Question

For a reversible reaction, the concentration of the reactants are doubled, then the equilibrium constant: (a) Becomes one-fourth (b) Is doubled (c) Is halved (d) Remains same

Short answer

(d) Remains same.

Step-by-step solution

Understand the Equilibrium Constant

The equilibrium constant, denoted as \( K \), of a reversible reaction is defined at a specific temperature and does not depend on the concentrations of reactants or products. It is generally expressed for a reaction \( aA + bB \rightleftharpoons cC + dD \) as:\[K = \frac{[C]^c[D]^d}{[A]^a[B]^b}\]where \([C]\), \([D]\), \([A]\), and \([B]\) are the equilibrium concentrations of the products and reactants.

Analyze the Change in Concentration

Since we are doubling the concentration of the reactants, let's consider the initial concentrations of \( [A]_0 \) and \( [B]_0 \). After doubling, these concentrations become \(2[A]_0\) and \(2[B]_0\).

Consider the Effect on Equilibrium

Even though the initial concentrations of reactants are changed, a reversible reaction will reach a new equilibrium state, and the concentrations will adjust accordingly. Importantly, the equilibrium constant \( K \) is dependent only on temperature. Therefore, no matter the initial concentrations, the value of \( K \) remains the same at a constant temperature.

Conclude on Equilibrium Constant Change

Given that the equilibrium constant \( K \) depends only on temperature and not on the concentrations of reactants or products, the act of doubling the concentrations of reactants does not affect \( K \).

Key concepts

Reversible Reactions

Reversible reactions are a fascinating concept in chemistry. They involve reactions where the products can revert to the original reactants under the right conditions. This process of going both ways—forward and backward—is what makes these reactions reversible. An example of this is the synthesis of ammonia in the Haber process, where nitrogen - and hydrogen react to form ammonia, but ammonia can break down back to nitrogen and hydrogen.

In reversible reactions, both reactants and products exist in a state of dynamic balance, which is an essential concept for understanding chemical equilibrium. You'll often see them represented with a double-headed arrow (\(\rightleftharpoons\)) in chemical equations. The conditions under which the reaction occurs, such as temperature and pressure, can shift the direction of the reaction, but it eventually stabilizes in an equilibrium state, as long as those conditions remain constant.

Concentration Changes

Concentration changes play a crucial role in chemical reactions, particularly in reversible reactions. When the concentration of reactants or products in a reaction is altered, it can cause the reaction to shift in order to regain equilibrium.

For example, if you increase the concentration of reactants in a reversible reaction, the system will try to counteract this change by producing more products, shifting the equilibrium towards the products. This is known as Le Chatelier's principle.

• Increasing reactant concentration typically shifts the equilibrium to the right, creating more product.
• Conversely, decreasing reactant concentration shifts the equilibrium to the left, favoring the reactants.
• Similarly, increasing the concentration of products can shift the equilibrium back towards the reactants.

Despite these changes in concentrations, the equilibrium constant remains unaffected as long as the temperature is unchanged.

Chemical Equilibrium

Chemical equilibrium is a state where the rates of the forward and reverse reactions become equal, leading to constant reactant and product concentrations. Despite seemingly being at rest, equilibrium is a highly dynamic state at the molecular level. Molecules of reactants are continuously converting into products and vice versa.

The concept of equilibrium also outlines that the macroscopic properties, such as concentration, appear stable even though molecular changes are constantly happening. An equilibrium constant, denoted as \( K \), describes the ratio of product concentrations to reactant concentrations at equilibrium and is specific for a given reaction at a specific temperature.

• The magnitude of \( K \) can tell us how far the reaction proceeds towards products.
• A large value of \( K \) indicates a reaction that favors product formation, while a small value suggests more reactants are present at equilibrium.
• It is crucial to remember that \( K \) is affected only by temperature—not by changes in concentration or pressure directly.

Effect of Concentration on Equilibrium

The effect of concentration changes on equilibrium may initially seem complex, but it's all about balance. According to Le Chatelier's principle, if a dynamic equilibrium is disturbed by changing the concentration of one of the components, the system will respond to restore a new equilibrium.

For example, when the concentration of a reactant is increased, the reaction will produce more products to offset the change, thus maintaining the balance. This mechanism allows the equilibrium position of a reaction to shift to either the right (favoring products) or to the left (favoring reactants).

• Adding a reactant pushes the equilibrium toward product formation.
• Removing a reactant will shift equilibrium back toward reactants.
• Likewise, adding a product shifts equilibrium toward reactants, while removing it favors the formation of more products.

Changes in concentration can effectively alter the quantities of products and reactants but remember—to keep the equilibrium constant unchanged, temperature must remain constant.