Question

The Deacon process for producing chlorine gas from hydrogen chloride is used in situations where \(\mathrm{HCl}\) is available as a by-product from other chemical processes. $$\begin{aligned} 4 \mathrm{HCl}(\mathrm{g})+\mathrm{O}_{2}(\mathrm{g}) \rightleftharpoons 2 \mathrm{H}_{2} \mathrm{O}(\mathrm{g})+2 \mathrm{Cl}_{2}(\mathrm{g}) & \\ \Delta H^{\circ}=&-114 \mathrm{kJ} \end{aligned}$$ A mixture of \(\mathrm{HCl}, \mathrm{O}_{2}, \mathrm{H}_{2} \mathrm{O},\) and \(\mathrm{Cl}_{2}\) is brought to equilibrium at \(400^{\circ} \mathrm{C}\). What is the effect on the equilibrium amount of \(\mathrm{Cl}_{2}(\mathrm{g})\) if (a) additional \(\mathrm{O}_{2}(\mathrm{g})\) is added to the mixture at constant volume? (b) \(\mathrm{HCl}(\mathrm{g})\) is removed from the mixture at constant volume? (c) the mixture is transferred to a vessel of twice the volume? (d) a catalyst is added to the reaction mixture? (e) the temperature is raised to \(500^{\circ} \mathrm{C} ?\)

Short answer

(a) The amount of \(Cl_{2}(g)\) would increase. (b) The amount of \(Cl_{2}(g)\) would decrease. (c) The amount of \(Cl_{2}(g)\) would decrease. (d) The amount of \(Cl_{2}(g)\) would remain the same. (e) The amount of \(Cl_{2}(g)\) would decrease.

Step-by-step solution

Additional O2 is added to the mixture at constant volume

Adding more \(O_{2}(g)\) would add more reactants to the left side of the equation, so the reaction would shift to the right to maintain equilibrium. This would increase the amount of \(Cl_{2}(g)\).

HCl is removed from the mixture at constant volume

Removing \(HCl(g)\) would reduce the amount of reactants on the left side of the equation, which would cause the equilibrium to shift to the left to compensate.The amount of \(Cl_{2}(g)\) would decrease as a result.

The mixture is transferred to a vessel of twice the volume

Increasing the volume decreases the pressure and therefore shifts the equilibrium towards the side with the greater number of moles. In this case, there are more moles on the left side of the equation, therefore the amount of \(Cl_{2}(g)\) would decrease.

A catalyst is added to the reaction mixture

Adding a catalyst would speed up both the forward and backward reactions equally. It does not shift the equilibrium, so the amount of \(Cl_{2}(g)\) remains the same.

The temperature is raised to 500C

Increasing the temperature would favor the endothermic reaction or the reaction that absorbs heat. However, the given reaction is exothermic as shown by an enthalpy change, \(\Delta H\), of -114 kJ indicating that heat is released. Therefore, an increase in temperature would shift the reaction to the left, decreasing the amount of \(Cl_{2}(g)\)

Key concepts

Equilibrium Reactions

Chemical reactions often reach a state called equilibrium, where the rate of the forward reaction equals the rate of the backward reaction. This means that the concentrations of reactants and products remain constant over time. However, it's important to note that equilibrium does not mean the reactants and products are equal in concentration. It simply means their ratios are stable. Understanding equilibrium is crucial for processes like the Deacon process, where multiple factors can shift the reaction toward more reactants or more products. Le Chatelier's Principle aids in predicting these shifts. It states that if an external change is applied to a system at equilibrium, the system adjusts to minimize that change.

• Adding a reactant will typically shift the equilibrium to produce more product.
• Removing a reactant or adding a product pushes the equilibrium in the opposite direction.
• Changes in pressure or volume affect gaseous equilibria based on the number of gas molecules on each side of the equation.
• Temperature changes can shift equilibria toward either endothermic or exothermic directions depending on whether heat is added or removed.

These principles help manage industrial processes to optimize product yield.

Chemical Thermodynamics

Chemical thermodynamics deals with the energy changes that occur during chemical reactions, particularly the principles that govern these changes. For the Deacon process, understanding these energy changes is vital to manipulate the reaction conditions effectively.Thermodynamics introduces important concepts such as enthalpy (\( \Delta H \)). In the exercise example, \( \Delta H = -114 \text{ kJ} \) indicates the reaction is exothermic, meaning it releases heat. This can significantly impact the reaction if the temperature conditions are altered.

• Heat released or absorbed during a reaction can shift equilibrium based on thermodynamic laws.
• Exothermic reactions release energy into the surroundings, usually favoring products.
• Endothermic reactions absorb energy, often favoring reactants unless additional heat is supplied.

Using thermodynamic concepts, students can predict how energy will flow in and out of reactions, helping optimize processes like the Deacon process under various conditions.

Deacon Process

The Deacon process is an industrial chemical reaction used to convert hydrogen chloride \((\text{HCl})\) into chlorine gas \((\text{Cl}_2)\) using oxygen \((\text{O}_2)\). This reaction is particularly useful in recycling \(\text{HCl}\) waste, turning it into valuable chlorine gas.In the Deacon process, the reaction involves both complex kinetics and equilibrium principles:

• Equilibrium reactions demonstrate how the reaction shifts to produce \(\text{Cl}_2\) when \(\text{O}_2\) is added, as seen in the exercise.
• Thermodynamics, through \(\Delta H\), illustrates why a rise in temperature slows chlorine production due to the exothermic nature of the reaction.
• A catalyst is generally used in this process to speed up the reaction without affecting the equilibrium.

Understanding these aspects allows industry professionals to adjust conditions—such as temperature, pressure, and reactant concentrations—to optimize the yield and efficiency of chlorine gas production, making the Deacon process a valuable method in chemical manufacturing.