Chemistry
Equilibrium Constant Kp
The equilibrium constant Kp is a measure of the extent to which a chemical reaction reaches equilibrium in terms of partial pressures of gases. It is defined as the ratio of the partial pressures of the products to the partial pressures of the reactants, each raised to the power of their respective stoichiometric coefficients, at equilibrium. A higher Kp value indicates a greater concentration of products at equilibrium.
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- Kevin Reel, Derrick C. Wood, Scott A. Best(Authors)
- 2014(Publication Date)
- Research & Education Association(Publisher)
.• If the reaction quotient is greater than the equilibrium expression, the ratio is too large. According to Le Chatelier’s principle, the reaction will shift toward the left and create reactants. In summary, when Q > K, the reaction proceeds to the left toward reactants. When Q < K, the reaction proceeds to the right toward products. When Q = K, the reaction is at equilibrium.EXAMPLE: See the following reaction:, Keq = 5.9 × 10–2 . The molar concentrations of the species are: [N2 ] = 0.40 M, [H2 ] = 0.80 M, and [NH3 ] = 0.20 M. In which direction will the reaction proceed as it begins to establish equilibrium?SOLUTION:Because Q > Keq , the reaction will shift left toward the reactants.Equilibrium Constants for Gaseous Reactions
• Concentrations of reactants and products in the gas phase may be expressed in units of molarity (moles/liter) or partial pressures (atm, Pa, mmHg).• The ideal gas law (PV = nRT) is used to convert between the equilibrium constant expressed in moles/liter (Kc ) to the equilibrium constant expressed in gas pressures (Kp ).• The equilibrium expression for Kp only contains those species that are in the gas phase. For example, following is the Kp for the formation of ammonia from nitrogen gas and hydrogen gas:EXAMPLE:All phase changes exist in a state of equilibrium. For example, at 100°C, water exists as both a liquid and a gas in equilibrium with one another.DID YOUKNOW?TEST TIPEXAMPLE:If pressures are given for products and reactants in an equilibrium, be sure to write the expression for Kp and not Kc .The value of Kp for the following reaction is 8.3 × 10–3 at 700 K. What is the value for Kc
eBook - ePubChemistry
Concepts and Problems, A Self-Teaching Guide
- Richard Post, Chad Snyder, Clifford C. Houk(Authors)
- 2020(Publication Date)
- Jossey-Bass(Publisher)
equilibrium constant is the ratio of the concentration of the products divided by the concentration of the reactants at equilibrium and at a specified temperature. The equilibrium constant (product concentrations divided by reactant concentrations) is valid only at a specified temperature after the reaction has gone to (completion, equilibrium) __________Answer: equilibriumBelow is a reversible reaction and the expression for the equilibrium constant for this reversible reaction.The symbol Keqrepresents the equilibrium constant and the brackets [] represent the concentration (usually in moles per liter) of each product and reactant. Look at the placement of each reactant and product in the equilibrium constant expression. In the equilibrium constant expression for a reversible reaction, the (products, reactants) ____________ are located in the numerator or upper part of the fraction and the (products, reactants) ________________ are located in the denominator or lower part of the fraction.Answer: products; reactantsThe standard equation, then, for Keqis as follows.Write the equilibrium constant expression for the following reversible reaction.Keq= ___________________Answer:(Since there are two products, they should be placed in the upper part of the fraction. The one reactant belongs in the lower part of the fraction.)
eBook - ePub- Adrian Dingle, Derrick C. Wood(Authors)
- 2014(Publication Date)
- Research & Education Association(Publisher)
concentrations of reactants and products. In fact, they are very rarely the same numerical value.
