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Catalysts

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Intros
Lessons
  1. The effect of catalysts
  2. Why are catalysts important to chemists?
  3. How catalysts affect mechanism and potential energy.
  4. Reaction profile with and without a catalyst.
  5. Heterogeneous catalysts.
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Examples
Lessons
  1. Explain the use of catalyst and their effect on the energy profile of a reaction.
    The decomposition of hydrogen peroxide is shown in the equation below:

    2H2O2→2H2O+O2\mathrm{2H_2O_2 \to 2 H_2O + O_2}

    1. This reaction is slow but is sped up with a catalyst such as potassium iodide. Show the uncatalyzed and catalyzed reaction pathways on a potential energy diagram. Assume the reaction is exothermic and the catalyzed pathway forms an intermediate.
  2. Explain the use of catalyst and their effect on the energy profile of a reaction.
    The Haber process is an important industrial process to produce ammonia. The equation is below:

    3H2(g)+N2(g)⇌2NH3(g)\mathrm{3H_{2 (g)} + N_{2 (g)} \rightleftharpoons 2NH_{3 (g)} }

    1. For practical reasons, this reaction is always performed with an iron-based catalyst. Explain why this is used and its effect on the position of equilibrium.
    2. Draw two energy profiles for this exothermic reaction, showing a catalyzed and a possible uncatalyzed pathway. Label the reactants and products on these profiles.
Topic Notes
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Introduction to Catalysts

Welcome to our exploration of catalysts and their crucial role in chemical reactions! Catalysts are remarkable substances that speed up reactions without being consumed in the process. They're like the unsung heroes of the chemical world, making reactions happen faster and more efficiently. In our introduction video, we'll dive into the fascinating world of catalysts and their impact on reaction kinetics. You'll see how these amazing substances work their magic, lowering activation energy and providing alternative reaction pathways. Understanding catalysts is key to grasping many industrial processes and even biological functions. As we progress, you'll discover how catalysts are used in everything from manufacturing to environmental protection. So, get ready to unlock the secrets of catalysts and revolutionize your understanding of chemical reactions! This video is your gateway to a deeper comprehension of reaction kinetics and the power of catalysts in shaping our world.

Importance and Examples of Catalysts in Industry

The Economic and Environmental Significance of Catalysts

Catalysts play a crucial role in numerous industrial processes, offering both economic and environmental benefits. These substances accelerate chemical reactions without being consumed, allowing for more efficient and sustainable production methods. The importance of catalysts in industry cannot be overstated, as they contribute to reduced energy consumption, lower production costs, and decreased environmental impact.

Economic Advantages of Catalysts

From an economic standpoint, catalysts offer several key advantages in industrial processes:

  • Increased reaction rates, leading to higher productivity
  • Lower energy requirements, resulting in reduced operational costs
  • Improved selectivity, minimizing waste and maximizing desired product yield
  • Extended equipment lifespan due to milder reaction conditions

Environmental Benefits of Catalysts

The environmental advantages of using catalysts in industrial processes are equally significant:

  • Reduced energy consumption, leading to lower greenhouse gas emissions
  • Decreased waste production through improved reaction efficiency
  • Minimized use of harmful chemicals and solvents
  • Enhanced ability to remove pollutants from waste streams

Lowering Activation Energy

One of the primary functions of catalysts is to lower the activation energy required for a chemical reaction to occur. By providing an alternative reaction pathway, catalysts enable reactions to proceed more quickly and under milder conditions. This reduction in activation energy translates to significant energy savings and increased efficiency in industrial processes.

Examples of Catalysts in Industrial Processes

The Haber-Bosch Process

The Haber-Bosch process is a prime example of the importance of catalysts in industry. This process, used for the production of ammonia, employs an iron-based catalyst to facilitate the reaction between nitrogen and hydrogen. Without this catalyst, the reaction would require extremely high temperatures and pressures, making it economically unfeasible. The catalyst allows the process to operate at more moderate conditions, significantly reducing energy consumption and production costs.

Hydrogenation of Ethene

The hydrogenation of ethene to produce ethane is another industrial process that benefits from catalysis. This reaction, important in the petrochemical industry, uses a nickel catalyst to promote the addition of hydrogen to ethene. The catalyst enables the reaction to occur at lower temperatures and pressures, improving energy efficiency and reducing the overall environmental impact of the process.

