Using your knowledge of periodicity and atomic structure, why does the first ionisation energy decrease moving down a group yet increase moving along a period in the periodic table?

Moving down a group means that the shielding of the outer electrons is increased due to the greater number of inner electron shells within the atom - these are between the nucleus and the outer electrons. As a result, the force of attraction between the positively charged protons and the negatively charged outer electrons is reduced - it requires less energy to remove the outer electrons, hence the first ionisation energy decreases down a group. Moving along a group, the trend in first ionisation energy is to increase due to a greater positive nuclear charge within the nucleus. The nuclear charge increases along a group due to more protons being in the nucleus of each atom. As a result, the outer electrons are more strongly attracted to the nucleus of the atom - more energy is required to remove the first electron. Another factor causing this trend is shielding. Along a period, nuclear charge increases due to more protons so the outer electrons are held tighter/closer to the nucleus and as such require more energy to be removed.