The Electron Configuration (Simplified) - Online Tutor, Practice Problems & Exam Prep

The periodic law dictates electron arrangements in elements, visualized through orbital diagrams. Degenerate orbitals, filled according to Hund's rule, illustrate that s, p, d, and f subshells can hold a maximum of 2, 6, 10, and 14 electrons, respectively. The Aufbau principle guides the electron configuration, filling lower energy orbitals first. Understanding paired and unpaired electrons is crucial, as paired electrons occupy orbitals with opposite spins, while unpaired electrons exist alone in an orbital. This knowledge is essential for predicting chemical behavior and reactivity.

The Electron Configuration (Simplified) Concept 1

As we stated earlier, the periodic law influences the electron arrangements of the elements, and the electron orbital diagrams are the visual representation within orbitals. Now we're going to say here we have what are called degenerate orbitals. These are electrons in the same set of orbitals having the same energy, and they're filled using Hund's rule. Now Hund's rule says that these degenerate orbitals are first half-filled before being totally filled. So if we take a look here, we have our s sublevel or s subshell. S can hold a maximum of 2 electrons. It has 1 orbital. Within that orbital, we have 2 electrons, one spins up, one spins down. So that would mean that the s sublevel has a maximum of 2 electrons.

For the p sublevel, we have 3 orbitals. Following Hund's rule, we would half-fill them first. So we go up, up, up, and each orbital we know can hold a maximum of 2 electrons, so we come back around, down, down, down. So the p sublevel holds a maximum of 6 electrons. For the d sublevel, we have 5 orbitals here. Hund's rule says we half-fill them first since they're all d set of orbitals, they all have similar energy. So then we come back around, down, down, down, down, down for a total of 10 electrons. And finally, the f sublevel has 7 of these orbitals, half-fill them again according to Hund's rule. So when we half-fill them according to Hund's rule, come back around to totally fill them in. When we do that we get a total of 14 electrons.

So just remember the periodic law influences the electron arrangement of elements, and it's these orbital diagrams that depict the visual representation of electrons within any given orbital, based on subshell level or subshell letter. So just keep that in mind, s can hold a maximum of 2 electrons, p can hold up to 6, d can hold up to 10, and f can hold up to 14.

The Electron Configuration (Simplified) Example 1

It says we need to properly fill in the orbitals of an atom that possesses 8 electrons within its d set of orbitals. Now we know that these are d set of orbitals because there are 5 of them. We have to fill in 8 electrons. Remember, since they're all the same set of orbital types being d, they all have the same energy and therefore are degenerate. We would half fill them based on Hund's rule. So we go up, up, up, up, up for our first 5. We need to fill in 8, so come back around. Down, down, down. So we stop there because again we only have 8 electrons to fill in, so here we don't need to completely fill in the last 2. They'll be left with just 1 electron in each. So this would be the way to properly fill in these d set of orbitals that have 8 total electrons.

Which electron configuration represents a violation of Hund's Rule? 

• A.
• B.
• C.
• D.

The Electron Configuration (Simplified) Concept 2

When writing the electron configuration of an element, it's important that it represents the distribution of electrons from s1 to s2 to p2 and so on within orbitals using the Aufbau principle. Now with the Aufbau principle, it says, starting with s1, we're going to have electrons filling lower energy orbitals before moving on to higher energy orbitals. We begin with s1 when we're doing the full ground state electron configuration of an element or an ion. To help us with this, you could take the Aufbau diagram approach. We start out with s1, then we move on, loop back around to s2, then we loop back around to p2, then to s3, then we loop back around to p3, and then s4, then continuously loop back around to d3, to p4, to s5. Now, this is our Aufbau diagram. In the Aufbau diagram, we have here s1 all the way down to s8. Then we have p2 to p7. Then we have d3 to d6, and then we have f4 and f5. Another way we can look at determining the electron configuration has to do more with the periodic table. So if you click on the next video, let's reimagine what the periodic table will look like when dealing with the electron configuration of elements and ions.

