Editing Ionization energy (section)

===Exceptions in ionization energies===
{{original research|section|date=December 2022}}
There are exceptions to the general trend of rising ionization energies within a period. For example, the value decreases from [[beryllium]] ({{nuclide|Be|&nbsp;}}: 9.3 eV) to [[boron]] ({{nuclide|B|&nbsp;}}: 8.3 eV), and from [[nitrogen]] ({{nuclide|N|&nbsp;}}: 14.5 eV) to [[oxygen]] ({{nuclide|O|&nbsp;}}: 13.6 eV). These dips can be explained in terms of electron configurations.<ref name=":Grandinetti">{{Cite web|url=https://www.grandinetti.org/ionization-energy-trends|title=Ionization Energy Trends {{!}} Grandinetti Group |last=Grandinetti |first=Philip J. |date=September 8, 2019 |access-date=2020-09-13|website=www.grandinetti.org}}</ref>

[[File:BerylliumVsBoronElectronConfiguration.jpg|thumb|200px|The added electron in boron occupies a [[atomic orbital|p-orbital]].]]
Boron has its last electron in a 2p orbital, which has its [[electron density]] further 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.<ref name=Miessler/>

{{Multiple images
| image1          = NitrogenVsOxygenElectronConfiguration.jpg
| alt1            = Nitrogen and oxygen's electron configuration
| caption1        = These electron configurations do not show the full and half-filled orbitals.
| image2          = NitrogenVsOxygenElectronConfigurationBoxAndArrows.jpg
| alt2            = Nitrogen and oxygen's electron configuration using box and arrows
| caption2        = Here the added electron has a spin opposed to the other 2p electrons. This decreases the ionization energy of oxygen}}
In oxygen, the last electron shares a doubly occupied p-orbital with an electron of opposing [[Spin (physics)|spin]]. The two electrons in the same orbital are closer together on average than two electrons in different orbitals, so that they [[Shielding effect|shield each other from the nucleus]] more effectively and it is easier to remove one electron, resulting in a lower ionization energy.<ref name=Miessler/><ref name=":0">{{cite web |url=https://www.kentchemistry.com/links/PT/PTIonE.htm |title=First Ionization Energy |last=Kent |first=Mr. |website=kentchemistry.com |publisher=KentChemistry |access-date=December 6, 2020 |quote=...The addition of the second electron into an already occupied orbital introduces repulsion between the electrons, thus it is easier to remove. that is why there is a dip in the ionization energy.}}</ref>

Furthermore, after every noble gas element, the ionization energy drastically drops. This occurs because the outer electron in the [[alkali metal]]s 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.<ref>{{Cite web|title=Group IA|url=https://chemed.chem.purdue.edu/genchem/topicreview/bp/ch9/alkali.php|access-date=2020-09-20|website=chemed.chem.purdue.edu}}</ref><ref>{{Cite web|title=Alkali Metals|url=http://hyperphysics.phy-astr.gsu.edu/hbase/pertab/alkmet.html|access-date=2020-09-13|website=hyperphysics.phy-astr.gsu.edu}}</ref><ref>{{Cite web|title=The Alkali Metals {{!}} Introduction to Chemistry|url=https://courses.lumenlearning.com/introchem/chapter/the-alkali-metals/|access-date=2020-09-13|website=courses.lumenlearning.com}}</ref>

{{multiple image
| image1            = ZincVsGalliumElectronConfiguration.jpg
| alt1              = Zinc and Gallium's respective electron configurations
| caption1          = Because of a single p-orbital electron in [[gallium]]'s configuration, makes the overall structure less stable, hence the dip in ionization energy values<ref name="Lang & Smith 2003"/>
| image2            = RadiumVsActiniumElectronConfiguration.jpg
| alt2              = Radium and Actinium's Electron Configuration (condensed)
| caption2          = [[Actinium]]'s electron configuration predetermines that it would require less energy to remove that single d-orbital electron, therefore even though it has a larger EC, [[radium]] still has the higher IE<ref>{{cite web |url=https://www.lenntech.com/periodic-chart-elements/ionization-energy.htm |title=Chemical elements listed by ionization energy |author=<!--Not stated--> |date=2018 |website=lenntech.com |publisher=Lenntech BV |access-date=December 6, 2020 |quote=The elements of the periodic table sorted by ionization energy click on any element's name for further information on chemical properties, environmental data or health effects. This list contains the 118 elements of chemistry.}}</ref>
}}