5. Graphically, the concentrations of reactants and products can be represented as: i. For a product-favored reaction: ii. For a reactant-favored reaction: C. Quantitative TreatmentAll concentrations being constant (not changing) is not the same thing as concentrations being equal (all having the same value). This is a common misconception and one that you must not have.1. Equilibrium constants for gaseous reactions: Kp , Kci. Equilibrium constant relates the concentrations of reactants and products at equilibrium at a given temperature.For the general reaction: aA + bB cC + dDThe equilibrium constantProduct molar concentrations are in the numerator.Reactant molar concentrations are in the denominator.Each concentration is raised to the power of its stoichiometric coefficient in the balanced chemical equation.Pure solids and pure liquids (e.g., water) are not placed into the expression.The expression Kc indicates that concentrations are used (moles per liter).The expression Kp indicates that partial pressures are used (pressure units, often atm) and only gases are included.Examples1. N2 (g) + 3H2 (g) 2NH3
eBook - ePub- Kevin R. Reel(Author)
- 2013(Publication Date)
- Research & Education Association(Publisher)
law of mass action .• Pure substances, such as water or solids, do not show up in the equilibrium expression; only molar solutions or, as with Kp , gaseous pressures.• The equilibrium constant for a multi-step process is equal to the product of the equilibrium constants for each step.Example: For a set of three reactions that add to equal a total reaction,Ktotal=K1×K2×K3• The equilibrium constant for a reverse reaction is the inverse of the equilibrium constant for a forward reaction.Example: • There are different equilibrium constants for different types of reactions.REACTION QUOTIENT, Q
• The reaction quotient, Q, can be used to calculate the direction and degree to which a reaction will shift when new products or reactants are added (Le Chatelier’s principle). • The reaction quotient for a reaction is found using the same ratio as the equilibrium constant, but at non-equilibrium conditions.For the reaction, aA + bB ↔ cC + dD• If the reaction quotient is greater than the equilibrium expression, there are more products than there would be at equilibrium. By Le Chatelier’s principle, the reaction will shift toward equilibrium by using products and producing more reactants.When Q > K, reaction proceeds to the left toward reactants. When Q < K, reaction proceeds to the right toward products. When Q = K, reaction is at equilibrium. Example: In the following reaction,N2 (g ) + 3 H2 (g ) ↔ 2 NH3 (g )The Kc is 5.9 × 10−2 . The molar concentrations of each reactant and product are: [N2 ] = 0.40 M, [H2 ] = 0.80 M, and [NH3
eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
This is because the reaction quotient does not describe the ratio of product to reactant concentrations of a reversible reaction at equilibrium as with K eq. Instead, it describes the ratio of the product to reactant concentrations for a reversible reaction at any point in the reaction. The determination of the value of the reaction quotient from the concentrations of the reactants and products at any particular point in time helps to determine the direction the reaction is likely to proceed from that point. A comparison of the value of the reaction quotient (Q) with the value of the equilibrium constant K eq tells which way the reaction will shift in order to reach chemical equilibrium as demonstrated in Fig. 7.6. Fig. 7.6 A comparison of the values of Q and K eq at any point in a chemical reaction tells how the concentrations of reactants and products will change for the reaction to achieve equilibrium. When Q is less than K eq, the ratio of [products] to [reactants] for Q is less than for K eq and there is an excess of reactants. This means that the reaction has not yet reached equilibrium. In order to reach equilibrium, more products must be formed from the available reactants. So the reaction proceeds to the right. But, when Q is greater than K eq, there are an excess of products. To achieve equilibrium, the reaction must proceed to the left using up the excess products and forming more reactants. In summary; • Q < K eq —The ratio of products to reactants is less than that for the system at equilibrium. More products will be formed from the excess reactants for the reaction to reach equilibrium. • Q > K eq —The ratio of products to reactants is larger than that for the system at equilibrium. More products are present than there would be at equilibrium. The reaction must produce more reactants from the excess products for the reaction to reach equilibrium. • Q = K eq —The reaction is already at equilibrium
