⏱️General Chemistry II Unit 2 – Chemical Equilibrium: Le Chatelier's Principle

Key Concepts

• Chemical equilibrium occurs when the rates of the forward and reverse reactions are equal and the concentrations of reactants and products remain constant over time
• Le Chatelier's principle states that when a system at equilibrium is disturbed by a change in temperature, pressure, or concentration, the system will shift its equilibrium position to counteract the disturbance and re-establish equilibrium
• Equilibrium constant ($K$) expresses the relationship between the concentrations of reactants and products at equilibrium and remains constant at a given temperature
• Reaction quotient ($Q$) represents the ratio of product concentrations to reactant concentrations at any given point in the reaction and can be used to determine the direction of the shift in equilibrium
• Factors affecting equilibrium include changes in concentration, pressure, volume, and temperature
• Catalysts do not affect the equilibrium position but can increase the rate at which equilibrium is reached by lowering the activation energy

Defining Le Chatelier's Principle

• Le Chatelier's principle, named after French chemist Henry Louis Le Chatelier, is a fundamental concept in chemistry that predicts the behavior of a system at equilibrium when subjected to a disturbance
• States that when a system at equilibrium is disturbed, it will shift its equilibrium position to counteract the disturbance and re-establish equilibrium
• Applies to changes in concentration, pressure, volume, and temperature
• Helps chemists understand and control chemical reactions by predicting the direction of the shift in equilibrium
• Principle is based on the idea that a system at equilibrium is in a state of dynamic balance, where the rates of the forward and reverse reactions are equal
• When a disturbance is introduced, the system responds by shifting the equilibrium position to minimize the effect of the disturbance
  • For example, if the concentration of a reactant is increased, the system will shift towards the products to consume the added reactant and re-establish equilibrium

Factors Affecting Equilibrium

• Changes in concentration affect equilibrium by shifting the position to consume the added species or replenish the removed species
  • Adding a reactant shifts equilibrium towards the products
  • Removing a product shifts equilibrium towards the products
• Changes in pressure or volume affect equilibrium in gaseous systems where the total number of moles of gas changes during the reaction
  • Increasing pressure (or decreasing volume) shifts equilibrium towards the side with fewer moles of gas
  • Decreasing pressure (or increasing volume) shifts equilibrium towards the side with more moles of gas
• Changes in temperature affect equilibrium by altering the equilibrium constant ($K$) and shifting the position to favor the endothermic or exothermic reaction
  • Increasing temperature shifts equilibrium towards the endothermic reaction
  • Decreasing temperature shifts equilibrium towards the exothermic reaction
• Addition of a catalyst does not affect the equilibrium position but increases the rate at which equilibrium is reached by lowering the activation energy of the reaction
• Inert gases do not participate in the reaction and do not affect the equilibrium position when added at constant volume

Applying Le Chatelier's Principle

• Identify the disturbance introduced to the system at equilibrium (change in concentration, pressure, volume, or temperature)
• Determine the direction of the shift in equilibrium based on the nature of the disturbance and the reaction
  • Changes in concentration: equilibrium shifts to consume the added species or replenish the removed species
  • Changes in pressure or volume: equilibrium shifts towards the side with fewer moles of gas when pressure is increased or volume is decreased, and towards the side with more moles of gas when pressure is decreased or volume is increased
  • Changes in temperature: equilibrium shifts towards the endothermic reaction when temperature is increased and towards the exothermic reaction when temperature is decreased
• Predict the effect of the disturbance on the concentrations of reactants and products and the value of the equilibrium constant ($K$)
  • Concentrations of species favored by the shift in equilibrium will increase, while concentrations of species on the opposite side will decrease
  • Equilibrium constant ($K$) remains unchanged for changes in concentration, pressure, or volume, but changes for temperature changes
• Consider the effect of multiple disturbances by applying Le Chatelier's principle sequentially
• Recognize that the system will re-establish equilibrium after the disturbance, and the rates of the forward and reverse reactions will once again be equal