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 rule|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 1s<sup>2</sup> 2s<sup>2</sup> 2p<sup>1</sup>. 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]].<ref>{{cite web |url=https://www.angelo.edu/faculty/kboudrea/periodic/trends_ionization_energy.htm |title=The Parts of the Periodic Table |last=Boudreaux |first=K.A. |date=August 13, 2020 |orig-date=July 26, 2006 |department=Department of Chemistry and Biochemistry |website=angelo.edu/faculty/kboudrea/<!--this is the real website, pls dont change--> |publisher=Angelo State University |location=2601 W. Avenue N, San Angelo, TX 76909, Texas |language=en |access-date=December 19, 2020 |via=angelo.edu |archive-date=July 10, 2022 |archive-url=https://web.archive.org/web/20220710025232/https://www.angelo.edu/faculty/kboudrea/periodic/trends_ionization_energy.htm |url-status=dead }}</ref>
* Occupying a p-subshell with its '''first''' electron with spin opposed to the other electrons: such as in nitrogen ({{nuclide|N|&nbsp;}}: 14.5 eV) to oxygen ({{nuclide|O|&nbsp;}}: 13.6 eV), as well as [[phosphorus]] ({{nuclide|P|&nbsp;}}: 10.48 eV) to [[sulfur]] ({{nuclide|S|&nbsp;}}: 10.36 eV). The reason for this is because oxygen, sulfur and selenium all have dipping ionization energies because of shielding effects.<ref>{{Cite web|date=2014-07-02|title=18.10: The Group 6A Elements|url=https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_(Zumdahl_and_Decoste)/18%3A_The_Representative_Elements/18.10%3A_The_Group_6A_Elements|access-date=2020-09-20|website=Chemistry LibreTexts|language=en}}</ref> 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]] ({{nuclide|Zn|&nbsp;}}: 9.4 eV) to [[gallium]] ({{nuclide|Ga|&nbsp;}}: 6.0 eV)
* Special case: decrease from [[lead]] ({{nuclide|Pb|&nbsp;}}: 7.42 eV) to [[bismuth]] ({{nuclide|Bi|&nbsp;}}: 7.29 eV). This cannot be attributed to size (the difference is minimal: lead has a covalent radius of 146 [[picometer|pm]] whereas [[bismuth]]'s is 148 pm<ref>{{Cite web|title=Covalent Radius for all the elements in the Periodic Table|url=https://periodictable.com/Properties/A/CovalentRadius.v.log.html|access-date=2020-09-13|website=periodictable.com}}</ref>). This is due to the spin-orbit splitting of the 6p shell (lead is removing an electron from the stabilised 6p<sub>1/2</sub> level, but bismuth is removing one from the destabilised 6p<sub>3/2</sub> level). Predicted ionization energies show a much greater decrease from [[flerovium]] to [[moscovium]], one row further down the periodic table and with much larger spin-orbit effects.
* Special case: decrease from radium ({{nuclide|Ra|&nbsp;}}: 5.27 eV) to [[actinium]] ({{nuclide|Ac|&nbsp;}}: 5.17 eV), which is a switch from an s to a d orbital. However the analogous switch from [[barium]] ({{nuclide|Ba|&nbsp;}}: 5.2 eV) to [[lanthanum]] ({{nuclide|La|&nbsp;}}: 5.6 eV) does not show a downward change.
* [[Lutetium]] ({{nuclide|Lu|&nbsp;}}) and [[lawrencium]] ({{nuclide|Lr|&nbsp;}}) both have ionization energies lower than the previous elements. In both cases the last electron added [[Electron configurations of the elements (data page)|starts a new subshell]]: 5d for Lu with electron configuration [Xe] 4f<sup>14</sup> 5d<sup>1</sup> 6s<sup>2</sup>, and 7p for Lr with configuration [Rn] 5f<sup>4</sup> 7s<sup>2</sup> 7p<sup>1</sup>. These dips in ionization energies for lutetium and especially lawrencium show that these elements belong in the d-block, and not lanthanum and actinium.<ref name="JensenLr">{{cite web|url=https://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/251.%20Lawrencium.pdf |title=Some Comments on the Position of Lawrencium in the Periodic Table |last1=Jensen |first1=W. B. |date=2015 |access-date=20 September 2015 |url-status=dead |archive-url=https://web.archive.org/web/20151223091325/https://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/251.%20Lawrencium.pdf |archive-date=23 December 2015 }}</ref>

====Ionization energy increases when====
* Reaching Group 18 [[noble gas]] elements: This is due to their complete electron subshells,<ref>{{cite book |last1=Singh |first1=Jasvinder |chapter=Inert Gases |page=122 |chapter-url=https://books.google.com/books?id=eKnrhryjqn0C&pg=PA122 |title=Sterling Dictionary of Physics |date=1999 |publisher=Sterling Publishers Pvt. Ltd |isbn=978-81-7359-124-2 }}</ref> so that these elements require large amounts of energy to remove one electron.
