What are the 3 types of catalysis?

The three types of catalysis are: Homogeneous catalysis Heterogeneous catalysis Autocatalysis

How do catalysts speed up reactions?

Catalysts speed up a reaction by lowering the required activation energy.

What is a heterogeneous catalyst?

A heterogeneous catalyst is in a different phase from the reactants. The reaction occurs at active sites on the catalyst surface.

What are the steps of surface adsorption theory?

Adsorption: the reactants stick to active sites of the catalyst surface. Reaction: the reaction takes place while the reactants are attached to the active sites. Desorption: the products detach from the catalyst surface.

Catalysts: Meaning, Example, Equation & Types

Catalysts are amazing, but how do they work exactly? Let's use Transition Metals to help us explain.

In this article, you'll discover how transition metals act as catalysts.
Then, you'll find out how heterogeneous and homogeneous catalysts work.
You'll then learn about the use of catalysts in industries, specifically in the Haber and Contact processes.
Finally, you'll discover autocatalysis.

Catalysts meaning

When catalysts take part in a chemical process, the reaction produces the same amount of catalyst that was added at the start of the reaction. Catalysts are essential in industry. They help speed up the rate of reaction and lower the activation energy.

Go to Factors Affecting Reaction Rates to learn more about this.

A catalyst is a substance that increases the rate of a chemical reaction without being changed in chemical composition or quantity.

Transition metals make great catalysts because of their variable oxidation states. In some cases, they adsorb substances on their surface and activate them.

So, now you know what catalysts do. But how do they work? Transition metal catalysts can either be heterogeneous or homogeneous. Keep reading to discover what this means!

Heterogeneous Catalyst

A heterogeneous catalyst is in a different phase from the reactants, and the reaction occurs at active sites on the surface.

In general, most heterogeneous catalysts are in the solid phase and do not get consumed in the reaction. At least one of the reactants gets adsorbed at active sites on the catalyst's surface in heterogeneous catalysis.

You may have heard of catalytic converters in cars. They are an example of heterogeneous catalysts at work. In a catalytic converter, a gaseous reactant passes over a solid catalyst.

Adsorption happens when the reactants stick to the catalyst's surface so that the reaction can take place. We call the location on the catalyst's surface where the reactants stick, the active site.

Transition metals are popular solid catalysts. The Haber and Contact processes use transition metal catalysts to increase the rate of reaction. How do they do this? Find out below.

The Haber Process

Factories around the world produce ammonia through the Haber process. We use most of this ammonia in fertilisers which farmers use to grow an abundance of crops. The Haber process produces ammonia through heterogeneous catalytic action.

The Haber process places hydrogen and nitrogen from the air in a reactor between 400-450ºC. The process takes place under a compromise pressure of 200 atm with iron pellets as a catalyst.

The surface of the iron attracts hydrogen and nitrogen gas molecules which become adsorbed onto the surface.
The reaction takes place while the hydrogen and nitrogen molecules are on the surface of the iron.
When ammonia forms, it desorbs from the surface.

Have a look at the chemical equation for the Haber process below.

Catalysts, chemical equation for The Haber process, StudySmarter The Haber process, StudySmarter Originals

You will notice that it is a reversible process. The purpose of iron in the Haber process is to speed up the reaction; it does not affect the equilibrium. The Haber process takes far too long to produce ammonia without the iron catalyst.

German scientist Fritz Haber invented this process in 1908. He revolutionised agriculture with his invention. However, plants do not assimilate half of the nitrogen in fertilisers. As a result of this, we now have high concentrations of ammonia in our water supply and the earth's atmosphere. Today, chemists are searching for another process that could be used to produce abundant crops without endangering our planet.

You can find out about compromise pressure in Le Chatelier's Principle.
Keep reading to learn about adsorption and desorption.

The Contact Process

What do dyes, detergents, paints, plastics, fertiliser and fabrics have in common? Sulfuric acid! Globally we manufacture 231 million tonnes of sulfuric acid every year. Most of the sulfuric acid produced industrially gets made into fertiliser, but we also use it to make paper, pigments, and fibres. How does sulfuric acid get made industrially? We use the Contact process, another example of heterogeneous catalysis.

The Contact process happens in three stages. We will consider the stage where a solid catalyst - vanadium pentoxide (also referred to as vanadium (V) oxide) - gets used to speed up the reaction rate. At this stage, sulfur dioxide reacts with oxygen to produce sulfuric acid. Look at the equation for the reaction below.

