Ionization energy is the least or smallest amount of energy needed to take away an atom's outermost valence electron. In the periodic table, as you progress towards the right side of it, the ionization energy increases since the nuclear charge induced by the protons (which pulls the electrons towards the nucleus) being added also increases. However, the ionization energy decreases as you progress downwards through the periods since more energy levels are being added and more energy levels means more electrons between the nucleus and valence electrons. All electrons possess a negative charge so essentially, the inner electrons are repelling the valence electrons being pulled inwards the nucleus. This phenomenon is known as electron shielding. So, basically, the element that's in the bottom left of the periodic table has the weakest ionization energy - other elements that are more electronegative can easily take away its valence electron. Meanwhile, the element that's in the top right corner of the periodic table has the highest ionization energy - no one can just simply take away its outer electrons.
I've attached a diagram hoping that it can help you understand the concepts I've explained much better.
