[H2 Chemistry] 2021 Kinetics 1

This paragraph introduces the topic of reaction kinetics, following a discussion on thermodynamics. It emphasizes the importance of enthalpy and entropy changes in determining whether a reaction is feasible or spontaneous, as described by the equation delta G = delta H - T delta S. The instructor clarifies that while exothermic reactions are often assumed to be spontaneous, entropy contributions are equally significant. The paragraph transitions into the concept of reaction kinetics, highlighting that a thermodynamically feasible reaction might be slow, using the example of diamond transforming into graphite. The activation energy, or kinetic barrier, is introduced as a key factor in reaction rates, and the instructor mentions that reaction kinetics will be explored in more detail, including rate equations and reaction orders.

The instructor underscores the significance of studying reaction kinetics for understanding the mechanisms by which reactions proceed. They explain that while chemical equations provide information about reactants and products, they do not reveal the intermediate steps or the 'what happens in between.' The paragraph delves into the curiosity of chemists to understand these intermediate steps, which can inform the design of reactions to achieve specific outcomes. The concept of reaction mechanisms is introduced as a fundamental aspect of study in organic chemistry, with five basic mechanisms to be aware of. The instructor encourages students to build a strong foundation in kinetics to ease their studies in later years, acknowledging that initial understanding may be shallow but will deepen over time.

This paragraph focuses on defining the rate of a chemical reaction, which is the change in concentration of a reactant or product per unit time. The instructor discusses the mathematical representation of reaction rates, emphasizing that the rate of disappearance of reactants is indicated by a negative slope due to decreasing concentration over time. However, calculated rates are positive numbers. The paragraph also addresses common misconceptions regarding rate units, explaining that they should reflect the change in concentration (e.g., moles per liter per second or per minute). An example of a simple reaction between H2 and I2 to form HI is used to illustrate how the rates of disappearance of reactants and the rate of product formation are related, highlighting the need to adjust for stoichiometric ratios.

The instructor introduces different types of reaction rates, particularly instantaneous and average rates. Instantaneous rate is likened to taking the derivative at a specific time on a concentration-time graph, reflecting the rate at that exact moment. Initial rate is the instantaneous rate at time equals zero, which is crucial for understanding how fast a reaction begins. Average rate is defined over a time interval, represented by a straight line on a graph, indicating the overall change in concentration over that period. The paragraph clarifies the mathematical concepts behind these rates, using graphs to illustrate the differences and emphasizing the importance of understanding the mathematical principles for studying chemistry.

This paragraph discusses a practical approach to determining reaction rates using data from a graph of a reaction between CH3Br and OH-. The reaction is identified as a nucleophilic substitution, where OH- acts as a nucleophile. The instructor guides students on how to analyze the graph to find the initial rates by drawing a tangent at time equals zero and explains the significance of the units for the rates. The paragraph also touches on the process of determining the instantaneous rate at a specific time and the average rate over a given period, highlighting the importance of careful data interpretation and graph analysis in understanding reaction kinetics.

The instructor introduces the technical aspects of rate equations, emphasizing that the order of reaction is not the same as the stoichiometric coefficient. The rate equation is presented as a product of the rate constant (k) and the concentrations of reactants raised to the power of their respective orders. The paragraph explains that the order of reaction must be determined experimentally and is not necessarily related to the stoichiometric coefficients. The rate constant is described as a proportionality constant that is sensitive to temperature and is also determined experimentally. The instructor warns against common misconceptions and stresses the importance of understanding the fundamental definitions and concepts before progressing to more complex topics.

This paragraph delves into the experimental determination of reaction orders, explaining that experiments are designed to test how changes in concentration affect reaction rates. The instructor introduces the concepts of zero-order, first-order, and second-order reactions, describing their characteristics and how they relate to the concentration of reactants. Zero-order reactions show rate independence from concentration, first-order reactions have a direct proportionality to the concentration, and second-order reactions show a rate proportional to the square of the concentration. The paragraph emphasizes the importance of understanding these orders through detailed study and practical exposure to different scenarios.

The instructor provides a clear example of writing a rate equation for a reaction between hydrogen peroxide and acidified iodine ions, given the orders with respect to H2O2 and H+ are first and zero, respectively. The rate equation is simplified, and the overall order of the reaction is determined. The paragraph also discusses the rate constant, explaining its dependence on temperature and presence of a catalyst, and clarifies misconceptions about catalysts and activation barriers. The rate constant's units are explored, demonstrating how they relate to the overall order of the reaction.

This paragraph outlines the process of determining rate equations using different types of graphs: concentration-time, rate against concentration, and rate against time. The instructor emphasizes the importance of understanding the shapes of these graphs and their mathematical implications for determining the order of reactions. The paragraph provides examples of how to interpret concentration-time graphs for zero-order, first-order, and second-order reactions, highlighting the significance of half-life and its relationship to reaction order. The instructor advises against memorization and instead encourages a mathematical understanding of the concepts.

The instructor continues the discussion on graph analysis for understanding reaction kinetics, focusing on rate against concentration graphs. For zero-order reactions, a horizontal line is expected, while first-order reactions yield a straight line passing through the origin. The paragraph introduces the technique of linearizing graphs by plotting log rates against log concentration to determine the order of reaction and the rate constant. The method is explained as a valuable tool when the reaction order is unknown. The instructor also briefly mentions rate against time graphs, noting their less intuitive nature and suggesting that understanding the mathematics behind the graphs is more important than memorizing their shapes.