* Group 12: The elements here, zinc ({{nuclide|Zn|&nbsp;}}: 9.4 eV), [[cadmium]] ({{nuclide|Cd|&nbsp;}}: 9.0 eV) and [[mercury (element)|mercury]] ({{nuclide|Hg|&nbsp;}}: 10.4 eV) all record sudden rising IE values in contrast to their preceding elements: [[copper]] ({{nuclide|Cu|&nbsp;}}: 7.7 eV), [[silver]] ({{nuclide|Ag|&nbsp;}}: 7.6 eV) and [[gold]] ({{nuclide|Au|&nbsp;}}: 9.2 eV), respectively. For mercury, it can be extrapolated that the [[relativistic quantum chemistry|relativistic]] stabilization of the 6s electrons increases the ionization energy, in addition to poor shielding by 4f electrons that increases the effective nuclear charge on the outer valence electrons. In addition, the closed-subshells electron configurations: [Ar] 3d<sup>10</sup> 4s<sup>2</sup>, [Kr] 4d<sup>10</sup>5s<sup>2</sup> and [Xe] 4f<sup>14</sup> 5d<sup>10</sup> 6s<sup>2</sup> provide increased stability.
* Special case: shift from [[rhodium]] ({{nuclide|Rh|&nbsp;}}: 7.5 eV) to [[palladium]] ({{nuclide|Pd|&nbsp;}}: 8.3 eV). Unlike other Group 10 elements, palladium has a higher ionization energy than the preceding atom, due to its electron configuration. In contrast to [[nickel]]'s [Ar] 3d<sup>8</sup> 4s<sup>2</sup>, and [[platinum]]'s [Xe] 4f<sup>14</sup> 5d<sup>9</sup> 6s<sup>1</sup>, palladium's electron configuration is [Kr] 4d<sup>10</sup> 5s<sup>0</sup> (even though the [[Aufbau principle#Exceptions to the rule in the transition metal|Madelung rule]] predicts [Kr] 4d<sup>8</sup> 5s<sup>2</sup>). Finally, [[silver]]'s lower IE ({{nuclide|Ag|&nbsp;}}: 7.6 eV) further accentuates the high value for palladium; the single added s electron is removed with a lower ionization energy than palladium,<ref>{{cite book |doi=10.1016/B978-0-7506-3365-9.50028-6 |chapter=Vanadium, Niobium and Tantalum |title=Chemistry of the Elements |year=1997 |pages=976–1001 |isbn=978-0-7506-3365-9 }}</ref> which emphasizes palladium's high IE (as shown in the above linear table values for IE)
* The IE of [[gadolinium]] ({{nuclide|Gd|&nbsp;}}: 6.15 eV) is somewhat higher than both the preceding ({{nuclide|Sm|&nbsp;}}: 5.64 eV), ({{nuclide|Eu|&nbsp;}}: 5.67 eV) and following elements ({{nuclide|Tb|&nbsp;}}: 5.86 eV), ({{nuclide|Dy|&nbsp;}}: 5.94 eV). This anomaly is due to the fact that gadolinium valence d-subshell borrows 1 electron from the valence f-subshell. Now the valence subshell is the d-subshell, and due to the poor shielding of positive nuclear charge by electrons of the f-subshell, the electron of the valence d-subshell experiences a greater attraction to the nucleus, therefore increasing the energy required to remove the (outermost) valence electron.
