Chemistry (9701) | Homogeneous and heterogeneous catalysts

AS/A Level Chemistry (9701) Study Notes: Catalysis

Introduction: Why Catalysts Matter

Welcome! This topic explores catalysis, which is crucial for understanding how we speed up reactions both in the lab and in industry. Catalysts are amazing because they help reactions happen much faster and often at lower temperatures, saving huge amounts of energy and money! When you understand the difference between homogeneous and heterogeneous catalysts, you unlock a key part of reaction kinetics.

Remember from Reaction Kinetics (Section 8.3) that a catalyst works by providing an alternative reaction pathway with a significantly lower activation energy (\(E_a\)).

Quick Review: The Effect on \(E_a\) and the Boltzmann Distribution

The rate of reaction increases because lowering the activation energy dramatically increases the number of particles that possess energy greater than or equal to \(E_a\). This is best visualized using the Boltzmann distribution curve:

• Without a catalyst, only a small fraction of molecules have enough energy (\(E_a\)).
• With a catalyst, the required energy barrier is lowered (\(E_{a, \text{cat}}\)), meaning a much larger proportion of molecules now have enough energy to react effectively.

Important Point: A catalyst speeds up the forward and reverse reactions equally. It does not change the position of equilibrium or the equilibrium constant (\(K_c\) or \(K_p\)).

1. Heterogeneous Catalysis (Different Phases)

1.1 Defining Heterogeneous Catalysis

A heterogeneous catalyst is one where the catalyst and the reactants are in different physical states or phases. The most common scenario is a solid catalyst reacting with gaseous or liquid reactants.

Analogy: Think of the catalyst surface like a busy car park. The reactants (cars) must land on an empty space (active site) to be assembled into the product (new cars) before driving away.

1.2 The Mechanism: The Three Core Steps

The process of heterogeneous catalysis happens entirely on the surface of the solid catalyst. It involves three key stages: Adsorption, Reaction, and Desorption.

Step 1: Adsorption

• Reactant molecules physically attach to the surface of the solid catalyst. This happens at specific points called active sites.
• This binding is often a form of chemisorption (chemical adsorption), where weak bonds form between the reactant molecules and the atoms on the catalyst surface.

Step 2: Reaction (Bond Weakening)

• The bonds that form between the reactant molecules and the catalyst atoms temporarily weaken the internal bonds within the reactant molecules.
• This weakening makes it easier for the reactants to collide, break their original bonds, and form new products. This is the stage where the activation energy is lowered.

Step 3: Desorption

• Once the product molecules form, they detach (desorb) from the active sites and diffuse away from the catalyst surface.
• The active sites are then ready for the next reactant molecules.

Key Takeaway: The magic happens in Step 2, where the catalyst surface acts like a stress point, pulling apart reactant bonds and holding them in the perfect position to react with their collision partner.

1.3 Real-World Examples of Heterogeneous Catalysis

Example 1: Iron in the Haber Process

The Haber process is used to manufacture ammonia (\(\text{NH}_3\)):

\( \text{N}_2(\text{g}) + 3\text{H}_2(\text{g}) \rightleftharpoons 2\text{NH}_3(\text{g}) \)

• Catalyst: Solid Iron (Fe).
• Mechanism Role: Both nitrogen (\(\text{N}_2\)) and hydrogen (\(\text{H}_2\)) molecules adsorb onto the iron surface. The strong triple bond in \(\text{N}_2\) and the single bond in \(\text{H}_2\) are weakened (or sometimes broken entirely) by this adsorption, allowing them to react much more readily than they would in the gas phase.

Example 2: Platinum, Palladium, and Rhodium in Catalytic Converters

Catalytic converters in cars reduce harmful pollutants like carbon monoxide (\(\text{CO}\)) and nitrogen oxides (\(\text{NO}_x\)) into less harmful gases (\(\text{CO}_2\) and \(\text{N}_2\)).

• Catalyst: A honeycomb coated with Platinum (Pt), Palladium (Pd), and Rhodium (Rh) (all solid metals).
• Reaction: Pollutant gases flow over the solid metal surface. For example:
  \( 2\text{CO}(\text{g}) + 2\text{NO}(\text{g}) \xrightarrow{\text{Pt/Pd/Rh}} 2\text{CO}_2(\text{g}) + \text{N}_2(\text{g}) \)
• Why it works: The metals adsorb the $\text{CO}$ and $\text{NO}$ molecules, holding them in position to react with each other efficiently, reducing the environmental impact of car exhaust.

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2. Homogeneous Catalysis (Same Phase)

2.1 Defining Homogeneous Catalysis

A homogeneous catalyst is one where the catalyst and the reactants are in the same physical state, usually as a liquid or gas mixture (or aqueous solutions).

Analogy: Unlike the car park (heterogeneous), homogeneous catalysis is like a matchmaker (the catalyst) moving between two separate groups (reactants) and helping them pair up by forming a temporary relationship.

2.2 The Mechanism: Intermediates and Regeneration

Homogeneous catalysts work via a multi-step reaction mechanism where the catalyst is used up in an early step, forming an intermediate compound, and then reformed in a later step. This ensures the catalyst is not consumed overall.

General Mechanism:

1. Step 1 (Usage): Reactant A reacts with Catalyst (C) to form an intermediate (I).
   \( \text{A} + \text{C} \rightarrow \text{I} \)
2. Step 2 (Regeneration): The intermediate (I) reacts with Reactant B, generating the Product (P) and reforming the original Catalyst (C).
   \( \text{I} + \text{B} \rightarrow \text{P} + \text{C} \)

The new pathway (via the intermediate) has a lower \(E_a\) than the direct, uncatalysed reaction (\( \text{A} + \text{B} \rightarrow \text{P} \)).

