Electron Configurations | General Chemistry

Dr. Bedard(Ph.D.) provides an in-depth explanation of electron configurations, emphasizing the Pauli Exclusion Principle, which dictates that each suborbital can hold a maximum of two electrons with opposite spins. This concept is visualized with half arrows pointing in opposite directions within a box representing each suborbital. As atomic numbers increase, each element gains an additional electron, affecting how electrons fill the orbitals—s, p, d, and f—each capable of holding 2, 6, 10, and 14 electrons respectively. The discussion spans from hydrogen, which fills the 1s suborbital, to neon, which completes the 2p orbitals. It introduces Hund's rule with carbon, where electrons fill degenerate orbitals singly as much as possible. The narrative then explains the use of condensed electron configurations, which simplify the representation of electron arrangements using noble gas notation, particularly useful for elements further down the periodic table. It also addresses the lower energy of the 4s orbital compared to the 3d, affecting elements in the fourth periodic row and beyond. Lastly, the placement of f-block elements at the bottom of the periodic table is explained as a design choice to maintain the table’s clarity and manageability, highlighting their unique properties and ensuring the table remains functionally concise.

Unit 6.5 - Electron Configurations(Atoms)

For more information please visit http://nocollegeneeded.org

#nocollegeneeded #science #chemistry #chemistryrocks #ncn #generalchemistry #collegechemistry #chemistrytips #chemistryhelp #stemeducation #education #stem #scienceexplained #learning #scienceeducation #sciencefacts

#electrons #orbitals #quantum #hydrogen #helium #hundsrule #noblegases #periodictable #atomicstructure #atomic #chemistry101 #quantumphysics #electronconfiguration
#elementaryparticles #chemicalbonds #learnchemistry #periodictrends #atomicnumber #energylevels #visuallearning