Chemical Equilibrium and Le Chatelier's Principle

A state where the forward and backward reaction rates are equal, resulting in constant reactant and product concentrations.

A chemical reaction where products can react to reform the original reactants.

When a stress is applied to a system in equilibrium, the system shifts to relieve the stress.

A change in reaction conditions such as concentration, pressure, or temperature.

Increased pressure favors the reaction producing fewer gaseous molecules.

Substance that speeds up the rate of a reaction, but without changing the equilibrium position.

A reaction that releases heat.

A reaction that absorbs heat.

When conditions change, the equilibrium of a reversible reaction shifts to favor either the forward or reverse reaction to re-establish balance.

Adding more of a reactant in a reversible reaction will cause the equilibrium to shift towards the product side to consume the excess reactant.

Increasing pressure in a reversible reaction favors the side with fewer gas molecules, as the system tries to relieve the pressure.

Raising temperature in an exothermic reaction shifts the equilibrium towards the reactants, as the system tries to absorb the excess heat.

Le Chatelier's Principle can be demonstrated by observing color changes in a solution of Fe^3+ and CNS^- ions.

Removing a product from a reversible reaction will cause the equilibrium to shift towards the product side to replenish the removed product.

The industrial production of ammonia from nitrogen and hydrogen gases utilizes Le Chatelier's principle to maximize ammonia yield.

When a change is imposed on a system in equilibrium, the system shifts to relieve the stress and re-establish equilibrium.

A value representing the ratio of product concentrations to reactant concentrations at equilibrium. It indicates the extent to which a reaction proceeds to completion.

A mathematical formula that shows the relationship between the concentrations of reactants and products at equilibrium. Kc equals the product of product concentrations raised to their stoichiometric coefficients, divided by the product of reactant concentrations raised to their stoichiometric coefficients.

The process of determining the equilibrium constant using the concentrations of reactants and products at equilibrium.

An increase in temperature favors the endothermic reaction by shifting the equilibrium to absorb heat. A decrease in temperature favors the exothermic reaction.

A device in car exhaust systems that uses catalysts to speed up the conversion of harmful pollutants into less harmful gases.

A type of catalysis where the catalyst is in a different phase than the reactants. For example, a solid catalyst in a gas-phase reaction.

A process used to manufacture ammonia by reacting nitrogen with hydrogen under high pressure and moderate temperature, typically around 200 atmospheres and 200°C.

The catalyst, typically iron (Fe), speeds up the reaction rate but doesn't affect the equilibrium position. It helps reach equilibrium faster by lowering the activation energy.

Le Chatelier's Principle states that the system will shift to relieve the stress applied. In the Haber process, high pressure favors ammonia production, while low temperature also favors production, but at a slower rate.

This process manufactures sulfuric acid. Sulfur dioxide is reacted with oxygen in the presence of a vanadium pentoxide (V₂O₅) catalyst to produce sulfur trioxide.

The Contact Process works best with a pressure slightly higher than atmospheric pressure and a temperature of around 450°C. These conditions allow for maximum conversion of sulfur dioxide to sulfur trioxide.

The formula for calculating Kc involves the concentration of products raised to the power of their stoichiometric coefficients, divided by the concentration of reactants raised to the power of their coefficients.

To calculate Kc, you need to know the equilibrium concentrations of all reactants and products. Substitute these values into the Kc formula to find the numerical value of Kc.

Adding KCNS to the Fe^3+^ (aq) + CNS^--^ (aq) Fe(CNS)^2+^ (aq) equilibrium mixture increases the concentration of CNS^-^, causing the equilibrium to shift forward to consume the added CNS^-^, thus decreasing the Fe^3+^ concentration.

Changing the pressure has no effect on the Fe^3+^ (aq) + CNS^--^ (aq) Fe(CNS)^2+^ (aq) equilibrium because it occurs in solution, not involving gases.

Cooling the Fe^3+^ (aq) + CNS^--^ (aq) Fe(CNS)^2+^ (aq) equilibrium mixture causes the red color to fade because the reaction is endothermic, and cooling shifts the equilibrium backwards, favoring reactants, thus reducing the red Fe(CNS)^2+^ concentration.