Catalyzed reactions are chemical reactions that happen faster because a catalyst gives them a lower-activation-energy pathway. In General Chemistry II, you use them to explain reaction rate, equilibrium time, and enzyme behavior.
Catalyzed reactions are reactions that proceed faster because a catalyst gives the reactants an easier path to products. In General Chemistry II, that means the catalyst changes the rate of the reaction, not the final equilibrium position. The reactants and products still have the same equilibrium relationship, but they get there more quickly.
The main idea is activation energy. Even if a reaction is thermodynamically allowed, the molecules may need a big energy push to reach the transition state. A catalyst lowers that barrier by offering a different mechanism, so more collisions lead to successful product formation at the same temperature.
A catalyst does not get used up overall. It may form temporary intermediates, bind to reactants, or provide a surface where molecules can react, but it is regenerated by the end of the mechanism. That is why a small amount of catalyst can speed up a large amount of reaction.
This is also why catalysts do not change equilibrium constants. In an equilibrium problem, the value of K depends on the relative energies of reactants and products, not on the speed of the path between them. A catalyst speeds up both the forward and reverse reactions, so equilibrium is reached sooner, but the ratio of concentrations at equilibrium stays the same.
Chemists often talk about catalyzed reactions in two big categories. Homogeneous catalysis means the catalyst is in the same phase as the reactants, like a dissolved acid catalyst in solution. Heterogeneous catalysis means the catalyst is in a different phase, such as a solid surface that adsorbs gases or liquids and helps them react. In both cases, the big idea is the same: lower activation energy, faster reaction, same equilibrium.
Enzymes are the biology version of this idea, but the chemistry is the same. They speed up metabolic reactions by binding specific substrates and stabilizing the transition state. That makes catalyzed reactions one of the clearest places where kinetics and equilibrium connect in Gen Chem II.
Catalyzed reactions connect directly to the parts of General Chemistry II where rate and equilibrium overlap. When you study equilibrium constants, you need to know why a reaction can have a perfectly valid K value and still be too slow to matter without a catalyst. That comes up in reaction mechanisms, lab work, and industrial chemistry.
This term also helps you separate thermodynamics from kinetics. A reaction may be spontaneous in the sense of having a favorable ΔG, but it can still crawl along if the activation energy is high. A catalyst fixes the speed problem, not the energy difference between reactants and products. That distinction shows up often in problem sets because it is easy to confuse “can happen” with “happens quickly.”
Catalyzed reactions also explain real-world processes students usually see in examples: enzymes in biology, catalytic converters in cars, and surface-catalyzed reactions in industry. Those examples make equilibrium less abstract because you can trace how a mechanism changes the pathway without changing the end result.
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Visual cheatsheet
view galleryCatalyst
A catalyst is the substance that makes the reaction faster, and catalyzed reactions are the reactions happening with that substance present. In Gen Chem II, you often describe the catalyst by what it does to the mechanism, especially whether it is homogeneous, heterogeneous, or an enzyme. The catalyst itself comes out regenerated, so it is not consumed overall.
Activation Energy
Catalyzed reactions work because the catalyst lowers activation energy. That means fewer molecules need enough energy to reach the transition state, so the reaction rate increases. On free-energy diagrams, you still see the same starting and ending energy levels, but the peak in the middle is lower.
Equilibrium
Catalysts do not change the position of equilibrium. They speed up both the forward and reverse reactions, which means the system reaches equilibrium faster but still settles at the same concentrations or pressures. This is why catalyzed reactions matter so much in equilibrium problems and industrial synthesis.
Temperature
Temperature and catalysts both affect reaction rate, but they do it in different ways. Raising temperature gives particles more kinetic energy and increases the fraction of successful collisions, while a catalyst lowers the activation barrier. In a problem set, you may need to tell whether a faster reaction came from heating the system or adding a catalyst.
A quiz or problem set will usually ask you to read a reaction profile, compare catalyzed and uncatalyzed pathways, or explain why equilibrium is reached faster without changing K. You may also be asked to identify whether a change in rate came from a catalyst, a temperature increase, or a concentration change. The move is to connect the observation to activation energy and to say whether the equilibrium position changes.
If you see a mechanism question, look for temporary catalyst steps, adsorption on a surface, or an enzyme-substrate complex. If the prompt mentions rate but not final amounts, the answer is probably about kinetics. If it asks about equilibrium concentrations, the catalyst should not change the calculation itself.
Catalysts and temperature both speed up reactions, but they are not the same change. Temperature gives particles more kinetic energy and changes the collision distribution, while a catalyst lowers the activation energy through an alternate pathway. In equilibrium questions, temperature can shift K for some reactions, but a catalyst does not.
Catalyzed reactions are faster because a catalyst lowers the activation energy needed for the reaction to proceed.
A catalyst is not used up overall, because it is regenerated after the mechanism finishes.
Catalysts change the rate at which equilibrium is reached, but they do not change the equilibrium constant or the equilibrium concentrations.
On an energy diagram, a catalyzed pathway has a lower peak than the uncatalyzed pathway, but the reactant and product energy levels stay the same.
In General Chemistry II, catalyzed reactions often show up in kinetics, equilibrium, enzyme examples, and reaction mechanism questions.
Catalyzed reactions are reactions that happen faster because a catalyst provides a lower-activation-energy pathway. In General Chemistry II, the big point is that the catalyst changes the reaction rate, not the final equilibrium position. You usually connect this idea to energy diagrams and mechanism steps.
No. A catalyst speeds up both the forward and reverse reactions, so the system reaches equilibrium faster, but the equilibrium constant stays the same. If a problem asks about final concentrations at equilibrium, the catalyst should not change the answer.
Look for a substance that appears in the mechanism, lowers the activation barrier, and is regenerated by the end. In lab or textbook examples, this might be a solid surface, a dissolved acid, or an enzyme. If the prompt shows a lower peak on a reaction coordinate diagram, that is another clue.
A catalyst changes the pathway, while temperature changes how much kinetic energy the particles have. Both can increase rate, but only temperature can also affect equilibrium constants for some reactions. That makes them different tools in kinetics and equilibrium problems.