2SO₂ + O₂ ⇌ 2SO₃

You will notice that the Contact process is a reversible reaction, like the Haber process. Vanadium pentoxide works to speed up the reaction. Otherwise, the process would be too slow!

Sulfur dioxide and oxygen enter the reactor as gases. In the reactor, they pass over a solid vanadium pentoxide catalyst. We use a mechanism called surface adsorption theory to explain how heterogeneous catalysts work.

First, adsorption happens when one of the reactants attaches to the catalyst surface.
The reaction takes place while the reactants are on the catalyst surface.
Later, desorption happens when the product of the reaction detaches from the catalyst surface.

Let us now examine the reaction stages of the Contact process. You can see surface adsorption theory at work here.

Stage 1: Sulfur dioxide adsorbs onto the vanadium(V) oxide. A redox reaction occurs when vanadium gets reduced from +5 to +4. The sulphur trioxide desorbs.
Stage 2: Another redox reaction occurs when oxygen reacts on the catalyst surface. Vanadium(IV) oxide gets oxidised back to +5. The original catalyst, V₂O₅, is regenerated.

Catalysts, diagram of the Contact process surface adsorption theory, StudySmarter Fig. 2 - The contact process

Surface adsorption theory has helped us design a solution for one of the biggest problems in the 21st century, pollution from vehicle exhaust gases. Find out more in the deep dive below!

Catalytic Converter

Cars produce pollutants that affect the environment and our health negatively. In some cities, the air is so dirty that humans cannot breathe safely. One of these pollutants is carbon monoxide, a toxic gas that restricts oxygen from getting to your vital organs. Another major harmful exhaust gas is nitrogen monoxide. Nitrogen monoxide oxidises easily in the air to nitrogen dioxide, a respiratory irritant that also contributes to acid rain in cities.

To deal with this, for decades now, automakers have fitted car exhausts with catalytic converters to reduce global emissions. Essentially, the inside of a catalytic converter is a ceramic honeycomb coated with a solid metal catalyst. Usually, the coating is a mixture of the transition metals platinum and rhodium. We also sometimes use palladium.

These transition metals act as catalysts through surface adsorption theory. Harmful gases adsorb onto active sites, where they react to produce harmless gases like carbon dioxide and nitrogen. The products desorb and get released through the car exhaust. The honeycomb interior increases the efficiency of the catalytic converter since it provides a greater surface area for the solid catalyst. This means there are more active sites where adsorption can take place. The honeycomb structure also helps to minimise the cost of catalytic converters since platinum and rhodium are expensive metals.

Catalytic converter, catalysts, studysmarter Fig. 2 - A catalytic converter

Catalytic 'poisoning'

While they have helped reduce pollution in cities, catalytic converters have a downside: heterogeneous catalysts can become 'poisoned' by impurities that block active sites. The platinum and rhodium get 'poisoned' by lead compounds in leaded petrol. So, cars with this kind of catalytic converter should not use leaded petrol. Catalytic poisoning reduces efficiency and increases the cost of the process.

Now you know how heterogeneous catalysts work. Let us discuss another type of catalyst - homogeneous catalysts.

Homogeneous Catalyst

A homogeneous catalyst is a catalyst which is in the same phase as the reactants.

Homogeneous catalysis often involves an aqueous catalyst and aqueous reactants, but this is not always the case. Sometimes, the catalyst and the reactants will be in the gas phase.

In homogeneous catalysis, the reaction proceeds through an intermediate species. What does this mean? You can see how this works in the reaction between the persulfate ion (peroxydisulfate) and iodine.

How Fe²⁺ ions catalyse the reaction between iodide and persulfate ions

Persulfate (or peroxodisulfate, its IUPAC name) acts as an oxidising agent when it reacts with iodide. The equation for the reaction is given below:

S₂O₈^2- + 2I^- ➔ 2SO₄^2- + I₂

At room temperature, the process is slow due to the negatively charged reactant ions repelling each other. However, in the presence of Fe²⁺ ions, the reaction is much faster. The Fe²⁺ ions and the reactants are in the aqueous phase, so this reaction is an excellent example of a homogeneous catalyst. Let us look at the steps below:

1. Since the Fe²⁺ ions and the S₂O₈^2- ions share opposite charges, they attract each other and react as follows:

S₂O₈^2- + 2Fe²⁺ ➔ 2SO₄^2- + 2Fe³⁺

2. The Fe³⁺ ions produced in the first reaction react with the I^- ions as follows:

2Fe³⁺ + 2I^- ➔ 2Fe²⁺ + I₂

As you can see, the original iron(II) ion catalyst gets regenerated, so the steps repeat themselves. The iron(III) ions that form in the first reaction act as the intermediate species.

