Define ionization energy and its trends in the periodic table
Ionization energy refers to the energy required to remove an electron from an atom or a positive ion in its gaseous state. It can also be defined as the minimum amount of energy needed to completely remove the most loosely held electron from an atom.
In terms of trends in the periodic table, ionization energy generally increases from left to right across a period and decreases from top to bottom within a group.
Across a period (from left to right): As you move from left to right across a period, ionization energy generally increases. This is because the number of protons in the nucleus and the number of electrons in the same energy level remain the same, resulting in increased attraction between the positively charged nucleus and the negatively charged electrons. As a result, it becomes more difficult to remove an electron, leading to higher ionization energy.
Within a group (from top to bottom): As you move down a group in the periodic table, ionization energy generally decreases. This is because the number of energy levels or electron shells increases as you move down the group. Electrons located in higher energy levels are further away from the positively charged nucleus, reducing the attraction between the nucleus and the outermost electron. Consequently, it becomes easier to remove an electron, resulting in lower ionization energy.