GCSE Transition metal electronic structure Watch

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Lauraa96
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Hey,

I've just done a practice paper for unit 3 chemistry and a question came up asking about the electronic structure of the transition elements. I can't find the topic in my revision guide and I was wondering if anyone could explain it to me or tell me a good website that can explain it.
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thegodofgod
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(Original post by Lauraa96)
Hey,

I've just done a practice paper for unit 3 chemistry and a question came up asking about the electronic structure of the transition elements. I can't find the topic in my revision guide and I was wondering if anyone could explain it to me or tell me a good website that can explain it.
For GCSE, don't you only have to know the electronic configuration of elements up until Z = 20, i.e. Calcium?

After Calcium, the electronic configuration (especially for transition metals) becomes confusing
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Lauraa96
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This was the question:

5 (c) (i) Transition elements have similar properties
Explain why in terms of electronic structure

5 (c) (ii) There are no transition elements between the group 2 element magnesium and the group 3 element aluminium.
Explain why in terms of electronic structure
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Booyah
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(Original post by Lauraa96)
This was the question:

5 (c) (i) Transition elements have similar properties
Explain why in terms of electronic structure

5 (c) (ii) There are no transition elements between the group 2 element magnesium and the group 3 element aluminium.
Explain why in terms of electronic structure

At the beginning of A-level the first thing we were taught was about sub shells, each shell can be divided up, for example He had a full S orbital written as 1S^2 kind of like that...

The proof for sub shells comes in first ionisation energy between across a group, as you know it generally increases but with some ionisations there is a large drop (Between groups 2 and 3, and between groups 5 and 6).

The periods in the periodic table can help you tell where an elements sub shells are filled, I suggest you label the first 2 columns S, the transition metals D and the last 6 columns P.

S sub shells hold 2 electrons, P sub shells hold 6 and transition metals contain 10.

Thats the basics now...

i) Transition metals have the same properties due to their outermost electrons being in the D block (As far as I am aware , I haven't learnt this yet it is just an educated guess... This is A2)

ii) There is no D block present between Mg and Al...

Hope that helped you... More important you know the background and then you can answer... As I said I haven't actually learnt this yet, I really just gave you an insight to subshells rather then answering your question to be honest... But I guess it relates into your question.
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Lauraa96
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Thank you so much
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qiaoyu.he
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(Original post by Lauraa96)
This was the question:

5 (c) (i) Transition elements have similar properties
Explain why in terms of electronic structure

5 (c) (ii) There are no transition elements between the group 2 element magnesium and the group 3 element aluminium.
Explain why in terms of electronic structure
As far as I know, you only need to know the electronic configurations up to Ca, i.e. 20. Interesting things happen after element 20.

First, you need to forget the 'stable octet' rules that you were taught. It is perfectly possible for shells to have more than 8 electrons in their shells (I got very confused when I was told this in A-Level Chemistry). The maximum number of electrons that can be accommodated in any shell is 2n^2 (i.e. twice a perfect square) where n is the 'principle quantum number' (basically what shell it is). So shell 1 can accommodate 2*1^2 = 2 electrons; shell 2 can accommodate 2*2^2 = 8 electrons; shell 3 can accommodate 2*3^2 = 18 electrons; shell 4 can accommodate 2*4^2 = 32 electrons etc.

Second, each shell is also divided up into sub-shells, called atomic orbitals, which are given letters to represent them. Don't worry about the letters used, they are just names. They are called s, p, d, f, g, h, i, j and are characterized by their different shapes. Each sub-shell can accommodate exactly 2 electrons. Shell 1 has 1 s orbital; shell 2 has 1 s orbital and 3 p orbitals; shell 3 has 1 s orbital, 3 p orbitals and 5 d orbitals; shell 4 has 1 s orbital, 3p orbitals, 5 d orbitals and 7 f orbitals etc. (you get the pattern). You can see that shell 1 has 1 atomic orbital, which can accommodate 2 electrons, so shell 1 can hold 1*2 electrons in total. Shell 2 has 4 orbitals in total (1 s and 3 p), each of which can accommodate 2 electrons, so shell 2 can hold 4*2 = 8 electrons. Shell 3 has 9 orbitals in total (1 s, 3 p, and 5d), which can accommodate 9*2 = 18 electrons in total. This follows the 2n^2 rule I mentioned previously.

