Short answers, pulled from the story.
Gilbert N. Lewis published his paper on the electron-pair bond in 1916. This publication fundamentally changed how humanity understands the universe by proposing that atoms share electrons rather than merely touching.
Richard Abegg proposed a rule in 1904 and Walther Kossel expanded upon it in 1916 to describe ionic bonding as the complete transfer of electrons. Covalent bonding involves the sharing of electron pairs between atoms, whereas ionic bonding creates electrostatic attraction between positive and negative ions.
Metallic bonding occurs when atoms donate electrons to a sea of electrons that flows freely through a lattice of atoms. This delocalization allows metals to exhibit high electrical and thermal conductivity while remaining malleable enough to be hammered into sheets or drawn into wires.
Øyvind Burrau derived the first mathematically complete quantum description of a simple chemical bond in 1927. Shortly thereafter, Walter Heitler and Fritz London developed the Heitler-London method which forms the basis of valence bond theory.
The strength of a chemical bond is determined by the electronegativity difference between the constituent elements. Bond strength is measured in kilojoules per mole, with values ranging from 146 kJ/mol for an oxygen-oxygen single bond to 1,072 kJ/mol for a carbon-oxygen triple bond.
Intermolecular forces including van der Waals forces and hydrogen bonds cause molecules to attract or repel each other without forming a permanent bond. These weak forces are responsible for the high boiling points of water and ammonia as well as the unique properties of biological molecules like DNA.