In physics and chemistry, ionization energy (IE) (American English spelling), or ionisation energy (British English spelling) is the minimum energy required to remove the most loosely bound electron(s) (the valence electron(s)) of an isolated gaseous atom, positive ion, or molecule. The first ionization energy is quantitatively expressed as X(g) + energy ⟶ X+(g) + e− where X is any atom or molecule, X+ is the resultant ion when the original atom was stripped of a single electron, and e− is the removed electron. Ionization energy is positive for neutral atoms, meaning that the ionization is an endothermic process. Roughly speaking, the closer the outermost electrons are to the nucleus of the atom, the higher the atom's ionization energy.
In physics, ionization energy (IE) is usually expressed in electronvolts (eV) or joules (J). In chemistry, it is expressed as the energy to ionize a mole of atoms or molecules, usually as kilojoules per mole (kJ/mol) or kilocalories per mole (kcal/mol).
Comparison of ionization energies of atoms in the periodic table reveals two periodic trends which follow the rules of Coulombic attraction:
Ionization energy generally increases from left to right within a given period (that is, row). Ionization energy generally decreases from top to bottom in a given group (that is, column).
The latter trend results from the outer electron shell being progressively farther from the nucleus, with the addition of one inner shell per row as one moves down the column.
The nth ionization energy refers to the amount of energy required to remove the most loosely bound electron from the species having a positive charge of (n − 1). For example, the first three ionization energies are defined as follows: 1st ionization energy is the energy that enables the reaction X ⟶ X+ + e−
2nd ionization energy is the energy that enables the reaction X+ ⟶ X2+ + e−
3rd ionization energy is the energy that enables the reaction X2+ ⟶ X3+ + e−
The most notable influences that determine ionization energy include: Electron configuration: This accounts for most elements' IE, as all of their chemical and physical characteristics can be ascertained just by determining their respective electron configuration (EC). Nuclear charge: If the nuclear charge (atomic number) is greater, the electrons are held more tightly by the nucleus and hence the ionization energy will be greater (leading to the mentioned trend 1 within a given period). Number of electron shells: If the size of the atom is greater due to the presence of more shells, the electrons are held less tightly by the nucleus and the ionization energy will be smaller. Effective nuclear charge (Zeff): If the magnitude of electron shielding and penetration are greater, the electrons are held less tightly by the nucleus, the Zeff of the electron and the ionization energy is smaller. Stability: An atom having a more stable electronic configuration has a reduced tendency to lose electrons and consequently has a higher ionization energy.
Minor influences include: Relativistic effects: Heavier elements (especially those whose atomic number is greater than about 70) are affected by these as their electrons are approaching the speed of light. They therefore have smaller atomic radii and higher ionization energies. Lanthanide and actinide contraction (and scandide contraction): The shrinking of the elements affects the ionization energy, as the net charge of the nucleus is more strongly felt. Electron pairing energies: Half-filled subshells usually result in higher ionization energies.
The term ionization potential is an older and obsolete term for ionization energy, because the oldest method of measuring ionization energy was based on ionizing a sample and accelerating the electron removed using an electrostatic potential.
Determination of ionization energies
The ionization energy of atoms, denoted Ei, is measured by finding the minimal energy of light quanta (photons) or electrons accelerated to a known energy that will kick out the least bound atomic electrons. The measurement is performed in the gas phase on single atoms. While only noble gases occur as monatomic gases, other gases can be split into single atoms. Also, many solid elements can be heated and vaporized into single atoms. Monatomic vapor is contained in a previously evacuated tube that has two parallel electrodes connected to a voltage source. The ionizing excitation is introduced through the walls of the tube or produced within.
When ultraviolet light is used, the wavelength is swept down the ultraviolet range. At a certain wavelength (λ) and frequency of light (ν=c/λ, where c is the speed of light), the light quanta, whose energy is proportional to the frequency, will have energy high enough to dislodge the least bound electrons. These electrons will be attracted to the positive electrode, and the positive ions remaining after the photoionization will get attracted to the negatively charged electrode. These electrons and ions will establish a current through the tube. The ionization energy will be the energy of photons hνi (h is the Planck constant) that caused a steep rise in the current: Ei = hνi.
When high-velocity electrons are used to ionize the atoms, they are produced by an electron gun inside a similar evacuated tube. The energy of the electron beam can be controlled by the acceleration voltages. The energy of these electrons that gives rise to a sharp onset of the current of ions and freed electrons through the tube will match the ionization energy of the atoms.
Atoms: values and trends
Generally, the (N+1)th ionization energy of a particular element is larger than the Nth ionization energy (it may also be noted that the ionization energy of an anion is generally less than that of cations and neutral atom for the same element). When the next ionization energy involves removing an electron from the same electron shell, the increase in ionization energy is primarily due to the increased net charge of the ion from which the electron is being removed. Electrons removed from more highly charged ions experience greater forces of electrostatic attraction; thus, their removal requires more energy. In addition, when the next ionization energy involves removing an electron from a lower electron shell, the greatly decreased distance between the nucleus and the electron also increases both the electrostatic force and the distance over which that force must be overcome to remove the electron. Both of these factors further increase the ionization energy.
