Ionization Energy of the Elements

Ionization Energy of the Elements The ionization energy, or ionization potential, is the energy required to completely remove an electron from a gaseous atom or ion. The closer and more tightly bound an electron is to the nucleus, the more difficult it will be to remove, and the higher its ionization energy will be. Key Takeaways: Ionization Energy Ionization energy is the amount of energy needed to completely remove an electron from a gaseous atom.Generally, the first ionization energy is lower than that required to remove subsequent electrons. There are exceptions.Ionization energy exhibits a trend on the periodic table. Ionization energy generally increases moving from left to right across a period or row and decreases moving top to bottom down an element group or column. Units for Ionization Energy Ionization energy is measured in electronvolts (eV). Sometimes the molar ionization energy is expressed, in J/mol. First vs Subsequent Ionization Energies The first ionization energy is the energy required to remove one electron from the parent atom. The second ionization energy is the energy required to remove a second valence electron from the univalent ion to form the divalent ion, and so on. Successive ionization energies increase. The second ionization energy is (almost) always greater than the first ionization energy. There are a couple of exceptions. The first ionization energy of boron is smaller than that of beryllium. The first ionization energy of oxygen is greater than that of nitrogen. The reason for the exceptions has to do with their electron configurations. In beryllium, the first electron comes from a 2s orbital, which can hold two electrons as is stable with one. In boron, the first electron is removed from a 2p orbital, which is stable when it holds three or six electrons. Both of the electrons removed to ionize oxygen and nitrogen come from the 2p orbital, but a nitrogen atom has three electrons in its p orbital (stable), while an oxygen atom has 4 electrons in the 2p orbital (less stable). Ionization Energy Trends in the Periodic Table Ionization energies increase moving from left to right across a period (decreasing atomic radius). Ionization energy decreases moving down a group (increasing atomic radius). Group I elements have low ionization energies because the loss of an electron forms a stable octet. It becomes harder to remove an electron as the atomic radius decreases because the electrons are generally closer to the nucleus, which is also more positively charged. The highest ionization energy value in a period is that of its noble gas. Terms Related to Ionization Energy The phrase ionization energy is used when discussing atoms or molecules in the gas phase. There are analogous terms for other systems. Work Function - The work function is the minimum energy needed to remove an electron from the surface of a solid. Electron Binding Energy - The electron binding energy is a more generic term for ionization energy of any chemical species. Its often used to compare energy values needed to remove electrons from neutral atoms, atomic ions, and polyatomic ions. Ionization Energy Versus Electron Affinity Another trend seen in the periodic table is electron affinity. Electron affinity is a measure of the energy released when a neutral atom in the gas phase gains an electron and forms a negatively charged ion (anion). While ionization energies may be measured with great precision, electron affinities are not as easy to measure. The trend to gain an electron increases moving from left to right across a period in the periodic table and decreases moving from top to bottom down an element group. The reasons electron affinity typically becomes smaller moving down the table is because each new period adds a new electron orbital. The valence electron spends more time further from the nucleus. Also, as you move down the periodic table, an atom has more electrons. Repulsion between the electrons makes it easier to remove an electron or harder to add one. Electron affinities are smaller values than ionization energies. This puts the trend in electron affinity moving across a period into perspective. Rather than a net release of energy when an electron is gain, a stable atom like helium actually requires energy to force ionization. A halogen, like fluorine, readily accepts another electron.