Ionization Energy

The ionization energy is the energy needed to remove an electron from an atom.  For an element X, it is written as:
                      X  +  energy ------>   X+  +  e-

The ionization energy is a measure of how tightly the electrons are held by the atom.

Successive Ionization Energies in kJ/mole for the First 12 Elements
               1             2             3             4               5               6               7                8
H          1312
He        2372      5250
Li           520   |  7297      11,810
Be          899      1757   |  14,845     21,000
B            800      2426        3659   |  25,020     32,820
C          1086      2352        4619        6221  |   37,820      47,260
N          1402      2855        4576        7473        9442   |   53,250      64,340
O          1314      3388        5296        7467     10,987       13,320  |   71,320     84,070
F           1680      3375        6045        8408     11,020       15,160      17,860  |  92,101
Ne         2080     3963        6130        9361     12,180       15,240       -------     -------
Na          496    |  4563        6913        9541     13,350       16,600       20,113     25,666
Mg         737      1450     |   7731     10,545     13,627      17,995       21,700      25,662
Al           580
Si            790
P           1067
S           1004
Cl          1260
Ar          1525
K            416     3082
Ca           592     1151    |   4960
Rb           403  |   2662
Cs           378  |   2432

Atoms with more than one electron have more than one ionization energy. These ionization energies correspond to the stepwise removal of electrons, one after the other.  Take for example Be. It is fairly easy to remove the first electron and about twice as much energy to remove a second.  To remove the third electron involves a major jump.

In general, successive ionization energies always increase because each subsequent electron is being pulled away from an increasingly more positive ion.

Periodic Table Trends
Ionization energy increases from bottom to top within a group, and increases from left to right within a period.

The trend within a group can be seen easily by observing how the ionization energies vary for either the alkali metals (Li through Cs) or the noble gases (He through Rn). The change in ionization energy from Li through Ne or from Na through Ar illustrates the trend across a period. These periodic trends in ionization energy are the opposite of the trends in atomic size within the periodic table.

      The reason for the left to right trend is the gradual increase in effective nuclear charge felt by the valence electrons. As this charge increases the valence electrons get pulled closer to the nucleus which causes the shrinking of the atom but it means that more energy will have to be expended to pull an electron away.

Noble Gas Stability
Look at the trends in the chart above. We always get a big increase in the ionization energy needed when we start to remove electrons from a new inner core of electrons.  The noble gases have electron configurations that have a full 'p' subshell.  In effect the noble gases have very stable cores of electrons.  For this reason they are considered to be inert and unreactive.

Homework Assignment:
Using the ionization energies chart from above:
1. Get a blank periodic table and fill in the first electron ionization energies in the appropriate element.   Verify for yourself that the ionization trends are as stated above.

2.  Using your periodic table from above, plot on a piece of graph paper the Ionization Energy vs Atomic number.  Connect dot-to-dot. Can you note any trends visually on the graph?