Physical Chemistry - Chemical Kinetics

Explore the fundamentals of chemical kinetics, defining spontaneity, extent, and rate, and examine factors like temperature, pressure, and concentration that control reaction rates.

Define the rate of reaction as the change in concentration per unit time, with fast, slow, and moderate examples and units such as moles per liter per second.

Explore the types of rate of reaction, including average rate and instantaneous rate, by linking concentration changes over time to slope on a graph and tangent methods in chemical kinetics.

Explore the rate law and rate constant, showing how the reaction rate depends on reactant concentrations, with the rate equation and the experimentally determined rate constant.

Learn to write the rate equation by applying sign conventions, with reactants negative and products positive, for average and instantaneous rates, illustrated by ammonia synthesis.

Explore how the rate constant units vary with reaction order—from zero, first, to second order—by linking concentration and time, with A, B, X, and Y examples.

Apply the law of mass action: the rate is proportional to the product of reactant concentrations raised to their stoichiometric coefficients, and it requires a balanced equation for single-step reactions.

Determine the reaction order from the rate law by identifying the exponents of reactant concentrations; sum them to get the overall order. Note that order is an experimental property.

Examine zero order reactions, where the rate remains constant regardless of concentration, with examples like ammonia decomposition on catalysts and photochemical reactions; derive the integrated rate expression.

Explore first order reactions by analyzing rate laws, concentration changes, and integration to derive ln C and log C relationships, with graphing interpretations.

This lecture presents four graphical representations of first-order reactions, deriving linear forms from ln[C], log10[C], and related plots, and interpreting slopes and intercepts.

Explore pseudo first order reactions, where a second-order process behaves as first order when one reactant is in excess, making the rate depend only on the limiting reactant.

Explore second order reactions by examining rate laws for a + b → products, derive the integrated rate law, and learn how time relates inversely to concentration.

Explore third-order reactions, derive the integrated rate form from an A to products example, link reactant decrease to product formation, and discuss limits and plotting.

Study n-order reactions and master the differential and integrated rate laws, using a simple formula to relate reactant concentration changes to time for first, second, and third order.

Explore the molecularity of reaction, defined as the number of molecules colliding at the same time; learn uni-, bi-, and tri-molecular cases and that molecularity derives from elementary steps, not complex reactions.

Explore activation energy and threshold energy, and how collision orientation and energy barriers govern the formation of products.

Learn collision theory and how only effective collisions with proper orientation and kinetic energy at or above activation energy yield product formation. Non-effective collisions fail to form products.

Describe transition state theory, including the activated complex, its equilibrium with reactants, and the activation energy barrier to form it, plus energy transfer from vibrational to translational modes.

Explore the five factors affecting reaction rate: concentration and rate law, nature of reactants, surface area, catalysts, and temperature with Arrhenius concepts.

Explore how a catalyst lowers activation energy by providing a new reaction pathway, increasing product formation and the rate of reaction without being consumed.

Temperature raises the average kinetic energy of reactants, increasing collision frequency and the fraction of molecules with energy at least activation energy, thereby speeding up the rate of reaction.

Present the Arrhenius equation, identify activation energy, the pre-exponential factor A, and the gas constant R, and show converting natural logs to base-10 logs for solving.

Derive and apply the Arrhenius equation to calculate activation energy from two rate constants at different temperatures, linking temperature rise to reaction rate.

Use the Arrhenius equation to calculate activation energy from two temperatures. Explore how rate constants relate to activation energy and temperature with the Arrhenius constant and universal gas constant.

Explain the Arrhenius equation and its simplified forms, then show how plotting natural or base-10 logarithms yields straight lines, revealing the slope and axis roles in the graphical representation.

This lecture explains the rate-determining step in complex reactions and how the slow step controls the overall rate, with example-driven derivations of the rate law.

Use the integration (substitution) method to determine reaction order from time intervals and remaining concentration, then apply first-, second-, and third-order formulas to find the rate constant.

Learn how to determine reaction order with the half-life method, comparing how half-life depends on initial concentration for zero, first, and second order reactions.

learn to determine reaction order using the initial rate method by varying concentrations of A and B, holding one in excess, solving for x and y to reveal order.

Use the van't Hoff differential method to determine reaction order by comparing rates at different concentrations and deriving n from logarithmic relationships of the rate law.

Discover graphical methods to determine reaction order with half-life plots across zero to third order, and use axes of concentration, one over concentration, and log plots.

Explore parallel reactions where one reactant forms two products, B and C, with rate constants k1 and k2, giving rate k1+k2 and yields B = k1/(k1+k2), C = k2/(k1+k2).

explores negative order reactions in chemical kinetics, where the rate is inversely proportional to reactant concentration, with examples including precipitation, aggregation, diffusion, and polymerization.

Explore photochemical reactions driven by light energy that excite molecules past energy barriers, with rates depending on light intensity and blue light, not temperature, including vision and photosynthesis.

Explain radioactivity as the decay of unstable nuclei into new elements via alpha, beta, or gamma radiation, linking decay rate to chemical kinetics and becquerel units.

Explore the rate law of radioactivity, deriving the integrated decay law using the decay constant λ, initial and remaining amounts, and natural logarithms.

Apply the decay law to show that the concentration halves at the half-life, given by 0.693 divided by lambda, and examine its graphical representation.

Explain graphical representation of radioactive decay by plotting activity and undecayed atoms over time. Show log(activity) versus time forms a straight line with negative slope, linking decay constant and half-life.

Explore nuclear fission through a uranium-235 example, demonstrating how neutron bombardment splits a radioactive nucleus into lighter products, releasing large energy used in the power plant.