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Why do reactions between ions often occur faster than reactions that break covalent bonds?

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A graph showing a potential energy curve for a chemical reaction, with the y-axis labeled potential energy and the x-axis labeled progress of reaction. The curve rises from the initial state (reactants) to a peak labeled transition state (activated complex), indicating activation energy, then drops to the final state (products). — chemical kinetics In the Arrhenius equation, the activation energy (E) represents the minimum amount of energy required to transform reactants into products in a chemical reaction. On a potential energy curve, the value of the activation energy is equivalent to the difference in potential energy between particles in an intermediate configuration (known as the activated complex, or transition state) and particles of reactants in their initial state. The activation energy thus can be visualized as a barrier that must be overcome by reactants before products can be formed.

Reactions between ions often occur faster than those involving the breaking of covalent bonds because of the fundamental differences in their reaction mechanisms and energy requirements.

In collision theory, a chemical reaction requires molecules to collide with sufficient kinetic energy to overcome an activation energy barrier. For reactions that involve breaking covalent bonds, this process is often slow because it requires significant energy to disrupt the shared electron pairs that hold the atoms together.12

Conversely, reactions between ions are often driven by electrostatic forces of attraction between oppositely charged species. Because these ions are already charged, they do not necessarily require the same high-energy bond-breaking steps to initiate the reaction. Instead, they can interact rapidly upon collision, often involving simple rearrangements or electron transfers that bypass the more complex, energy-intensive bond-breaking processes required for covalent molecules.34