Chemistry 105
The Chemistry of Color

Course Lecture Topic Information

Chem 105 Web Site
Dr. Kimberly Lawler-Sagarin
Elmhurst College

Discussion 4:
How do neon lights work?

Discussion 4 will run for two weeks.

This discussion will take place February 20 - March 5

Assignments for the discussion board

This week's discussion questions are as follows:

  • Q1: Based on your student number, post an answer to the following question.
    • (Students 1,2,9,10) Choose one of the following questions (try to pick on that no one has written about yet on the discussion board): What are the basic components of a neon light? Why is neon used in the lights? When/how were they discovered? How are these related to energy levels in atoms?
    • (Students 3,4,11,12,16) Look up information about your assigned element below. All of these are noble or "rare" gases and some are used in "neon lights". (see the periodic table references from discussion 3 - go to list).
      • student 3: neon
      • student 4: argon
      • student 11: krypton
      • student 12: xenon
      • student 16: radon
    • (Students 5,6,13,14) Choose one way of making neon lights different colors and share this with the class. What do you think are the advantages/disadvantages of this method?
    • (Students 7,8,15,17) Choose some aspect of the manufacturing of neon lights to share with the class. (filling the tubes, bending glass, obtaining the neon, etc.)

  • Q2: All students: Choose two elements from the periodic table that are in the same group (same column - see lab meeting 2 lecture) other than the noble gas elements. Write out the electron configuration of each element and share with everyone how the elements are similar and how they are different from one another. For example: what similar chemical or physical properties do they have? How do their chemical or physical properties differ?

  • Q3: Choose one of:
    • Comment on the answer of another student, either by adding to the answer or asking a followup question.
    • Answer another student's question.
    • Look up information regarding other types of lights (incandescent, fluorescent, sodium vapor, halogen, etc.) and describe how they work to the class.

Text Readings
  • Electron Configuration: Chapter 3: Sections 3.7-3.8
  • Ion Formation: Chapter 5: Section 5.1-5.5
  • Lewis Dot Structures: Chapter 5: Section 5.7, 5.11
  • See also: lab meeting 2 lecture)

Topic Background

Electrons and their Configuration in Atoms

As we learned last week, electrons occupy different energy levels or "shells" in an atom. This week, we will learn a little about how the electrons are arranged in the atom in more detail.

Atomic energy levels are discrete (or not continuous) - that is, only certain values for the energy are allowed. Electrons in the lower energy levels are generally closer to the nucleus. Each energy "shell" is given a number (1, 2, 3, etc.) designated by the letter "n". Each electron energy shell can only hold a certain number of electrons.

The higher the energy level, the larger the number of electrons it can hold.

Electron Shell (n) Maximum # Electrons (2n2)
1 2
2 8
3 18
4 32

Electron Shells and Subshells

The number n, associated with each electron energy shell is known as the principle quantum number or energy quantum number.

Each electron shell can be further broken down into one or more smaller units called subshells.

Each of these subshells also has an associated quantum number, l.

  • For a given shell, l can take on any integer value from 0 to n-1.

  • Each value of l is associated with a subshell.

  • We also use letters (s, p, d, f...) to identify these different subshells. s is used for l=0, p for l=1, etc.

  • Each subshell can be represented by a symbol which combines the numerical value of n (1,2,3, etc.), with the symbolic value for l (s,p,d,f), as in:
    3p    for n=3   and  l= 1 = symbol p

One can determine the number of subshells for a given shell from the value of n for that shell.


The first, or lowest, energy shell is the n=1 shell.

Given that l can take on only integer values from 0 to n-1:

    l=0 is the only allowed value for l for this shell (n-1=0 when n=1). The symbol for l=0 is s

Thus, the n=1 shell is composed of only one subshell, with the symbol: 1s

Another Example:

The next highest energy shell is the n=2 shell.

Given that l can take on only integer values from 0 to n-1:

  • There are two allowed values for lL l=0 and l=1 (n-1=1 when n=2).

Thus, the n=2 shell is composed of two subshells, with the symbols: 2s and 2p.

Here's how things break down for the first four energy shells. Note that every time you increase "n", an additional subshell is added.

Shell (n) allowed l
(0 to n-1)

Energy Ordering of the Subshells

Within a given shell, the ordering of the energy of the subshells is given by:

s < p < d < f

Each subshell consists of one or more orbitals. These orbitals can be thought of as representing, in an average way, the region of space that the electrons inhabit. We'll be talking about these more later in the course. The type of subshell determines the number of orbitals in the subshell and the total number of electrons the subshell can hold, as follows:

Subshell Type # orbitals max # electrons
s 1 2
p 3 6
d 5 10
f 7 14

You may notice the the number of electrons is twice the number of orbitals. This is because each orbital can hold two electrons.

Putting this information together with the number of subshells allowed in each shell, we come up with a more detailed picture of how many electrons are allowed in each energy shell.

  • Shell 1 (n=1): one subshell, the 1s. An s subshell holds 2 electrons, so a maximum of 2 electrons fit into the n=1 level.

  • Shell 2 (n=2): two subshells, the 2s and 2p. An s subshell holds 2 electrons, a p holds 6, so a maximum of 8 electrons fit into the n=2 level.