Common Examples and Reactions

• Haber-Bosch process for ammonia synthesis: $\ce{N2(g) + 3H2(g) <=> 2NH3(g)}$
  • Increasing pressure shifts equilibrium towards the products (fewer moles of gas)
  • Increasing temperature shifts equilibrium towards the reactants (endothermic reaction)
• Synthesis of hydrogen iodide: $\ce{H2(g) + I2(g) <=> 2HI(g)}$
  • Increasing temperature shifts equilibrium towards the products (endothermic reaction)
  • Adding a catalyst (platinum) increases the rate of reaching equilibrium without affecting the equilibrium position
• Dissolution of carbon dioxide in water: $\ce{CO2(g) + H2O(l) <=> H2CO3(aq)}$
  • Increasing pressure shifts equilibrium towards the products (fewer moles of gas)
  • Decreasing temperature shifts equilibrium towards the products (exothermic reaction)
• Formation of nitrogen dioxide: $\ce{2NO(g) + O2(g) <=> 2NO2(g)}$
  • Removing $\ce{NO2}$ shifts equilibrium towards the products (replenishing the removed species)
  • Increasing volume shifts equilibrium towards the reactants (more moles of gas)

Calculations and Problem-Solving

• Calculate the equilibrium constant ($K$) using the equilibrium concentrations of reactants and products
  • $K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$ for the general reaction $\ce{aA + bB <=> cC + dD}$
• Determine the direction of the shift in equilibrium by comparing the reaction quotient ($Q$) to the equilibrium constant ($K$)
  • If $Q < K$, the reaction proceeds towards the products to reach equilibrium
  • If $Q > K$, the reaction proceeds towards the reactants to reach equilibrium
  • If $Q = K$, the system is at equilibrium, and no net change occurs
• Calculate the equilibrium concentrations of reactants and products using the equilibrium constant ($K$) and initial concentrations
  • Set up an ICE table (Initial, Change, Equilibrium) to organize the information
  • Express the change in concentrations using the stoichiometric coefficients and a variable (e.g., $x$)
  • Substitute the equilibrium concentrations into the equilibrium constant expression and solve for the variable
• Determine the effect of changes in concentration, pressure, volume, or temperature on the equilibrium concentrations and the value of the equilibrium constant ($K$)
  • Changes in concentration, pressure, or volume do not affect $K$, but alter the equilibrium concentrations
  • Changes in temperature affect both $K$ and the equilibrium concentrations

Real-World Applications

• Hemoglobin-oxygen binding in blood: $\ce{Hb(aq) + O2(g) <=> HbO2(aq)}$
  • Increasing altitude (lower pressure) shifts equilibrium towards the reactants, reducing oxygen binding
  • Increasing carbon dioxide concentration (Bohr effect) shifts equilibrium towards the reactants, promoting oxygen release in tissues
• Ocean acidification due to increased atmospheric carbon dioxide: $\ce{CO2(g) + H2O(l) <=> H2CO3(aq)}$
  • Rising atmospheric $\ce{CO2}$ levels shift equilibrium towards the products, forming more carbonic acid and lowering ocean pH
  • Lower pH affects marine life, particularly organisms with calcium carbonate shells or skeletons (e.g., corals, mollusks)
• Nitrogen fixation in soil by bacteria: $\ce{N2(g) + 8H+(aq) + 8e- <=> 2NH3(g) + H2(g)}$
  • Presence of nitrogenase enzyme (catalyst) in nitrogen-fixing bacteria increases the rate of ammonia formation
  • Ammonia produced is converted to other nitrogen compounds (e.g., nitrates) that plants can use for growth
• Exhaust gas recirculation (EGR) in internal combustion engines: $\ce{2NO(g) + O2(g) <=> 2NO2(g)}$
  • EGR reduces the concentration of oxygen in the combustion chamber, shifting equilibrium towards the reactants and reducing $\ce{NO2}$ formation
  • Lower $\ce{NO2}$ levels help reduce smog and air pollution

Review and Practice

• Summarize the key concepts of Le Chatelier's principle, including the factors affecting equilibrium and the direction of the shift in response to disturbances
• Practice applying Le Chatelier's principle to various reactions and disturbances, predicting the direction of the shift in equilibrium and the effect on concentrations and the equilibrium constant
• Work through sample problems involving calculations of equilibrium constants, reaction quotients, and equilibrium concentrations
  • Emphasize the use of ICE tables and equilibrium constant expressions
  • Practice solving for the equilibrium concentrations given the initial concentrations and the equilibrium constant
• Analyze real-world applications of Le Chatelier's principle, identifying the disturbances and predicting the effects on the equilibrium system
  • Discuss the implications of these effects on the environment, biological systems, or industrial processes
• Review the common examples and reactions discussed in class, ensuring a clear understanding of how Le Chatelier's principle applies to each case
• Collaborate with classmates to create and solve additional practice problems, focusing on areas of difficulty or confusion
• Consult with the instructor or teaching assistants for further clarification or guidance on challenging concepts or problems