identify and explain trends in ionisation energies of atoms andtheir relationship with the Periodic Table

Ionisation energy is the energy required to remove an electron from a gaseous atom or ion. The trend in ionisation energies across the periodic table is influenced by several factors:

1. Atomic Radius: As you move from left to right across a period, the atomic radius decreases. This is because electrons are being added to the same energy level at the same time the nucleus is increasing in protons. The increase in the number of protons causes an increase in the positive charge of the nucleus which attracts the electrons on the same energy level closer to the nucleus thereby decreasing the atomic radius. As the atomic radius decreases, it becomes harder to remove an electron, hence the ionisation energy increases.

2. Nuclear Charge: As you move from left to right across a period, the nuclear charge increases. This is because the number of protons in the nucleus increases. The increase in nuclear charge attracts the electrons more strongly, making it harder to remove an electron, hence the ionisation energy increases.

3. Electron Shielding: As you move down a group, the electron shielding effect increases. This is because there are more energy levels between the nucleus and the outermost energy level. The inner energy levels 'shield' the outermost energy level from the positive charge of the nucleus. This makes it easier to remove an electron, hence the ionisation energy decreases.

4. Electron Pair Repulsion: Electrons are negatively charged and therefore repel each other. In a multi-electron atom, the repulsion between electrons can make an electron easier to remove. For example, the ionisation energy of oxygen is less than that of nitrogen. In nitrogen, the outermost electrons are all in different orbitals, whereas in oxygen, two of the electrons are paired in the same orbital. The repulsion between these two electrons makes one of them easier to remove, hence the ionisation energy decreases.

In summary, ionisation energy generally increases across a period (from left to right) and decreases down a group. However, there are exceptions to these trends due to electron shielding and electron pair repulsion.