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Catalysis is the process of increasing the rate of a chemical reaction by adding a substance known as a 触媒. Catalysts are not consumed in the reaction and remain unchanged. They work by providing an alternative reaction pathway with a lower activation energy ([latex]E_a[/latex]), thereby accelerating both the forward and reverse reactions without altering the overall thermodynamics ([latex]\Delta H[/latex]).
The fundamental principle of catalysis lies in its ability to alter the kinetics of a chemical reaction without affecting its thermodynamics. A 触媒 introduces a new reaction mechanism, often involving one or more intermediate steps. For a reaction A + B → C, a catalyst C’ might participate as follows: A + C’ → AC’ and then AC’ + B → C + C’. The catalyst C’ is regenerated at the end of the process. This alternative pathway has a lower transition state energy compared to the uncatalyzed reaction. The Arrhenius equation, [latex]k = Ae^{-E_a/(RT)}[/latex], shows that a lower activation energy ([latex]E_a[/latex]) leads to an exponential increase in the reaction rate constant (k). Importantly, the catalyst does not change the Gibbs free energy change ([latex]\Delta G[/latex]) or the equilibrium constant ([latex]K_{eq}[/latex]) of the reaction. It only affects the speed at which equilibrium is reached. This concept was first formally described by Jöns Jacob Berzelius in 1835, who observed that certain substances could accelerate reactions without being consumed, coining the term “catalysis” from Greek words meaning “to dissolve” or “to break down”.
この原理は反応座標図を用いて視覚化され、触媒反応経路では非触媒反応経路よりも低いエネルギーピーク（遷移状態）が示されます。反応物と生成物の間の全体的なエネルギー差は同じままですが、乗り越えるべきエネルギー障壁は大幅に低下します。これにより、より多くの反応物分子が衝突時に反応するのに十分なエネルギーを持つようになり、一定の温度において反応速度が速くなります。
触媒作用
（日付が不明または関連性がない場合、例えば「流体力学」などでは、その注目すべき出現時期の概算値が提示されます。）
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