Chemistry
Properties of Equilibrium Constant
The properties of equilibrium constant include being temperature-dependent, providing information about the extent of a chemical reaction at equilibrium, and being independent of initial concentrations. It is a ratio of the concentrations of products to reactants at equilibrium and is used to predict the direction of a reaction under different conditions.
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6 Key excerpts on "Properties of Equilibrium Constant"
eBook - PDF- Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
- 2013(Publication Date)
- Wiley(Publisher)
Chapter Six GENERAL CONCEPTS OF CHEMICAL EQUILIBRIUM “The worst form of inequality is to try to make unequal things equal.” — Aristotle Learning Objectives WHAT ARE SOME OF THE KEY THINGS WE WILL LEARN FROM THIS CHAPTER? ● The equilibrium constant (key equations: 6.12, 6.15), pp. 194 ● Calculation of equilibrium concentrations, p. 195 ● Using Excel Goal Seek to solve one-variable equations, p. 197 ● The systematic approach to equilibrium calculations: mass balance and charge balance equations, p. 204 ● Activity and activity coefficients (key equation: 6.19), p. 211 ● Thermodynamic equilibrium constants (key equation: 6.23), p. 217 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 general 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 potentiometric calculations (Chapter 13). 6.1 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 equilibrium as the condition when the rates of the forward and reverse reactions are equal.
eBook - PDF- 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.
eBook - PDFPhysical Chemistry
Thermodynamics
- Horia Metiu(Author)
- 2006(Publication Date)
- Taylor & Francis(Publisher)
20 CHEMICAL EQUILIBRIUM: THE DEPENDENCE OF THE EQUILIBRIUM CONSTANT ON TEMPERATURE AND PRESSURE §1. Introduction. Of all the quantities calculated here, the equilibrium composi-tion is the most interesting to the practicing chemist. To calculate it, the equilibrium constant K must be known. This puts K at the center of the stage: G 0 and S 0 are of interest only as intermediate quantities for computing K . Given this central role of K , it would be useful to have equations that give the vari-ation of K with temperature and pressure. In this chapter, I derive such equations. They have three uses. 1. They lead to Le Chatelier’s principle, which allows us to predict the direction in which the equilibrium composition changes when we change T or p . These predictions are qualitative (e.g. if you raise the temperature, you make more ammonia) and require no calculations. 409 410 Dependence of Equilibrium Constant on T and p 2. The equations allow me to calculate the equilibrium constant at any temperature and pressure, if I know its value at one temperature and pressure. 3. The equations allow me to calculate the heat of reaction if I measure the equilib-rium composition at several temperatures. Since concentration measurements are accurate and heat measurements are not, this is a good method for obtaining heats of reaction. The Change of the Equilibrium Constant K with Temperature and Pressure: the Equations §2. The Change of Equilibrium Constant with Temperature. In Chapter 17, §19, I derived the equation − RT ln K = c i = 1 ν ( i ) µ 0 ( i ; T , p ) ≡ G 0 ( T , p ) (20.1) This is Eq. 17.35 for A = 0, which is the equilibrium condition. µ 0 ( i ; T , p ) is the chemical potential of the pure component i . To proceed I need to remind you of the equation (see Chapter 14, §12, Eq. 14.24) ∂ ∂ T µ 0 ( i ) T p , n = − h 0 ( i ) T 2 (20.2) Here h 0 ( i ) is the molar enthalpy of the pure component i .
eBook - PDF- John A. Olmsted, Gregory M. Williams, Robert C. Burk(Authors)
- 2020(Publication Date)
- Wiley(Publisher)
As the LiF example illustrates, the most direct way to determine the value of an equi- librium constant is to mix substances that can undergo a chemical reaction, wait until the system reaches equilibrium, and measure the concentrations of the species present once equilibrium is established. Although the calculation of an equilibrium constant requires knowledge of the equilibrium concentrations of all species for which concentrations appear in the equilibrium constant expression, stoichiometric analysis often can be used to deduce the concentration of one species from the known concentration of another species. Example 14.9 shows how to approach an equilibrium constant problem from a molecular perspective. 14. 5 Working with Equilibria 701 EX AMPLE 14.9 K eq from a Molecular View The figure represents a molecular view of a gas-phase reaction that has reached equilibrium. Assuming that each molecule in the molecular view represents a partial pressure of 1.0 bar, deter- mine K eq for this reaction. Strategy: Our seven-step approach to equilibrium problems will lead to the correct result. The first four steps are part of the strategy: View at equilibrium 2 AB AB 2 + A + + 1. Determine what is asked for. The problem asks us to calculate an equilibrium constant, K eq . 2. Identify the major chemical species. The chemical species present are the hypothetical gas-phase molecules, AB 2 , A, and AB. 3. Determine what chemical equilibria exist. The forward reaction is provided. At equilib- rium the reverse reaction proceeds at an equal rate: 2 AB ⟶ ⟵ AB 2 + A 4. Write the equilibrium constant expression. We use the chemical reaction to determine the equilibrium constant expression: K eq = ( p AB 2 ) eq ( p A ) eq ______________ ( p AB ) eq 2 Solution: 5.
