Dynamic Equilibrium

9 Key excerpts on "Dynamic Equilibrium"

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  eBook - PDF

  Foundations of College Chemistry

  • Morris Hein, Susan Arena, Cary Willard(Authors)
  • 2016(Publication Date)
  • Wiley

    (Publisher)

  EXAMPLE 16.2 Is a reversible chemical reaction at equilibrium a static or a dynamic system? Explain. SOLUTION A reversible chemical reaction is a dynamic system in which two opposing reactions are taking place at the same time and at the same rate of reaction. P R A C T I C E 1 6 . 2 What symbolism is used in a chemical equation to indicate that a chemical reaction is reversible? 372 CHAPTER 16 • Chemical Equilibrium 16.3 Le Châtelier’s Principle Use Le Châtelier’s principle to predict the changes that occur when concentration, temperature, or volume is changed in a system at equilibrium. In 1888, the French chemist Henri Le Châtelier (1850–1936) set forth a simple, far- reaching generalization on the behavior of equilibrium systems. This generalization, known as Le Châtelier’s principle, states: If a stress is applied to a system in equilibrium, the system will respond in such a way as to relieve that stress and restore equilibrium under a new set of conditions. The application of Le Châtelier’s principle helps us predict the effect of changing con- ditions in chemical reactions. We will examine the effect of changes in concentration, temperature, and volume. Effect of Concentration on Equilibrium The manner in which the rate of a chemical reaction depends on the concentration of the reactants must be determined experimentally. Many simple, one-step reactions result from a collision between two molecules or ions. The rate of such one-step reactions can be altered by changing the concentration of the reactants or products. An increase in concentration of the reactants provides more individual reacting species for collisions and results in an increase in the rate of reaction. An equilibrium is disturbed when the concentration of one or more of its components is changed. As a result, the concentration of all species will change, and a new equilibrium mix- ture will be established.

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  Chemistry 2e

  • Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
  • 2019(Publication Date)
  • Openstax

    (Publisher)

  INTRODUCTION CHAPTER 13 Fundamental Equilibrium Concepts 13.1 Chemical Equilibria 13.2 Equilibrium Constants 13.3 Shifting Equilibria: Le Châtelier’s Principle 13.4 Equilibrium Calculations Imagine a beach populated with sunbathers and swimmers. As those basking in the sun get too hot, they enter the surf to swim and cool off. As the swimmers tire, they return to the beach to rest. If the rate at which sunbathers enter the surf were to equal the rate at which swimmers return to the sand, then the numbers (though not the identities) of sunbathers and swimmers would remain constant. This scenario illustrates a dynamic phenomenon known as equilibrium, in which opposing processes occur at equal rates. Chemical and physical processes are subject to this phenomenon; these processes are at equilibrium when the forward and reverse reaction rates are equal. Equilibrium systems are pervasive in nature; the various reactions involving carbon dioxide dissolved in blood are examples (see Figure 13.1). This chapter provides a thorough introduction to the essential aspects of chemical equilibria. 13.1 Chemical Equilibria LEARNING OBJECTIVES By the end of this section, you will be able to: • Describe the nature of equilibrium systems • Explain the dynamic nature of a chemical equilibrium The convention for writing chemical equations involves placing reactant formulas on the left side of a reaction Figure 13.1 Transport of carbon dioxide in the body involves several reversible chemical reactions, including hydrolysis and acid ionization (among others). CHAPTER OUTLINE arrow and product formulas on the right side. By this convention, and the definitions of “reactant” and “product,” a chemical equation represents the reaction in question as proceeding from left to right. Reversible reactions, however, may proceed in both forward (left to right) and reverse (right to left) directions.

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  Chemistry

  Concepts and Problems, A Self-Teaching Guide

  • Richard Post, Chad Snyder, Clifford C. Houk(Authors)
  • 2020(Publication Date)
  • Jossey-Bass

    (Publisher)

  12 Chemical Equilibrium

  You have just learned several properties of solutions (mixtures of solids, liquids, and gases). We have discussed reactions that go to completion (reactants totally consumed, leaving only new products) in Chapter 5 and electrolytes that dissociate completely in water in Chapter 11 . Both of these concepts imply a one-way reaction, continuous movement toward the product side. However, in Chapter 10 we discussed a Dynamic Equilibrium where the rate of evaporation equals the rate of condensation, that is, the reactions are “reversible.”

  Many chemical reactions are reversible. The products formed react to give back the original reactants, even as the reactants are forming more products. After some time, both the forward and reverse reactions will be going on at the same rate. When this occurs, the reaction is said to have reached equilibrium. There is no further change in the amount of any reactant or product, though both reactions still go on (forever). Since there are many such reactions that appear to go only partway to completion, their study is of major importance to the chemist.