2.3 Real-World Examples of Homogeneous Catalysis

Example 1: Atmospheric Oxides of Nitrogen (\(\text{NO}/\text{NO}_2\))

Oxides of nitrogen, often produced by lightning or high-temperature combustion, can catalyse the oxidation of atmospheric sulfur dioxide (\(\text{SO}_2\)) to sulfur trioxide (\(\text{SO}_3\)) which contributes to acid rain.

Overall reaction (Uncatalysed, very slow):
\( 2\text{SO}_2(\text{g}) + \text{O}_2(\text{g}) \rightarrow 2\text{SO}_3(\text{g}) \)

Catalytic Cycle (Using \(\text{NO}_2\) and \(\text{NO}\) as intermediates/catalyst):

Step 1 (Usage): The $\text{SO}_2$ reacts with $\text{NO}_2$, transferring oxygen to form the product, $\text{SO}_3$, and reducing the catalyst to $\text{NO}$ (nitric oxide).

\( \text{SO}_2(\text{g}) + \text{NO}_2(\text{g}) \rightarrow \text{SO}_3(\text{g}) + \mathbf{NO}(\text{g}) \)

Step 2 (Regeneration): The $\text{NO}$ immediately reacts with atmospheric oxygen to reform the original catalyst, $\text{NO}_2$.

\( 2\mathbf{NO}(\text{g}) + \text{O}_2(\text{g}) \rightarrow 2\mathbf{NO}_2(\text{g}) \)

The $\text{NO}$ and $\text{NO}_2$ cycle rapidly, speeding up the overall conversion.

Example 2: Transition Metal Ions (\(\text{Fe}^{2+}\) or \(\text{Fe}^{3+}\))

Transition metal ions are excellent homogeneous catalysts because they have variable oxidation states (Section 28.1). They can be easily oxidised or reduced, making them perfect intermediates.

Consider the reaction between peroxodisulfate ions (\(\text{S}_2\text{O}_8^{2-}\)) and iodide ions (\(\text{I}^-\)):

Overall Reaction: \( \text{S}_2\text{O}_8^{2-}(\text{aq}) + 2\text{I}^-(\text{aq}) \rightarrow 2\text{SO}_4^{2-}(\text{aq}) + \text{I}_2(\text{aq}) \)

This reaction is very slow uncatalysed because both reactants are negatively charged and repel each other.

Catalytic Cycle (Using \(\text{Fe}^{2+}\) ions):

Step 1 (Oxidation of Catalyst): The $\text{Fe}^{2+}$ reduces the highly oxidising $\text{S}_2\text{O}_8^{2-}$, being oxidised itself to $\text{Fe}^{3+}$.

\( \text{S}_2\text{O}_8^{2-}(\text{aq}) + 2\mathbf{Fe^{2+}}(\text{aq}) \rightarrow 2\text{SO}_4^{2-}(\text{aq}) + 2\mathbf{Fe^{3+}}(\text{aq}) \)

Step 2 (Regeneration of Catalyst): The $\text{Fe}^{3+}$ then oxidises the $\text{I}^-$ ions, converting back to $\text{Fe}^{2+}$.

\( 2\mathbf{Fe^{3+}}(\text{aq}) + 2\text{I}^-(\text{aq}) \rightarrow 2\mathbf{Fe^{2+}}(\text{aq}) + \text{I}_2(\text{aq}) \)

Since $\text{Fe}^{2+}$ is regenerated, it acts as a catalyst. The benefit is that the steps involving the iron ions have a much lower \(E_a\) than the original reaction.

(Note: You could also start the cycle with $\text{Fe}^{3+}$ ions, which would be reduced in Step 1 and regenerated (oxidised) in Step 2.)

3. Comparison and Review

3.1 Summary of Catalyst Types

	Heterogeneous Catalyst	Homogeneous Catalyst
Phase	Different phases (e.g., solid catalyst, gaseous reactants)	Same phase (e.g., aqueous catalyst, aqueous reactants)
Mode of Action	Surface interaction (Adsorption, bond weakening, desorption)	Forms an intermediate that is then consumed and reformed
Example	Iron in Haber process	\(\text{Fe}^{2+}/\text{Fe}^{3+}\) ions in redox reactions
Advantages	Easily separated from products; reusable.	High activity; active sites are more available.

3.2 Sketching Reaction Pathway Diagrams

The syllabus requires you to construct and interpret reaction pathway diagrams (enthalpy profiles) for catalysed and uncatalysed reactions (Section 8.3/26.2).

• Both pathways start at the same reactant energy level and end at the same product energy level. This means the overall enthalpy change (\(\Delta H\)) is identical for both pathways.

• The uncatalysed pathway shows one large energy hump representing the high original activation energy (\(E_a\)).

• The catalysed pathway shows a pathway with two or more smaller energy humps, representing the intermediate steps (formation and breakdown of the intermediate/surface complex). The highest point of this new pathway is significantly lower than the uncatalysed \(E_a\).

(Imagine drawing an exothermic reaction: the products are lower than the reactants. The catalysed route simply has a shorter hill to climb, achieved in multiple stages.)

Key Takeaway: Catalysts do not change the start or end points (thermodynamics); they only change the route taken (kinetics).