“equilibrium constant”的概念、定义、翻译、参考文献-科学参考

equilibrium constant

Chemistry

For a reversible reaction of the type
$x A + y B ⇌ z C + w D$
chemical equilibrium occurs when the rate of the forward reaction equals the rate of the back reaction, so that the concentrations of products and reactants reach steady-state values. It can be shown that at equilibrium the ratio of concentrations
${[C]}^{z} {[D]}^{w} / {[A]}^{x} {[B]}^{y}$
is a constant for a given reaction and fixed temperature, called the equilibrium constant K_c (where the c indicates concentrations have been used). Note that, by convention, the products on the right-hand side of the reaction are used on the top line of the expression for equilibrium constant. This form of the equilibrium constant was originally introduced in 1867 by Cato Maximilian Guldberg and Peter Waage using the law of mass action. They derived the expression by taking the rate of the forward reaction
$k_{f} {[A]}^{x} {[B]}^{y}$
and that of the back reaction
$k_{b} {[C]}^{z} {[D]}^{w}$
Since the two rates are equal at equilibrium, the equilibrium constant K_c is the ratio of the rate constants k_f/k_b. The principle that the expression is a constant is known as the equilibrium law or law of chemical equilibrium.
The equilibrium constant shows the position of equilibrium. A low value of K_c indicates that [C] and [D] are small compared to [A] and [B]; i.e. that the back reaction predominates. It also indicates how the equilibrium shifts if concentration changes. For example, if [A] is increased (by adding A) the equilibrium shifts towards the right so that [C] and [D] increase, and K_c remains constant.
For gas reactions, partial pressures are used rather than concentrations. The symbol K_p is then used. Thus, in the example above
$K_{p} = p_{C^{z}} p_{D^{w}} / p_{A^{x}} p_{B^{y}}$
It can be shown that, for a given reaction K_p = K_c(RT)^Δν‎, where Δν‎ is the difference in stoichiometric coefficients for the reaction (i.e. z + w – x – y). Note that the units of K_p and K_c depend on the numbers of molecules appearing in the stoichiometric equation. The value of the equilibrium constant depends on the temperature. If the forward reaction is exothermic, the equilibrium constant decreases as the temperature rises; if endothermic it increases (see also van’t Hoff’s isochore).
The expression for the equilibrium constant can also be obtained by thermodynamics; it can be shown that the standard equilibrium constant K^⦵ is given by exp(−Δ‎G^⦵/RT), where Δ‎G^⦵ is the standard Gibbs free energy change for the complete reaction. Strictly, the expressions above for equilibrium constants are true only for ideal gases (pressure) or infinite dilution (concentration). For accurate work activities are used.

Chemical Engineering

A reversible process, chemical or physical, in a closed system will eventually reach a state of equilibrium. The equilibrium is dynamic and may be considered as a state at which the rate of the process in one direction exactly balances the rate in the opposite direction. For a chemical reaction, the equilibrium concentrations of the reactants and products will remain constant providing the conditions remain unchanged. For the homogenous system:
$w A + x B \leftrightarrow y C + z D$

the ratio of the molar concentrations of products to reactants remains constant at a fixed temperature:
$K_{C} = \frac{{[C]}^{y} {[D]}^{z}}{{[A]}^{w} {[B]}^{x}}$

where K_C is the equilibrium constant and the square brackets indicate equilibrium concentrations. The relationship is known as the equilibrium law. For example, for the Haber process for the synthesis of ammonia, nitrogen is reacted with hydrogen as:
$N_{2} (g) + 3 H_{2} (g) \leftrightarrow 2 N H_{3} (g)$

The equilibrium constant is expressed as partial pressures as:
$K_{C} = \frac{{[N H_{3}]}^{2}}{[N_{2}] {[H_{2}]}^{3}} = \frac{p_{N H_{3}}^{2}}{p_{N_{2}} . p_{H_{2}}^{3}}$

单词	equilibrium constant
释义	equilibrium constant Chemistry For a reversible reaction of the type $x A + y B ⇌ z C + w D$ chemical equilibrium occurs when the rate of the forward reaction equals the rate of the back reaction, so that the concentrations of products and reactants reach steady-state values. It can be shown that at equilibrium the ratio of concentrations ${[C]}^{z} {[D]}^{w} / {[A]}^{x} {[B]}^{y}$ is a constant for a given reaction and fixed temperature, called the equilibrium constant K_c (where the c indicates concentrations have been used). Note that, by convention, the products on the right-hand side of the reaction are used on the top line of the expression for equilibrium constant. This form of the equilibrium constant was originally introduced in 1867 by Cato Maximilian Guldberg and Peter Waage using the law of mass action. They derived the expression by taking the rate of the forward reaction $k_{f} {[A]}^{x} {[B]}^{y}$ and that of the back reaction $k_{b} {[C]}^{z} {[D]}^{w}$ Since the two rates are equal at equilibrium, the equilibrium constant K_c is the ratio of the rate constants k_f/k_b. The principle that the expression is a constant is known as the equilibrium law or law of chemical equilibrium. The equilibrium constant shows the position of equilibrium. A low value of K_c indicates that [C] and [D] are small compared to [A] and [B]; i.e. that the back reaction predominates. It also indicates how the equilibrium shifts if concentration changes. For example, if [A] is increased (by adding A) the equilibrium shifts towards the right so that [C] and [D] increase, and K_c remains constant. For gas reactions, partial pressures are used rather than concentrations. The symbol K_p is then used. Thus, in the example above $K_{p} = p_{C^{z}} p_{D^{w}} / p_{A^{x}} p_{B^{y}}$ It can be shown that, for a given reaction K_p = K_c(RT)^Δν‎, where Δν‎ is the difference in stoichiometric coefficients for the reaction (i.e. z + w – x – y). Note that the units of K_p and K_c depend on the numbers of molecules appearing in the stoichiometric equation. The value of the equilibrium constant depends on the temperature. If the forward reaction is exothermic, the equilibrium constant decreases as the temperature rises; if endothermic it increases (see also van’t Hoff’s isochore). The expression for the equilibrium constant can also be obtained by thermodynamics; it can be shown that the standard equilibrium constant K^⦵ is given by exp(−Δ‎G^⦵/RT), where Δ‎G^⦵ is the standard Gibbs free energy change for the complete reaction. Strictly, the expressions above for equilibrium constants are true only for ideal gases (pressure) or infinite dilution (concentration). For accurate work activities are used. Chemical Engineering A reversible process, chemical or physical, in a closed system will eventually reach a state of equilibrium. The equilibrium is dynamic and may be considered as a state at which the rate of the process in one direction exactly balances the rate in the opposite direction. For a chemical reaction, the equilibrium concentrations of the reactants and products will remain constant providing the conditions remain unchanged. For the homogenous system: $w A + x B \leftrightarrow y C + z D$ the ratio of the molar concentrations of products to reactants remains constant at a fixed temperature: $K_{C} = \frac{{[C]}^{y} {[D]}^{z}}{{[A]}^{w} {[B]}^{x}}$ where K_C is the equilibrium constant and the square brackets indicate equilibrium concentrations. The relationship is known as the equilibrium law. For example, for the Haber process for the synthesis of ammonia, nitrogen is reacted with hydrogen as: $N_{2} (g) + 3 H_{2} (g) \leftrightarrow 2 N H_{3} (g)$ The equilibrium constant is expressed as partial pressures as: $K_{C} = \frac{{[N H_{3}]}^{2}}{[N_{2}] {[H_{2}]}^{3}} = \frac{p_{N H_{3}}^{2}}{p_{N_{2}} . p_{H_{2}}^{3}}$
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