  We will discuss several types of equilibrium in this chapter, along with their associated problems and concepts. You will use the concept of molarity you just learned in Chapter 11

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  General Chemistry for Engineers

  • Jeffrey Gaffney, Nancy Marley(Authors)
  • 2017(Publication Date)
  • Elsevier

    (Publisher)

  Chapter 7

  Chemical Equilibrium

  Abstract

  This chapter reviews the principles of chemical equilibrium beginning with the concept that many chemical reactions are reversible. These reversible reactions achieve a Dynamic Equilibrium where the rate of the forward reaction equals the rate of the reverse reaction. This Dynamic Equilibrium is described by an equilibrium constant, which is determined from the concentrations of the reactants and products at equilibrium. The magnitude of the equilibrium constant is presented as measure of the extent that the forward and reverse reactions take place. Le Chatelier's principle is introduced and the changes in reaction conditions that can disturb a chemical equilibrium are reviewed. The response of the chemical reaction to these changes in reaction conditions is explained in detail. The reaction quotient is presented as a means of determining the direction the reaction is likely to proceed.

  Keywords

  Reversible reactions; Dynamic Equilibrium; Equilibrium constant; Partial pressures; Le Chatelier's principle; Exothermic; Endothermic; Reaction quotient

  Outline

  7.1  

  Reversible Reactions

  7.2  

  The Equilibrium Constant

  7.3  

  Relationships Between Equilibrium Constants

  7.4  

  Le Chatelier's Principle: Disturbing a Chemical Equilibrium

  7.5  

  The Reaction Quotient

  Important Terms

  Study Questions

  Problems

  7.1 Reversible Reactions

  The idea that a chemical reaction can be reversible was introduced by Claude Louis Berthollet in 1803 when he observed the formation of sodium carbonate crystals at the edge of a limestone salt lake in Egypt (Fig. 7.1 ). Since the salt lake was a landlocked body of water with a very high concentration of dissolved sodium chloride (> 3 g/L) and other minerals, he knew that the formation of these crystals must be a result of the following chemical reaction;

  Fig. 7.1

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  Chemistry: Atoms First 2e

  • Edward J. Neth, Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
  • 2019(Publication Date)
  • Openstax

    (Publisher)

  INTRODUCTION CHAPTER 13 Fundamental Equilibrium Concepts 13.1 Chemical Equilibria 13.2 Equilibrium Constants 13.3 Shifting Equilibria: Le Châtelier’s Principle 13.4 Equilibrium Calculations Imagine a beach populated with sunbathers and swimmers. As those basking in the sun get too hot, they enter the surf to swim and cool off. As the swimmers tire, they return to the beach to rest. If the rate at which sunbathers enter the surf were to equal the rate at which swimmers return to the sand, then the numbers (though not the identities) of sunbathers and swimmers would remain constant. This scenario illustrates a dynamic phenomenon known as equilibrium, in which opposing processes occur at equal rates. Chemical and physical processes are subject to this phenomenon; these processes are at equilibrium when the forward and reverse reaction rates are equal. Equilibrium systems are pervasive in nature; the various reactions involving carbon dioxide dissolved in blood are examples (see Figure 13.1). This chapter provides a thorough introduction to the essential aspects of chemical equilibria. We now have a good understanding of chemical and physical change that allow us to determine, for any given process: 1. Whether the process is endothermic or exothermic 2. Whether the process is accompanied by an increase of decrease in entropy 3. Whether a process will be spontaneous, non-spontaneous, or what we have called an equilibrium process Recall that when the value ∆G for a reaction is zero, we consider there to be no free energy change—that is, no Figure 13.1 Transport of carbon dioxide in the body involves several reversible chemical reactions, including hydrolysis and acid ionization (among others). CHAPTER OUTLINE free energy available to do useful work. Does this mean a reaction where ΔG = 0 comes to a complete halt? No, it does not.

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  Chemistry

  With Inorganic Qualitative Analysis

  • Therald Moeller(Author)
  • 2012(Publication Date)
  • Academic Press

    (Publisher)

  Section 25.6 .)

  Reactions in the gas phase are particularly easy to study as examples of equilibrium. Suppose some nitrogen dioxide, NO2 , is confined in a closed vessel of fixed volume at a constant temperature. Immediately, the formation of dinitrogen tetroxide, N2 O4 , begins. We choose to call this the “forward” reaction:

  As soon as some N2 O4 has formed the “reverse” reaction begins—the breakdown of N2 O4 molecules to give NO2 molecules. The rates of the forward and reverse reactions and the amounts of the two gases continue to change until equilibrium is reached. At this point the “opposing forces”—the forward and reverse reactions—are balanced.

  The forward and reverse reactions proceed at equal rates, with the net result that the amounts of the two gases are constant. It is a dynamic state. Each time an N2 O4 molecule is formed, two molecules of NO2 disappear; each time an N2 O4 molecule decomposes, two new NO2 molecules appear. As long as the temperature and pressure remain constant and nothing is added to or taken from the mixture, the equilibrium state remains unchanged.

