We know that an element is defined by the number of protons it contains. Now we will learn how electrons and their particular arrangement within an atom can affect how that atom reacts, chemically and physically.


What is it, why is it important to our discussion about atoms and the arrangement of their electrons?

Electromagnetic radiation is energy – we describe it as a wave – visible light is only a small portion of the electromagnetic spectrum. Each type of energy (x-rays, gamma-rays, microwaves, radiowaves) as well as each type of light (red, blue, green) has different wave characteristics. The characteristics which distinguish different types of light are the electromagnetic radiation's



Wavelength – the distance between two peaks (or two troughs) of the wave.

Frequency – the number of wavelengths passing a given point in one second.

Energy -- is directly proportional to the frequency-- if the frequency increases, so does the energy of the radiation and vice versa.


When atoms are excited by energy (such as heat) – they emit energy in the form of light (sometimes colored light) and different atoms emit different energies.

These energy spectra are how scientists explore the universe today--to learn the content of stars, whether the universe is expanding, etc.

Bohr’s Theory:

When white light is passed through a prism, it is separated into its components of the spectrum, or the "rainbow" of colors. What, then, is white light made of? This "rainbow" is called a continuous spectrum, as shown below (look familiar?):


When atoms are excited they emit light--but not as a continuous spectrum. The light is emitted as a specific pattern of visible, colored lines which is particular to each element. Hence these line spectra can be used to identify individual elements. The hydrogen line spectrum is shown below:


Why do elements not emit a continuous spectra and why does each element emit a particular set of lines which make up its line specturm?

The answer lies in Bohr's Theory of the Atom and its structure. Bohr found that:

Following are two graphics which show this effect:




When electrons are excited they absorb energy and move to a higher energy level (gold arrow). When they emit light, they move to a lower energy level. Violet light (violet arrow) is produced when electrons move from level 6 to level 2. Blue light (blue arrow) is produced when electrons move from level 5 to level 2. Green light (green arrow) is produced when electrons move from level 4 to level 2. Red light (red arrow) is produced when electrons move from level 3 to level 2. This series of lines which is the visible hydrogen line spectrum is called the Balmer series. The energies of these emissions just happen to be in the visible range so we can see the colors. There are many other transitions which occur, but since they are not in the visible range we cannot see them. Some have wavelengths less than 400 nm and some have wavelengths greater than 750 nm (remember, the visible range is very limited, only 400-750 nm).


Visualize the atom as an onion with different levels of structure and sub-structure, all of which make up electron energy levels – each major energy level is designated by a number (called the quantum number):

Below is shown the relative energy levels and structure of the hydrogen atom. Remember, n = 1 is the major shell closest to the nucleus and it has only one subshell which, in turn, contains only one orbital, which can contain two electrons. What is the maximum number of electrons that can reside in the first major shell? In the diagram, the major shells are designated by the black lines coming from the energy arrow. The different types of subshells are indicated by the differently colored boxes. The orbitals themselves are the individual boxes. Remember, each box (orbital) can contain only a maximum of two electrons. Consequently, the energy of an electron is going to depend upon "which box it is sitting in" and what the energy of that box is. The lower the box (the closer it is to the nucleus) the lower (more negative) the energy and the more tightly an electron in that box is held.

Although the major shell primarily determines the energy of an electron, there are minor differences between the energies of the subshells within any major shell, particularly in an atom with more than one electron (actually, any atom other than hydrogen):

Once we have many electrons (which means many shells and subshells) in an atom, we begin to see some overlap of subshells--particularly, the s subshell from the fourth level (4s) overlaps with the d subshell from the third level (3d) such as shown below:

Shapes of the subshells:

For many reasons, we cannot define exactly where an electron is within an atom. We must, instead, define the volume which will most likely contain electron density. Consequently, the shapes of those volumes which can contain electron density are somewhat "fuzzy". We can depict volumes of high electron density (many dots) or volumes of low electron density (few dots), as shown in the diagrams below.

Cross-sections of an s-subshell orbital

A spherical s subshell orbital with high (very blue)

and decreasing (less blue) electron density




All s subshell orbitals are spherical, the main difference

between different major shells is in the size of the sphere

electron density model                                       p subshell orbitals


d subshell orbitals

Electron Configurations are a particular arrangement or distribution of electrons among different subshells and shells within an atom.


1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 5d10 6p6

lower energy higher energy

strongest held e- less tightly held e-

You can see this pattern emerging from the energy levels in the diagrams above.


Pauli Exclusion Principle: an orbital can contain no more than two electrons and those two electrons must be paired, in other words, they must have opposite spin (usually indicated by one up-arrow and one down-arrow, as we will see shortly).

Hund’s Rule: in a subshell where there are multiple orbitals with the same energy, electrons will enter the each orbital singly until all orbitals are half filled before pairing with other electrons in the subshell.

Orbital Diagrams: For ground state electron configurations, electrons enter the electron configuration (orbital diagram) pattern beginning at the lowest energy and filling the pattern to the highest energy while obeying the Pauli Exclusion Principle and Hund’s Rule.

