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Ionization Energy

 

 

From-http://www.sparknotes.com/testprep/books/sat2/chemistry/chapter4section6.rhtml
The ionization energy of an atom is the energy required to remove an electron from the atom in the gas phase. 

M(g) ----> M+(g) + e-

Although removing the first electron from an atom requires energy, the removal of each subsequent electron requires even more energy. This means that the second IE is usually greater than the first, the third IE is greater than the second, and so on. The reason it becomes more difficult to remove additional electrons is that they’re closer to the nucleus and thus held more strongly by the positive charge of the protons.

 

Ionization energies differ significantly, depending on the shell from which the electron is taken. For instance, it takes less energy to remove a p electron than an s electron, even less energy to extract a d electron, and the least energy to extract an f electron. As you can probably guess, this is because s electrons are held closer to the nucleus, while f electrons are far from the nucleus and less tightly held. You’ll need to remember two important facts about ionization energy for the test. The first is that ionization energy increases as we move across a period.

 

The reason for this, as is the case with periodic trends in atomic radii, is that as the nucleus becomes more positive, the effective nuclear charge increases its pull on the electrons and it becomes more difficult to remove an electron. The second thing you’ll need to remember is that ionization energy decreases as you move down a group or family. The increased distance between electrons and the nucleus and increased shielding by a full principal energy level means that it requires less energy to remove an electron. Shielding occurs when the inner electrons in an atom shield the outer electrons from the full charge of the nucleus. Keep in mind that this phenomenon is only important as you move down the periodic table! Here are the values for the first ionization energies for some elements:

As for Bonding, positive ions (cations) are formed because they have low ionization energy. They tend to give up their electrons easily. So they become positive ions. Nonmetals become positive because of electron affinity.

on to Electron Affinity

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