Pre-Test
Q&A	Elements of Physics: Matter: Atoms and Molecules

Instruction 1-8

Connection Among the Location in the Table, the Atomic Number, and Mass | How to Identify Metals, Semimetals, Nonmetals, and Halogens | How to Identify Alkaline Metals, Alkaline Earth Metals, and Transition Metals | Lanthanide, Actinide, Transactinide, and Transuranium Elements | Ionization Energy, Electronegativity, Relative Sizes | How Many Electrons Can Bond? | Size and Mass | Location and Quantum Electron Configuration | Summary

The format for writing it is this: 1s²

The number 1 refers to the energy level (in this case, energy level 1). The “s” stands for subshell that the electrons are occupying. The superscript number 2, (raised) tells us the number of electrons in that subshell.

For example: 3d⁴. This notation indicates that in the 3rd shell of a particular atom, its d subshell contains 4 electrons.

You learned about electron configuration in Instruction 1-1, but it might be helpful to repeat just a little of that Instruction.

You know about the electron shells of atoms (k, l, m, n, o, p and q or 1, 2, 3, 4, 5, 6 and 7). And you also know that each shell (energy level) is divided into subshells (sublevels). Those subshells are designated s, p, d and f. The number of subshells an energy level can contain is the same as its number. For example, the first principal shell, the k shell, has one subshell, which is designated the 1s shell. The second principal shell, the l shell, has two subshells, which are designated 2s and 2p. The third shell, the m shell, has three subshells: 3s, 3p and 3d. And the fourth shell, the n shell, has four: 4s, 4p, 4d and 4f. There are more subshells in larger shells, but they are not currently used.

To summarize -- shell 1 has 1 sublevel, shell 2 has 2 sublevels, shell 3 has 3 sublevels, and shell 4 has 4 sublevels.
But even that isn’t enough!

These subshells (sublevels) are further divided into ortibals, each of which can hold a maximum of two electrons. Therefore an s subshell, which is made up of one orbital, can contain two electrons. A p subshell has three orbitals, so it can contain six electrons. A d subshell has five orbitals so it can contain 10 electrons. And an f subshell has seven orbitals so it can contain 14 electrons (7 orbitals x 2 electrons per orbital).

Let’s put it in chart form:

In order to simplify writing about the different configurations, a notation is used to specify the information call the electron quantum numbers. There are 4 numbers and named Principle Quantum Number (n), Azimuthal Quantum Number (l), Magnetic Quantum Number (m), and Spin Quantum Number (s). Each electron in an atom is given a different set of these 4 numbers which specifies its location. A summary of these numbers is provided here: http://www.saskschools.ca/curr_content/chem30/modules/module2/lesson2/quantumandelectronst.htm

The Principle Quantum Number (n) describes the energy level of the electron. It ranges in value from 1 to infinity. But since we usually work with the electrons in the ground state, it ranges from 1 to 4. It is sometimes call the shell number. The energy levels are associated with a certain shell as outlined in the table.

The Azimuthal Quantum Number (l) describes the shape of the subshell orbital. It ranges from 0 to (n - 1). The shapes of the orbitals are associated with the s, p, d, f designation as outlined in the table. This means that when n = 1, l can only be 0; only the 1s orbital exist. When n = 2, l can be either 0 or 1; the orbitals can be either 2s or 2p.

The Magnetic Quantum Number (m) describes the orientation of the orbital in relation to an arbitrary grid. Since the s orbital is spherical, there is only need for 1 value of the Magnetic Quantum Number. Since there are three p orbitals, there needs to be 3 values for this number. The number m takes on the values from -l to l as shown in the table.

The Spin Quantum Number (s) describes the orientation of the electron in the orbital. There are only 2 ways the electron can be oriented, so there are only 2 values: -½, and ½.

Let’s look at the outer most electron in cesium. It is in the 6th energy level, so n = 6. It is in the s orbital, so l = 0. The Magnetic Quantum Number can only be 0 when l = 0, so m = 0. The Spin Quantum Number can be assigned either value, lets give s = ½. It can also be designed as 6s¹.

Let’s review the shell filling order for electrons that was presented in Instruction 1-6. The figure is shown on the right and following the arrows tells us the order the shells are filled. The designation for all the electrons in cesium is 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s²4d¹⁰5p⁶6s¹. Counting all the electrons, there are 2 + 2 + 6 + 2 + 6 + 2 + 10 + 6 + 2 + 10 + 6 + 1 = 55 electrons, which is the atomic number of cesium and how many protons there are in the atom.

The reason for this peculiar order is that electrons fill up low energy orbitals before they fill higher energy ones. Low energy levels are the ones closest to the nucleus. The 4s (5s, 6s) orbital will have a slightly lower energy level than the previous 3d (4d, 5d) orbital because of the pull from the protons in the nucleus. This means the s orbital of the next shell will fill up before the d orbitals.

To write the electronic configuration, we need to follow some rules. Using the diagram above to give us the order of the subshells to fill and knowing that there can be a maximum of 2 electrons in the s subshell, 6 electrons in the p subshell, 10 electrons in the d subshell and 14 electrons in the f subshell. We just need to have as many electrons as protons in the atom (this is the atomic number).

So let’s learn to write a formula. Let’s start with element #1, hydrogen (H). Hydrogen has only one electron in the 1s orbital – 1s¹. So writing the electron configuration for hydrogen would be just that – 1s¹.
Now let’s go to the second Period and start filling the second level. Lithium’s 3rd electron goes into the 2s orbital because 2s has lower energy than 2p. So lithium’s electronic configuration would be 1s²2s¹. You see that the superscript (raised) number equals lithium’s atomic number: 3.

Let’s try one more – a slightly more complicated one: phosphorus (P). Phosphorus has an atomic number of 15, so the superscript numbers should total 15. That makes the notation for phosphorus: 1s²2s²2p⁶3s²3p³.

Sure, it’s a little hard to work out -- but aren’t you glad you’ve mastered it?

Because, believe us, you will be having test questions on writing electron configurations!

for Students, Parents and Teachers

Now let's do Practice Exercise 1- 8 (top). 

Summary

You have now completed Lesson 1 on Chemistry and are ready to do the Problem and Test sections.

You may wish to review any or all of the topics before answering the questions that follow. You may also wish to obtain additional material from the links in any of the Instructions or the links below
before answering the questions.

Good luck!