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Ionisation Energies HELP PLEASE!! AQA AS LEVEL

Hi guys

I'm finding it hard to understand ionisation energies, I understand the equations but i do not understand the trends across groups 1-6?
Original post by AlphaMolecule
There are four main factors which affect the amount of energy needed to ionise an atom:

•

Proton number - the more protons, the greater the effective nuclear charge, meaning that it requires more energy to displace an electron by overcoming the forces of electrostatic attraction between the positively-charged nucleus and the electrons

•

Electron pairing - if two electrons are paired in an orbital, the fact that their charges are like results in them repelling each other which contributes to the energy required to displace the electron and ionise the atom - meaning that the ionisation energy required to displace the electron is lower

•

Electron shielding shells - if there are more electron shells between the highest energy level electron and the atomic nucleus, the shells exert a dampening effect on the size of the force of electrostatic attraction - so the ionisation energy decreases as the effective nuclear charge is weaker

•

Orbitals - if the electron is located in an orbital further away from the nucleus (ie a higher energy orbital), the force of electrostatic attraction experienced by the electron is lower and the electron subsequently has more kinetic energy. Therefore it requires less energy to displace it from the atomic radius



Across a period, these three factors vary. The proton number always increases, however electron shielding and pairing doesn't. You need to write out the electronic configuration for atoms in a period to assess whether or not the ionisation energy increases or decreases.

So take note of when orbitals and shells change, while acknowledging the fact that the proton number increases constantly - however the orbitals/electron shielding factors are more effective than an increase in nuclear charge due to more protons.


Thank you can you help me with this please?

understand how ionisation energies in Period 3 (Na – Ar) and in Group 2 (Be – Ba) give evidence for electron arrangement insub-levels and in levels.
Crossing the period the number of protons in the nucleus increases but the electrons are being added to the same quantum shell and so there is no increase in shielding this results in a greater attraction between the outermost electrons and the nucleus and so more energy is needed to remove the outermost electron and there is a general increase in 1st ionisation energy.
However for period 3 Al is less than Mg and S is less than P. This is because the outermost electron in Al is in a different subshell (P subshell) compared to Mg (S subshell) which is at a higher energy level and experiences more shielding which overcomes the extra attractive force of the additional proton so it needs less energy to remove it. This is evidence for the S and P sub shells going across the periods as the same occurs in period 2 with Be and B. (You don't need to explain the drop back from P to S or N to O but this is because in P and N the 3 P electrons are in separate orbitals but in O and S one of the P orbitals has a pair of electrons in it which results in some repulsion and makes the extra electron easier to remove)
Evidence for levels is from going down the groups because the ionisation energies decrease (a lot) but the proton number increases so there must be a factor which is overcoming the extra attraction due to having more protons. This is evidence for shells because if the outer electron is in a shell further from the nucleus it has more shielding due to the inner shells and so there is less attraction from the nucleus and it needs less energy to remove the outermost electron.

Hope this helps.

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