/*! This file is auto-generated */ .wp-block-button__link{color:#fff;background-color:#32373c;border-radius:9999px;box-shadow:none;text-decoration:none;padding:calc(.667em + 2px) calc(1.333em + 2px);font-size:1.125em}.wp-block-file__button{background:#32373c;color:#fff;text-decoration:none} Problem 81 Many metallic catalysts, particu... [FREE SOLUTION] | 91影视

91影视

Log out

Learning Materials

Features

Discover

Chapter 14: Problem 81

Many metallic catalysts, particularly the precious-metal ones, are often deposited as very thin films on a substance of high surface area per unit mass, such as alumina \(\left(\mathrm{Al}_{2} \mathrm{O}_{3}\right)\) or silica \(\left(\mathrm{SiO}_{2}\right) .\) (a) Why is this an effective way of utilizing the catalyst material compared to having powdered metals? (b) How does the surface area affect the rate of reaction?

Short Answer

Expert verified
Thin films maximize surface area, enhancing efficiency and reaction rates. Increased surface area allows more contact with reactants, thus accelerating the process.

Step by step solution

01

Understand the Catalyst Film Deposition

Catalysts are often deposited as thin films on substrates like alumina or silica because these materials offer a high surface area per unit mass. This means that a small amount of catalyst material can cover a large area, ensuring that a greater area is available for reactants to interact with the catalyst. This is more efficient than using the catalyst in a bulk or powdered form, which would not maximize the surface available for the reaction.
02

Compare Thin Films with Powdered Metals

When metallic catalysts are used in powdered form, they might clump together, reducing the accessible surface area. In contrast, thin films ensure that the catalytic material is extended over a large surface area, maximizing exposure to reactants and thus enhancing the catalytic efficiency.
03

Explore Surface Area's Effect on Reaction Rate

The rate of reaction in catalysis is highly dependent on surface area. A greater surface area allows more reactant molecules to come into contact with the catalyst simultaneously, increasing the number of active sites available for the reaction, thus accelerating the reaction rate.
04

Synthesize the Concept

By depositing catalysts as thin films over a large surface area substrate, the efficiency of the catalyst is increased, and the amount of catalyst needed is reduced. Increased surface area enhances the rate of reaction, making this method both cost-effective and efficient in industrial applications.

Unlock Step-by-Step Solutions & Ace Your Exams!

  • Full Textbook Solutions

    Get detailed explanations and key concepts

  • Unlimited Al creation

    Al flashcards, explanations, exams and more...

  • Ads-free access

    To over 500 millions flashcards

  • Money-back guarantee

    We refund you if you fail your exam.

Over 30 million students worldwide already upgrade their learning with 91影视!

Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Thin Films
Thin films in catalysis involve the application of a very thin layer of catalyst material onto a support substrate. The use of thin films is an effective strategy because it allows for the optimal use of expensive catalyst materials, such as precious metals. Thin films enable:
  • Maximized catalyst exposure to reactants.
  • Efficient utilization of material, as only minimal amounts are needed to cover large areas.
  • Avoidance of clumping, which is common in powdered catalysts, ensuring consistent and optimal performance.
This method ensures that the active catalyst sites are evenly distributed over the support surface, providing high efficiency in catalysis applications.
Surface Area
Surface area plays a crucial role in the effectiveness of catalytic processes. Increasing the surface area of a catalyst allows more molecules of reactants to come into contact with the catalyst. This is because:
  • A larger surface has more available sites for reactions.
  • More interaction points for reactant molecules mean that more reactions can occur simultaneously.
  • The efficiency of the catalyst is significantly enhanced as a result.
Materials such as alumina and silica are commonly used to support catalyst films because they offer a high surface area to mass ratio, which is optimal for maintaining a high level of catalytic activity.
Reaction Rate
The rate at which a chemical reaction proceeds can be significantly impacted by the surface area of the catalyst. A higher surface area increases the reaction rate owing to more active sites being available for reactants to engage with the catalyst. Specifically, this leads to:
  • A greater number of successful collisions between reactant molecules and the catalyst surface.
  • Faster conversion rates of reactants to products.
  • Higher overall process efficiency.
This principle is vital in industrial chemistry, where speed and efficiency translate to time and cost savings.
Metallic Catalysts
Metallic catalysts, particularly those using precious metals, benefit greatly from being deployed as thin films on substrates with high surface area. These metals are efficient in catalysis due to their unique ability to lower activation energy in chemical reactions. By using them as thin films:
  • The cost of precious metals is reduced since less material is required.
  • The likelihood of the catalytic material becoming inactive due to clumping is minimized.
  • The surface becomes more accessible, increasing interaction with reactants.
This efficient tactic maximizes the catalytic potential of these metals, ensuring that industrial processes remain both cost-effective and environmentally friendly.

