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Group Activity:
Electron Configuration

Here is a template for electron configurations:
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 etc.

Introduction

The world of the atom is a very strange world. It is quite unlike the world we are used to. The usual relatively simple laws of physics do not apply. Instead the world of the atom is one dominated by waves

and probability. The speed and even the location of objects in this world are subject to uncertainty.


Electron Orbitals

Werner Heisenberg, one of the founders of Quantum Mechanics, was the first to describe this uncertainty mathematically. The Heisenberg Uncertainty Principle states that it is impossible to know with high levels of certainty both the location and the velocity of an electron. The better you know where an electron is the worse your knowledge of where it is going.

electrons.not.on.circles (2K)

The practical meaning of this is that the common picture given of atoms that looks like a solar system is not just a little inaccurate. It is dead wrong. Diagrams such as the one at right are useful for counting electrons and illustrating different energy levels. But they are worthless as picture of what atoms really look like.

The solar system model of the atom uses orbits to show where the electrons are. Electrons do not really orbit the nucleus. They do live in regions of space called orbitals. Orbitals are organized by the energy level or shell that they belong to.

1s (18K)
The 1s Orbital

Orbitals are clouds of probability. That is, an orbital is the shape of the region in space where an electron is most likely to be found. At left you see what is called the 1s orbital. This ball-shaped orbital is gives a better idea of what a hydrogen atom actually looks like. The dots in the picture represent possible positions of the one electron the hydrogen has. The nucleus (too small to see) is in the center of the picture. Notice that there are more dots near the center of the sphere than near the edges. This shows that the electron is most likely to be near the nucleus. But that does not mean it is impossible to find the electron very far away from the nucleus. There is an infinitesimal probability (less than the probability of winning the lottery) that the electron in a hydrogen atom will be on the Moon.

1s (18K)
The 2s Orbital

It might help to think of it this way. Picture a bird feeder on a pole in someone’s backyard. A huge flock of small birds has just descended into the yard on its way to a summer habitat in Canada. If you take a picture of the yard at different moments in time the birds are most likely to be found as close to the birdfeeder as they can get. The shape of the region of space around the feeder where a bird is most likely to be found will be roughtly spherical, just like the 1s orbital of an atom. Each bird represent the location where an electron might be at any given time.

The 2s orbital of an atom is shaped just like the 1s orbital. The difference is that is larger and has a higher energy than the 1s orbital. The 2s orbital also has a small part that is closer to the nucleus so that it looks a bit like a sphere within a sphere.

Generally, s orbitals all look alike but get bigger at higher energies. The same is true for the other types of orbitals: at higher energies they look similar but are bigger.



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Orbitals and Energy

The orbitals of an electron have a variety of shapes. The shapes fall into certain categories that are known by letters. The s orbitals are all spherical and there is only one s orbital in each shell. The p orbitals all look like barbells made out of balloons and there are always three p orbitals per shell. Each p orbital points in a direction perpendicular to the others (px, py, pz). Most of the d orbitals look like various combinations of four balloons. There are always five d orbitals in a shell that contains d orbitals. There are also f, g, h, i, (etc.) orbitals but their shapes are very complicated and are really of no importance at this point. Incidentally, there are seven f orbitals in shells which can contain f orbitals.

2s 2px 3dxy
The 2s Orbital

The 2px Orbital

The 3dxy Orbital

The orbitals are organized by their energies. Each shell of the electron structure of an atom is organized by its energy: the more energy, the farther a shell is from the nucleus. The first shell can only have one kind of orbital: the s orbital. The second shell has both an s orbital and a p orbital. The third shell has s, p, and also d orbitals. The fourth shell has s, p, d, and f orbitals. After that, as the energy of the shells increases the types of orbitals go by the letter of the alphabet.

electron.orbitals.relative.energy (5K)

Look at the diagram above. It shows how the orbitals are organized by their relative energy. In this activity you will learn how to determine where the electrons in atoms and ions are placed in such a diagram. You will also learn how to write a shorthand notation for such diagrams.

The energy levels or shells are distinguished by their number. This number is called the principal quantum number and is the most important quantum number for determining how much energy an electron has.

In each energy level the orbitals have specific names. For electrons with principal quantum number 1 there is only a ‘1s’ orbital. In the shell with principal quantum number 2 there are both a ‘2s’ orbital and three ‘2p’ orbitals. In the third shell there are one ‘3s’, 3 ‘3p’, and 5 ‘3d’ orbitals. The pattern continues this way and at each level a new type of orbital is added.



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Apartment House Analogy

One way to approach the very abstract task of filling electron orbitals is to use an analogy to make it a bit more concrete. Imagine you are the landlord of a very strange apartment building. Your job is to fill the apartments in the building in the most efficient way possible. You are required by the owner of the building to fill the rooms in a certain way. The rules you have to follow are as strange as the building because quantum mechanics is not like anything you might have expected. The rules are summarized in the table below.