* Moving into d-block elements: The elements Sc with a 3d<sup>1</sup> electronic configuration has a ''higher'' IP ({{nuclide|Sc|&nbsp;}}: 6.56 eV) than the preceding element ({{nuclide|Ca|&nbsp;}}: 6.11 eV), contrary to the decreases on moving into s-block and p-block elements. The 4s and 3d electrons have similar shielding ability: the 3d orbital forms part of the n=3 shell whose average position is closer to the nucleus than the 4s orbital and the n=4 shell, but electrons in s orbitals experience greater penetration into the nucleus than electrons in d orbitals. So the mutual shielding of 3d and 4s electrons is weak, and the effective nuclear charge acting on the ionized electron is relatively large. Yttrium ({{nuclide|Y|&nbsp;}}) similarly has a higher IP (6.22 eV) than {{nuclide|Sr|&nbsp;}}: 5.69 eV.
* Moving into f-block elements; The elements ({{nuclide|La|&nbsp;}}: 5.18 eV) and ({{nuclide|Ac|&nbsp;}}: 5.17 eV) have only very slightly lower IP's than their preceding elements ({{nuclide|Ba|&nbsp;}}: 5.21 eV) and ({{nuclide|Ra|&nbsp;}}: 5.18 eV), though their atoms are anomalies in that they add a d-electron rather than an f-electron. As can be seen in the above graph for ionization energies, the sharp rise in IE values from ({{nuclide|Cs|&nbsp;}}: 3.89 eV) to ({{nuclide|Ba|&nbsp;}}: 5.21 eV) is followed by a small increase (with some fluctuations) as the f-block proceeds from {{nuclide|Ba|&nbsp;}} to {{nuclide|Yb|&nbsp;}}. This is due to the [[lanthanide contraction]] (for lanthanides).<ref name=Housecroft>{{cite book |last1=Housecroft |first1=C.E. |last2=Sharpe |first2=A.G. |date=November 1, 1993 |title=Inorganic Chemistry |url=https://www.pearson.com/us/higher-education/program/Housecroft-Inorganic-Chemistry-5th-Edition/PGM2178749.html |type=eBook |language=en |volume=3 |edition=15th |location=Switzerland |publisher=Pearson Prentice-Hall |publication-date=November 1, 1993 |pages=536, 649, 743 |doi=10.1021/ed070pA304.1 |isbn=978-0-273-74275-3 |archive-url=https://web.archive.org/web/20210414235943/https://www.pearson.com/us/higher-education/program/Housecroft-Inorganic-Chemistry-5th-Edition/PGM2178749.html |archive-date=April 14, 2021 |access-date=December 14, 2020 |url-status=bot: unknown }}</ref><ref name=Cotton>{{Cotton&Wilkinson5th|pages=776, 955}}</ref><ref name=Jolly>{{cite journal |doi=10.1021/ed062pA137.1 |title=Modern Inorganic Chemistry (Jolly, William L.) |year=1985 |last1=Billo |first1=E. J. |journal=Journal of Chemical Education |volume=62 |issue=4 |pages=A137 |bibcode=1985JChEd..62..137B |doi-access=free }}</ref> This decrease in ionic radius is associated with an increase in ionization energy in turn increases, since the two properties correlate to each other.<ref name=":1" /> As for d-block elements, the electrons are added in an inner shell, so that no new shells are formed. The shape of the added orbitals prevents them from penetrating to the nucleus so that the electrons occupying them have less shielding capacity.

