Quantity of ionization energy

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Ionization energy is the quantity of energy that a neutral atom (or an ion in case of second and third ionization energy) must absorb to discharge an electron. Each electron that is extracted from an atom requires a different amount of energy. Each succeeding ionization energy is larger than the preceding ionization energy. In other words, an atom could have first, second, or third ionization potential, and the third one will be greater than the second and the second will be greater than the first one(1).


In general, the more further away an electron is from the nucleus, the easier it is for it to be expelled. In other words, ionization energy is a function of atomic radius; the larger the radius the smaller the amount required to remove the electron from the outer most orbital. In a chemical reaction this characteristic is significantly important to understand the behavior of an element to make either ionic or covalent bond. For instance, the ionization energy of Sodium (alkaline earth metal) is 496KJ/mol(2)whereas Chlorine's first ionization potential is 1251.1 KJ/mol(3). Due to this difference in their ionization potential, when they chemically combine they make ionic bond. Elements that reside close to each other in the periodic table or the ones that do not have much of a difference in ionization energy make covalent bonds. For example, Carbon and Oxygen make CO2 (Carbon dioxide). Carbon and Chlorine make CCl4 (Carbon tetrachloride), molecules that are covalently bonded.

Periodic Table and Trend of Ionization Energies

As described above, ionization energies are dependent upon theatomic radius. Since goingfrom left to righton the periodic table,the atomic radius decreases therefore the ionization energy increases across the groups (please see diagram 1). Earth metals (IA group) have small ionization energies as compared to halogens or VIIA group (see diagram 2). In addition to radius (distance between nucleus and the electrons in outermost orbital), number of electrons in a given orbital also make a difference. The more the number of electrons the higher the shielding effect of electrons is. Shielding is inability of electrons not to see the positively charged nucleus. The higher theshielding effectthe lower the ionization energy. For the very reason the ionization potential decreases from top to bottom within a group. From this trend, Cesium is said to have the lowest ionization energy and Fluorine is said to have the highest ionization energy (with the exception of Helium and Neon).


The symbolI1stands for thefirst ionization energy(energy required to take away an electron from a neutral atom) and the symbolI2stands for thesecond ionization energy(energy required to take away an electron from an atom with a +1 charge. Each succeeding ionization energy is larger than the preceding energy. This means thatI1<I2<I3< ...

Example of how ionization energy increases as succeeding electrons are taken away.

Mg(g)->Mg+(g)+e-I1= 738kJ/mol

Mg+(g)->Mg2+(g) +e-I2= 1451kJ/mol

Metallic and Non-metallic Characteristics:

Ionization potential of a neutral atom or ion also describes the tendency of an electron to lose the electron during a chemical reaction. For example, Halogens can capture an electron easily as compared to elements in the first and second group. This tendency is termed aselectronegativityand along the periods of periodic table it follows the same trend as ionization potential does(4).

As indicated above, the elements to the right side of periodic table have tendency to receive the electron while the one at the left are more electropositive. Also, from left to right, the metalic characteristics of elements decrease(5).

Prediction of Covelent and Ionic Bonds:

The difference of electronegativity or ionization potential between two reacting elements determine the fate of the bond. For example, sodium and chlorine has a big difference between their ionization energies and electronegativities. Therefore, sodium completely removes the electron from its outermost orbital and chlorine completely accepts the electron, and as a result we have anionic bond(5). However, in cases where there is no differencein electronegativites, the sharing of electrons produces acovalent bond. For example, electronegativity of Hydrogen is 2.1 andthe combination of two Hydrogen atoms will definitely make a covalent bond (by sharing of electrons). The combination of Hydrogen and Fluorine (electronegativity=3.96) will produce apolarcovalent bondbecause they have small differences between electronegativities(6).


Consider the ground state electronic configuration of elements that are given below:

(A) 1s22s22p6(B) 1s22s22p4(C) 1s22s22p63s2(D) 1s22s22p63s1(E) 1s22s22p5.

Which element has the lowest and which one is considerd to have highest ionization energy?(7).


Element Dhas the lowest energy of ionization (because it has to remove only one electron)

Element Ahas the largest energy of ionization (because it has to remove 6 electrons and each successive electron removal requires more energy than the preceeding electron).

Outside links

  • Kaufman, Myron J.; Trowbridge, C. G. "The Ionization Energy of Helium ."J. Chem. Educ.19997688.
  • Rioux, Frank; DeKock, Roger L. "The Crucial Role of Kinetic Energy in Interpreting Ionization Energies."J. Chem. Educ.199875537.


  1. Petrucci, Ralph H.General Chemistry. 9th ed. New Jersey: Pearson Prentice Hall, 2005.
  2. http://hyperphysics.phy-astr.gsu.edu.../bondd.html#c1
  3. http://www.chemicool.com/elements/chlorine.html)
  4. http://hyperphysics.phy-astr.gsu.edu.../bondd.html#c1
  5. http://www.tutorvista.com/content/sc...nts/trends.php
  6. http://www.tutor-homework.com/Chemis...egativity.html
  7. http://www.tutorvista.com/content/ch...ion-energy.php