1.2 ELECTRONIC STRUCTURE AND IONISATION ENERGIES

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Hi again! My name is Aneta and today we are going to learn about orbitals and subshells and ionisation energies. To start off, you have probably learned at school that each atom has shells and you can hold 8 electrons in each shell (apart from the first one where you only have two). However, for A-levels you need to know that this is actually more complicated. Ok, let me explain.

Each shell has subshells. These are called subshell s,p and d. There is also a subshell f but you don't need to worry about it. On the image above you can see how they look. Subshell s is a sphere and subshell p looks like a dumbell (There are three different p orbitals that's what the x,y and z represent). In your exam you may be asked to draw these subshells however you don't need to worry about drawinng subshell d).

As you probably realised I said something about orbitals. In each subshell there is a certain amount of atomic orbitals that can be present. For example in subshell s there is one orbital and in subshell p there are three orbitals and in subshell d there are 5 orbitals. You need to know that in each orbital there can be up to 2 electrons. Therefore to sum this up:

s subshell -> 1 atomic orbital -> max. 2 electrons

p subshell -> 3 atomic orbitals -> max. 6 electrons

d subshell -> 5 atomic orbitals -> max. 10 electrons

Atomic orbital- region in an atom that can hold up to two electrons with opposite spins.

Now, we need to know how we represent the electronic configuration (how electrons are arranged in atoms). There are actually two ways but lets start with the drawing one. So, you have one orbital in s subshell therefore we only draw one box. ( Each box represents one atomic orbital). Inside that box we have two arrows one going up and one down. This is because in each orbital there is a max of 2 electrons with opposite spins ( one up and one down). Also, remember that you only put one electron in each box of a given subshell first before filling the box fully with two electrons.

So, you probably think: I know there are three subshells and I know how many orbitals and electrons there should be in each BUT How do I know which subshell to fill in first? How are they arranged in an atom? Let me come with an explanation. Electrons fill the atomic orbitals in order of increasing energy. I got you a chart in the above image. Of course chemists came up with a way of representing the electrons in subshells and this is what I drawn at the top of the image above. The big number in front represents the atomic orbital, the letter shows what subshell it is and the small number shows how many elecrons are present in each subshell.

The reason why we fill in the 4s box first before the 3d is because 4s has a lower energy level.

As always what a better way to explain something than give an example.

Staring with the easy examples:

1. Carbon - this example show what I told you before about putting one electron in each box for a given subshell first. You can see there are two single electrons in the 2p subshell. This is a rule that you need to follow.

2. Aluminium - again whenever you get a question like this firstly draw the boxes for yourself, then go and see on the periodic table how many electrons that atom has and simply just start filling the boxes.

Now, we have the medium difficulty examples:

1. Potassium- do everything I told you in the prevoius example, draw boxes, check how many electrons an atom has and fill the boxes. However, this time you need to apply the rule I explained you before, the 4s comes in first before the 3d.

2. Titanium- after you fill in the 4s then you go back to the 3d and put one electron in each box first before adding the second one.

Lastly, you have the difficult examples. I called them difficult because they are the exceptions and you will see why now:

1. Chromium- remember I told you to fill in the 4s before 3d, yes you kind of do this here too BUT 3d is much more stable (and happier) if you fill it half way so 4s donates one electron to it.

2. Manganese- this time you basically just fill in the boxes and start with 4s before 3d.

3. Copper- again 3d is much more happier with a full shell therefore 4s donates one electron for it to be stable.

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Moving on to the other part of this topic which is ionisation energies. How about some definitions for a start.

Ionisation energy- all atoms have electrons and you need energy to remove an electron. Ionisation energy is that energy needed to remove an electron from an atom. After you remove an electron from an atom this atom is called an ion (because it has been ionised)

First molar ionisation energy- it mentions molar therefore we are going to remove one mole of electrons from one mole of its GASEOUS atom. It has to be in gaseous state.

You can wirte equations to represent the ionisation energies and some tips from me that I found helpful:

1. See which ionisation energy they want you to write (1st, 2nd, 3rd, etc.)

2. Start with the right hand side. Whatever the ionisation energy it is you have to write that number with a plus on the right hand side next to the element symbol. (we put a plus because every time an electron is removed the atom becomes a positve ion).

3. Always add an electron on the right hand side.

4. Balance the equation .

5. Put gaseous state next to the elements.

Electrons are held very closely to the nucleus due to it is being positive (protons and neutrons in the nucleus and only protons have a positive charge as neutrons have no charge). The greater the attraction from the nucleus the higher the ionisation energy and the harder it is to remove an electron. However, you need to know that this attraction and ionsation energy depends on three main factors:

1. Size- this means the size of an atom. A bigger atom means that there are more protons in the nucleus and therefore a bigger attraction for the electrons. Concluding in higher ionisation energy.

2. Distance- this is the distance of the electrons from the nucleus. As you probably can imagine the attraction will be bigger for the electrons that are closer to the nucleus. If the elctrons are far away from the nucleus the ionisation energy will be lower.

3. Shielding- this is the repulsion between electrons in different shells. The electrons from inner shell will repel the electrons from outer shell.

And now using all of the information that we just learned about, we are going to look at two evidence that subshells exist.

This is the first one. It is a graph with first ionisation energies for the first 20 elements on the periodic table. As you can see, across period the I.E. increases. This is because there is more attraction from the nucleus as there are more protons. Across the group on the other hand the ionisation energy decreases and this is because there are more shells present and shielding happens.

In the exam they usually ask you to compare the ionisation energy of two atoms and I prepared the 5 most common examples for you. To do this you would use the periodic table.

1. B is shielded by the 2s electrons, therefore Be has a higher first ionisation energy.

2. Oxygen has one paired electron in the 2p subshell and therefore repulsion happens. Meaning Nitrogen has a higher first ionisation energy.

3. Lithium has more shielding than Helium and therefore Helium has a higher first ionisation energy.

4. Hydrogen has only one proton while Helium has two, so more attraction from the nuclueus. Meaning He has a higher first I.E.

5. This is shielding again. Neon has one more shell than Helium and therefore has extra shielding, so Helium has a higher first ionisation energy.

Now, the evidence for subshells number 2:

This time we only have one element that we are going to explore. It is a graph of successive ionisation energies for Sodium. Clearly you can see from the graph that the line goes up and this is because the energy needed to remove an electron increases. As you remove electrons there are less and less of them in the atom meaning the shielding is lower and the attraction is greater and I.E. increases.

This graph is an evidence for subshells as the energy needed to remove the 10th electron is much higher than the energy to remove the 2nd electron. This suggests that the atom has shells. Otherwise, the energy to remove every electron would be the same.

And this is all for this topic. I hope you understood it and now I invite you to my youtube channel where you can practice some past paper questions.

PS. Please remember, I am only a student, and as anyone, I can make mistakes. If you think you can see one, don't hesitate and comment (either here on on my youtube channel) Thank you!