One App. One Place for Learning.

All the tools & learning materials you need for study success - in one app.

Get started for free

Most popular questions from this chapter

In solution, chemical species as simple as \(\mathrm{H}^{+}\) and \(\mathrm{OH}^{-}\) can serve as catalysts for reactions. Imagine you could measure the \(\left[\mathrm{H}^{+}\right]\) of a solution containing an acidcatalyzed reaction as it occurs. Assume the reactants and products themselves are neither acids nor bases. Sketch the \(\left[\mathrm{H}^{+}\right]\) concentration profile you would measure as a function of time for the reaction, assuming \(t=0\) is when you add a drop of acid to the reaction.

Based on their activation energies and energy changes and assuming that all collision factors are the same, rank the following reactions from slowest to fastest. (a) \(E_{a}=75 \mathrm{~kJ} / \mathrm{mol} ; \Delta E=-20 \mathrm{~kJ} / \mathrm{mol}\) (b) \(E_{a}=100 \mathrm{~kJ} / \mathrm{mol} ; \Delta E=+30 \mathrm{~kJ} / \mathrm{mol}\) (c) \(E_{a}=85 \mathrm{~kJ} / \mathrm{mol} ; \Delta E=-50 \mathrm{~kJ} / \mathrm{mol}\)

The rate of a first-order reaction is followed by spectroscopy, monitoring the absorbance of a colored reactant at \(520 \mathrm{nm}\). The reaction occurs in a \(1.00-\mathrm{cm}\) sample cell, and the only colored species in the reaction has an extinction coefficient of \(5.60 \times 10^{3} \mathrm{M}^{-1} \mathrm{~cm}^{-1}\) at \(520 \mathrm{nm}\). (a) Calculate the initial concentration of the colored reactant if the absorbance is 0.605 at the beginning of the reaction. (b) The absorbance falls to 0.250 at \(30.0 \mathrm{~min}\). Calculate the rate constant in units of \(\mathrm{s}^{-1}\), (c) Calculate the half-life of the reaction. (d) How long does it take for the absorbance to fall to \(0.100 ?\)

(a) Consider the combustion of ethylene, \(\mathrm{C}_{2} \mathrm{H}_{4}(g)+\) \(3 \mathrm{O}_{2}(\mathrm{~g}) \longrightarrow 2 \mathrm{CO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(g)\). If the concentra- tion of \(\mathrm{C}_{2} \mathrm{H}_{4}\) is decreasing at the rate of \(0.025 \mathrm{M} / \mathrm{s}\), what are the rates of change in the concentrations of \(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2} \mathrm{O} ?(\mathbf{b})\) The rate of decrease in \(\mathrm{N}_{2} \mathrm{H}_{4}\) partial pressure in a closed reaction vessel from the reaction \(\mathrm{N}_{2} \mathrm{H}_{4}(g)+\mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{NH}_{3}(g)\) is \(10 \mathrm{kPa}\) per hour. What are the rates of change of \(\mathrm{NH}_{3}\) partial pressure and total pressure in the vessel?

What is the molecularity of each of the following elementary reactions? Write the rate law for each. (a) \(2 \mathrm{NO}(g) \longrightarrow \mathrm{N}_{2} \mathrm{O}_{2}(g)\) (b) \(\mathrm{H}_{2} \mathrm{C}-\mathrm{CH}_{2}(g) \longrightarrow \mathrm{CH}_{2}=\mathrm{CH}-\mathrm{CH}_{3}(g)\) (c) \(\mathrm{SO}_{3}(g) \longrightarrow \mathrm{SO}_{2}(g)+\mathrm{O}(g)\)
See all solutions

Recommended explanations on Chemistry Textbooks

Inorganic Chemistry

Read Explanation

Chemical Analysis

Read Explanation

Organic Chemistry

Read Explanation

Nuclear Chemistry

Read Explanation

Physical Chemistry

Read Explanation

Chemical Reactions

Read Explanation
View all explanations

What do you think about this solution?

We value your feedback to improve our textbook solutions.

Study anywhere. Anytime. Across all devices.