Synthesis Gas Production

Synthesis gas, a mixture of carbon monoxide and hydrogen, is a crucial intermediate in many industrial processes. The production of synthesis gas through steam reforming of natural gas relies heavily on catalysts, typically nickel-based. These catalysts allow the reaction to proceed efficiently at lower temperatures, saving energy and reducing carbon dioxide emissions. The resulting synthesis gas serves as a feedstock for various products, including methanol, ammonia, and synthetic fuels.

Resource Conservation through Catalysis

By enhancing reaction efficiency and selectivity, catalysts play a vital role in conserving resources. They enable industries to:

  • Utilize raw materials more effectively, reducing waste
  • Minimize energy consumption in production processes
  • Extend the lifespan of equipment and reduce maintenance needs
  • Develop more sustainable production methods

Conclusion

The economic and environmental importance of catalysts in industrial processes cannot be overstated. From the Haber-Bosch process to the production of synthesis gas, catalysts enable more efficient, cost-effective, and environmentally friendly manufacturing methods. By lowering activation energy, improving selectivity, and conserving resources, catalysts continue to drive innovation and sustainability in the chemical industry. As we face growing environmental challenges

Effects of Catalysts on Reaction Mechanisms

Catalysts are remarkable substances that play a crucial role in chemical reactions by providing alternative reaction pathways with lower activation energy. To understand how catalysts work, let's dive into the fascinating world of reaction mechanisms and explore the concepts of intermediates, transition states, and reaction profiles.

Imagine you're hiking up a mountain. The direct path to the summit is steep and requires a lot of energy. This is similar to a chemical reaction without a catalyst - it needs a high activation energy to proceed. Now, picture a winding trail that takes a longer route but has a gentler slope. This represents the alternative pathway provided by a catalyst, which requires less energy to reach the same destination.

In chemical terms, the activation energy is the minimum energy required for a reaction to occur. Catalysts work by lowering this energy barrier, making it easier for reactants to transform into products. They achieve this by interacting with the reactants and forming temporary chemical bonds, creating what we call "intermediates."

Intermediates are temporary species formed during a reaction that are neither the initial reactants nor the final products. They're like pit stops on our hiking trail, allowing the reaction to progress through a series of smaller, more manageable steps rather than one giant leap. These intermediates are typically unstable and short-lived, quickly transforming into the next species along the reaction pathway.

As the reaction progresses, it passes through "transition states." These are the highest energy points along the reaction pathway, representing the moment when bonds are breaking and forming. In our hiking analogy, transition states are like the steepest parts of the trail - the most challenging points to overcome. Catalysts help by lowering the energy of these transition states, making them easier to surpass.

To visualize how catalysts affect a reaction, chemists use reaction profiles or energy diagrams. These graphs show the energy changes as reactants transform into products. On a reaction profile, the uncatalyzed reaction appears as a high energy hill, while the catalyzed reaction shows a series of smaller hills with lower peaks. This illustrates how the catalyst breaks down the reaction into multiple steps with lower energy barriers.

Let's consider a real-world example: the catalytic converter in your car. This device uses precious metals like platinum as catalysts to convert harmful exhaust gases into less harmful emissions. Without these catalysts, the reactions would be too slow at normal temperatures, and your car would release more pollutants. The catalysts provide alternative reaction pathways that allow these conversions to happen quickly and efficiently at lower temperatures.

It's important to note that while catalysts lower the activation energy and speed up reactions, they don't change the overall energy difference between reactants and products. This means they don't affect the thermodynamics of the reaction - only the kinetics. In other words, catalysts can't make an unfavorable reaction favorable; they can only make a favorable reaction happen faster.

Catalysts also remain unchanged at the end of the reaction. They participate in the process but are regenerated in their original form, ready to catalyze the next reaction. This is why a small amount of catalyst can facilitate the conversion of a large amount of reactants.

Understanding how catalysts work is crucial in many fields, from industrial chemistry to biochemistry. In our bodies, enzymes act as biological catalysts, enabling the countless reactions necessary for life to occur at body temperature. In industry, catalysts are used to produce everything from plastics to pharmaceuticals, making processes more efficient and environmentally friendly.