eBook - PDF- Juan J. de Pablo, Jay D. Schieber(Authors)
- 2014(Publication Date)
- Cambridge University Press(Publisher)
10 Reaction equilibrium Reactions are an essential component of chemical engineering, and reaction engineering itself is a very broad discipline [43 , 107 ]. However, before we study the time evolution of chemical reactions, it is important to know the final equilibrium state that a system can reach. The equilibrium state of a chemical reaction is determined by thermodynamic principles. In this chapter we examine the thermodynamic principles that govern chemical reactions, and find methods for calculating the final composition of a mixture that results from chemical reactions. The methods consist of two general steps: determination of the equilibrium constant(s) from ther-modynamic properties of the constituent chemical compounds; and determination of the chemical composition of the equilibrium mixture from the initial composition and the equilibrium constant(s). In general, finding this composition requires the use of models with good estimates for the activities (or fugacities) of the constituent species, as shown in the previous two chapters. The equilibrium constant depends only on the temperature and the species present, not on the composition. We also consider how thermodynamic driving forces can influence reaction rates. However, few details are given in this book, since that is a topic best left for other courses and textbooks. Denat-uration of DNA strands is discussed in Section 10.9, where denaturation is applied to polymerase chain reactions (PCR). PCR is a technique that can be used to amplify small amounts of genetic material. Finally, in the last section, Section 10.10, connections are made to statistical mechanics. 10.1 A SIMPLE PICTURE: THE REACTION COORDINATE ......................................................................................................................
eBook - PDFIntroductory Chemistry
An Active Learning Approach
- Mark Cracolice, Edward Peters, Mark Cracolice(Authors)
- 2020(Publication Date)
- Cengage Learning EMEA(Publisher)
720 Chapter 18 Chemical Equilibrium 18.9 The Significance of the Value of K Goal 12 Given an equilibrium equation and the value of the equilibrium constant, identify the direction in which the equilibrium is favored. By definition, an equilibrium constant is a ratio—a fraction. The numerical value of an equilibrium constant may be very large, very small, or anyplace in between. Even though there is no defined intermediate range, equilibria with constants between 0.01 and 100 (10 22 to 10 2 ) will have appreciable quantities of all species present at equilibrium. To see what is meant by “very large” or “very small” K values, consider an equi- librium similar to the hydrogen iodide system studied in Section 18.8. If we substi- tute chlorine for iodine, the equilibrium equation is H 2 sgd 1 Cl 2 sgd (Figure 18.19) m 2 HCl(g). At 25°C, K 5 fHClg 2 fH 2 g fCl 2 g 5 2.4 3 10 33 This is a very large number—10 billion times larger than the number of particles in a mole! The only way an equilibrium constant ratio can become so huge is for the concentration of one or more reacting species to be very close to zero. If the denom- inator of a ratio is nearly zero, the value of the ratio will be very large. A near-zero denominator and large K mean the equilibrium is favored overwhelmingly in the forward direction. By contrast, if the equilibrium constant is very small, it means the concentra- tion of one or more of the species on the right-hand side of the equation is nearly zero. This puts a near-zero number in the numerator of K, and the equilibrium is strongly favored in the reverse direction. You improved your skill at writing equilibrium constant expressions. What did you learn by solving this Active Example? Practice Exercise 18.9 Write the equilibrium constant expressions for the following: a) 2 H 2 O(O) m 2 H 2 (g) 1 O 2 (g) b) CoCl 2 (s) 1 6 H 2 O(g) m CoCl 2 ? 6 H 2 O(s) Figure 18.19 Chlorine gas has a greenish-yellow color.
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