In the process between iodide ions and persulfate ions, we can also use iron(III) ions as the original catalyst. In this case, Fe²⁺ is the intermediate species. The reaction would take place in the reverse order:

2Fe³⁺ + 2I^- ➔ 2Fe²⁺ + I₂
S2O₈^2- + 2Fe²⁺ ➔ 2SO₄^2- + 2Fe³⁺

Nitrogen dioxide (NO₂) as a catalyst.

An example of the use of a homogenous catalyst can be seen with nitrogen dioxide. We can explore this through formation of acid rain which contains H₂SO₄ (sulphuric acid).

When sulphur dioxide (SO₂), a pollutant in the atmosphere, is oxidised, it turns into SO_3.This can then react with rain water to produce H₂SO_4.

_{This can be shown as:}

SO₃(g) + H₂O(l) → H₂SO₄(aq)

So what role does nitrogen dioxide play in this reaction? It not only leads to acid rain, but also acts as a catalyst. This is due to nitrogen dioxide catalysis of SO₂ to SO₃.

This can be shown as:

SO₂(g) + NO₂(g) → SO₃(g) + NO(g)

The NO in this reaction can be regenerated, like most catalysts, back to NO₂where it can go on to catalyse the reaction of SO₂to SO_3.

This can be presented as:

NO(g) + 1/2 O₂(g) → NO₂(g)

Before we conclude this discussion on catalysts, let us look at one last type of catalysis, namely, autocatalysis.

Autocatalysis

In the following reaction, negative manganate(VII) ions react with negative ethanedioate ions.

2MnO₄^-_(aq) + 5C₂O₄^2-_(aq) + 16H⁺_(aq) ➔ 2Mn²⁺_(aq) + 10CO_{2 (g)} + 8H₂O _(l)

This reaction is fascinating, as the reaction rate increases as more Mn²⁺ ions get produced. We call it autocatalysis when a process is catalysed by one of the products of the reaction. In this case, Mn²⁺_(aq) acts as an autocatalyst. The process starts slow, but as more manganese (II) ions form, it becomes faster and faster. Eventually, the reaction slows down when the catalyst gets used up.

Similar to the previous reaction where iron acts as a catalyst, the Mn(II) gets regenerated in a redox cycle.

The Mn²⁺ ions react with the MnO₄^- ions to produce Mn³⁺ ions.
The Mn³⁺ ions react with the ethanedioate ions to regenerate Mn²⁺ ions.

We can show this redox cycle using the following equations:

4Mn²⁺_(aq) + MnO₄^-_(aq) + 8H⁺_(aq) → 5Mn³⁺_(aq) + 4H₂O _(aq)

2Mn³⁺_(aq) + C₂O₄^2-_(aq) → 2CO₂_(g) + 2Mn²⁺_(aq)

And there you have it: the wonderful world of transition metal catalysts!

Catalysts - Key takeaways

A catalyst is a substance that increases the rate of a chemical reaction without being changed in chemical composition or quantity.
A heterogeneous catalyst is in a different phase from the reactants. The reaction occurs at active sites on the catalyst surface.
The Haber process produces ammonia through a heterogeneous catalyst, iron.
Heterogeneous catalysts work through surface adsorption theory. First, the reactants get adsorbed on the surface of the catalyst. Then the reaction takes place. Lastly, the products desorb from the surface of the catalyst.
The Contact process produces Sulfuric acid through a heterogeneous catalyst, vanadium pentoxide.
Catalytic converters reduce vehicle emissions through surface adsorption theory.
A catalytic converter is a ceramic honeycomb coated with a metal catalyst mixture of platinum and rhodium or palladium.
Catalytic poisoning reduces the efficiency and increases the cost of catalytic converters.
A homogeneous catalyst is in the same phase as the reactants.
In homogeneous catalysis, the reaction proceeds through an intermediate species.
Autocatalysis occurs when a process is catalysed by one of the reaction products.

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