As we continue to face global challenges like climate change and resource scarcity, the development of new and improved catalysts becomes increasingly important. These molecular helpers allow us to design cleaner, more efficient chemical processes, paving the way for sustainable technologies and a greener future.

In conclusion, catalysts are powerful tools in the chemist's arsenal, providing alternative reaction pathways with lower activation energies. By forming intermediates and lowering the energy of transition states, they allow reactions to proceed more quickly and efficiently. Understanding reaction profiles helps us visualize these effects and appreciate the elegant way catalysts guide chemical transformations. As we continue to unlock the secrets of catalysis, we open doors to new possibilities in science, technology, and environmental stewardship.

Potential Energy Curves and Catalyzed Reactions

Potential energy curves are powerful tools for visualizing and understanding chemical reactions, particularly when comparing catalyzed and uncatalyzed reaction pathways. These curves provide a graphical representation of how the potential energy of a system changes as the reaction progresses, offering valuable insights into reaction mechanisms, energy barriers, and the role of catalysts.

In an uncatalyzed reaction, the potential energy curve typically shows a single, high-energy peak representing the transition state. This peak corresponds to the activation energy required for the reaction to occur. The height of this peak determines the reaction rate, with higher activation energies resulting in slower reactions.

When a catalyst is introduced, the potential energy curve changes dramatically. Catalysts work by providing an alternative reaction pathway with lower activation energy. This is represented on the curve as a series of smaller peaks and valleys, rather than a single high peak. Each of these smaller peaks corresponds to the formation of intermediates temporary molecular structures that exist briefly during the reaction process.

To illustrate this concept, let's consider a simple example of the decomposition of hydrogen peroxide (H2O2) into water and oxygen. In the uncatalyzed reaction, the potential energy curve would show a single high peak, representing the energy required to break the O-O bond in H2O2. However, when a catalyst like manganese dioxide (MnO2) is added, the curve changes.

The catalyzed reaction pathway might involve the following steps:

  1. H2O2 adsorbs onto the MnO2 surface
  2. The O-O bond in H2O2 weakens and breaks
  3. One oxygen atom remains bound to the catalyst
  4. A second H2O2 molecule reacts with the bound oxygen
  5. O2 is released, and the catalyst is regenerated

Each of these steps would be represented by a small peak on the potential energy curve, with valleys in between representing the intermediate states. The overall activation energy the highest point on the curve is significantly lower than in the uncatalyzed reaction.

The formation of intermediates is crucial in catalyzed reactions. These temporary structures allow the reaction to proceed through a series of smaller, more easily surmountable energy barriers, rather than one large barrier. Intermediates often involve chemical species bonded to the catalyst surface, which can stabilize reactive molecules or bring reactants into close proximity.

Energy profiles, which are essentially potential energy curves plotted against the reaction coordinate, provide a clear visual comparison between catalyzed and uncatalyzed reactions. In these profiles, the uncatalyzed reaction appears as a single smooth curve with a high peak, while the catalyzed reaction shows a more complex path with multiple smaller peaks and troughs.

The significance of intermediates extends beyond just lowering activation energy. They can also:

  • Provide insight into the reaction mechanisms
  • Offer opportunities for selective product formation
  • Allow for the isolation of valuable compounds
  • Help in understanding enzyme function in biochemical reactions

Understanding potential energy curves and energy profiles is essential for chemists and chemical engineers. These visual representations help in:

  • Designing more efficient catalysts
  • Optimizing reaction conditions
  • Predicting reaction rates and equilibria
  • Developing new synthetic routes for complex molecules

In conclusion, potential energy curves and energy profiles are invaluable tools for comparing catalyzed and uncatalyzed reactions. They clearly illustrate how catalysts lower activation energy by providing alternative reaction pathways involving intermediates. By understanding these concepts, scientists can develop more efficient and selective chemical processes, leading to advancements in fields ranging from industrial manufacturing to pharmaceutical development and environmental protection.

Heterogeneous Catalysts and Their Applications

Heterogeneous catalysts are a cornerstone of modern industrial processes, playing a crucial role in making chemical reactions more efficient and environmentally friendly. These catalysts are called "heterogeneous" because they exist in a different phase from the reactants they help transform. Typically, heterogeneous catalysts are solid materials that interact with gases or liquids, facilitating reactions without being consumed in the process.