In the building the different floors are like the different energy levels (or shells) in an atom. The energy levels are numbered starting from one, just like the floors in an apartment house. Each room corresponds to one orbital: one box on the diagram on the previous page. The rooms have a capacity of two electrons (two people) each. In each room only a man and a woman may be paired together. In the strange world of quantum mechanics there are no same-gender room mates.


Apartment House Rules Electron Rules
From the Bottom Up: Rooms must be filled from the ground floor up. Fill the one room on the first floor before starting to put new tenants on the second floor. Then fill the s room before the p rooms. At higher floors the order might change a bit. Aufbau Principle: the electrons fill the available orbitals from lowest energy to highest energy. In the ground state all the electrons are in the lowest possible energy level.
Singles First: the owner of the building wants to have the tenants spread out as much as possible. For that reason singles are placed in rooms before couples. If couples must be placed into a room then all of the other rooms on that floor must already have a single in them. Hund’s Rule: The electrons must be placed into the orbitals in such a way that no pairs are put together unless absolutely necessary. That is, single electrons must be placed into boxes first and then paired up if necessary.
Opposite Gender Only: When two people are placed in a room they must be of opposite genders. No men may room together and no women may room together. This is an arbitrary rule on the part of the owners: in a just world we wouldn’t have to follow it. But quantum mechanics has nothing to do with justice. Pauli Exclusion Principle: Electrons come in two varieties based on the direction they are ‘spinning’. There is an Up spin and a Down spin. Up and Down spins are always paired together and Up-Up or Down-Down combinations are not allowed. No two electrons can ever be in the same place at the same time.

Here is an example of the rooms in the apartment house having been filled in by the rules. The element Sulfur has 16 electrons and starting from the bottom up you place two (1 Up and 1 Down) in the 1s orbital.

e.Config.Arrows.Sulfur (6K)

Next, electrons are placed in each of the 2p orbitals until there is one in each orbital. Then the rest of the electrons are filled into the 2p orbitals. After that the 3s orbital is filled and then the 3p. Electrons are placed in the boxes of the 3p orbitals one at a time until there is one in each orbital. After that the remaining electron is placed as part of a pair in the first 3p box.

One key thing that is of great help with this is that the business with the boxes can be abbreviated the way you see it next to the word Sulfur in the figure at right. The short hand works like this:

[principal quantum number][orbital type][# of e– ]


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The shorthand notation for the electron configuration of Sulfur is 1s2 2s2 2p6 3s2 3p4. In that notation 1s2 means principal quantum number 1, s-type orbital, 2 e-; 2s2 means principal quantum number 2, s-type orbital, 2 e-; 2p6 means principal quantum number 2, p-type orbital, 6 e-; 3s2 means principal quantum number 3, s-type orbital, 2 e-; and, finally, 3p4 means principal quantum number 3, p-type orbital, 4 e-.


Connection to Valence Electrons

Remember that when we first wrestled with the idea of the electron configuration of atoms we only looked at the shell number. Now you know that the principal quantum number is the real identity of that shell number. The valence shell of electrons is the outermost shell of electrons. Here is why: the inner shells of electrons are smaller than the outer shells. For this reason, only the outermost shell of electrons can interact with the rest of the world. The reason elements in the same column of the periodic table (the so-called Groups) have similar properties is not just because they have the same number of electrons in their valence shell. They have similar properties because their outermost orbitals have the same number of electrons and the same shapes..

Oxygen 1s2 2s2 2p4
Sulfur 1s2 2s2 2p6 3s2 3p4

Look at the electron configurations of Oxygen and Sulfur. Notice that in both cases the last orbitals to fill are p orbitals. Notice also that there are 2 electrons in an s orbital and 4 electrons in p orbital for both Oxygen and Sulfur.

For these reasons Oxygen and Sulfur belong to the p-block or elements in the periodic table. In fact, all of the elements from Groups 13 to 18 are part of the p-block. All of the elements in this block fill a set of p orbitals last and all of them have their valence electrons in s and p orbitals.

Here you can finally understand the reason that having eight electrons is so important. Each s orbital holds 2 electrons and each p orbital holds 6 electrons. Together they can hold the eight electrons it takes to finish off the orbitals and make the valence shell full.

