Ionization energy trends | Periodic table | Chemistry | Khan Academy

The video script begins with an exploration of the concept of ions, which are atoms or molecules with an unequal number of protons and electrons, resulting in a net charge. Positive ions, or cations, have more protons than electrons, while negative ions, or anions, have more electrons. The script uses hydrogen as an example of a positive ion and fluorine for a negative ion. It then introduces ionization energy, which is the energy required to remove an electron from an atom, and explains how this concept relates to the ease of turning elements into cations. The periodic table trends are discussed, highlighting that elements in group one, such as alkali metals, have a low ionization energy because they readily lose an electron to achieve a stable electron configuration similar to noble gases.

This paragraph delves into the trends of ionization energy as one moves across and down the periodic table. It explains that ionization energy generally increases from left to right and from bottom to top, indicating that it becomes progressively harder to remove an electron from an atom. Alkali metals, positioned on the left side, have low ionization energies because they easily lose an electron to achieve a noble gas configuration. Conversely, noble gases on the right side have high ionization energies due to their stable electron configurations. The paragraph also discusses the impact of additional electron shells on ionization energy, noting that electrons in outer shells are less tightly bound and thus easier to remove, as exemplified by cesium compared to lithium. The video script concludes by mentioning that even large atoms like radon, despite being noble gases, can have lower ionization energies than smaller atoms like hydrogen, due to the increased distance of their outermost electrons from the nucleus.