One of the most familiar examples of heterogeneous catalysts in action is the catalytic converter found in automobiles. This device is a prime illustration of how these catalysts work to solve real-world problems. Catalytic converters contain precious metals like platinum, palladium, and rhodium, which act as catalysts to convert harmful exhaust gases into less harmful emissions. As exhaust fumes pass through the converter, these metals catalyze reactions that transform carbon monoxide into carbon dioxide, hydrocarbons into carbon dioxide and water, and nitrogen oxides into nitrogen and oxygen.

The effectiveness of heterogeneous catalysts largely depends on a process called adsorption. Adsorption occurs when molecules from a gas or liquid adhere to the surface of the solid catalyst. This process is crucial because it brings reactant molecules into close proximity with each other and with the catalyst surface, significantly increasing the likelihood of a reaction occurring. The more molecules that can adsorb onto the catalyst's surface, the higher the chance of successful reactions.

To maximize the potential for adsorption and subsequent reactions, heterogeneous catalysts are often designed with extremely high surface areas. This is achieved by creating materials with complex, porous structures or by using nanoparticles. The increased surface area provides more sites for reactant molecules to adsorb and interact, enhancing the catalyst's efficiency. In the case of catalytic converters, the precious metal catalysts are typically dispersed as tiny particles across a honeycomb-structured ceramic support, creating a vast surface area within a compact device.

The importance of heterogeneous catalysts in industrial processes cannot be overstated. They enable the production of a wide range of essential products, from fuels and plastics to pharmaceuticals and fertilizers. By lowering the activation energy required for chemical reactions, these catalysts allow processes to occur at lower temperatures and pressures, reducing energy consumption and operational costs. Moreover, they often improve selectivity, meaning they can guide reactions towards desired products while minimizing unwanted by-products.

In the context of environmental protection, heterogeneous catalysts are invaluable. Beyond catalytic converters in vehicles, they are used in power plants to reduce emissions, in water treatment facilities to remove contaminants, and in the production of clean fuels. The ability of these catalysts to facilitate gas phase reactions is particularly important for addressing air pollution and developing cleaner energy technologies.

As we continue to face global challenges related to energy efficiency, resource conservation, and environmental protection, the role of heterogeneous catalysts becomes increasingly critical. Researchers are constantly working to develop new catalytic materials and improve existing ones, aiming for higher activity, better selectivity, and increased durability. Advances in nanotechnology and materials science are opening up new possibilities for catalyst design, potentially leading to breakthroughs in fields such as renewable energy production and carbon dioxide conversion.

Understanding the principles behind heterogeneous catalysis is not just important for chemists and engineers; it's relevant to anyone interested in how technology can address environmental and energy challenges. The next time you start your car, remember that the catalytic converter under your vehicle is a prime example of heterogeneous catalysis in action, silently working to make our air cleaner. This everyday application is just one of many that demonstrate the profound impact of heterogeneous catalysts on our modern world, highlighting their indispensable role in shaping a more sustainable and efficient future.

Catalysts in Environmental Applications

Catalysts play a crucial role in environmental applications, particularly in reducing pollution and improving air quality. These remarkable substances accelerate chemical reactions without being consumed, making them invaluable tools in addressing various environmental challenges. One of the most significant and well-known applications of catalysts in environmental protection is their use in vehicle exhaust systems.

In the automotive industry, catalytic converters are essential components that utilize catalysts to convert harmful exhaust gases into less harmful emissions. These devices contain precious metals such as platinum, palladium, and rhodium, which act as catalysts to facilitate the conversion of toxic pollutants. For instance, carbon monoxide (CO) is oxidized to carbon dioxide (CO2), while nitrogen oxides (NOx) are reduced to nitrogen gas (N2) and oxygen (O2). Additionally, unburned hydrocarbons are converted into CO2 and water vapor. This process significantly reduces the environmental impact of vehicle emissions, contributing to improved air quality in urban areas and along busy highways.