The Electron Configuration (Simplified) Concept 3

So here if we reimagine the periodic table, we'll see it in this depiction. Now realize here we have blue, we have yellow, we have purple, and we have red sectors. Now, each of these is called a block. The s block is what's in blue, this is the p block, the d block, and the f block. Remember the s sublevel has 1 orbital and that one orbital can hold a maximum of 2 electrons, which is why the s block is these 2 columns, to represent the 2 maximum electrons that the s sublevel can hold. The p sublevel has 3 orbitals, right, and each one again can hold a maximum of 2 electrons. So p theoretically can hold a maximum of 6 electrons. That's why the p block has 1, 2, 3, 4, 5, and 6 slots for it. D can have up to 10 electrons because it has 5 orbitals, and if you count you'll see in here there are 10 spots. And then for the f sublevels, they can hold up to 14 electrons, and if you were to count these rows in red, you'd see they come out to 14.

Now, realize here that in this periodic table, the first slot here which represents hydrogen starts off our electron configuration as s11. Then as we move to the next one we add another electron, so helium is s12. Then when we get to the 2nd row, since we're in the 2nd row, we now have 2. This is still the s block, so this is s12s21, and it continues onward and onward. Over here, we'd go s12s22p21. So following this pattern, this would be s3s4s5s6s7. And then here, this would be p3p4p5p6p7.

When we go to the d block, realize that it's gonna drop down by 1. So here this is s4, but then when we go to d block, it drops down by 1 number, so now it's d3d4d5d6. Now, notice how this number goes from 57 to 72, that's because 58 to 71 are here. Remember this red line here says that this entire red row exists between 57 and 72. And then we go between 89 and 104 because 90 to 103 is down here. These are your f blocks. They also drop down by one number. So this is f4f5. So basically, if you can reimagine the periodic table in this fashion, you can use it to figure out the electron configuration of any element or ion given to you. So we're gonna put this periodic table to use in order to do future electron configurations.

The Electron Configuration (Simplified) Example 2

To help us with this example question, I decided to leave this periodic table up so we can see how to best use it. Here it says to write the ground state electron configuration for the following element. Here we're dealing with fluorine, and they tell us that it has an atomic number of 9, which means it has 9 electrons. On this periodic table, we find fluorine right here. Ground state means that we're going to start out with 1s orbital and work our way up to fluorine. So we're going to count to fluorine. So we'd say 1s², 1s², 2s² because of 1, 2 slots. And then we have to count to f, so that would be 2p, we're in the 2p row. And how many slots do we have to count to get to fluorine? 1, 2, 3, 4, 5. 2p⁵. So here, this would be the ground state electron configuration of the fluorine atom. And realize here that these are the number of electrons. So when you add them up, 2 plus 2 plus 5, that gives me 9 electrons, which is related to the atomic number here of Fluorine. Remember, when an element is neutral, its atomic number tells us both the number of protons and the number of electrons. So again, rely on this depiction of the periodic table to help guide you to the right electron configuration of any element or ion given to you.

Which electron configuration represents a violation of the Auf Bau Principle? 

• A.
• B.
• C.
• D.

Identify the element with the given electron orbital diagram. 

• A. Silicon
• B. Fluorine
• C. Sulfur
• D. Chlorine
• E. Phosphorus

Write the electron configuration and electron orbital diagram for the following element:

Sulfur   (Z = 16)

• A.
• B.
• C.
• D.
• E.

Write the ground state electron configuration for the following element:

Magnesium   (Z = 12)  

• A.
• B.
• C.
• D.
• E.

The Electron Configuration (Simplified) Concept 4

Recall an orbital can hold a maximum of 2 electrons that pair up with opposite spins. One spins up and one spins down. Now with this, we have the ideas of unpaired and paired electrons. An unpaired electron is when an orbital contains 1 electron with its own spin. So an example of an unpaired electron would be here. This orbital has only 1 electron in it pointing up. A paired electron is when an orbital contains 2 electrons, each with its own spin, so one spins up and one spins down. So this would be a paired up orbital. But now, if we're looking at an entire electron orbital diagram, we can apply this idea of unpaired and paired electrons. The image to the left represents an unpaired electron orbital diagram because there is at least 1 orbital with an electron by itself. So as long as at least one of the orbitals has one electron, the whole thing is considered an unpaired electron orbital diagram. Now with paired, every single electron is paired up. Every single orbital we can see has a partner, and they have opposite spins. So, keep this in mind when you're doing electron orbital diagrams. Are any of the electrons not paired up? If so, you would say that's an unpaired electron orbital diagram that's being displayed.

