Ionization Energy

allows the outer electrons to be more easily borrowed to other atoms. The effective nuclear charge is the atomic number with the number of shielding electrons separated from it. In other words:

Z(effective)= Z (atomic number) minus S (number of shielding electrons)

An example of this would be chlorine: It has an atomic number of 17 and 10 shielding electrons (those protected in stable orbitals). The effective Z value is +7 because there are seven electrons left over in an outer orbital that is not complete with a filled octet. The value of +7 is the effective nuclear charge. This means that the nucleus is pulling the outer electrons closer to the nucleus, leading to a smaller atomic radius. Sodium has a much lesser nuclear charge at +1 because its atomic number is 11 (having 10 shielding electrons). Its pull on the electrons is less.

Moving from left to light across the periodic table, the nucleus has a greater pull on the outer electrons and the atomic radius goes down. Moving down the periodic table means larger atomic numbers and a decreased pull on the outer valence electrons and a larger atomic radius.

IONIZATION ENERGY

This is an important feature of the periodic table. In order to remove an electron from an atom, you need enough energy to remove an electron from the positive charge of the nucleus. The IE or I or Ionization energy is the energy necessary to get rid of an electron from the atom or ion. This is always a positive number. The “first IE” is the energy needed to remove the first valence electron, while the “second IE” is the energy required to remove the second valence electron. There is a third and fourth IE as well.

An example is sodium. The first IE is that to make a sodium ion Na+ (Sodium Plus) and an electron, while the second IE is the energy to take Na+ to make Na2+ and the second electron. Going up a group increases the ionization energy because the atom is smaller and has fewer shielding electrons and a stronger ionization energy.