====Ionization energy anomalies in groups====

Ionization energy values tend to decrease on going to heavier elements within a group<ref name=":Grandinetti" /> as shielding is provided by more electrons and overall, the valence shells experience a weaker attraction from the nucleus, attributed to the larger covalent radius which increase on going down a group<ref>{{Cite web|title=Patterns and trends in the periodic table - Periodicity - Higher Chemistry Revision|url=https://www.bbc.co.uk/bitesize/guides/zxc99j6/revision/6|access-date=2020-09-20|website=BBC Bitesize|language=en-GB}}</ref> Nonetheless, this is not always the case. As one exception, in Group 10 palladium ({{nuclide|Pd|&nbsp;}}: 8.34 eV) has a higher ionization energy than nickel ({{nuclide|Ni|&nbsp;}}: 7.64 eV), contrary to the general decrease for the elements from technetium {{nuclide|Tc|&nbsp;}} to xenon {{nuclide|Xe|&nbsp;}}. Such anomalies are summarized below:
* Group 1:
** [[Hydrogen]]'s ionization energy is very high (at 13.59844 eV), compared to the alkali metals. This is due to its single electron (and hence, very small [[electron cloud]]), which is close to the nucleus. Likewise, since there are not any other electrons that may cause shielding, that single electron experiences the full net positive charge of the nucleus.<ref>{{Cite web|date=2013-10-03|title=Ionization Energies|url=https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Physical_Properties_of_Matter/Atomic_and_Molecular_Properties/Ionization_Energy/Ionization_Energies|access-date=2020-09-20|website=Chemistry LibreTexts|language=en}}</ref>
** [[Francium]]'s ionization energy is higher than the precedent [[alkali metal]], [[cesium]]. This is due to its (and radium's) small ionic radii owing to relativistic effects. Because of their large mass and size, this means that its electrons are traveling at extremely high speeds, which results in the electrons coming closer to the nucleus than expected, and they are consequently harder to remove (higher IE).<ref>{{Cite web|date=2019-11-06|title=IYPT 2019 Elements 087: Francium: Not the most reactive Group 1 element|url=https://www.compoundchem.com/2019/11/06/iypt087-francium/|access-date=2020-09-20|website=Compound Interest|language=en-GB}}</ref>
* Group 2: [[Radium]]'s ionization energy is higher than its antecedent [[alkaline earth metal]] [[barium]], like francium, which is also due to relativistic effects. The electrons, especially the 1s electrons, experience ''very high effective nuclear charges''. To avoid falling into the nucleus, the 1s electrons must move at very high speeds, which causes the special relativistic corrections to be substantially higher than the approximate classical momenta. By the [[uncertainty principle]], this causes a relativistic contraction of the 1s orbital (and other orbitals with electron density close to the nucleus, especially ns and np orbitals). Hence this causes a cascade of electron changes, which finally results in the outermost electron shells contracting and getting closer to the nucleus.
* Group 4:
** [[Hafnium]]'s near similarity in IE with [[zirconium]]. The effects of the lanthanide contraction can still be felt ''[[lanthanide contraction#Influence on the post-lanthanides|after the lanthanides]]''.<ref name=Cotton/> It can be seen through the former's smaller atomic radius (which contradicts the [https://www.chem.tamu.edu/class/fyp/stone/tutorialnotefiles/fundamentals/trends.htm#:~:text=WHY%3F%20%2D%20The%20number%20of%20energy,a%20period%2C%20atomic%20radius%20decreases. observed periodic trend] {{Webarchive|url=https://web.archive.org/web/20181011230430/http://www.chem.tamu.edu/class/fyp/stone/tutorialnotefiles/fundamentals/trends.htm#:~:text=WHY%3F%20%2D%20The%20number%20of%20energy,a%20period%2C%20atomic%20radius%20decreases. |date=2018-10-11 }}) at 159 pm<ref>{{cite web |url=https://www.gordonengland.co.uk/elements/hf.htm |title=Hafnium |author=<!--Not stated--> |date=2020 |website=gordonengland.co.uk |publisher=Gordon England |access-date=December 7, 2020 |quote=...Atomic Radius 159 pm...}}</ref> ([[atomic radius#Notes|empirical value]]), which differs from the latter's 155 pm.<ref>{{cite web |url=https://pubchem.ncbi.nlm.nih.gov/element/Zirconium#section=Atomic-Radius |title=Zirconium (Element) - Atomic Radius |author=<!--Not stated-->|website=pubchem.ncbi.nlm.nih.gov |publisher=PubChem |access-date=December 8, 2020 |quote=155 pm (Empirical)}}
</ref><ref>{{cite journal |last1=Slater |first1=J. C. |title=Atomic Radii in Crystals |journal=The Journal of Chemical Physics |date=15 November 1964 |volume=41 |issue=10 |pages=3199–3204 |doi=10.1063/1.1725697 |bibcode=1964JChPh..41.3199S }}</ref> This in turn makes its ionization energies increase by 18 kJ/mol<sup>−1</sup>.