Examples:                   Orbital Diagrams or Electron Configurations for: H, He                       Row 1

Orbital Diagrams or Electron Configurations for: Li, Be, B, N, F, Ne   Row 2

Below are shown electronic configurations for the above elements in three forms of symbolism:

I     Complete Electron Configuration
II   Shorthand Electron Configuration
III Shorthand Valence Electron Configuration

I                                                              II                                       III

Periodic Table Row 1:

1H    ­                                                             1H     1s1                             1H   1s1

2He   ­¯                                              2He   1s2                           2He   1s2  

Periodic Table Row 2:

3Li  ­¯   ­                                                        3Li   1s2 2s1                       3Li [He] 2s1
       1s  2s

4Be ­¯   ­¯                                                     4Be 1s2 2s2                      4Be [He] 2s2
       1s   2s

5B ­¯  ­¯     ­                                                5B 1s2 2s2 2p1                  5B [He] 2s2 2p1
     1s   2s   2p

6C ­¯  ­¯    ­   ­                                        6C 1s2 2s2 2p2                  6C [He] 2s2 2p2
     1s   2s      2p

 7N   ­¯   ­¯  ­  ­  ­                          7N 1s2 2s2 2p3                    7N [He] 2s2 2p3
        1s    2s      2p

8O ­¯  ­¯     ­¯  ­  ­                         8O 1s2 2s2 2p4                         8O [He] 2s2 2p4   
     1s   2s         2p

9F ­¯  ­¯   ­¯ ­¯ ­                           9F 1s2 2s2 2p5                      9F [He] 2s2 2p5
     1s   2s       2p

10Ne ­¯  ­¯    ­¯ ­¯ ­¯                 10Ne 1s2 2s2 2p6                    10Ne [He] 2s2 2p6
        1s   2s          2p


Please Notice:


Periodic Table Row 3:

11Na  ­¯  ­¯    ­¯ ­¯ ­¯     ­                   11Na  1s2 2s2 2p6 3s1                  11Na [Ne] 3s1
         1s   2s          2p        3s

12Mg  ­¯ ­¯    ­¯ ­¯ ­¯   ­¯                   12Mg 1s2 2s2 2p6 3s2                    12Mg [Ne] 3s2
          1s  2s         2p        3s

13Al ­¯ ­¯     ­¯ ­¯ ­¯    ­¯  ­               13Al  1s2 2s2 2p6 3s2 3p1              13Al [Ne] 3s2 3p1
       1s  2s           2p         3s 3p

14Si  ­¯ ­¯   ­¯ ­¯ ­¯    ­¯     ­ ­            14Si 1s2 2s2 2p6 3s2 3p2                    14Si [Ne] 3s2 3p2
        1s  2s         2p         3s     3p

15­¯ ­¯     ­¯ ­¯ ­¯   ­¯   ­ ­ ­        15P  1s2 2s2 2p6 3s2 3p3                  15P [Ne] 3s2 3p3
      1s  2s           2p        3s      3p

16S  ­¯ ­¯   ­¯ ­¯ ­¯   ­¯  ­¯ ­ ­          16 1s2 2s2 2p6 3s2 3p4                      16S [Ne] 3s2 3p4
       1s  2s         2p        3s      3p

17Cl  ­¯ ­¯   ­¯ ­¯ ­¯    ­¯    ­¯  ­        17Cl  1s2 2s2 2p6 3s2 3p5                17Cl [Ne] 3s2 3p5
        1s  2s         2p         3s     3p

­¯ ­¯  ­¯ ­¯ ­¯   ­¯   ­¯ ­¯ ­¯     18Ar  1s2 2s2 2p6 3s2 3p6                18Ar [Ne] 3s2 3p6
        1s  2s        2p         3s        3p

Once again, please notice that the configurations in blue are the inner or core electrons (represented by the noble gas element of that configuration, Ne for row 3 elements) and the configurations in pink are the valence electrons. The electron configurations of the valence electrons are very important.

Now, let's compare the electron configurations (complete and valence) of elements in a group, say, group IA:

3Li   ­¯  ­                                                               3Li [He] 2s1
       1s  2s

­¯ ­¯   ­¯ ­¯ ­¯   ­                                    11Na [Ne] 3s1
        1s   2s        2p        3s

19­¯ ­¯   ­¯ ­¯ ­¯    ­¯   ­¯ ­¯ ­¯   ­             19K [Ar] 4s1
      1s  2s          2p        3s         3p       4s

Well, I think you can see the pattern. Within a group, elements have very similar valence electron configurations--same number of electrons in the same type of subshells but the only difference is in the major shell. What other similarities do elements in the same group exhibit? (see chapter four on the periodic table)


Notice that in the orbital diagrams above the outer electrons of elements in the same group are similar: the same number of electrons, in the same types of orbitals but in different major shells.

This type of periodicity can be used to quickly determine the electron configuration or orbital diagram of any element based upon its position in the periodic table.

Valence Electrons and Valence Shells:


10Ne 1s2        2s2 2p6 valence shell containing the maximum number of electrons

11Na 1s2 2s2 2p6      3s1 valence shell containing one valence electron

13Al 1s2 2s2 2p6       3s2 3p1 valence shell containing three valence electrons




Many physical properties depend upon the electronic properties of the atoms. Since the valence electronic structures are periodic, those physical properties also reflect that periodicity. Properties such as atomic and ionic radii and ionization energy are periodic and can be predicted based upon an atoms relative position in the periodic table.