And so, ( \text{SO}_4^{2-} ) was born. It looked like a perfect pyramid: Sulfur in the center, four Oxygens at the points.
He formed four double bonds (S=O). But to the Electron Geometry, those double bonds count as just of electron density each. So, looking at the electron clouds only: Sulfur had four regions of high electron density pushing away from him. so4 2 electron geometry and molecular geometry
But here was the twist. Because the ion had a ( 2- ) charge, the Oxygens were slightly jealous—they wanted even more negative attention. So they began to delocalize . The double bonds started switching places so fast that, if you looked at the molecule, every bond looked identical: 1.5 bonds (a resonance hybrid). And so, ( \text{SO}_4^{2-} ) was born
"No lone pairs to hide," Sulfur said. "What you see is what you get." But to the Electron Geometry, those double bonds
Sulfur made a decision. He would use his d-orbital expansion. He promoted one of his 3s electrons to a higher energy level, creating six unpaired electrons. Then, he borrowed two extra electrons from the universe (giving the ion its ( 2- ) charge). Now, with eight electrons to allocate, he invited the four Oxygens to bond.
Deep in the valley of the Periodic Table lived a large, charismatic atom named Sulfur. Sulfur was unique. Unlike his neighbor, the rigid Carbon, Sulfur had an expanded wardrobe—empty d-orbitals that allowed him to dress up in more than eight electrons. Today, Sulfur faced a dilemma. He had four Oxygen atoms asking for his attention. Each Oxygen needed two electrons to complete its own valence shell.
The four Oxygens stood at the corners of a tetrahedron, repelling each other equally. The molecule was symmetric, stable, and perfectly non-polar in its charge distribution, despite carrying a net negative charge.