Question

Consider the reaction \(\mathrm{A}(g)+\mathrm{B}(g) \rightleftharpoons \mathrm{C}(g)+\mathrm{D}(g) .\) A friend asks the following: "I know we have been told that if a mixture of \(\mathrm{A}, \mathrm{B}, \mathrm{C},\) and \(\mathrm{D}\) is at equilibrium and more of \(\mathrm{A}\) is added, more \(\mathrm{C}\) and \(\mathrm{D}\) will form. But how can more \(\mathrm{C}\) and \(\mathrm{D}\) form if we do not add more \(\mathrm{B} ?^{\prime \prime}\) What do you tell your friend?

Short answer

Adding more reactant A to the system at equilibrium according to Le Châtelier's principle will shift the reaction in the direction that favors the formation of products C and D. As a result, the reaction will consume more A and B to restore equilibrium, leading to higher concentrations of C and D, even without adding more reactant B.

Step-by-step solution

Explain Chemical Equilibrium

Chemical equilibrium is a state in a chemical reaction where the rate of the forward reaction (formation of products) is equal to the rate of the reverse reaction (breakdown of products). In other words, the concentrations of reactants and products no longer change, even though the reaction is still happening at the molecular level, just at equal rates in both directions.

Describe Le Châtelier's Principle

Le Châtelier's Principle states that if a system at equilibrium is subjected to a change in conditions, such as concentration, temperature, or pressure, the system will adjust itself to counteract the change and restore a new equilibrium. Essentially, the reaction will shift in the direction that opposes the change.

Adding More Reactant A

Now, let's consider the case where we add more of reactant A to the system at equilibrium. According to Le Châtelier's principle, the reaction will shift in a direction to counteract the increase in A's concentration. In this case, the forward reaction (A + B ⇌ C + D) will increase its rate in order to consume the added A and restore equilibrium.

No Need to Add More Reactant B

Because only A has been added, the concentration of B initially remains the same. However, as the forward reaction increases its rate and consumes more A and B to form C and D, the concentration of B will start to decrease as well. As a result, the equilibrium will shift to accommodate the decrease in B by further favoring product formation and consume more A to restore the balance. This process will continue until a new equilibrium is reached, with higher concentrations of C and D than before. So, even though we did not add more B to the system, the reaction's shift to restore equilibrium will lead to the formation of more products C and D, as predicted by Le Châtelier's principle.

Key concepts

Le Châtelier's Principle

Imagine you have a balance scale perfectly leveled. When you add more weight to one side, the scale tips. Le Châtelier's Principle is quite similar. It describes how a chemical reaction at equilibrium responds to changes such as concentration, temperature, or pressure. Consider a reaction at equilibrium, like our example \[\mathrm{A}(g)+\mathrm{B}(g) \rightleftharpoons\mathrm{C}(g)+\mathrm{D}(g) \]. When this equilibrium is disturbed, say by adding more \( \mathrm{A} \), the system behaves like the scale, trying to rebalance itself.

• The reaction will "shift" towards the side that counteracts the change. In our case, this means producing more \( \mathrm{C} \) and \( \mathrm{D} \) to consume the excess \( \mathrm{A} \).
• This shift continues until a new state of balance (equilibrium) is established.

Understanding Le Châtelier's Principle helps predict how a chemical reaction will adapt to different changes in conditions.

Reaction Rates

In any chemical reaction, the rate at which reactants turn into products can be dynamic. Reactants \( \mathrm{A} \) and \( \mathrm{B} \) need energy to collide and react, turning into products \( \mathrm{C} \) and \( \mathrm{D} \). When the system is at equilibrium, these rates are equal, meaning the forward reaction forms products as quickly as the reverse reaction breaks them back down.
When more \( \mathrm{A} \) is added, the forward reaction rate increases because there are more \( \mathrm{A} \) molecules available to collide with \( \mathrm{B} \).

• This means more products will form, shifting equilibrium towards \( \mathrm{C} \) and \( \mathrm{D} \).
• Even though \( \mathrm{B} \) isn’t added directly, it will get consumed faster due to the increased forward reaction rate.

Eventually, when the extra \( \mathrm{A} \) is sufficiently reacted, a new equilibrium is reached. Reaction rates are crucial in understanding how quickly and in what direction a system approaches equilibrium.

Chemical Reactions

Chemical reactions involve the transformation of reactants into products. Consider different molecules, such as \( \mathrm{A}(g) \) and \( \mathrm{B}(g) \), that come together to create \( \mathrm{C}(g) \) and \( \mathrm{D}(g) \) through a chemical reaction. This process involves breaking and forming bonds, where energy changes play a pivotal role. The initial reaction rate can be altered by changing conditions or concentrations.

• In our example, adding more \( \mathrm{A} \) doesn't just mean more reactants; it means those extra molecules are likely to find \( \mathrm{B} \) faster to form \( \mathrm{C} \) and \( \mathrm{D} \).
• This interaction between molecules highlights how changes in a reaction's conditions can influence the progress of the reaction.

Chemical reactions give us insight into how substances change and why some products might be favored over others under different conditions. Understanding these interactions helps predict the outcomes in various scenarios, such as the one described.