Beyond automotive applications, catalytic technology has broader implications in addressing environmental challenges. In industrial settings, catalysts are employed to reduce emissions from power plants and manufacturing facilities. Selective catalytic reduction (SCR) systems use catalysts to convert NOx emissions from industrial processes into harmless nitrogen and water. This technology has been instrumental in helping industries comply with stringent environmental regulations and reduce their carbon footprint.

Catalysts also play a vital role in water treatment processes. Advanced oxidation processes utilizing catalysts can effectively remove organic pollutants and pathogens from wastewater, making it safer for release into the environment or for reuse. In the realm of renewable energy, catalysts are essential for the production of clean fuels such as hydrogen through water splitting reactions, offering promising solutions for sustainable energy storage and transportation.

The development of new and more efficient catalysts continues to drive innovation in environmental applications. Researchers are exploring nanocatalysts and bio-inspired catalysts that offer enhanced performance and selectivity. These advancements hold the potential to further improve pollution reduction techniques, increase energy efficiency, and support the transition to more sustainable industrial processes. As environmental concerns continue to grow globally, the role of catalysts in providing effective solutions for pollution control and environmental remediation becomes increasingly critical, underscoring their importance in shaping a cleaner, more sustainable future.

Conclusion: The Future of Catalytic Chemistry

Catalysts play a crucial role in chemical reactions and industrial processes, accelerating reactions without being consumed. They reduce energy requirements, increase efficiency, and enable the production of vital materials. The introduction video provides a foundational understanding of catalytic chemistry, highlighting its importance across various sectors. Catalysts are essential in manufacturing, environmental protection, and energy production. They contribute to sustainable practices by reducing waste and energy consumption. The future of catalytic chemistry holds immense potential, with ongoing research into more efficient and selective catalysts. These advancements could revolutionize industries, leading to greener technologies and novel materials. To fully grasp the significance of catalysts and their future applications, further engagement in this field is crucial. Explore additional resources, participate in research, or consider a career in catalytic chemistry. The potential for innovation and positive impact on society through catalytic chemistry is boundless, making it an exciting and vital area of study.

The Effect of Catalysts

Why are catalysts important to chemists?

Step 1: Introduction to Catalysts

In this section, we will explore the fundamental concept of catalysts and their significance in chemical reactions. Catalysts are substances that increase the rate of a chemical reaction without being consumed in the process. They achieve this by lowering the activation energy required for the reaction to proceed, making it easier for reactants to convert into products. This property of catalysts is crucial in both industrial and laboratory settings.

Step 2: The Role of Catalysts in Chemical Reactions

Catalysts play a vital role in chemical reactions by affecting the kinetics of the reaction. By lowering the activation energy, catalysts enable reactions to occur under milder conditions, such as lower temperatures and pressures. This not only makes the reactions safer and more efficient but also reduces the energy and resources required. For chemists, this means that reactions can be carried out more economically and with less environmental impact.

Step 3: Economic and Environmental Benefits

The use of catalysts offers significant economic and environmental advantages. Economically, catalysts help in saving money by reducing the need for high temperatures and pressures, which are energy-intensive and costly. This is particularly important in industrial processes where large-scale production is involved. Environmentally, catalysts contribute to sustainability by minimizing resource consumption and reducing waste generation. This makes chemical processes more eco-friendly and aligns with the goals of green chemistry.

Step 4: Examples of Industrial Catalysts

Several industrial processes rely heavily on catalysts to enhance efficiency and productivity. For instance, the Haber-Bosch process, which produces ammonia for fertilizers, uses an iron-based catalyst. Another example is the hydrogenation of ethene to produce ethane, a process that utilizes a nickel catalyst. These processes are essential for producing large quantities of chemicals that are vital for various applications, from agriculture to manufacturing.

Step 5: Specific Effects on Reaction Mechanisms

Catalysts not only lower the activation energy but also influence the reaction mechanisms. They provide alternative pathways for the reaction to occur, which can lead to different intermediate species and transition states. Understanding these mechanisms is crucial for chemists as it allows them to design more efficient catalysts and optimize reaction conditions. Potential energy curves are often used to illustrate these catalyzed reaction pathways, showing how the energy barrier is reduced in the presence of a catalyst.

