Si on observe sur une certaine durée l'équilibre NH4 I(s) ⇄NH3( g)+HI(g) établi à 600 K, à part

Si on observe sur une certaine durée l'équilibre NH4 I(s)...

Si on observe sur une certaine durée l'équilibre $$ \mathrm{NH}_4 \mathrm{I}(\mathrm{s}) \rightleftarrows \mathrm{NH}_3(\mathrm{~g})+\mathrm{HI}(\mathrm{g}) $$ établi à $600 \mathrm{~K}$, à partir de $0,25 \mathrm{~mol}$ de $\mathrm{NH}_4 \mathrm{I}$, dans un réacteur de 3 litres [Ex. 14.20], on constate une lente variation de la pression totale, due à une dissociation de l'iodure d'hydrogène ; l'équation stæchiométrique de cette dissociation s'écrit : $$ 2 \mathrm{HI}(\mathrm{g}) \stackrel{1}{\stackrel{1}{\rightleftharpoons}} \mathrm{H}_2(\mathrm{~g})+\mathrm{I}_2(\mathrm{~g}) $$ À $600 \mathrm{~K}$ la constante d'équilibre de la réaction (1) est égale à 1,69 , et celle de la réaction (-1) $\mathrm{H}_2+\mathrm{I}_2 \rightleftarrows 2 \mathrm{HI}$ est égale à 64 [Ex. 14.19]. a) Vers quelles valeurs d'équilibre les pressions partielles des quatre constituants gazeux et la pression totale tendent-elles? b) Vers quelle valeur la masse restante d'iodure d'ammonium tend-elle?

Si on observe sur une certaine durée l'équilibre
$$
\mathrm{NH}_4 \mathrm{I}(\mathrm{s}) \rightleftarrows \mathrm{NH}_3(\mathrm{~g})+\mathrm{HI}(\mathrm{g})
$$
établi à $600 \mathrm{~K}$, à partir de $0,25 \mathrm{~mol}$ de $\mathrm{NH}_4 \mathrm{I}$, dans un réacteur de 3 litres [Ex. 14.20], on constate une lente variation de la pression totale, due à une dissociation de l'iodure d'hydrogène ; l'équation stæchiométrique de cette dissociation s'écrit :
$$
2 \mathrm{HI}(\mathrm{g}) \stackrel{1}{\stackrel{1}{\rightleftharpoons}} \mathrm{H}_2(\mathrm{~g})+\mathrm{I}_2(\mathrm{~g})
$$

À $600 \mathrm{~K}$ la constante d'équilibre de la réaction (1) est égale à 1,69 , et celle de la réaction (-1) $\mathrm{H}_2+\mathrm{I}_2 \rightleftarrows 2 \mathrm{HI}$ est égale à 64 [Ex. 14.19].
a) Vers quelles valeurs d'équilibre les pressions partielles des quatre constituants gazeux et la pression totale tendent-elles?
b) Vers quelle valeur la masse restante d'iodure d'ammonium tend-elle?

Key Concepts

Ideal Gas Law in Equilibrium Calculations

The ideal gas law (PV = nRT) links the physical properties of a gas (pressure, volume, and temperature) with the amount in moles. In equilibrium problems involving gases, it allows one to convert between moles and partial pressures, providing a way to apply the equilibrium constant (expressed in pressure terms) to determine the individual pressures of the components in the mixture.

Coupled Equilibria

Coupled equilibria occur when two or more reactions share one or more species and are interdependent. A shift in one equilibrium influences the other through the common component. Analyzing these systems requires simultaneous consideration of all equilibrium expressions and the conservation of matter across the interconnected reactions.

Reaction Stoichiometry

Reaction stoichiometry describes the quantitative relationships between reactants and products as dictated by a balanced chemical equation. It enables the prediction of how changes in the quantity of one species, due either to reaction progress or simultaneous reactions, affect the amounts of other species present in the system.

Chemical Equilibrium

Chemical equilibrium is the state in which the forward and reverse reactions occur at the same rate so that the concentrations or partial pressures of the reactants and products remain constant over time. In heterogeneous equilibria, pure solids (or liquids) are not included in the equilibrium constant expression because their activity is defined as one, which simplifies the analysis of the system.

Equilibrium Constant Expressions

The equilibrium constant (K) quantitatively relates the concentrations or partial pressures of reactants and products at equilibrium according to the balanced reaction equation. For gaseous reactions, the expression is written in terms of partial pressures. This concept is crucial when evaluating how changes in a system (for example, additional reactions) affect the distribution of species.

Transcript

00:01 Okay, so in this reaction, we're starting with a solid, okay, and we're producing two gases.

00:10 And we want to figure out what the total pressure is in the system.

00:14 Well, our gases are in a one -to -one ratio, okay, this hi.

00:21 So i can say i'm going to form equal amounts of them.

00:25 I'll call them both x, and since this is a solid, i won't worry about it right now.

00:29 It's not going to contribute to the partial pressure or be involved in our k expression so k equals the partial pressure of n h3 times the partial pressure of h i and we know k is 0 .215 so that just equals x squared okay so x is going to equal 0 .464 atmospheres and that's the partial pressure of both of them.

01:02 So that's your partial pressure of nh3, and it's also the partial pressure of the hi.

01:12 So our total pressure, we'll just go ahead and add those together, and we'll get 0 .9 -28 -atm, and that'll be our total pressure.

01:24 Okay? and then what we want to do is try to figure out how much of this solid is left when we're done.

01:31 Okay? so i know i formed, point, 0 .464 atm of ammonia.

01:42 So i'm going to use that to find how many moles of ammonia were produced.

01:45 I'm going to use the ideal gas law.

01:48 Pv equals nrt.

01:50 I'm going to rearrange it a little bit for moles...