Shells are filled with electrons in the following order (with a few exceptions). To find any element's electronic structure, just add electrons to each orbital, down the group until that orbital is filled and then move to the next one:

Shell 1 s orbitals (2 electrons)
Shell 2 s orbitals (2 electrons)
Shell 2 p orbitals (6 electrons as there are 3 p orbitals)
Shell 3 s orbitals (2 electrons)
Shell 3 p orbitals (6 electrons)
Shell 4 s orbitals (2 electrons)
Shell 3 d orbitals (10 electrons as there are 5 d orbitals)
Shell 4 p orbitals (6 electrons)
Shell 5 s orbitals (2 electrons)
Shell 4 d orbitals (10 electrons)

There is an order to the filling, but its a bit complicated to explain. For example, carbon has 6 electrons, so it will have 2 electrons in the shell 1 s orbital, 2 electron in the shell 2 s orbital and 2 electrons in the shell 2 p orbital. Sulfur, having 16 electrons would have 2 electrons in the shell 1 s orbital, 2 electrons in the shell 2 s orbital, 6 electrons in the shell 2 p orbitals, 2 electrons in the shell 3 s orbitals and 4 electrons in the shell 3 p orbitals.

Thus you can see that Calcium, having 20 electrons will fill up to the 4 s orbital, and if you just look at shells, will have electronic structure 2,8,8,2 - exactly what you learnt. But Scandium, the first transition element has 21 electrons and the extra electron would enter the 3rd shell, in the d orbitals, giving it electronic structure 2,8,9,2.

If you continue along the periodic table, titanium will have 2,8,10,2 etc. To answer your first question, there are 10 transition metals in the 3rd period and they account for the filling of the 3rd shell d orbitals (which can accommodate 10 electrons). As they all have electrons in the 3 d shell, they all have similar properties. Interestingly, the period 4 transition metals have electrons in the shell 4 d orbitals (try to work it out and see). To answer your second question, there are no d orbitals in the second shell so there can't be transition elements in the second period.

I hope this helps (don't worry if you don't understand, I don't think transition metal chemistry should come up at GCSE). You will learn more detail about all this if you take A level chemistry. Quote me if you have any questions.
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usycool1
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(Original post by qiaoyu.he)
As far as I know, you only need to know the electronic configurations up to Ca, i.e. 20. Interesting things happen after element 20.

First, you need to forget the 'stable octet' rules that you were taught. It is perfectly possible for shells to have more than 8 electrons in their shells (I got very confused when I was told this in A-Level Chemistry). The maximum number of electrons that can be accommodated in any shell is 2n^2 (i.e. twice a perfect square) where n is the 'principle quantum number' (basically what shell it is). So shell 1 can accommodate 2*1^2 = 2 electrons; shell 2 can accommodate 2*2^2 = 8 electrons; shell 3 can accommodate 2*3^2 = 18 electrons; shell 4 can accommodate 2*4^2 = 32 electrons etc.

Second, each shell is also divided up into sub-shells, called atomic orbitals, which are given letters to represent them. Don't worry about the letters used, they are just names. They are called s, p, d, f, g, h, i, j and are characterized by their different shapes. Each sub-shell can accommodate exactly 2 electrons. Shell 1 has 1 s orbital; shell 2 has 1 s orbital and 3 p orbitals; shell 3 has 1 s orbital, 3 p orbitals and 5 d orbitals; shell 4 has 1 s orbital, 3p orbitals, 5 d orbitals and 7 f orbitals etc. (you get the pattern). You can see that shell 1 has 1 atomic orbital, which can accommodate 2 electrons, so shell 1 can hold 1*2 electrons in total. Shell 2 has 4 orbitals in total (1 s and 3 p), each of which can accommodate 2 electrons, so shell 2 can hold 4*2 = 8 electrons. Shell 3 has 9 orbitals in total (1 s, 3 p, and 5d), which can accommodate 9*2 = 18 electrons in total. This follows the 2n^2 rule I mentioned previously.

Shells are filled with electrons in the following order (with a few exceptions). To find any element's electronic structure, just add electrons to each orbital, down the group until that orbital is filled and then move to the next one:

Shell 1 s orbitals (2 electrons)
Shell 2 s orbitals (2 electrons)
Shell 2 p orbitals (6 electrons as there are 3 p orbitals)
Shell 3 s orbitals (2 electrons)
Shell 3 p orbitals (6 electrons)
Shell 4 s orbitals (2 electrons)
Shell 3 d orbitals (10 electrons as there are 5 d orbitals)
Shell 4 p orbitals (6 electrons)
Shell 5 s orbitals (2 electrons)
Shell 4 d orbitals (10 electrons)

There is an order to the filling, but its a bit complicated to explain. For example, carbon has 6 electrons, so it will have 2 electrons in the shell 1 s orbital, 2 electron in the shell 2 s orbital and 2 electrons in the shell 2 p orbital. Sulfur, having 16 electrons would have 2 electrons in the shell 1 s orbital, 2 electrons in the shell 2 s orbital, 6 electrons in the shell 2 p orbitals, 2 electrons in the shell 3 s orbitals and 4 electrons in the shell 3 p orbitals.

