[H2 Chemistry] 2021 Topic 7 Chemical Equilibria 2

The lecturer introduces the concept of the reaction quotient (Qc) in chemical equilibrium. The reaction quotient is used to calculate the concentration of reactants and products before equilibrium is reached. If Qc equals the equilibrium constant (Kc), the system is at equilibrium. If Qc is less than Kc, the reaction will proceed in the forward direction, increasing the product concentration. Conversely, if Qc is greater than Kc, the reaction will shift to the left, favoring the formation of reactants.

The lecturer explains how to apply the reaction quotient (Qc) using the Haber process as an example. The process involves calculating Qc for a given set of initial concentrations of nitrogen, hydrogen, and ammonia. By comparing Qc to the equilibrium constant (Kc), one can predict whether the reaction will shift to the right or left to achieve equilibrium. A detailed calculation shows that Qc is less than Kc, indicating the reaction will proceed to the right to produce more ammonia.

The lecturer discusses the limitations of Le Chatelier's Principle and introduces the use of Qc in specific scenarios. For example, when water is added to an aqueous solution at equilibrium, the concentration of water remains constant. Le Chatelier's Principle may not accurately predict changes, but using Qc helps to solve these issues by showing that the equilibrium will shift based on the reaction quotient's value.

The impact of adding inert gases to a system at equilibrium is explored. Two scenarios are considered: adding inert gas at constant pressure and constant volume. At constant pressure, the total number of moles increases, affecting Qp (reaction quotient for partial pressures) and shifting the equilibrium. At constant volume, the addition of inert gas does not affect the equilibrium position. Detailed calculations illustrate these principles.

Several classic equilibrium reactions, including the dimerization of NO2 and the Haber process, are used to demonstrate how changes in conditions affect the equilibrium position. For example, adding inert gas at constant pressure increases the total number of moles, affecting the reaction quotient (Qp) and shifting the equilibrium. Conversely, adding inert gas at constant volume has no effect on the equilibrium position.

The lecturer demonstrates how to use ICE (Initial, Change, Equilibrium) tables to calculate equilibrium concentrations for various chemical reactions. Detailed examples, such as the dissociation of ethanoic acid in a liquid phase, show how to systematically approach these problems. By calculating the changes in concentrations and using equilibrium constants, one can determine the equilibrium concentrations of reactants and products.

A step-by-step approach to calculating equilibrium constants for different chemical reactions is presented. Examples include the reaction of ethanoic acid with ethanol, the equilibrium between H2, I2, and HI, and the dissociation of carboxylic acids. Detailed calculations show how to set up ICE tables, calculate changes in concentration, and determine equilibrium constants (Kc) for these reactions.

The lecturer explains how to determine equilibrium constants (Kc) from given initial conditions and equilibrium concentrations. Examples include reactions involving carboxylic acids, ethanoic acid, and esters. By using ICE tables and incorporating initial amounts and changes, the equilibrium constants for these reactions can be calculated accurately.

Quadratic equations often arise in equilibrium calculations, especially when determining equilibrium concentrations from initial amounts and equilibrium constants. The lecturer provides detailed examples, such as the equilibrium between H2, I2, and HI, to illustrate how to solve these quadratic equations and accurately calculate equilibrium concentrations.

The relationship between temperature and equilibrium constants (Kc and Kp) is explored. The Van't Hoff equation, which relates changes in temperature to changes in equilibrium constants, is introduced. Detailed examples show how to calculate equilibrium constants at different temperatures, considering the effects of enthalpy and entropy changes on the equilibrium position.

The lecturer presents more practical examples of equilibrium calculations, focusing on determining equilibrium constants (Kc and Kp) for various reactions. Examples include the dissociation of PCl5 and the reaction between H2 and I2. By using ICE tables and solving quadratic equations, the equilibrium constants for these reactions are calculated accurately.

Advanced equilibrium calculations are discussed, including scenarios where initial amounts and equilibrium constants are given. Examples involve reactions like the dissociation of PCl5 and the equilibrium between H2 and I2. Detailed step-by-step calculations demonstrate how to use ICE tables and solve quadratic equations to determine equilibrium concentrations and constants.

The lecturer explains how changes in initial conditions and equilibrium constants affect the equilibrium position. Using examples such as the reaction between H2 and I2, the impact of adding reactants or products on the equilibrium position is explored. Detailed calculations show how to predict the direction of the equilibrium shift and calculate new equilibrium concentrations.

The concept of dynamic equilibrium and the use of reaction quotients (Qc) to predict equilibrium shifts are revisited. Examples include the decomposition of NO2 and the reaction between H2 and I2. By comparing Qc to the equilibrium constant (Kc), the direction of the equilibrium shift can be predicted accurately.

Complex equilibrium calculations are discussed, including reactions involving multiple steps and components. Examples include the decomposition of NO2 and the equilibrium between CO, H2, and methanol. Detailed step-by-step calculations show how to determine equilibrium constants and predict equilibrium positions for these complex reactions.

The practical application of equilibrium concepts in industrial processes, such as the Haber process for ammonia synthesis, is discussed. The interplay between thermodynamics and kinetics in optimizing reaction conditions is explored. Detailed examples show how to balance temperature, pressure, and catalyst use to maximize product yield while maintaining economic feasibility.

Methods for experimentally determining equilibrium constants (Kc and Kp) are presented. Examples include titration of equilibrium mixtures and using colorimetry for complex ion equilibria. The importance of maintaining constant temperature and accurately measuring concentrations is emphasized. Practical tips for designing experiments to determine equilibrium constants are provided.

The relationship between equilibrium constants (K) and Gibbs free energy (ΔG) is explored. The lecturer explains how ΔG determines the direction of the reaction and the position of equilibrium. The mathematical relationship between ΔG and K, including the use of the Van't Hoff equation, is discussed. Examples show how changes in temperature affect ΔG and K, influencing the equilibrium position.

The lecturer provides a summary of key concepts in chemical equilibrium, including the use of reaction quotients (Qc), equilibrium constants (Kc and Kp), and the relationship between ΔG and K. Advanced topics, such as the Van't Hoff equation and the impact of temperature on equilibrium constants, are revisited. The importance of understanding these concepts for future studies in chemistry, including acid-base and solubility equilibria, is emphasized.

Practical tips for performing accurate equilibrium calculations are provided. The lecturer emphasizes the importance of using ICE tables, carefully tracking changes in concentrations, and accurately solving quadratic equations. Examples from previous sections are revisited to reinforce key points and demonstrate best practices for equilibrium calculations.