has an infinite number of possible complete ground state electron configurations. The electron configuration of an atom associated with the lowest energy level of the atom corresponds to a quantum-mechanical state (zero vibrational level) called the ground state of an atom.
Other ground state electron configurations correspond to excited states of the same atom, associated with energy levels other than the lowest energy levels. Ground state is one in the atom. But there are many more excited states in the same atom.
For example, the ground state of the sodium (Na) atom is known from spectroscopy experiments to have the ground state electron configuration 1s2 2s2 2p6 3s1 . The electron configuration 1s2 2s2 2p6 3p1 represents an excited state of the sodium (Na) atom. The physical and chemical properties of an atom are related primarily to the electron configuration of its ground state.
Building-Up Principle (Aufbau Principle)
Most of the ground state electron configurations in the atom can be explained in terms of the building-up principle (or Aufbau principle), a scheme used to reproduce the electron configurations of the ground states of atoms by successively filling orbital subshells with electrons in a specific order (the building-up order in the ground state electron configuration).
Following this Aufbau principle, you can obtain the ground state electron configuration of an atom by successively filling orbital subshells in the following order: 1s,–> 2s,–> 2p,–> 3s,–> 3p,–> 4s, –>3d, –>4p, –>5s, –>4d, –>5p, –>6s, –>4f, –>5d, –>6p, –>7s, –>5f.
As you will see, you can very easily obtain it from the periodic table of ground state electron configuration. The building-up order in Aufbau principle corresponds for the most part to increasing energy of the orbital subshells. You might not wonder about this. By filling in an orbital subshell of lowest energy first, you usually get the lowest total energy (ground state) of the atom in the periodic table.
Remember that the energy of an orbital depends only on the principal quantum numbers n and angular quantum number l (The energy of the Hydrogen atom, however, depends only on n. Because l is zero). Orbitals subshell with the same principal quantum number n and angular quantum number l but different m(l) —that is, different orbitals of the same subshell—have the same energy level.
The energy depends primarily on principal quantum number n, increasing with its value. For example, a 3s orbital subshell has greater energy than a 2s orbital subshell. Except for the Hydrogen atom, the energies of orbitals with the same n increase with the l quantum number.
A 3p orbital subshell has slightly greater energy than a 3s orbital subshell. The subshell orbital of lowest energy is 1s; next higher are 2s and 2p, then 3s and 3p go on. The order of these subshells orbital by energy, from 1s to 3p, follows the building-up order(Aufbau principle) as listed earlier.
When orbital subshells have nearly the same energy level, however, the building-up order (Aufbau’s principle) is not strictly determined by the order of their energies. The ground state electron configuration of atoms, which we are trying to predict by the building-up order, are determined by the total electronic energies of the atoms.
The total electronic energy of an atom depends not only on the energies of the orbital subshells but also on the energies of interaction among the different orbital subshells. It so happens that for all elements with Z = 21 or greater, the energy of the 3d subshell orbital is lower than the energy of the 4s subshell orbital, which is opposite to the building-up order.
You need the building-up order (Aafbau principle) to predict the electron configurations of the ground states of atoms. Now you can see how to reproduce the ground state electron configurations using the building-up principle. Remember that the number of electrons in a neutral atom equals the atomic number Z. (The nuclear charge is +Z.)
In the case of the simplest atom, hydrogen atom (Z = 1), you obtain the ground state by placing the single electron into the 1s orbital subshell, giving the configuration 1s1 (this is read as “oneess-one”).
Now you go to the helium atom (Z = 2). The first electron goes into the 1s orbital subshell, as in hydrogen, followed by the second electron, because any orbital subshell can hold two electrons. The configuration is 1s2 . Filling the n = 1 shell creates a very stable configuration, and as a result, helium is chemically unreactive. It is the first noble gas configuration.
How to write ground state electron configuration
You continue this way through the other atomic elements in the periodic table, each time increasing Z by 1 and adding another one electron. You can obtain the ground state electron configuration of an atom from that of the preceding atomic element by adding an electron into the next available orbital, following the building-up order using the Aufbau principle.
In a lithium atom (Z = 3), the first two electrons give the configuration 1s2 , like a helium atom, but the third electron goes into the next higher orbital subshell in the building-up order, because the 1s orbital is now filled. This gives the ground state electron configuration 1s2 2s1 .