  What happens if we pump more NO2 molecules into the vessel? With more NO2 available, whatever collisions precede the formation of N2 O4 occur more often. If the formation of N2 O4 proceeds faster than its decomposition, the amount of N2 O4 increases. As the concentration of N2 O4 increases, the rate of its decomposition also increases. Changes in the concentrations and in the forward and reverse reaction rates continue until equilibrium has been reached once more (Figure 16.1

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  Introduction to Geochemistry

  Principles and Applications

  • Kula C. Misra(Author)
  • 2012(Publication Date)
  • Wiley-Blackwell

    (Publisher)

  This rearrangement of atoms from one configuration to another is called a ­ chemical reaction. It is possible that the forward ­reaction X + Y → C + D will result in complete conversion of X and Y into C and D. Many chemical reactions, however, remain incomplete because of the reverse reaction C + D → X + Y. The reaction is considered to have attained chemical ­equilibrium if and when the rate of the forward reaction becomes equal to the rate of the reverse reaction. If conditions of the experiment remain unchanged, then at equilibrium all the four substances would coexist without any further change in their concentrations with time. For a chemical reaction, the terms reversibility and equilibrium are often used interchangeably because a reversible process is conceptualized as a ­process that proceeds in such infinitely small steps that the system is at equilibrium for every step. Real geochemical systems seldom, if ever, attain a state of equilibrium, but the equilibrium model, because of its simple mathematical relationships, serves as a useful reference for evaluating chemical reactions. An irreversible chemical reaction is unidirectional and so can never achieve equilibrium. A reaction may also not attain equilibrium either because the rate of the forward or the reverse reaction is too slow or because one or more of the products are removed from the system. If a chemical reaction occurs as written at the molecular level (i.e., without involving intermediate steps), then its rate is proportional to the concentrations of the reacting substances (see Chapter 9 for a more elaborate discussion of rates of chemical reactions)

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  Analytical Chemistry

  • Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
  • 2020(Publication Date)
  • Wiley

    (Publisher)

  Chapter 5 GENERAL CONCEPTS OF CHEMICAL EQUILIBRIUM “The worst form of inequality is to try to make unequal things equal.” —Aristotle KEY THINGS TO LEARN FROM THIS CHAPTER The equilibrium constant (key Equations: 5.12, 5.15) Calculation of equilibrium concentrations Using Excel Goal Seek to solve one-variable equations The systematic approach to equilibrium calculations: mass balance and charge balance equations Activity and activity coefficients (key Equation: 5.19) ThermoDynamic Equilibrium constants (key Equation: 5.23) Even though in a chemical reaction the reactants may almost quantitatively react to form the products, reactions never go in only one direction. In fact, reactions reach an equilibrium in which the rates of reactions in both directions are equal. In this chapter we review the equilibrium concept and the equilibrium constant and describe gen- eral approaches for calculations using equilibrium constants. We discuss the activity of ionic species along with the calculation of activity coefficients. These values are required for calculations using thermoDynamic Equilibrium constants, that is, for the diverse ion effect, described at the end of the chapter. They are also used in potentio- metric calculations (Chapter 20). 5.1 Chemical Reactions: The Rate Concept In 1863 Guldberg and Waage described what we now call the law of mass action, which states that the rate of a chemical reaction is proportional to the “active masses” of the reacting substances present at any time. The active masses may be concentrations or pressures. Guldberg and Waage derived an equilibrium constant by defining equilib- rium as the condition when the rates of the forward and reverse reactions are equal.

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  Biosimulation

  Simulation of Living Systems

  • Daniel A. Beard(Author)
  • 2012(Publication Date)
  • Cambridge University Press

    (Publisher)

  5 Chemical reaction systems: thermodynamics and chemical equilibrium Overview This and the following two chapters are focused on analyzing and simulating chem- ical systems. These chapters will introduce basic concepts of thermodynamics and kinetics for application to biochemical systems, such as biochemical synthesis, cel- lular metabolism and signaling processes, and gene regulatory networks. Although we have seen examples of chemical kinetics in previous chapters, notably in Sections 2.3 and 2.4, in those examples we developed the expressions governing the chemistry more from intuition than from a physical theory. One of the primary goals here will be to develop a formal physical/chemical foundation for analyzing and simulating complex biochemical systems. As is our practice throughout this book, these concepts will be applied to analyze real data (and understand the behavior of real systems) later in this chapter and elsewhere. Yet, because the rules governing the behavior of biochemical systems are grounded in thermodynamics, we must begin our investigation into chemical systems by establishing some fundamental concepts in chemical thermo- dynamics. The concept of free energy is particularly crucial to understanding thermodynamic driving forces in chemistry. We will see that both a physical definition and an intuitive understanding of free energy require physical definitions and intuitive understandings of temperature and entropy. All of this means that this chapter will begin with some abstract thought experiments and derivations of physical concepts. 5.1 Temperature, pressure, and entropy 5.1.1 Microstates and macrostates All thermodynamic theory arises from the fact that physical systems composed of many atoms and/or molecules attain a large number (often a practically infinite 146 Chemical reaction systems: thermodynamics and chemical equilibrium number) of microstates under defined macroscopic conditions, such as temper- ature, pressure, and volume.

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