The first ionization energy IE1 is the energy required to remove the one electron from a gaseous atom. With this the effective nuclear charge increases and thus it takes more energy to remove the next electron. IE2 is always greater than IE1. Each subsequent electron becomes harder to remove. In the plot below it is apparent that there are periodic patterns to the first ionization energies. Within a given row the ionization energy increases with increasing atomic number. Atoms with filled or half-filled subshells have higher energies than the trend. The IE1 of a row increases until a maximum is reached with a noble gas. The IE1 drops dramatically as the next shell begins to fill, starting with an alkali metal. It is obviously more difficult to remove an electron from a noble gas than form a metal. The transition metals tend to be similar in their properties. It is apparent that the most significant chemical properties are due to the s and p orbital electrons. The filling of the d and particularly the f orbitals make far less difference. In both the noble gases and the alkali metals there is a gradual decrease in IE1 with increasing atomic number. In the smaller atoms the electrons are closer to the nucleus and are thus held more firmly. The nuclear charge is shielded by the inner electrons in larger atoms.