1. Ionisation energy

2. Definitions The first ionisation energy is the energy required to remove the first (or outermost) electron from each atom in a mole of gaseous atoms. Eg; Na (g) → Na + (g) + e - NB You must use state symbols in these equations.

3. Similarly the second ionisation energy is that required to remove the second electron from each ion in a mole of gaseous, single charged ions. Eg; Na + (g) → Na 2+ (g) + e

4. Q: Write an equation, including state symbols, to represent the fifth ionisation energy of sodium. The fifth electron can only be removed after the first four have gone, so start with the 4+ ion. Na 4+ → Na 5+ + e - Then add the state symbols; Na 4+ (g) → Na 5+ (g) + e -

5. Values of ionisation energies Ionisation energies depend on; 1) The nuclear charge. Electrons are negative, opposite charges attract, so they will be attracted by the positive nucleus. The bigger the atomic number the greater will be the attraction between the protons and the electrons and the higher the ionisation energy.

6. 2) The distance of the electrons from the nucleus The force of attraction between the nucleus and the electrons decreases with distance. This means that the further an electron is away from the nucleus the weaker the force of it will feel. Less energy will be required to remove it, so the ionisation energy will be lower.

7. 3) Shielding Electrons are arranged in shells. Like charges repel each other. Electrons are all negatively charged. This means that the electrons in the outer shells will be repelled from those in the inner shells. They are said to be shielded, or screened. This means that the outer electrons will not experience the full force of attraction of the nucleus. Less energy will be required to remove them and so their ionisation energies will be lower.

8. Ionisation Energy

9. Ionisation energies and electron shells The outermost electrons are easier to remove from an atom, as their ionisation energies are the lowest. Electrons in the same shell have similar ionisation energies. But, due to shielding, there are big differences between the ionisation energies of the shells. So successive ionisation energies allow us to predict how many shells of electrons are present in an atom, and how many electrons are in each shell.

10. Arrangement of electrons in oxygen Six electrons with similar IEs = same shell. Big jump = change of shell. Two electrons with similar IEs = inner shell.

11. Which element is this? Electronic configuration = 2,8,8,1 I electron in outer shell = Group 1. N o electrons = 19 = Atomic Number = K

12. Ionisation energy generally increases across a period. Atomic Number increases, nuclear charge increases, so greater attraction for the electron, so more energy is needed to remove it.

13. Moving from group 2 to 3 the first electron is lost from a p orbital, rather than an s orbital. p orbital is higher energy than s orbital, so easier to lose electron.

14. 5 ↑↑↑ 6↑↓↑↑ In group 5 the electron is removed from a p orbital with one electron. In group 6 it is removed from one with two electrons. Due to repulsion it is easier to remove the latter. So IEs are slightly lower.

15. Ionisation energies decrease down a group Electrons are further from the nucleus and there is more shielding, so they are easier to remove.

16. Periodicity of first ionisation energy A repeating pattern is seen going along the periods. NB A group of 2 electrons

17. Followed by a group of 6, divided into two groups of 3.

18. Then a group of 10.

19. Big drops in IE allow the electron shells to be identified. First shell = 2 Second shell = 8 Third shell = 8 Second part of 3 rd shell = 10 Fourth shell = 2=6=10 =18