electron configurations

 

Topics discussed:

·        Electron Clouds

·        The Schrodinger model of the atom (versus the Bohr Model)

·        How to write electron configurations, including what order to write them in

·        Why we write electron configurations and what we actually do with them

 

Student:    Hey, professor, I think I have come across the most confusing subject I’ve ever seen in chemistry thus far.

Professor:  What subject is that?

Student:    Electron configurations.

Professor:  Can you tell me more about what you find so confusing about them?

Student:    Well, yeah, everything basically.  I think we should start from the beginning.  I haven’t the faintest idea what they are, how to write them, or what they have to do with anything, or what you do with them.

Professor:  Well, tell me what you know about atoms and electrons.

Student:    I know that an atom has a proton and a neutron in the middle, and the electrons orbit around the nucleus like little planets around a sun.

Professor:  Mmm, not quite.

Student:    What do you mean not quite?  That’s what every science book I’ve read tells me!

Professor:  I agree with the part about the proton and the neutron being in the middle (the nucleus).  And I agree that the electrons are out there away from the nucleus, flying around in a sense.  But I disagree with the part about the electrons being like little planets.  That doesn’t quite jive with scientists current understanding of the electrons.

Student:    Oh, I guess some new discovery has been made about atoms since my textbook was written.

Professor:  Actually, yes and no.  Lots of new information has been learned about atoms in almost every year in the 20^th and 21^st century, but physicists and chemists knew as early as the 1930’s and 1940’s that the electrons didn’t orbit around the nucleus just like little planets.

Student:    They have?  Why do science books and shows still portray atoms this way?

Professor:  The “electrons-are-like-planets-around-a-sun” is called the “Bohr” model of the atom and it’s easy to understand (albeit a little incorrect), so the image has stuck in the mind of the public.

Student:    Well, if the electrons don’t orbit around the nucleus like little planets, what do they do?

Professor:  Ever since the 1930’s and 1940’s scientists have believed in the Schrodinger model of the atom.  The Schrodinger model of the atom says that electrons can go wherever they want, but they are most likely to be found in regions called “electron clouds” that hover around outside of the nucleus.

Student:    Hmm, I think I’ve heard of electron clouds at some point, but that’s about it.  Can you tell me more about them?

Professor:  In the 1930’s, a guy names Erwin Schrodinger discovered that electrons do float around outside the nucleus, but that they do not follow well defined orbits like planets around a sun.  What he also discovered is that we can’t really predict exactly where we will find the electrons – we can only give probabilities of where we will find them.

Student:    Oh, yeah, sounds like the weather man.  He can’t tell if when and where it’s going to rain – he can only tell you the probability that it will rain in a particular area.

Professor:  Exactly!  The regions where electrons are likely to be found are called “electron clouds.”

Student:    How did Schrodinger figure out all this?

Professor:  That’s beyond the scope of first semester chemistry.  It’s something that you will study when and if you ever take a course in Quantum Mechanics.

Student:    Oh, okay – so basically I should stop asking how we know about electron clouds until I get to that level.

Professor:  Well, I’ll put it this way – it’s difficult to discuss how we know about electron clouds until you’ve had three semesters of calculus and two semesters of differential equations.

Student:    Heh, yeah I’ll try to hold off on my curiosity then!

Professor:  So anyways, we have names for these electron clouds, and a few numbers that go with them. 

                The first electron cloud is called 1s.  Two electrons can live in it at a time, and it’s shaped like a sphere around the nucleus.  The second electron cloud is called 2s.  It’s also shaped like a sphere (albeit a larger one) around the nucleus and two electrons can live in it as well.

 

Student:    I suppose the third electron cloud is called 3s?!

Professor:  No, the third electron cloud is called 2p, and six electrons can live there at a time.  The 2p cloud has a bit of a lobe shape.  There should be a picture of it in your chemistry book.

Student:    Why is there a jump in logic like that?  If there’s a 2p, shouldn’t there be a 1p as well?  And why all of a sudden can six electrons live in one cloud?

Professor:  The reasons for all that have to do with more of the Schrodinger stuff that you’ll study when you take Quantum Mechanics – for the time being it’s best if you just accept it.

Student:    Oh, okay.

Professor:  So far we have 1s, 2s, and 2p.  The s electron clouds can hold two at a time, and the p clouds can hold up to 6 at a time.

Student:    Okay, so what’s after 2p?

Professor:  The next electron cloud is called 3s.

Student:    Um, I’m getting a bit confused.  There doesn’t seem to be any logic to the order of these electron clouds?  Do they even really go in order?

Professor:  The answer to your second question is yes, they go in order.  Now, with regard to your first question, it might be easier if I draw the following table for you:

 

                               

                               

Student:    What’s up with the arrows and the little superscripts?

Professor:  The arrow tells you what order the electron clouds go in.  Although there are exceptions to the arrow rules, they should suit you just fine for general chemistry.  The subscripts tell you the maximum number of electrons that can live in each cloud.

Student:    I know that there’s a long list of questions that I’m not supposed to ask because the reasons are too complicated to explain to someone who’s never had differential equations or quantum mechanics, but is there any rhyme or reason at all to the order of these things?