In the beryllium atom (Z = 4), the fourth electron fills the 2s orbital, giving the total ground state electron configuration 1s2 2s2 . Using the abbreviation [He] for 1s2 , the configurations are
Z = 3 lithium atom 1s2 2s1 or [He]2s1
Z = 4 beryllium atom 1s2 2s2 or [He]2s2
With boron atom (Z = 5), the electrons begin filling the 2p subshell. You get
Z = 5 boron atom 1s2 2s2 2p1 or [He]2s2 2p1
Z = 6 carbon atom 1s2 2s2 2p2 or [He]2s2 2p2
Z = 7 nitrogen atom 1s2 2s2 2p3 or [He]2s2 2p3
Z = 8 oxygen atom 1s2 2s2 2p4 or [He]2s2 2p4
Z = 8 Fluorine atom 1s2 2s2 2p5 or [He]2s2 2p5
Z = 10 neon atom1s2 2s2 2p6 or [He]2s2 2p6
Having filled the 2p subshell, you again find a particularly stable noble gas electron configuration. Neon is chemically unreactive as a result.
With sodium atom (Z = 11), the 3s orbital begins to fill. Using the abbreviation of neon noble gas electron configuration [Ne] for 1s2 2s2 2p6 , you have the following ground state electron configuration
Z = 11 sodium 1s2 2s2 2p6 3s1 or [Ne]3s1
Z = 12 magnesium 1s2 2s2 2p6 3s2 or [Ne]3s2
Then the 3p subshell orbital begins to fill the electron in their electron configuration.
Z = 13 aluminum 1s2 2s2 2p6 3s2 3p1 or [Ne]3s2 3p1
Z = 18 argon 1s2 2s2 2p6 3s2 3p6 or [Ne]3s2 3p6
With the 3p subshell orbital filled fully, a stable noble gas electron configuration has been attained; argon is an unreactive element.
Now the 4s orbital subshell begins to fill its electron configuration. You get following ground state electron configurations
[Ar]4s1 for potassium (Z = 19) and [Ar]4s2 for calcium (Z = 20) ([Ar] = 1s2 2s2 2p6 3s2 3p6 ).
At this point the 3d subshell begins to fill its electron configuration. You get following ground state electron configurations[Ar]3d1 4s2 for scandium (Z = 21),
[Ar]3d 2 4s2 for titanium (Z = 22),
and [Ar]3d 3 4s2 for vanadium (Z = 23).
Note that we have written the ground state electron configurations with orbital subshells arranged in order by shells. This generally places the subshells in order by energy and puts the orbital subshells involved in chemical reactions at the far right. Let us skip to the zinc atom (Z = 30). The 3d subshell has filled fully; the ground state electron configuration is [Ar]3d104s2 .
Now the 4p orbital subshell begins to fill its ground state electron configuration, starting with gallium (Z = 31), configuration [Ar]3d104s2 4p1 , and ending with krypton (Z = 36), configuration [Ar]3d104s2 4p6.
Electron Configurations and the Periodic Table
By this time you can understand a pattern developed among the ground-state electron configurations of the atoms. This pattern explains the periodic table, which was briefly described in this article. Consider helium, neon, argon, and krypton, elements in Group 8A of the periodic table, which are called noble gas elements.
Neon, argon, and krypton elements have electron configurations in which a p subshell orbital has just filled. (Helium has a filled 1s subshell orbital; no 1p subshell orbital is possible.)
neon 1s2 2s2 2p6
argon 1s2 2s2 2p6 3s2 3p6
krypton 1s2 2s2 2p6 3s2 3p6 3d104s2 4p6
These elements are the first members of the group called noble gases because of their relative unreactivity as compared to other elements in the periodic table. Look now at the electron configurations of beryllium, magnesium, and calcium, members of the group of alkaline earth metals (Group 2A), which are similar, moderately reactive elements in the left side of the periodic table.
beryllium 1s2 2s2 or [He]2s2
magnesium 1s2 2s2 2p6 3s2 or [Ne]3s2
calcium 1s2 2s2 2p6 3s2 3p6 4s2 or [Ar]4s2
Each of these electron configurations consists of a noble-gas core, that is, an inner-shell electron configuration corresponding to one of the noble gases, plus two outer electrons with an ns2 configuration.