Some values for elements of the third period are given in the following table:
+Successive ionization energy values / kJ mol−1 (96.485 kJ mol−1 ≡ 1 eV) Element First Second Third Fourth Fifth Sixth Seventh Na 496 4,560 Mg 738 1,450 7,730 Al 577 1,816 2,881 11,600 Si 786 1,577 3,228 4,354 16,100 P 1,060 1,890 2,905 4,950 6,270 21,200 S 1,000 2,295 3,375 4,565 6,950 8,490 27,107 Cl 1,256 2,260 3,850 5,160 6,560 9,360 11,000 Ar 1,520 2,665 3,945 5,770 7,230 8,780 12,000
Large jumps in the successive molar ionization energies occur when passing noble gas configurations. For example, as can be seen in the table above, the first two molar ionization energies of magnesium (stripping the two 3s electrons from a magnesium atom) are much smaller than the third, which requires stripping off a 2p electron from the neon configuration of Mg2+. That 2p electron is much closer to the nucleus than the 3s electrons removed previously.
Ionization energy is also a periodic trend within the periodic table. Moving left to right within a period, or upward within a group, the first ionization energy generally increases, with exceptions such as aluminium and sulfur in the table above. As the nuclear charge of the nucleus increases across the period, the electrostatic attraction increases between electrons and protons, hence the atomic radius decreases, and the electron cloud comes closer to the nucleus because the electrons, especially the outermost one, are held more tightly by the higher effective nuclear charge.
On moving downward within a given group, the electrons are held in higher-energy shells with higher principal quantum number n, farther from the nucleus and therefore are more loosely bound so that the ionization energy decreases. The effective nuclear charge increases only slowly so that its effect is outweighed by the increase in n.
Exceptions in ionization energies
There are exceptions to the general trend of rising ionization energies within a period. For example, the value decreases from beryllium (: 9.3 eV) to boron (: 8.3 eV), and from nitrogen (: 14.5 eV) to oxygen (: 13.6 eV). These dips can be explained in terms of electron configurations.
Boron has its last electron in a 2p orbital, which has its electron density farther away from the nucleus on average than the 2s electrons in the same shell. The 2s electrons then shield the 2p electron from the nucleus to some extent, and it is easier to remove the 2p electron from boron than to remove a 2s electron from beryllium, resulting in a lower ionization energy for B.
In oxygen, the last electron shares a doubly occupied p-orbital with an electron of opposing spin. The two electrons in the same orbital are closer together on average than two electrons in different orbitals, so that they shield each other from the nucleus more effectively and it is easier to remove one electron, resulting in a lower ionization energy.
Furthermore, after every noble gas element, the ionization energy drastically drops. This occurs because the outer electron in the alkali metals requires a much lower amount of energy to be removed from the atom than the inner shells. This also gives rise to low electronegativity values for the alkali metals.
The trends and exceptions are summarized in the following subsections:
Ionization energy decreases when Transitioning to a new period: an alkali metal easily loses one electron to leave an octet or pseudo-noble gas configuration, so those elements have only small values for IE. Moving from the s-block to the p-block: a p-orbital loses an electron more easily. An example is beryllium to boron, with electron configuration 1s2 2s2 2p1. The 2s electrons shield the higher-energy 2p electron from the nucleus, making it slightly easier to remove. This also happens from magnesium to aluminium. Occupying a p-subshell with its first electron with spin opposed to the other electrons: such as in nitrogen (: 14.5 eV) to oxygen (: 13.6 eV), as well as phosphorus (: 10.48 eV) to sulfur (: 10.36 eV). The reason for this is because oxygen, sulfur and selenium all have dipping ionization energies because of shielding effects. However, this discontinues starting from tellurium where the shielding is too small to produce a dip. Moving from the d-block to the p-block: as in the case of zinc (: 9.4 eV) to gallium (: 6.0 eV) Special case: decrease from lead (: 7.42 eV) to bismuth (: 7.29 eV). This cannot be attributed to size (the difference is minimal: lead has a covalent radius of 146 pm whereas bismuth's is 148 pm). This is due to the spin-orbit splitting of the 6p shell (lead is removing an electron from the stabilised 6p1/2 level, but bismuth is removing one from the destabilised 6p3/2 level). Predicted ionization energies show a much greater decrease from flerovium to moscovium, one row farther down the periodic table and with much larger spin-orbit effects. Special case: decrease from radium (: 5.27 eV) to actinium (: 5.17 eV), which is a switch from an s to a d orbital. However the analogous switch from barium (: 5.2 eV) to lanthanum (: 5.6 eV) does not show a downward change.