Step 6: Practical Applications and Future Directions

The practical applications of catalysts are vast and continue to expand with ongoing research and development. In addition to traditional industrial processes, catalysts are being explored for use in renewable energy technologies, such as fuel cells and biofuel production. The development of new and more effective catalysts holds the promise of further improving the efficiency and sustainability of chemical processes. For chemists, staying informed about these advancements is essential for driving innovation and addressing global challenges related to energy and the environment.

FAQs

  1. What is a catalyst and how does it work?

    A catalyst is a substance that increases the rate of a chemical reaction without being consumed in the process. It works by providing an alternative reaction pathway with lower activation energy, allowing the reaction to proceed more quickly. Catalysts form temporary bonds with reactants, creating intermediates that facilitate the conversion to products.

  2. What is the difference between homogeneous and heterogeneous catalysts?

    Homogeneous catalysts are in the same phase as the reactants (usually liquid), while heterogeneous catalysts are in a different phase (typically solid catalysts with liquid or gas reactants). Heterogeneous catalysts, like those in catalytic converters, often work through surface adsorption of reactants.

  3. How do catalysts contribute to environmental protection?

    Catalysts play a crucial role in environmental protection by reducing pollution and improving air quality. They are used in catalytic converters to convert harmful vehicle emissions into less toxic substances, in industrial processes to reduce emissions, and in water treatment to remove contaminants. Catalysts also enable the production of cleaner fuels and support renewable energy technologies.

  4. What are some common industrial applications of catalysts?

    Catalysts are widely used in industry for processes such as the Haber-Bosch process for ammonia production, petroleum refining, synthesis gas production, and the manufacture of plastics and pharmaceuticals. They enable more efficient and cost-effective production methods, often under milder conditions with reduced energy consumption.

  5. What are the future prospects for catalytic chemistry?

    The future of catalytic chemistry is promising, with ongoing research into more efficient and selective catalysts. Advancements in nanotechnology and materials science are opening up new possibilities for catalyst design. Future developments may lead to breakthroughs in renewable energy production, carbon dioxide conversion, and the creation of more sustainable industrial processes.

Prerequisite Topics for Understanding Catalysts

To fully grasp the concept of catalysts in chemistry, it's crucial to have a solid foundation in several key areas. Understanding activation energy is paramount, as catalysts work by lowering this energy barrier, allowing reactions to proceed more quickly. This concept is intrinsically linked to the broader field of reaction kinetics, which explores how chemical reactions occur and the factors that influence their rates.

A thorough comprehension of chemical reactions forms the bedrock of catalyst studies. Catalysts are substances that accelerate chemical reactions without being consumed in the process, so understanding the basics of how reactions occur is essential. This knowledge helps in appreciating how catalysts can dramatically alter reaction pathways and speeds.

Delving into reaction mechanisms is also vital when studying catalysts. These mechanisms provide a step-by-step breakdown of how reactants transform into products, and catalysts often introduce alternative pathways with lower energy requirements. By understanding reaction mechanisms, students can better visualize and predict how catalysts influence chemical processes.

The interplay between activation energy and reaction kinetics is particularly significant in the context of catalysts. Catalysts function by providing an alternative reaction pathway with a lower activation energy, which directly impacts the kinetics of the reaction. This relationship underscores why a solid grasp of both concepts is crucial for anyone studying catalysts.

Moreover, the study of chemical reactions provides the necessary context for understanding catalytic processes. Catalysts don't change the overall thermodynamics of a reaction; they simply facilitate the process. Therefore, a strong foundation in chemical reactions helps students appreciate the role of catalysts within the broader scope of chemical transformations.

Reaction mechanisms become especially relevant when examining how catalysts operate at a molecular level. Many catalysts work by forming intermediate compounds or by providing surfaces for reactions to occur more easily. Understanding these mechanisms allows students to comprehend the intricate dance of molecules and energy that catalysts orchestrate.

In conclusion, mastering these prerequisite topics creates a robust framework for exploring the fascinating world of catalysts. Each concept builds upon the others, forming a comprehensive understanding of how catalysts function and their critical role in chemistry and industry. By investing time in these foundational areas, students will be well-equipped to tackle more advanced concepts in catalysis and chemical kinetics.

In this lesson, we will learn:

  • To recall some important examples of catalysts in chemical industry.
  • To explain the specific effect of catalysts on reaction mechanisms and potential energy.
  • How to use potential energy curves to show catalyzed reaction pathways.