** [[Titanium]]'s IE is smaller than that of both hafnium and zirconium. Hafnium's ionization energy is similar to zirconium's due to lanthanide contraction. However, why zirconium's ionization energy is higher than the preceding elements' remains unclear; we cannot attribute it to atomic radius as it is higher for zirconium and hafnium by 15 pm.<ref>{{Cite web|title=WebElements Periodic Table » Titanium » radii of atoms and ions|url=https://www.webelements.com/titanium/atom_sizes.html|access-date=2020-09-20|website=www.webelements.com}}</ref> We also cannot invoke the ''condensed'' ionization energy, as it is more or less the same ([Ar] 3d<sup>2</sup> 4s<sup>2</sup> for titanium, whereas [Kr] 4d<sup>2</sup> 5s<sup>2</sup> for zirconium). Additionally, there are no half-filled nor fully filled orbitals we might compare. Hence, we can only invoke zirconium's ''full'' electron configuration, which is 1s<sup>2</sup>2s<sup>2</sup>2p<sup>6</sup>3s<sup>2</sup>3p<sup>6</sup>'''3d<sup>10</sup>'''4s<sup>2</sup>4p<sup>6</sup>4d<sup>2</sup>5s<sup>2</sup>.<ref>{{Cite web|last=Straka |first=J. |title=Periodic Table of the Elements: Zirconium - Electronic configuration|url=https://www.tabulka.cz/english/elements/configuration.asp?id=40|access-date=2020-09-20|website=www.tabulka.cz}}</ref> The presence of a full 3d-block sublevel is tantamount to a higher shielding efficiency compared to the 4d-block elements (which are only two electrons).{{efn|Nonetheless, further research is still needed to corroborate this mere inference.}}
* Group 5: akin to Group 4, [[niobium]] and [[tantalum]] are analogous to each other, due to their electron configuration and to the lanthanide contraction affecting the latter element.<ref>{{Cite web|title=Tantalum {{!}} chemical element|url=https://www.britannica.com/science/tantalum|access-date=2020-09-20|website=Encyclopedia Britannica|language=en}}</ref> Ipso facto, their significant rise in IE compared to the foremost element in the group, [[vanadium]], can be attributed due to their full d-block electrons, in addition to their electron configuration. Another intriguing notion is niobium's half-filled 5s orbital; due to repulsion and exchange energy (in other words the ''"costs"'' of putting an electron in a low-energy sublevel to completely fill it instead of putting the electron in a high-energy one) overcoming the energy gap between s- and d-(or f) block electrons, the EC does not follow the Madelung rule.
* Group 6: like its forerunners groups 4 and 5, group 6 also record high values when moving downward. [[Tungsten]] is once again similar to [[molybdenum]] due to their electron configurations.<ref>{{cite book |doi=10.1002/0471435139.tox038 |chapter=Chromium, Molybdenum, and Tungsten |title=Patty's Toxicology |year=2015 |last1=Langård |first1=Sverre |isbn=978-0-471-12547-1 }}</ref> Likewise, it is also attributed to the full 3d-orbital in its electron configuration. Another reason is molybdenum's half filled 4d orbital due to electron pair energies violating the aufbau principle.
* Groups 7-12 6th period elements ([[rhenium]], [[osmium]], [[iridium]], [[platinum]], [[gold]] and [[mercury (element)|mercury]]): All of these elements have extremely high ionization energies compared to the elements preceding them in their respective groups. The essence of this is due to the lanthanide contraction's influence on post lanthanides, in addition to the relativistic stabilization of the 6s orbital.
* Group 13:
** Gallium's IE is higher than aluminum's. This is once again due to d-orbitals, in addition to scandide contraction, providing weak shielding, and hence the effective nuclear charges are augmented.
** Thallium's IE, due to poor shielding of 4f electrons<ref name="Lang & Smith 2003">{{cite journal |last1=Lang |first1=Peter F. |last2=Smith |first2=Barry C. |title=Ionization Energies of Atoms and Atomic Ions |journal=Journal of Chemical Education |date=August 2003 |volume=80 |issue=8 |pages=938 |doi=10.1021/ed080p938 |bibcode=2003JChEd..80..938L }}</ref> in addition to lanthanide contraction, causes its IE to be increased in contrast to its precursor [[indium]].
* Group 14: [[Lead]]'s unusually high ionization energy ({{nuclide|Pb|&nbsp;}}: 7.42 eV) is, akin to that of group 13's thallium, a result of the full 5d and 4f subshells. The lanthanide contraction and the inefficient screening of the nucleus by the 4f electrons results in slightly ''higher'' ionization energy for lead than for [[tin]] ({{nuclide|Sn|&nbsp;}}: 7.34 eV).<ref>{{Cite web|date=2015-12-02|title=The Group 14 elements|url=https://www.webelements.com/nexus/the-group-14-elements/|access-date=2020-09-13|website=Chemistry Nexus|language=en-US}}</ref><ref name="Lang & Smith 2003"/>