Professor:  Schrodinger determined the energy of the electrons in each cloud.  The electrons fill up clouds in order of ascending energy.  This is called the aufbau principle.  The 1s cloud has the lowest energy, and fills up first.  The 2s cloud has the next highest energy and fills up second, etc.

Student:    So what about the fact that certain electron clouds can hold more electrons than others.  Why is that?

Professor:  It’s more Schrodinger stuff.

Student:    How did I guess?! 

Professor:  So let’s do an example.  Let’s write out the electron configuration of Phosphorus, element number 15!

Student:    Well, I’m sure you would start with 1s, I suppose.

Professor:  Agreed.  Now how many electrons can the 1s cloud hold?

Student:    I think you said 2 earlier.

Professor:  Right.  So the first thing you would write is 1s².  So according to the table I gave you earlier, which electron cloud fills up next?

Student:    The 2s cloud fills up next, and it can hold 2 electrons, just like all the s clouds.

Professor:  Exactly!  So far we have 1s²2s².  What’s next?

Student:    According to the chart above, the next cloud to fill up is 2p, and it can hold 6 electrons.  So far we have 1s²2s²2p⁶.

Professor:  Excellent.  So how many electrons have you used up so far?

Student:    What do you mean?

Professor:  Well, you have two electrons in the 1s cloud, you have two electrons in the 2s cloud, and 6 in the 2p cloud.  What is the total number of electrons that you have used?

Student:    Oh, 2 + 2 + 6 = 10.  I have used ten electrons so far.

Professor:  So how many electrons total do you have to play with?

Student:    The atomic number of phosphorus is 15, so I guess I have 15 to play with.  I’ve used 10 so far, so I have 5 left.

Professor:  Agreed.  Now where do those remaining 5 electrons go?

Student:    According to the chart, the next orbital to fill up is 3s, which holds two electrons.  So far I have 1s²2s²2p⁶3s².  I just used two more so I have 3 left to play with.

Professor:  Right.  Now where do the remaining 3 electrons go?

Student:    Well, according to the chart they go in the 3p cloud.  But the 3p cloud holds 6 electrons, and I only have three left.  What do I do?

Professor:  Go ahead and stick the remaining three in the 3p cloud by writing 3p³.

Student:    But the 3p cloud holds 6, not 3!

Professor:  It would be more correct to say that the 3p cloud can hold up to 6.  That’s not to say that it has to hold 6.  It can hold any number of electrons up to 6, inclusive.

Student:    Okay, so I guess the overall electron configuration of phosphorus would be 1s² 2s² 2p⁶ 3s² 3p³.

Professor:  Right!  That’s it!  You’re done!  You’ve taken all of the electrons of phosphorus and shown what electron clouds they will fill up.

Student:    That’s nice and all, but what do you actually do with an electron configuration?  I mean, what did we just accomplish?

Professor:  Well, let me ask you this: how many atoms can bond to phosphorus?

Student:    As many as it wants, I suppose.

Professor:  Actually, no.  Under ordinary circumstances, phosphorus can form three bonds – because it is three electrons shy of having a full orbital.

Student:    Oh?!

Professor:  Yes.  The purpose of drawing out electron configurations is to see how many more electrons are needed to give the atom a full outer orbital.  The electron configuration will also tell you how many valence electrons the atom has.

Student:    So how do we tell how many valence electrons phosphorus has from looking at its electron configuration?

Professor:  Well, the highest cloud number of the electrons in phosphorus is 3.  And in the number three cloud there are s electrons and p electrons, totaling 2 and 3 respectively – so phosphorus has 2+3=5 valence electrons.

Student:    Oh so electron configurations are used to determine the number of valence electrons that an atom has.

Professor:  Right.  There are other uses, but that is the main one that you will come across in general chemistry.  Also, there are situations (in advanced chemistry, mostly), where we need to know what cloud or clouds the valence electrons are in.  Sometimes, that makes a difference in how the atom reacts.  They’re also used in assigning formal charges, which you’ll learn about later in the semester.

Student:    Oh, okay!

Professor:  A couple of last things.  From now on, don’t call them “clouds”, call them “orbitals”.  Chemists usually call them orbitals.  And second, the level one orbitals can hold up to 2*1² = 2 electrons, the level two orbitals (as in 2s and 2p) can altogether hold 2*2²=8, the level 3 orbitals (as in 3s, 3p, and 3d) can altogether hold a total of 2*3² = 18 electrons, and so forth.  In general an n level orbital can hold a total of 2n² electrons, where n is the number of the orbital level.  Also, the letters s, p, d, and so on are referred to as subshells.

                By the way, just for your reference, you can look up the electron configuration of just about any element on the table by looking online.  Just remember that when you get into the elements that are really big, sometimes there are exceptions to the arrow chart I gave you above.  And another thing you may come across is:

                [Xe] 6s² 4f⁵.

                This means that the electron configuration is the same as for Xenon (Xe), and then add a 6s and a 4f cloud, to get the electron configuration.  In this case, the electron configuration is for Promethium.

Student:    Wow.  Thanks professor.  The electron configurations make much more sense now.

Professor:  I’m always glad to help.