  • Shell 3 (n=3): three subshells, the 3s, 3p and 3d. An s subshell holds 2 electrons, a p holds 6, and a d holds 10, so a maximum of 18 electrons fit into the n=3 level.

Here's the above in tabular form with the addition of the n=4 level.

Shell (n)
e-'s in subshell
total e-'s in shell

Writing Electron Configurations

The electron configuration for an atom is a way of writing down how electrons of an atom occupy the different energy shells and subshells.

Electron configurations for atoms are written by first grouping electrons into subshells. Next, write the designation of each subshell (nl) followed by a superscript indicating the number of electrons in that subshell.

n l (number of electrons in subshell)

(n = 1,2,3...., l = s,p,d,f....)

Terms for each subshell are placed after one another until all subshells containing electrons are included. An electron configuration looks something like this:


But, how do we get there? Use the following rule:

  • Electrons fill the lowest energy subshell available. Once that is full, electrons are placed into the next lowest subshell available.

Let's start with hydrogen (H). Hydrogen, with atomic number 1, has 1 electron. The lowest energy subshell is the 1s level, so the electron is placed there, and the configuration looks like:


Helium, atomic number 2, has 2 electrons. The second electron fills the 1s subshell (remember, s subshells can only hold 2 electrons).


Lithium, atomic number 3, has 3 electrons. The first two filled the 1s orbital, so the third one has to go into the next subshell, the 2s:


Next in line is Beryllium (Be), with 4 electrons.


When we get to boron (B), with 5 electrons, we have to begin a new subshell because the 2s is filled. The 2p in next in line:


As we proceed up in atomic number, we fill he 2p subshell and eventually have to start filling the 3s level. The electrons for the first 18 elements are placed into subshells in the following order: 1s, 2s, 2p, 3s, 3d. Their configurations are below:



Electron configurations for the first 18 elements
Atomic Number
Electron Configuration
H 1 1s1
He 2 1s2
Li 3 1s22s1
Be 4 1s22s2
B 5 1s22s22p1
C 6 1s22s22p2
N 7 1s22s22p3
O 8 1s22s22p4
F 9 1s22s22p5
Ne 10 1s22s22p6
Na 11 1s22s22p63s1
Al 13 1s22s22p63s23p1
Cl 17 1s22s22p63s23p5
Ar 18 1s22s22p63s23p6

Noble Gas Notation

A simplified way of writing electron configurations makes use of the noble gas notation.

Consider the electron configuration of Li:


Noting that 1s2 is the electron configuration of He, this can also be written as:


Similarly, for sodium, we can replace 1s22s22p6 with [Ne]:


Atomic Number
Electron Configuration
Be 4 [He]2s2
F 9 [He]2s22p5
Ne 10 [He]2s22p6 = [Ne]
Cl 17 [Ne]3s23p5
Ar 18 [Ne]3s23p6 = [Ar]


Beyond the first 18 elements

Beyond the first 18 elements, things get a little more complicated. Basically, the energy shells overlap. For example, the 4s subshell actually has a lower energy than the 3d. Thus, the n=3 and n=4 levels overlap a bit. Electrons will fill the lowest energy subshell, so they will fill the 4s BEFORE the 3d. Below is a table of the energy ordering for the various subshells. As you go up vertically in the table, the subshells increase in energy. Note the overlapping shells.

Energy Level Ordering for Subshells




Electron Configurations After Element 18

The 4s subshell lies lower in energy than the 3d subshell, thus it is filled before the 3d.

The electron configuration of potassium (K), element 19 is:


Using the simplified ''noble gas notation'' we can write this in the simplified way:


where [Ar] represents the electron configuration of an argon atom.  

The electron configuation of calcium (Ca), element 20 is:




Starting with scandium (Sc, atomic number 21), we begin to fill the 3d:


Once the 3d is filled, we can return to the n=4 shell by beginning to fill the 4p subshell with the element gallium (Ga, atomic number 31):


A table of the elements K through Kr is given below. Note the there are a few exceptions due to the fact that 1/2 filled subshells are especially stable, hoever this is not something you will be held responsible for.

Atomic Number
Electron Configuration
K 19 [Ar]4s1
Ca 20 [Ar]4s2
Sc 21 [Ar]4s23d1
Ti 22 [Ar]4s23d2
V 23 [Ar]4s23d3
Cr 24 [Ar]4s13d5 - note: 1/2 filled subshells
Mn 25 [Ar]4s23d5
Fe 26 [Ar]4s23d6
Co 27 [Ar]4s23d7
Ni 28 [Ar]4s23d8
Cu 29 [Ar]4s13d10 - note: 1/2 filled s subshell
Zn 30 [Ar]4s23d10
Ga 31 [Ar]4s23d104p1
Ge 32 [Ar]4s23d104p2
As 33 [Ar]4s23d104p3
Se 34 [Ar]4s23d104p4
Br 35 [Ar]4s23d104p5
Kr 36 [Ar]4s23d104p6 = [Kr]

* 1/2 filled subshells are especially stable  

Neon Lights

Find out more about neon lights for your discussion question Q1 at the following websites:

Chemistry 105 Home
Blackboard Home
Elmhurst Chemistry Home
Elmhurst College Home

 K. Lawler-Sagarin, 2004