Problem 120 - Chemistry Matter and Change

Question

Hydrogen and fluorine react to form HF according to the following equilibrium equation. $$\mathrm{H}_{2}(\mathrm{g})+\mathrm{F}_{2}(\mathrm{g}) \rightleftharpoons 2 \mathrm{HF} \Delta H=-538 \mathrm{kJ}(\mathrm{g})$$ Will raising the temperature cause the amount of product to increase? Explain. (Chapter 17)

Step-by-Step Solution

Answer

Raising the temperature will decrease \(\mathrm{HF}\) as equilibrium shifts toward reactants.

Step 1: Understand the Reaction

The reaction given is \( \mathrm{H}_{2}(\mathrm{g}) + \mathrm{F}_{2}(\mathrm{g}) \rightleftharpoons 2 \mathrm{HF} \). This is an exothermic reaction, as indicated by the negative \( \Delta H = -538 \text{ kJ/mol} \). This means heat is released during this reaction.

Step 2: Apply Le Chatelier's Principle

According to Le Chatelier's Principle, if an external change is applied to a system at equilibrium, the system adjusts in such a way as to counteract that change. For an exothermic reaction, increasing the temperature is akin to adding more 'product' in the form of heat.

Step 3: Predict System Response to Temperature Increase

As temperature is increased, the equilibrium will shift to oppose the change, meaning it will shift towards the endothermic direction (the reverse reaction). Thus, instead of producing more of the product \( \mathrm{HF} \), the system will adjust by favoring the formation of the reactants \( \mathrm{H}_{2} \) and \( \mathrm{F}_{2} \).

Key Concepts

Exothermic Reaction

An exothermic reaction is a chemical reaction that releases energy by light or heat. In the context of the given reaction, \( \mathrm{H}_{2}(\mathrm{g}) + \mathrm{F}_{2}(\mathrm{g}) \rightleftharpoons 2 \mathrm{HF} \), the negative \( \Delta H = -538 \text{ kJ/mol} \) indicates that it is exothermic, meaning heat is released when hydrogen and fluorine form hydrogen fluoride (HF).

Exothermic reactions are often spontaneous because they release energy, allowing the products to reach a lower energy state. This release of heat may be felt physically, such as a flask warming up during a laboratory experiment. Exothermic reactions are common and include various processes such as combustion, respiration, and certain oxidation reactions.

• Release energy as the reaction progresses.
• Have products that are more stable due to lower energy levels than reactants.
• Are often spontaneous as they give off heat.

A clear understanding of exothermic reactions helps to predict how changing conditions can affect a system at equilibrium.

Equilibrium

Chemical equilibrium occurs when the rates of the forward and reverse reactions are equal, resulting in no net change in the concentration of reactants and products. In the reaction \( \mathrm{H}_{2}(\mathrm{g}) + \mathrm{F}_{2}(\mathrm{g}) \rightleftharpoons 2 \mathrm{HF} \), equilibrium is achieved when the production of HF from hydrogen and fluorine occurs at the same rate as its decomposition back into the reactants.

Equilibrium in a chemical reaction is dynamic, meaning that even though the concentrations remain constant, the forward and reverse reactions continue to occur. It is not a static state but a balance between opposing processes.

• The concentrations of reactants and products remain constant, though reactions continue to happen.
• The position of equilibrium can be altered by changing concentration, pressure, or temperature.
• The system strives to minimize the effect of any imposed changes.

Understanding equilibrium is crucial for predicting how a system will react to external changes, such as temperature alterations, which is where Le Chatelier's Principle becomes applicable.

Temperature Effect on Equilibrium

Temperature changes can have significant effects on the equilibrium of a chemical reaction, particularly for an exothermic reaction like \( \mathrm{H}_{2}(\mathrm{g}) + \mathrm{F}_{2}(\mathrm{g}) \rightleftharpoons 2 \mathrm{HF} \). According to Le Chatelier's Principle, if the temperature of an exothermic reaction is increased, the equilibrium will shift in the direction that absorbs this added heat, essentially the endothermic or reverse direction.

For the hydrogen-fluorine reaction, raising the temperature means the system will favor the decomposition of HF back into \( \mathrm{H}_{2} \) and \( \mathrm{F}_{2} \), hence decreasing the amount of HF produced. Conversely, decreasing the temperature would shift the equilibrium toward the exothermic direction, thereby favoring the synthesis of more HF.

• Increasing temperature shifts equilibrium toward the endothermic direction.
• Decreasing temperature shifts equilibrium toward the exothermic direction.
• The principle helps to predict and control the outcomes of reactions.

By manipulating temperature, chemists can control the yields of chemical processes involving equilibrium reactions.