Reaction Rates and Rate Laws

This page delves deeper into the quantitative aspects of reaction rates and introduces rate laws. It explains how reaction rates are measured and how they change over time.

The relationship between concentration and rate is explored, introducing the concept of initial rates and instantaneous rates. The page emphasizes that reaction rates typically decrease over time as reactant concentrations diminish.

Vocabulary: Instantaneous rate is the slope of a line tangent to the concentration-time curve at any given point.

Rate laws are introduced as mathematical expressions relating reaction rates to reactant concentrations. The general form of a rate law is presented:

rate = k[A]^a[B]^b

Where:

• k is the rate constant
• [A] and [B] are reactant concentrations
• a and b are reaction orders

Highlight: The overall order of a reaction is the sum of the individual orders with respect to each reactant.

The page stresses that stoichiometric ratios in balanced equations do not necessarily correspond to the exponents in rate laws. Methods for determining rate laws experimentally are outlined.

Example: To determine a rate law, one can compare reaction rates under conditions where one reactant concentration is held constant while the other is varied.

Kinetics and Equilibrium: Fundamentals of Reaction Rates

This page introduces the core concepts of chemical kinetics and equilibrium. It explores the factors that influence reaction rates and the principles of collision theory. The enduring understandings highlight how concentration, pressure, and environmental factors affect reaction rates.

Factors affecting reaction rates are explained in detail:

1. Physical state: More homogeneous mixtures lead to faster reactions due to increased collisions.
2. Concentration: Higher concentrations increase the likelihood of molecular collisions.
3. Temperature: Elevated temperatures increase kinetic energy, resulting in more frequent and energetic collisions.
4. Surface area: Greater surface area enhances the probability of collisions.

The concept of activation energy is introduced, emphasizing its role in activation energy barriers restriction to chemical reactions.

Definition: Activation energy (Ea) is the minimum amount of energy required for a reaction to occur.

Reaction coordinate diagrams are explained, illustrating the energy changes during a reaction, including the formation of the activated complex.

Highlight: The top vertex of a reaction coordinate diagram represents the formation of the activated complex, a crucial intermediate in chemical reactions.

Maxwell-Boltzmann distributions are presented to show how temperature affects the distribution of molecular energies and, consequently, reaction rates.

Example: As temperature increases, the Maxwell-Boltzmann distribution curve flattens and broadens, indicating that more molecules have sufficient energy to overcome the activation energy barrier.