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Chapter 7: Problem 45

What is a catalyst, and what effect does it have on the activation energy of a reaction?

Short Answer

Expert verified
A catalyst lowers the activation energy of a reaction, increasing its rate without being consumed.

Step by step solution

01

Understanding Catalysts

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.
02

Alternative Reaction Pathway

The alternative pathway provided by the catalyst requires lower activation energy compared to the uncatalyzed pathway.
03

Effect on Activation Energy

By lowering the activation energy, the catalyst makes it easier for reactants to convert to products, leading to an increased reaction rate.

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Key Concepts

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

Activation Energy
Activation energy is the minimum energy required for a chemical reaction to occur. Imagine it as an energy barrier that reactants need to overcome to transform into products. In the absence of sufficient energy, a chemical reaction cannot proceed. Activation energy (denoted as \(E_a\)) is crucial because it determines the speed and feasibility of a reaction. Factors that can alter activation energy include:
  • Temperature: Higher temperatures increase the energy of reactant molecules, potentially lowering the observable activation energy.
  • Catalysts: They provide an alternative pathway with a lower activation energy, which we'll discuss further.
Reducing activation energy doesn't change the nature of the reactants or products. It simply makes the conversion more achievable under given conditions.
Reaction Rate
The reaction rate is a measure of how quickly reactants are converted into products during a chemical reaction. Various factors influence this rate, including temperature, concentration, physical state, and the presence of a catalyst.
  • Temperature: Generally, raising the temperature increases the reaction rate by providing energy that helps overcome the activation energy barrier.
  • Concentration: Reactants that are more concentrated often collide more, leading to faster reactions.
  • Surface Area: For reactions involving solids, greater surface area allows more collisions and speeds up the reaction.
  • Catalysts: Catalysts are especially effective as they speed up reactions without being consumed, offering a "shortcut" that bypasses a portion of the energy barrier required for the reaction to proceed.
Adjusting these factors can help control and optimize reactions for industrial or laboratory processes.
Chemical Reaction Mechanics
Chemical reaction mechanics involves understanding how reactions progress from reactants to products. This concept includes analyzing energy barriers, intermediate steps, and transition states within a reaction. A catalyst plays a significant role in reaction mechanics by providing a lower energy pathway. It participates by stabilizing the transition state or forming an intermediate that reacts more readily than the original reactants. Here鈥檚 how it works:
  • In its role, a catalyst lowers the activation energy needed for the reaction to proceed. This enables the reaction to occur more easily and often at a lower temperature.
  • By altering the reaction mechanism, catalysts can change how quickly products form. However, they are not used up in the process. This characteristic makes catalysts essential in many industrial processes, like the synthesis of ammonia in the Haber process.
Understanding chemical reaction mechanics, including the use of catalysts, is vital for developing efficient chemical processes.

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Most popular questions from this chapter

The following reaction is used in the industrial synthesis of polyvinyl chloride (PVC) polymer: \(\mathrm{Cl}_{2}(g)+\mathrm{H}_{2} \mathrm{C}=\mathrm{CH}_{2}(g) \longrightarrow \mathrm{ClCH}_{2} \mathrm{CH}_{2} \mathrm{Cl}(l)\) $$ \Delta H=-218 \mathrm{~kJ} / \mathrm{mol} $$ (a) Is \(\Delta S\) positive or negative for this process? (b) Is this process spontaneous at all temperatures? Explain.

Why do catalysts not alter the amounts of reactants and products present at equilibrium?

What effect do the listed changes have on the position of the equilibrium in the reaction of carbon with hydrogen? $$ \mathrm{C}(s)+2 \mathrm{H}_{2}(g) \rightleftarrows \mathrm{CH}_{4}(g) \quad \Delta H=-75 \mathrm{~kJ} / \mathrm{mol} $$ (a) Increasing temperature (b) Increasing pressure by decreasing volume (c) Allowing \(\mathrm{CH}_{4}\) to escape continuously from the reaction vessel

For the evaporation of water, \(\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{H}_{2} \mathrm{O}(g),\) at \(373 \mathrm{~K}, \Delta H=+40.7 \mathrm{~kJ} / \mathrm{mol}\) (a) How many kilojoules are needed to vaporize \(10.0 \mathrm{~g}\) of \(\mathrm{H}_{2} \mathrm{O}(l) ?\) (b) How many kilojoules are released when \(10.0 \mathrm{~g}\) of \(\mathrm{H}_{2} \mathrm{O}(g)\) is condensed?

Hydrogen chloride can be made from the reaction of chlorine and hydrogen: $$ \mathrm{Cl}_{2}(g)+\mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{HCl}(g) $$ For this reaction, \(K=26 \times 10^{33}\) and \(\Delta H=-184 \mathrm{~kJ} / \mathrm{mol}\) at \(298 \mathrm{~K}\). (a) Is the reaction endothermic or exothermic? (b) Are the reactants or the products favored at equilibrium? (c) Explain the effect on the equilibrium of (1) Increasing pressure by decreasing volume (2) Increasing the concentration of \(\mathrm{HCl}(g)\) (3) Decreasing the concentration of \(\mathrm{Cl}_{2}(g)\) (4) Increasing the concentration of \(\mathrm{H}_{2}(g)\) (5) Adding a catalyst
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