15B.29-15B.32 - Heterogeneous Catalysis

Why Transition Metals Make Useful Catalysts

A gas-phase reaction can be slow even when it is thermodynamically possible. In an industrial reactor or a car exhaust, that slow rate matters: unreacted gases keep flowing past, useful product is not made quickly enough, or harmful emissions leave before they can react.

Transition metals and their compounds are often useful catalysts because they can interact with reactant molecules without being used up overall. Two features are especially important in this lesson:

• many transition metals can adsorb gases onto their surfaces
• many transition-metal compounds can change oxidation state and then be regenerated

The specification names both heterogeneous and homogeneous catalysts.

Heterogeneous does not mean "unrelated"

For a solid catalyst and gas reactants (catalyst solid, reactants gases), the phases are different. That difference in phase is not just a label. It tells you where the reaction happens: at the exposed surface of the solid catalyst. The bulk of the solid is not the important region; the accessible surface atoms and active sites are.

Adsorption, Bond Weakening And Desorption

The core model for heterogeneous catalysis is a surface pathway:

[DIAGRAM: asset_name: Adsorption surface reaction and product desorption; asset_slug: edexcel_a_level_chemistry_l074_catalyst_surface; recommended_method: deterministic_chart; description: Adsorption surface reaction and product desorption.]

gas reactants adsorb → bonds weaken and surface reaction occurs → products desorb

Adsorption is not the same as absorption. In catalytic converters and many industrial reactors, adsorption is temporary and useful: it brings reactant molecules onto active sites where reaction is easier.

A good catalyst binds reactants strongly enough to weaken bonds, but not so strongly that products remain stuck on the surface. If a product or impurity binds too strongly, it can block active sites and reduce catalytic activity.

Explanation route

When asked to explain a heterogeneous catalyst, use a clear chain:

different phase → surface active sites → adsorption → bond weakening/surface reaction → desorption → catalyst regenerated

A catalyst provides an alternative route with lower activation energy; it does not supply energy. At a fixed temperature it speeds forward and reverse routes without changing the equilibrium composition. This separates reaching equilibrium faster from changing the eventual yield.

Vanadium Oxide In The Contact Process

In the Contact process, sulfur dioxide is oxidised to sulfur trioxide:

$\ce{2SO2(g) + O2(g) -> 2SO3(g)}$

The catalyst is solid vanadium(V) oxide, $\ce{V2O5}$. The reacting gases and the solid catalyst are in different phases, so this is heterogeneous catalysis. The role of $\ce{V2O5}$ is explained in terms of oxidation number.

Oxidation numbers first

Use oxygen as $-2$ in these oxides.

For $\ce{V2O5}$:

$5(-2) = -10$, so the two vanadium atoms total $+10$: vanadium is $+5$.

For $\ce{V2O4}$:

$4(-2) = -8$, so the two vanadium atoms total $+8$: vanadium is $+4$.

For sulfur:

sulfur is $+4$ in $\ce{SO2}$ and $+6$ in $\ce{SO3}$.

So sulfur is oxidised from $+4$ to $+6$ when $\ce{SO2}$ becomes $\ce{SO3}$.

The catalytic cycle

One useful way to show the catalyst is:

In the first step, $\ce{V2O5}$ oxidises $\ce{SO2}$ to $\ce{SO3}$. Vanadium is reduced from $+5$ in $\ce{V2O5}$ to $+4$ in $\ce{V2O4}$, while sulfur is oxidised from $+4$ in $\ce{SO2}$ to $+6$ in $\ce{SO3}$.