Thus you can see that Calcium, having 20 electrons will fill up to the 4 s orbital, and if you just look at shells, will have electronic structure 2,8,8,2 - exactly what you learnt. But Scandium, the first transition element has 21 electrons and the extra electron would enter the 3rd shell, in the d orbitals, giving it electronic structure 2,8,9,2.

If you continue along the periodic table, titanium will have 2,8,10,2 etc. To answer your first question, there are 10 transition metals in the 3rd period and they account for the filling of the 3rd shell d orbitals (which can accommodate 10 electrons). As they all have electrons in the 3 d shell, they all have similar properties. Interestingly, the period 4 transition metals have electrons in the shell 4 d orbitals (try to work it out and see). To answer your second question, there are no d orbitals in the second shell so there can't be transition elements in the second period.

I hope this helps (don't worry if you don't understand, I don't think transition metal chemistry should come up at GCSE). You will learn more detail about all this if you take A level chemistry. Quote me if you have any questions.
I'm not the OP, but is this why Newland's law of octaves only worked for the first 16 or so elements?
thegodofgod
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(Original post by usycool1)
I'm not the OP, but is this why Newland's law of octaves only worked for the first 16 or so elements?
:yes:

After about the 4s orbital, anything can happen (e.g. sulfur has 12 electrons instead of 8 in sulfuric acid). This is because all of the orbitals are so close in energy that the electrons can 'jump' from one to another. This explains why the 3d sub-shell fills after the 4s orbital (usually ).

As the principle quantum number increases, the sub-shells begin to have more and more overlapping energies, so the electrons 'jump' around more easily, confusing the electronic configuration of some ions / elements.

From the diagram below, you can see that there's a (relatively) massive gap between the energies of the 1s and 2s orbitals, whereas afterwards, the sub-shells become closer together in terms of energy differences.

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Booyah
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(Original post by thegodofgod)
:yes:

After about the 4s orbital, anything can happen (e.g. sulfur has 12 electrons instead of 8 in sulfuric acid). This is because all of the orbitals are so close in energy that the electrons can 'jump' from one to another. This explains why the 3d sub-shell fills after the 4s orbital (usually ).

As the principle quantum number increases, the sub-shells begin to have more and more overlapping energies, so the electrons 'jump' around more easily, confusing the electronic configuration of some ions / elements.

From the diagram below, you can see that there's a (relatively) massive gap between the energies of the 1s and 2s orbitals, whereas afterwards, the sub-shells become closer together in terms of energy differences.

you have mind ****ed me harder then my Chemistry teacher... Jeez you make it look so complicated!
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Mcrfan
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I am actually doing this question now and I'm in year 9 (13 year old)
Soooo can anyone explain it in English please?

Q- there are no transition elements between group 2 element magnesium and the group 3 element aluminium. Explain why in terms of electronic structure.
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AB12345678
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(Original post by Mcrfan)
I am actually doing this question now and I'm in year 9 (13 year old)
Soooo can anyone explain it in English please?

Q- there are no transition elements between group 2 element magnesium and the group 3 element aluminium. Explain why in terms of electronic structure.
Hi!

I would guess it's because the definition of a transition element is one whose atoms have a "partially filled d sub-shell" and, across Period 3, there are insufficient electron shells to achieve this. Or something like that. How much detail does the question want?
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r-sharpox
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Hi, I did exactly the same paper yesterday and I'm in year 9... Do you know where to find the paper online?
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Stevenson Dude
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Yegh guys that's too complicated :/
I made a video about it here https://www.youtube.com/watch?v=ahyRGN9Ke1k
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1999266
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Idea of same number of electrons in outer shell therefore similar reactivity/properties. Inner shells incomplete (this is the GCSE level of understanding required in 2016)

Next question - Aluminium follows magnesium in proton number. You couldn't fit an element between 2 consecutive proton numbers without halving a proton.
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