Course Lecture Topic Information
Discussion 4 will run for two weeks.
This discussion will take place February 20 - March 5Assignments for the discussion board
This week's discussion questions are as follows:
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 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.
One can determine the number of subshells for a given shell from the value of n for that shell.
Here's how things break down for the first four energy shells. Note that every time you increase "n", an additional subshell is added.
Energy Ordering of the Subshells
Within a given shell, the ordering of the energy of the subshells is given by:
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:
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.
Here's the above in tabular form with the addition of the n=4 level.
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 = 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:
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:
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]:
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.
* 1/2 filled subshells are especially stable
Neon LightsFind out more about neon lights for your discussion question Q1 at the following websites: