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Chapter 13: Problem 105

Why doesn't a catalyst appear in the overall chemical equation for a reaction?

Short Answer

Expert verified
Catalysts aren't consumed and appear unchanged after the reaction, so they don't appear in the overall equation.

Step by step solution

01

Introduction to Catalysts

Catalysts are substances that increase the rate of a chemical reaction without being consumed in the process. They provide an alternative pathway for the reaction with a lower activation energy.
02

Role in the Reaction Mechanism

While catalysts participate in a reaction by forming intermediate complexes, they are regenerated at the end of the reaction mechanism and thus are not consumed or permanently altered.
03

Exclusion from Overall Equation

The overall chemical equation represents only the net change in reactants and products. Since catalysts are not part of the net change but rather facilitate it, they do not appear in the overall equation.
04

Conservation of Mass and Energy

Since the catalyst is regenerated and not used up, its presence in the overall equation would conflict with the principle of conservation of mass, which requires the equation to only reflect substances that permanently change.

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Key Concepts

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

Rate of Chemical Reactions
The rate of a chemical reaction refers to how quickly reactants are converted into products. It signifies the speed at which a reaction occurs. In general, several factors can affect this rate, including temperature, pressure, the concentration of reactants, and the physical state of the reactants.
Catalysts play a crucial role in altering reaction rates. A catalyst increases the reaction rate by offering an alternative pathway with lower activation energy. It's like providing a shortcut for the reaction to proceed at a faster pace.
  • Increased reaction rate means products are formed more rapidly.
  • Catalysts do not alter the overall equilibrium position of the reaction. They simply help attain equilibrium more swiftly.

Despite the faster reaction rate, the catalyst itself remains unchanged at the end of the process, ready to act again.
Reaction Mechanisms
Understanding reaction mechanisms is essential to grasping how catalysts work. A reaction mechanism is a detailed step-by-step description of the sequence of events that take place during a chemical reaction. Each step in the mechanism is called an elementary step.
Catalysts participate by forming temporary intermediate complexes in these mechanisms. This means that they interact with the reactants to form a new intermediate state, lowering the activation energy needed to reach the final product.
  • These intermediates are key components of a reaction mechanism.
  • The catalyst is regenerated at the end of the mechanism, ready to facilitate another reaction cycle.

Ultimately, catalysts do not appear in the overall reaction equation because their presence and role is only visible when inspecting the reaction mechanism closely.
Activation Energy
Activation energy is the minimum energy that reactant molecules must have to successfully collide and form products. It's the barrier that reactions must overcome.
Lowering the activation energy is where catalysts have a significant impact. They provide an alternative path for the reaction with a reduced activation energy, allowing more molecules sufficient energy to react and speed up the process.
  • With lower activation energy, more molecules possess the required energy to reach the transition state.
  • The more frequent collisions with enough energy increase the rate of the reaction.

Thus, by reducing the activation energy, catalysts optimize the efficiency of chemical reactions without being included in the chemical equation.

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Most popular questions from this chapter

Butadiene \(\left(\mathrm{C}_{4} \mathrm{H}_{6}\right)\) reacts with itself to form a dimer with the formula \(\mathrm{C}_{8} \mathrm{H}_{12} .\) The reaction is second order in \(\mathrm{C}_{4} \mathrm{H}_{6}\). Assume the rate constant at a particular temperature is \(4.0 \times 10^{-2} \mathrm{M}^{-1} \mathrm{~s}^{-1}\) and the initial concentration of \(\mathrm{C}_{4} \mathrm{H}_{6}\) is \(0.0200 \mathrm{M}\). (a) What is its molarity after a reaction time of \(1.00 \mathrm{~h}\) ? (b) What is the time (in hours) when the \(\mathrm{C}_{4} \mathrm{H}_{6}\) concentration reaches a value of \(0.0020\) M?

You wish to determine the reaction order and rate constant for the following thermal decomposition reaction: $$ 2 \mathrm{AB}_{2} \longrightarrow \mathrm{A}_{2}+2 \mathrm{~B}_{2} $$ (a) What data would you collect? (b) How would you use these data to determine whether the reaction is zeroth order, first order, or second order? (c) Describe how you would determine the value of the rate constant.

In aqueous solution the oxidation of nitric oxide occurs according to the equation: $$ 4 \mathrm{NO}(a q)+\mathrm{O}_{2}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 4 \mathrm{HNO}_{2}(a q) $$ The reaction is second order in nitric oxide and first order in oxygen. (a) Write the rate law. (b) What is the overall reaction order? (c) How does the reaction rate change when the concentration of oxygen is doubled? (d) How does the reaction rate change when the concentration of nitric oxide is doubled and the concentration of oxygen is halved?

For the one-step reaction \(\mathrm{NO}_{2}(g)+\mathrm{CO}(g) \longrightarrow\) \(\mathrm{NO}(g)+\mathrm{CO}_{2}(g)\), draw four possible ways that the two reactant molecules can collide. Which is most likely to result in a successful reaction?

Sulfur dioxide is oxidized to sulfur trioxide in the following sequence of reactions: $$ \begin{aligned} &2 \mathrm{SO}_{2}(g)+2 \mathrm{NO}_{2}(g) \longrightarrow 2 \mathrm{SO}_{3}(g)+2 \mathrm{NO}(g) \\ &2 \mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{NO}_{2}(g) \end{aligned} $$ (a) Write the chemical equation for the overall reaction. (b) Identify any molecule that acts as a catalyst or intermediate in this reaction.
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