12.7 Catalysis

Catalysis

Catalysts speed up chemical reactions by offering an alternative pathway with lower activation energy. They don't get consumed in the process, and they don't change the final products or equilibrium of a reaction. Understanding catalysis ties together several kinetics concepts: activation energy, reaction rates, and collision theory.

Catalysts and Activation Energy

Every reaction has an energy barrier that reactants must overcome to form products. That barrier is the activation energy ($E_a$), the minimum energy needed to reach the transition state. A catalyst lowers this barrier by providing a different reaction pathway.

Why does a lower $E_a$ speed things up? At any given temperature, molecules have a range of kinetic energies. When you lower the activation energy, a larger fraction of molecules now have enough energy to react when they collide. More successful collisions per second means a faster rate.

A few critical points about catalysts:

• They are not consumed by the reaction. A catalyst participates in the mechanism but is regenerated at the end of each cycle, so it can be reused.
• They do not change the equilibrium constant ($K_{eq}$) or the overall energy change ($\Delta G$) of the reaction. The thermodynamics stay the same.
• They speed up both the forward and reverse reactions equally, so the system reaches equilibrium faster but at the same position.

Homogeneous vs. Heterogeneous Catalysis

The two main categories of catalysis are based on whether the catalyst is in the same phase as the reactants.

Homogeneous catalysis means the catalyst and reactants are in the same phase, most often all dissolved in a liquid solution.

• Enzymes are biological homogeneous catalysts. For example, pepsin breaks down proteins in the aqueous environment of your stomach.
• Acid catalysis is common in organic reactions. Sulfuric acid ($H_2SO_4$) catalyzes esterification reactions, such as producing methyl salicylate (oil of wintergreen) from methanol and salicylic acid.

Heterogeneous catalysis means the catalyst is in a different phase from the reactants, typically a solid catalyst interacting with liquid or gas reactants.

• Catalytic converters in cars use solid metals like platinum, palladium, and rhodium to convert toxic exhaust gases (carbon monoxide, nitrogen oxides) into less harmful substances.
• The Haber-Bosch process uses a solid iron ($Fe$) catalyst to combine gaseous $N_2$ and $H_2$ into ammonia ($NH_3$), one of the most important industrial reactions in the world.
• Hydrogenation of vegetable oils uses solid nickel ($Ni$) catalysts to add hydrogen across double bonds in unsaturated fats, producing saturated fats like margarine.

Steps in Heterogeneous Catalysis

Heterogeneous catalysis follows a three-step cycle on the catalyst surface:

1. Adsorption: Reactant molecules bind to the surface of the solid catalyst. This can happen through weak van der Waals attractions (physisorption) or through actual chemical bond formation (chemisorption). Either way, adsorption concentrates the reactants near each other on the surface.

2. Surface reaction: The adsorbed reactants react at specific locations called active sites. The catalyst surface lowers $E_a$ in several ways:

   • It holds reactants in the right orientation for a successful collision
   • It forms temporary bonds with reactants, which weakens bonds within the reactant molecules and makes them more reactive
   • It can facilitate electron transfer between reactants

3. Desorption: The product molecules detach from the catalyst surface. This frees up the active sites so new reactant molecules can adsorb, and the cycle repeats.

The catalyst surface is regenerated after each cycle, which is why a small amount of catalyst can process a large amount of reactant over time.

Catalyst Efficiency and Selectivity

• Turnover number measures how many reactant molecules a single catalyst molecule (or active site) can convert before it becomes inactive. A higher turnover number means a more efficient catalyst.
• Selectivity describes a catalyst's ability to favor one specific product when multiple products are possible. This matters a lot in industrial chemistry, where you want to maximize the yield of the desired product and minimize waste.
• An inhibitor (sometimes called a catalyst poison) is a substance that reduces a catalyst's effectiveness. It can do this by blocking active sites or by interfering with the reaction mechanism. For example, lead poisons the platinum catalyst in a catalytic converter, which is one reason leaded gasoline was phased out.