Notes:

  • We saw in Factors affecting rate of reaction that a catalyst speeds up a chemical reaction by providing an alternative route of lower activation energy for a reaction. The catalyst is important to meet the activation energy, otherwise the reaction may never even begin.
    • Most reactions of fuels burning (combustion) are very exothermic. Overall, they release a lot of heat energy because the products are so much lower energy than the reactants. If you think of a reaction as going up a hill and then back down the other side, they would be like climbing up a staircase to fall off of a building.
      But if the activation energy of the reaction isnt met, it does not matter; the reaction will not start. Thats why fuels need a source of ignition, like a spark, to begin burning
    • Catalysts are incredibly useful for financial and environmental reasons when running industrial scale chemical reactions.
      With no catalyst to lower the activation energy, very harsh conditions like high temperature and pressure must provide this energy the reaction needs. Harsh conditions costs energy to make and maintain, and the longer you run a reaction, the longer you must hold these conditions. This costs companies a lot of money and the environment more of its resources!
    • The Haber process (iron), hydrogenation of ethene (nickel) and production of synthesis gas all use catalysts. They are very important industrial processes and catalysts make them less environmentally damaging.

    This lesson looks in more detail at how catalysts have their effect.

  • The catalysts alternative route is the reaction going by a different reaction mechanism than it does when no catalyst is present. The catalyst might lead to a lower energy intermediate substance that the reactants cant make without it.
    This would be like trying to climb over a high wall with a chair/stool, or without one. Which would be easier, and which would take less energy out of you?
    • The reaction mechanism is the sequence of steps in how the reactant bonds break and the product bonds form. Intermediate(s) and transition state(s) are temporary molecules that form for a short time during the reaction, as the bonds rearrange. More on this in Reaction mechanisms.
    • The reaction mechanism depends on the reactants and the catalyst used, but generally catalysts provide other possible reaction mechanisms. Some will be useless, higher energy than the uncatalyzed mechanism, while other mechanisms are lower energy and allow more reactant particles to become products, increasing the reaction rate as a result.
      • Think about this with the chair/stool analogy: when a stool is on its side on the floor or upside down, its useless. When it is upright, its a lot more useful. Catalysts are the same it just provides more options, only one needs to be lower energy than without it though.
    • The whole description of the reaction procedure, including potential energy involved and the stages in the mechanism is called the reaction pathway. These are often described using a potential energy diagram.
    • A typical reaction with a catalyzed and uncatalyzed pathway will appear different on potential energy curves. The catalyzed pathway will have a lower activation energy and may have a dip at the crest of the peak this is where an intermediate is formed.

  • Many industrial processes, such as the Haber process, have their reactants in the gas phase.
    While this means the reactant molecules are high energy, it also means that the chances of collision between particles is lower because gas particles are not very compact.
    To counteract this, most industrial processes use a solid phase catalyst for their gas phase reactants. This is called a heterogeneous catalyst, which is any catalyst in a different phase to the reactants.
    • Using a solid phase catalyst provides more surface area for the reactant particles to adsorb onto a surface. This is an attractive interaction that temporarily holds particles in place on a surface; this holding increases the chance of collisions between reactant particles happening at the surface. The larger the surface area of a catalyst, the more adsorption can happen, so the chance of successful collisions between reactant particles increases.
    • An example of this is a catalytic converter in a car, which uses a solid platinum or rhodium catalyst to convert toxic carbon monoxide and NOx gases into harmless N2 and CO2 before the gases are released to the environment. It has a thin honeycomb-like structure, so we get to use the most surface area for the smallest possible quantity of the expensive catalyst.
    • Sometimes, particles adsorb to a surface and do not come off the holding is permanent. Chemists call this change poisoning of the catalyst, because that site where those molecules are bound permanently cant be used for the reaction anymore. Eventually, catalysts need replacing as more and more of their surface gets poisoned and they become less and less effective.

  • When dealing with reactions at equilibrium, remember that activation energy barriers and reaction mechanisms exist and work for reverse reactions too! A catalyst will speed up forward and reverse reactions in a chemical process; equilibrium position is unchanged. Only the time taken to reach equilibrium is reduced.