eBook - ePub- Mark E. Davis, Robert J. Davis(Authors)
- 2013(Publication Date)
- Dover Publications(Publisher)
APPENDIX AReview of ChemicalEquilibria
A.1 | Basic Criteria for Chemical Equilibriumof Reacting Systems
The basic criterion for equilibrium with a single reaction is:where ΔG is the Gibbs function, NCOMP is the number of components in the system,viis the stoichiometric coefficient of species i , and i is the chemical potential of species i . The chemical potential is:where R g is the universal gas constant, is the standard chemical potential of species i in a reference state such that a i = 1, and a i is the activity of species i . The reference states are: (1) for gases (i.e., 0 = 1) (ideal gas, P = 1 atm) where is the fugacity, (2) for liquids, the pure liquid at T and one atmosphere, and (3) for solids, the pure solid at T and one atmosphere. If multiple reactions are occurring in a network, then Equation (A.1.1) can be extended to give:where NRXN is the number of independent reactions in the network.In general it is not true that the change in the standard Gibbs function, ΔG 0 , is zero. Thus,Therefore, or by using Equation (A.1.2): Now consider the general reaction:Application of Equation (A.1.6) to Equation (A.1.7) and recalling that ΔG = 0 at equilibrium gives:Thus, the equilibrium constant K a is defined as:Differentiation of Equation (A.1.8) with respect to T yields:Note that ΔG 0 = ΔH 0 − T ΔS 0 , where ΔH 0 and ΔS0 are the standard enthalpy and entropy, respectively, and differentiation of this expression with respect to T gives:Equating Equations (A. 1.10) and (A.1.l1) provides the functional form for the temperature dependence of the equilibrium constant:or after integration (assume ΔH 0 is independent of T ):Notice that when the reaction is exothermic (ΔH 0 is negative), K a increases with decreasing T
eBook - ePubEnzyme Kinetics
Rapid-Equilibrium Applications of Mathematica
- Robert A. Alberty(Author)
- 2011(Publication Date)
- Wiley(Publisher)
equation 1.2-1 becomes(1.2-3)At constant temperature and pressure, the Gibbs energy of reaction is given by(1.2-4)Equation 1.2-2 for a chemical species can be written as(1.2-5)where Δf G j ° is the standard Gibbs energy of formation of species j . In chemical thermodynamics, this equation is used for ideal solutions and activity coefficients are introduced, but in biochemical thermodynamics this equation is used at the specified ionic strength and the Debye-Huckel equation is used to account for the effects of ionic strength. Substituting equation 1.2-5 into equation 1.2-4 yields(1.2-6)where Δr G ° is the standard Gibbs energy of reaction Σv j Δf G j °, and Q is the reaction quotient. At equilibrium, Δr G = 0 and Q becomes the chemical equilibrium constant K .(1.2-7)Chemists have developed tables of values of Δf G j ° for species with respect to the elements. In other words, the Δf G j ° for elements are taken to be zero in a defined reference state. These tables can be used to calculate a number of chemical equilibrium constants K that are of biochemical interest, but the Natonal Bureau of Standards Tables [10] are limited to C2 .More information about chemical thermodynamics is given in text books on chemical thermodynamics like Beattie and Oppenheim [8], and in my two books on biochemical thermodynamics [20,23].The first publication that applied chemical thermodynamics to biochemical reactions was by Burton and Krebs [2], and the first table of thermodynamic properties was published by Burton in Krebs and Kornberg, Energy Transformations in Living Matter [4]. Burton recognized that the equilibrium constants for enzyme-catalyzed reactions together with the standard Gibbs energies Δf G j ° of species determined with chemical methods can yield Δf G j ° biochemical species. He made a table that could be used to calculate equilibrium constants of biochemical reactions that have not been studied. But he ran into problems with reactants like ATP that are sums of species at pH 7. Wilhoit [5] extended these tables, but ATP remained a problem. I became involved with ATP through electrophoresis, and my group determined acid dissociation constants and magnesium complex dissociation constants of ATP [1,3]. I worked on the thermodynamics of petroleum processing in the 1980–1990 period and learned that when the concentration of a species (like H+