Other sections of the periodic table are also known by the type of orbital that fills last in that section. Broadly, this way of looking at the periodic table can explain why it has just the shape that it has. In Groups 1 and 2 the s orbitals are filling. In Groups 3 through 12 the d orbitals are the last to be filled. At a principal quantum of one (n = 1) there is only an s orbital. At n = 2 there are both s and p orbitals available. At n = 3 there are s, p, and d orbitals available. This is why there are no transition metals between Magnesium and Aluminum. Incidentally, one important fact is the the 4s orbital fills before the 3d orbitals start to fill. This is because the 4s orbital is actually slightly lower in energy than the 3d orbitals. See the diagram on the second page of this packet. You will not be much concerned with any orbitals higher than the 4s orbital since there are some interesting but complicated exceptions to the rules when it comes to filling the 3d orbitals.



Shell.Filling.n1-n7.PTOE (8K)
This image is modified from the PDF file produced by the Chemistry Dept. at USM to teach this topic. See it here.


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Questions and Problems

Comprehension Questions

  1. In your own words explain what the Heisenberg Uncertainty Principle means.
  2. What is the difference between an orbit and an electron orbital?
  3. Quantum Mechanics relies on probability to say what region of space an electron is most likely to occupy. In your own words, describe how this principle results in the seemingly solid shapes of the orbitals. Use an analogy.
  4. In your own words describe the shapes of the s, p, and d orbitals. Draw a picture.
  5. What is the meaning of the term principal quantum number?
  6. How does the Aufbau Principle work? That is, what does it mean to say that orbitals are filled from the bottom up?
  7. What is Hund’s Rule? That is, how do you go about filling up electron orbitals according to this rule?
  8. Explain what is meant by the Pauli Exclusion Principle?
  9. In the following shorthand notation for the configuration of electrons tell what each part means: 1s2 2s2 2p6 3s2 3p6 4s2 3d5. (For example: 1s2 means that in the shell with principal quantum number 1 there are two electrons in an s orbital.)
  10. Why does the 4s orbital get filled with electrons before any electrons can be placed into a 3d orbital?


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Electron Configurations

The image at right is borrowed from Hyperphysics.
order.of.orbital.filling (7K)
Orbital Filling
Order
Note for Instructors: I have made a sheet with three problems in electron configuration for use as practice. It does not specifiy the atom or ion and goes up to the 4p level. You can see it here.
  1. Fill the boxes below with the arrow notation for electrons showing the correct ground state electron configuration for the element He.
electron.boxes.to.fill (3K)
  1. Write out the electron configuration in the short-hand notation.
  2. Fill the boxes below with the arrow notation for electrons showing the correct ground state electron configuration for the element Ne.
electron.boxes.to.fill (3K)
  1. Write out the electron configuration in the short-hand notation.
  2. Fill the boxes below with the arrow notation for electrons showing the correct ground state electron configuration for the element Ar.
electron.boxes.to.fill (3K)
  1. Write out the electron configuration in the short-hand notation.


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  1. Fill the boxes below with the arrow notation for electrons showing the correct ground state electron configuration for the element Na.
electron.boxes.to.fill (3K)
  1. Write out the electron configuration in the short-hand notation.
  2. Fill the boxes below with the arrow notation for electrons showing the correct ground state electron configuration for the element Si.
electron.boxes.to.fill (3K)
  1. Write out the electron configuration in the short-hand notation.
  2. Fill the boxes below with the arrow notation for electrons showing the correct ground state electron configuration for the element N.
electron.boxes.to.fill (3K)
  1. Write out the electron configuration in the short-hand notation.


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  1. Write out the shorthand notation for the electron configuration of B.
  2. Write out the shorthand notation for the electron configuration of Cl.
  3. Write out the shorthand notation for the electron configuration of F.
  4. Write out the shorthand notation for the electron configuration of Ca.
  5. Write out the shorthand notation for the electron configuration of Kr.
  6. Write out the shorthand notation for the electron configuration of O2-. Notice that this is an anion!
  7. Write out the shorthand notation for the electron configuration of Na+. Notice that this is a cation!
  8. Why are Groups 1 and 2 referred to as the s-block of the periodic table?
  9. Why are Groups 3 through 12 referred to as the d-block of the periodic table?
  10. Using what you now know about electron configurations explain the notion that elements in the same column in the periodic table have similar chemical and physical properties.
  11. Draw Lewis diagrams for all of the elements whose electron configuration you have figured out. Remember, the only electrons you draw dots for are the electrons in the valence shell. The valence shell is the shell of highest principal quantum number that contains electrons. As an example, Chlorine is done for you. Be careful with the ions!
lewis.dots.chlorine (1K) Instructors: Here is a link to a periodic table showing all of the s-block and p-block elements with their Lewis dots. I borrowed it from HyperPhysics.





  1. Look at the electron configurations for O2- and Na+ which you figured out above. Compare them to the other electron configurations you did. Do they have anything in common with any of them? Why?
  2. Look at the electron configurations of the noble gases that you have done. What do they have in common? Why do think that makes them more stable than the other elements?
Here is the homework assignment that goes with this group activity.
Last updated: Mar 20, 2007       Home