In simple terms
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Catalysts: The Reaction Accelerators
Catalysts are chemical matchmakers that speed up reactions by providing an easier route, without getting used up themselves. They come in two main types: homogeneous (in the same state as reactants) and heterogeneous (in a different state).
Imagine you need to get to a village on the other side of a tall mountain. Walking over the top requires a huge amount of energy (high activation energy). A catalyst is like a tunnel through the mountain. It provides an alternative, much easier path (lower activation energy), so you can get to the village much faster. The tunnel itself isn't changed by you passing through it, and the starting and ending points (reactants and products) are at the same heights as before, so the overall height difference (enthalpy change) is the same.
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A catalyst provides an alternative reaction pathway with a lower activation energy (), increasing the reaction rate. It is not consumed in the overall reaction.
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In homogeneous catalysis, the catalyst is in the same physical state (phase) as the reactants, often forming an intermediate species.
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In heterogeneous catalysis, the catalyst is in a different phase from the reactants. The reaction occurs on the catalyst's surface via adsorption.
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A catalyst does not change the overall enthalpy change () or the position of equilibrium (). It increases the rate of both the forward and reverse reactions equally.
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Full topic notes
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The Role of a Catalyst
A catalyst is a substance that increases the rate of a chemical reaction but is not consumed in the overall process. It achieves this by providing an alternative reaction pathway with a lower activation energy (). Think of it as creating a shortcut for the reaction. Because more reactant particles will have energy equal to or greater than this new, lower activation energy, the frequency of effective collisions increases, leading to a faster rate of reaction.
A catalyst provides an alternative reaction route with a lower activation energy ().
It is not used up in the overall reaction; it is regenerated.
It does not affect the enthalpy change () of the reaction.
It does not change the position of equilibrium but allows it to be reached more quickly.
Homogeneous Catalysis
In homogeneous catalysis, the catalyst and the reactants are in the same physical state, or phase. For example, a reaction between two aqueous solutions might be catalysed by another substance dissolved in the same solution. The mechanism typically involves the catalyst reacting with one of the reactants to form an intermediate. This intermediate then goes on to react with the other reactant, forming the final product and regenerating the catalyst. Transition metals are often excellent homogeneous catalysts because their ability to exist in variable oxidation states allows them to be easily oxidised and reduced during the catalytic cycle.
Heterogeneous Catalysis
In heterogeneous catalysis, the catalyst is in a different phase from the reactants. Most commonly, a solid catalyst is used for reactions involving gases or liquids. The process occurs on the surface of the catalyst at locations called active sites. The mechanism can be broken down into three main stages: adsorption of reactants onto the surface, reaction on the surface, and desorption of products from the surface. The effectiveness of a heterogeneous catalyst is highly dependent on its surface area; a larger surface area means more active sites are available.
Haber Process: Solid iron (Fe) catalyses the reaction between N₂(g) and H₂(g) to form NH₃(g).
Contact Process: Solid vanadium(V) oxide (V₂O₅) catalyses the oxidation of SO₂(g) to SO₃(g).
Catalytic Converters: A ceramic honeycomb coated with platinum (Pt), palladium (Pd), and rhodium (Rh) catalyses the conversion of toxic exhaust gases (CO, NOx, unburnt hydrocarbons) into less harmful substances (CO₂, N₂, H₂O).
Worked examples
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The reaction between peroxodisulfate(VI) ions and iodide ions is very slow: S₂O₈²⁻(aq) + 2I⁻(aq) → 2SO₄²⁻(aq) + I₂(aq). The reaction is catalysed by iron(II) ions, Fe²⁺(aq). Explain, with the aid of equations, the catalytic role of Fe²⁺(aq).
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The uncatalysed reaction is slow because it involves a collision between two negatively charged ions (S₂O₈²⁻ and I⁻), leading to strong electrostatic repulsion which results in a high activation energy.
In the Contact process, vanadium(V) oxide, V₂O₅, is used as a heterogeneous catalyst to oxidise sulfur dioxide. The overall equation is: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g). Write two equations to show how V₂O₅ acts as a catalyst.
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Vanadium(V) oxide is a solid, while the reactants are gases, so this is heterogeneous catalysis. The catalyst works by being reduced and then re-oxidised.
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What is a catalyst?
A substance that increases the rate of a chemical reaction by providing an alternative pathway with a lower activation energy, without being chemically changed at the end of the reaction.
Key takeaways
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A catalyst provides an alternative reaction route with a lower activation energy ().
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It is not used up in the overall reaction; it is regenerated.
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It does not affect the enthalpy change () of the reaction.
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It does not change the position of equilibrium but allows it to be reached more quickly.
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