Atoms naturally seek to lower their overall potential energy by rearranging their outermost electrons. They achieve this lower energy state by forming different types of chemical bonds. In covalent bonds, nonmetal atoms share pairs of valence electrons to achieve chemical stability.
In ionic bonds, atoms completely transfer electrons to create oppositely charged atomic ions. These oppositely charged ions then strongly attract each other to form a rigid crystal lattice. In metallic bonds, valence electrons roam freely across a large network of metal cations.
The specific type of bond completely determines a substance’s melting point, hardness, and chemical reactivity. Stronger chemical bonds generally result in much higher melting points and harder solid materials. Chemists accurately measure the strength of a bond using a concept known as bond energy.
Bond length describes the average physical distance between the nuclei of two bonded atoms. Bond polarity occurs when two bonded atoms share their valence electrons unequally between them. Shorter chemical bonds with higher bond orders are generally much stronger and harder to break.
Many students mistakenly believe that breaking chemical bonds releases energy into the environment. In reality, breaking any chemical bond always requires an input of outside energy. Energy is only released when new chemical bonds form during a chemical reaction.