Electron Configuration


2. Electron Configuration

2.1 The One-Electron Atom: Energy levels

When matter receives energy from its surroundings, the internal structure of its atoms is modified. The total energy of an atom is quantified,that is, it can only take certain specific values.These values are commonly called Energy levels of an atom.An energy level En of order n is expressed by the following relation:

En = -13.6/n2 e.v where n is a whole number >=1 and 1e.v = 1.6x 10 -19 J.

The minimum energy state (n=1) is the ground state of the hydrogen atom. The corresponding energy level is called K level.When an atom receives energy from an external source, the electron moves to a higher energy level ( n > 1).We say that the atom is in an excited state. The electron has to occupy one of these levels by an abrupt jump.The absorption of a definite energy E takes the atom from the ground state to the excited state. Returning to the ground state involves emission of the same energy E.

The following diagram is the energy diagram of the hydrogen atom is used to represent its different energy levels:


Notice that successive energy levels are distinctly spaced for small values of n,become closer together for increasing values of n, and are nearly superimposed when n is large. To each value of n is associated a letter, according to the following table:


N Alphabetical Representation
1 K
2 L
3 M
4 N
5 O
6 P
7 Q

Atoms are normally found in their ground state which is their most stable state. The atom is in the most stable state when the electron occupies the lowest possible energy level.


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2.2 Atoms containing more than one electron: Energy Sublevels

The study of the spectra of different elements reveals that each energy level consists of many closely spaced levels.We say that each energy level(shell) is made up of one or many sublevels (subshells).The number of energy sublevels in each energy level En is equal to the numerical value of n.


The level K ( n=1) has one energy sublevel called 1s
The level L (n=2 ) has two energy sublevels called 2s and 2p

The level M ( n =3) has three energy sublevels,3s,3p,3d

The level N ( n=4) has four sublevels 4s,4p,4d and 4f.


Electrons in the same level of a multi-electron atom do not have exactly the same energy.For example, in the L level, the electrons are divided into two groups, where one possesses an energy slightly higher than the other does. The maximum number of electrons in an energy level of order n is 2n2.


The K level ( n=1) may contain 2 electrons
The L level ( n=2) may contain 8 electrons

The M level ( n=3) may contain 18 electrons

The N level ( n=4) may contain 32 electrons


As there is a limit to the number of electrons that can occupy an energy level, so there is a limit to the number of electrons which are capable of occupying an energy sublevel.

A sublevel s may hold a maximum of 2 electrons
A sublevel p may hold a maximum of 6 electrons

A sublevel d may hold a maximum of 10 electrons

A sublevel f may hold a maximum of 14 electrons


We can thus draw the diagram of levels and sublevels of a multi-electron atom, and indicate the maximum number of electrons in each level and sublevel.Note that beyond level 3, the energy level 4s precedes energy level 3d, although it should normally come after it. This change in the order of energy sublevels is called inversion, and the order of energy sublevels becomes more complicatedwhen we reach the levels 5 and 6.


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2.3 Filling of energy levels

The electron configuration of an atom is the distribution of electrons in the different energy sublevels.To write the electron configuration of an atom zX , we fill the energy sublevels successively with electrons, in increasing order of energy using the diagram of energy sublevels, until the number of electrons placed in these sublevels is Z. We may go on to the next energy sublevel only when the preceding sublevel has been completely filled.Often, the last energy sublevel is not completely filled with electrons.For example the element (Z=25) has the following electron configuration:

1s2 2s2 2p6 3s2 3p6 4s2 3d5

To write the electron configuration of an element we order the sublevels of an atom in the ground state according to the diagonal rule shown below.Click on the gray ball to start:



There are, however, a few exceptions to this rule : these are some possessing partially filled , d energy sublevels. Example:

The electron configuration of the element Z =29 is: 1s2 2s2 2p6 3s2 3p6 4s1 3d10
The atoms are not necessarily always in their ground state. They could be in their excited state.


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2.4 Lewis dot symbol

Of all the electrons of an atom, only the electrons that are located on the outer shell are responsible for chemical reactions ( formation of bonds and ions). The outer shell is called valence shell. The electrons that belong to this shell are called valence electrons. To determine the shape of a molecule, two things are important. The first is a correct Lewis Dot Structure; the second is the application of the Valence-Shell Electron-Pair Repulsion Theory ( also known as the VSEPR Theory).

The symbolic representation of valence electrons on an atom as pairs and unpaired electrons is called Lewis dot symbol . We can represent an electron pair by a dashed line and an unpaired electron by a single dot. The Lewis dot symbol represents atoms when they are ready to react. In Lewis dot symbol, an electron pair can be represented by two adjacent dots. The maximum number of unpaired electrons that could be placed is four. There are four rules to writing a good Lewis Dot Structure.

1.Add up all of the valence shell electrons in the molecule. (do not forget that ions have either gained or lost electrons

2.Pick the central atom. The central atom will be the atom with the largest bong capacity (or bond order).

3.Arrange the other atoms around the central atom. If the total number of electrons is even, the outer atoms will have a noble gas structure. Any extra electrons are put around the central atom.

4.Check the structure that you have drawn using Formal Charges.


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2.5 Valence shell electron pair repulsion theory (VSEPR)

A simple method of predicting shapes of covalent molecules is the VSEPR. In the theory, valence shell electron pairs are assumed to repel each other, assuming orientations to minimize repulsions and establish certain groups of molecular shapes.



As shown above, these shapes can be classified according to how many electron pairs surround the central atom. Two pairs - linear, three - trigonal planar, four-tetrahedral, five-trigonal bipyrimidal, six-octahedral.

VSEPR notation:

To help translate the Lewis Structure to a molecular structure, one can use the VSEPR notation. The VSEPR notation looks very much like the formula for a compound, with a few changes.

First, the central atom is called "A." Secondly, all the outer atoms are designated with an "X." This is true even if the atoms are different (it only matters that these are atoms that surround the central atom. Also, molecules with double bonds are treated as if the double bond was merely a single bond, because there is still only one outer atom in the bond.) Finally, any lone pair electrons are designated with an "E." For instance, the water atom has 2 hydrogen atoms and 2 lone pair electrons surrounding the central atom. Therefore, its VSEPR notation is AX2E2.

Now, let's look at the different molecular configurations. The first Electron configuration is linear:



The second electron configuration is trigonal planar. It has three electron pairs surrounding the central atom. The possible molecular configurations are either trigonal planar (if all of the electron pairs are bonding pairs) or bent (If one is a lone pair).



If our molecule has a total of 4 pairs of electrons its electron configuration is tetrahedral. Below are the possible molecular structures stemming from this.





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