Editing Ionization energy (section)

====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"/>