The elements boron, aluminum, and gallium (Group 3A) in the periodic table also have similarities. Their electron configurations are
boron 1s2 2s2 2p1 or [He]2s2 2p1
aluminium 1s2 2s2 2p6 3s2 3p1 or [Ne]3s2 3p1
gallium 1s2 2s2 2p6 3s2 3p6 3d104s2 4p1 or [Ar]3d104s2 4p1
Boron and aluminum have noble-gas cores plus three electrons with the general ground state electron configuration ns2 np1 . Gallium has an additional filled 3d subshell. The noble-gas core together with (n – 1)d10 electrons is often referred to as a pseudo-noble-gas core of the atom, because these electrons usually are not involved in chemical reactions.
An electron in an atom outside the noble-gas or pseudo-noble-gas core of the atom is called a valence electron. Such electrons are primarily involved in chemical reactions, and similarities among the configurations of valence electrons (the valence-shell electron configurations) account for similarities of the chemical properties among groups of elements.
This similarity explains what chemists since Mendeleev have known—the physical and chemical properties of elements in any group are similar. The main-group (or representative) elements of periodic table have valence-shell electron configurations nsa npb , with some atoms of a and b. (b could be equal to 0.)
In other words, the outer s or p subshell orbital is being filled. Similarly, in the d-block transition elements (often called simply transition elements or transition metals), a d subshell orbital is being filled. In the f-block transition elements (or inner transition elements), an f subshell orbital is being filled.
The polarity of the molecules
The polarity of the molecules are listed as follows
- Polarity of BeCl2
- Polarity of SF4
- Polarity of CH2Cl2
- Polarity of NH3
- Polarity of XeF4
- Polarity of BF3
- Polarity of NH4+
- Polarity of CHCl3
- Polarity of BrF3
- Polarity of BrF5
- Polarity of SO3
- Polarity of SCl2
- Polarity of PCl3
- Polarity of H2S
- polarity of CS2
- Polarity of NO2+
- Polarity of HBr
- Polarity of HCl
- Polarity of CH3F
- Polarity of SO2
- Polarity of CH4
Lewis Structure and Molecular Geometry
Lewis structure and molecular geometry of molecules are listed below
- CH4 Lewis structure and CH4 Molecular geometry
- BeCl2 Lewis Structure and BeCl2 Molecular geometry
- SF4 Lewis Structure and SF4 Molecular geometry
- CH2Cl2 Lewis Structure and CH2Cl2 Molecular geometry
- NH3 Lewis Structure and NH3 Molecular geometry
- XeF4 Lewis Structure and XeF4 Molecular geometry
- BF3 Lewis Structure and BF3 Molecular geometry
- NH4+ Lewis Structure and NH4+ Molecular geometry
- CHCl3 Lewis Structure and CHCl3 Molecular geometry
- BrF3 Lewis Structure and BrF3 Molecular geometry
- BrF5 Lewis Structure and BrF5 Molecular geometry
- SO3 Lewis Structure and SO3 Molecular geometry
- SCl2 Lewis structure and SCl2 Molecular Geometry
- PCl3 Lewis structure and PCl3 Molecular Geometry
- H2S Lewis structure and H2S Molecular Geometry
- NO2+ Lewis structure and NO2+ Molecular Geometry
- HBr Lewis structure and HBr Molecular Geometry
- CS2 Lewis structure and CS2 Molecular Geometry
- CH3F Lewis structure and CH3F Molecular Geometry
- SO2 Lewis structure and SO2 Molecular Geometry
- HCl Lewis structure and HCl Molecular Geometry
- HF Lewis structure and HF Molecular Geometry
- HI Lewis structure and HI Molecular Geometry
- CO2 Lewis structure and CO2 Molecular Geometry
- SF2 Lewis structure and SF2 Molecular Geometry
- SBr2 Lewis structure and SBr2 Molecular Geometry
- PF3 Lewis structure and PF3 Molecular Geometry
- PBr3 Lewis structure and PBr3 Molecular Geometry
- CH3Cl Lewis structure and CH3Cl Molecular Geometry
- CH3Br Lewis structure and CH3Br Molecular Geometry
- CH3I Lewis structure and CH3I Molecular Geometry
- SCl4 Lewis structure and SCl4 Molecular Geometry
- SBr4 Lewis structure and SBr4 Molecular Geometry
- CH2F2 Lewis structure and CH2F2 Molecular Geometry
- CH2Br2 Lewis structure and CH2Br2 Molecular Geometry
- XeCl4 Lewis structure and XeCl4 Molecular Geometry
- BCl3 Lewis structure and BCl3 Molecular Geometry