The Electron Configuration (Simplified) Example 3

Here it says to determine the number of unpaired electrons in vanadium. So vanadium is V and it has an atomic number of 23. If we were to write out its electron configuration, we get 1s², 2s², 2p⁶, 3s², 3p⁶, 4s², 3d³. Now, Jules, do I have to look through every single one of these to find unpaired electrons? No. Just remember that s orbitals can hold a maximum of 2 electrons. And if you're hitting your maximum, that means all the electrons there are paired up. P, remember, can hold up to 6, so all those are reaching their maximum. D, though, can hold up to 10 electrons. So if it isn't hitting its maximum number of electrons, we got to check it out. So here we're gonna look at 3d, and remember d has 5 sets of orbitals. Following Hund’s rule, we'd half fill them because they have the same energy and there's 3 of them. Right? So we go up, up, up. There's no more electrons we can add because there's only 3 electrons here in 3d. And if we look, all 3 of them are unpaired, all of them are by themselves. That means that the answer here would have to be option D. There are 3 unpaired electrons within the vanadium element.

Which of the following atoms has no unpaired electrons?

• A. Ca
• B. N
• C. C
• D. F

Which of the following atoms has the most unpaired electrons?

• A. B
• B. Si
• C. P
• D. O
• E. Cl

Here’s what students ask on this topic:

What is the Aufbau principle in electron configuration?

The Aufbau principle states that electrons fill lower energy orbitals before moving to higher energy orbitals. This principle helps in determining the electron configuration of an element. Starting with the 1s orbital, electrons fill in the order of increasing energy levels: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on. This systematic filling ensures that the atom is in its lowest energy state, known as the ground state. The Aufbau diagram is often used to visualize this process, showing the sequence in which orbitals are filled.

How does Hund's rule apply to electron configurations?

Hund's rule states that electrons will fill degenerate orbitals (orbitals with the same energy) singly before pairing up. This means that for orbitals within the same subshell, one electron will enter each orbital with parallel spins before any orbital gets a second electron with an opposite spin. For example, in the p subshell with three orbitals, electrons will fill as ↑, ↑, ↑ before pairing as ↑↓, ↑, ↑. This minimizes electron-electron repulsions and stabilizes the atom.

What is the significance of paired and unpaired electrons in an orbital diagram?

Paired electrons in an orbital diagram have opposite spins and occupy the same orbital, which stabilizes the atom. Unpaired electrons, on the other hand, are single electrons in an orbital and can make the atom more reactive. The presence of unpaired electrons is crucial in determining the magnetic properties and chemical reactivity of an element. For instance, elements with unpaired electrons are often paramagnetic, while those with all paired electrons are diamagnetic.

How do you determine the electron configuration of an element using the periodic table?

The periodic table can be divided into blocks (s, p, d, f) that correspond to the electron subshells. The s block includes Groups 1 and 2, the p block includes Groups 13-18, the d block includes the transition metals, and the f block includes the lanthanides and actinides. To determine the electron configuration, start from hydrogen and move across the table, filling electrons according to the Aufbau principle. For example, carbon (atomic number 6) has the configuration 1s² 2s² 2p².

What are degenerate orbitals and how are they filled?

Degenerate orbitals are orbitals within the same subshell that have the same energy level. According to Hund's rule, these orbitals are filled singly with parallel spins before any orbital gets a second electron. For example, in the p subshell, which has three degenerate orbitals, electrons will fill as ↑, ↑, ↑ before pairing as ↑↓, ↑, ↑. This minimizes electron repulsion and stabilizes the atom.