No. of electrons in each energy level
• Electrons occupy the lowest possible energy level in an atom.
Bohr Model – electrons travelling in ﬁxed circular paths (INACCURATE)
• It is not possible to specify exact loca2on of electron in atom –
• Only the probability of loca2ng an electron in a given region of space.
• An ion is a charged atom!
• Atoms form charged ions by the gain or loss of e‐’s
• e‐’s can be removed from atoms by using energy. Ionisa2on Energy
The ﬁrst ionisa2on energy of an element is the minimum energy required to remove the most loosely bound electron from an isolated atom of that element in its gaseous state.
Unit 7, 12
Na to Na+
Ionisation Energy required to remove an electron from an atom of Na
• Na has one e‐ in its outer energy level. It wants to lose this e‐ so the energy required to remove it would be very low. Therefore a small ionisa2on energy would be required to ionise an atom of Na.
• There are as many IEs as there are electrons • Subsequent IEs are higher than the first because you are removing a negative charge (electron) from an increasingly positive atom/ion
• Subsequent IEs make a huge jump after the electrons in the outer shell are lost.
• It is not difficult for Mg to lose 12th and
11th electron, but very difficult for it to lose it’s 10th electron.
Ionisa2on Energies of Magnesium
Bohr diagram explaining IE’s of Mg
Graph of the ﬁrst Twenty Ionisa2on Energies
Unit 7, 19
Explaining the Graph
1. The maximum values are for the noble gases. Reason: Their atoms are very stable because of their electronic conﬁgura2on [full outer (sub) level], so it is diﬃcult to remove an electron.
2. The minimum values are for the group one metals (alkali metals). Reason: Their atoms have only one electron in their outer level, so it is easily removed (as when this is lost it will have noble gas conﬁgura2on.) This is why group one are so reac2ve.
3. In general, ionisa'on energies increase in moving across a period from the alkali metal to the next noble gas. Reason: 1. Increase in nuclear charge (greater pull for electrons) 2. Decrease in atomic radius.
4. Ionisa'on energies gradually decrease in moving down a group. Reason: 1. Increase in atomic radius. 2. Screening eﬀect. (This is where the inner shell or shells of electrons help to shield the outer electrons from the posi2ve charge in the nucleus.
Excep2ons to Rule to 3 – Across a Period
There are two excep'ons to this generalisa'on:
(a) Group two elements (e.g. Be, Mg) have abnormally high values. This is because the most loosely bound electron comes from a full orbital which is a rela2vely stable state.
When the next element in each case (B, Al) is be ionised, the electron being removed is the single electron in the next orbital
(b) Group ﬁve elements also show abnormally high values (e.g. N and P). The reason here is that the electrons being removed are from exactly half – ﬁlled orbitals, and the half ﬁlled orbitals are the next most stable state a^er that of completely ﬁlled orbitals.
Values of Ionisa2on Energy decrease down the groups in the Periodic Table
• Increasing Atomic Radius – electrons further from a`rac2ve force easier to remove.
• Screening eﬀect – outermost electrons shielded from a`rac2ve force of +ve nucleus.
Ionisa2on Energy Trends ‐ Summary
Increase going across a period.
Decrease going down a group.
Increase in atomic radius.
Excep'ons, Group 2.
Increase in nuclear charge.
Decrease in atomic radius.
Full (outer) sublevel.
Excep'ons, Group 5.
Half full (outer) sublevel.
Unit 7, 23
Trends in Ionisa2on Energies
Ionisa2on Energy Trends
Unit 7, 24