# Chemistry AS Level - Ionisation energies

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#1
The question says:
Electron configurations for atoms of different elements are shown below.
Which electron configuration represents the element with the largest first ionisation energy?
A 1s2 2s2
B 1s2 2s2 2p4
C 1s2 2s2 2p6
D 1s2 2s2 2p6 3s2

The answer is C but I don’t understand how?? I thought it would be A because the electrons are closer to the nucleus so more energy would be required to remove them.
0
1 month ago
#2
(Original post by Pringlewingle)
The question says:
Electron configurations for atoms of different elements are shown below.
Which electron configuration represents the element with the largest first ionisation energy?
A 1s2 2s2
B 1s2 2s2 2p4
C 1s2 2s2 2p6
D 1s2 2s2 2p6 3s2

The answer is C but I don’t understand how?? I thought it would be A because the electrons are closer to the nucleus so more energy would be required to remove them.
A is Be
B is Oxygen
C is Neon
D is calcium
C is the right answer because we know 1st IE increases across a group due to increased nuclear charge, and Neon is the element with the highest charge.
0
1 month ago
#3
A = Beryllium (Be) / B = Oxygen (O) / C = Neon (Ne) / D = Magnesium (Mg)

Yes, the trend in first ionisation energy increases across a period.
The atomic radius (ie the distance from the outermost electrons and the nucleus) decreases (across period 2) because the nuclear charge increases, due to the increase in protons. This increase in protons will make the atom's nucleus more positive, increasing the attraction for electrons. This means that from A to B to C, the attraction of electrons increases. (This increased attraction means that more energy will be required to remove the first electron, thus increasing the first IE).

Have you learnt the graph of first IEs across a period? Just because I think sometimes it helps if you know that pattern.

I don't know if I explained that very well, but I hope it helps!
1
#4
(Original post by simxne_)
A = Beryllium (Be) / B = Oxygen (O) / C = Neon (Ne) / D = Magnesium (Mg)

Yes, the trend in first ionisation energy increases across a period.
The atomic radius (ie the distance from the outermost electrons and the nucleus) decreases (across period 2) because the nuclear charge increases, due to the increase in protons. This increase in protons will make the atom's nucleus more positive, increasing the attraction for electrons. This means that from A to B to C, the attraction of electrons increases. (This increased attraction means that more energy will be required to remove the first electron, thus increasing the first IE).

Have you learnt the graph of first IEs across a period? Just because I think sometimes it helps if you know that pattern.

I don't know if I explained that very well, but I hope it helps!
Wow thanks so much!! Makes so much sense
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