Reaction Kinetics

Introduction

Studying chemical reactions has been the endeavor of human societies since time immemorial. Through the ages, our understanding of the subject has increased step-by-step, and just as important as the products of the reactions has been the exploration of the steps involved therein. Reaction kinetics, also known as chemical kinetics, is the branch of chemistry that deals with our quest to know at what rates the various reactions occur and what impact does it have on the final product. It is also an important field for the industrial sector where the efficiency in production time and conditions matter a lot. Reactions in industry are carried out in reactors, where substances are mixed together, perhaps heated and agitated for a while, and then transported to the next stage of the process. It is critical to understand how long it takes to maintain the reaction at a particular step before going on in order to ensure that the previous reaction has been completed before beginning the next.

We can infer what is going on at the molecular level by changing the reaction conditions and measuring the effect of the changes on the rate of reaction. Numerous processes may be enhanced by being aware of how reactions function. For instance, if a specific intermediate is understood to be involved in a process, circumstances (such as some solvents) that are incompatible with that intermediate may be avoided. Additionally, chemicals that facilitate particular phases in the process may be introduced.

Rates of chemical reactions

The rate at which chemical reactions occur varies substantially. Some are quite quick, while others may take millions of years to stabilize. The Reaction Rate of a chemical reaction is the rate at which the concentration of the reactants or of the products changes per unit time. A chemical reaction's speed may be defined as the change in concentration of a compound divided by the interval of time during which this variation is observed. In the form of equations, the rate of a chemical reaction may be represented as:

The rate of disappearance of R = decline in concentration of R/time taken = -[R]/t, whereas the rate of emergence of P = rise in concentration of P/time taken = + [P]/t

R stands for reactants, while P stands for products.

Because the concentration of the reactants decreases, [R] is a -ve quantity and since  the concentration of the products grows, [P] is a +ve quantity.

However, the actual process of experimentally measuring a reaction rate is not so straightforward. As the reactants are consumed, most reactions slow down. As a result, when measured over longer time intervals t, the rates produced by the formulas described above tend to lose their relevance. For such situations, instantaneous rates (also known as differential rates) are important.

Reactants A and B are consumed throughout the reaction indicated below, while the concentration of product AB grows. 

The slope of a tangent to the concentration-vs.-time graph gives the immediate rate of a reaction. These instantaneous rates are actually limiting rates, which are described as:

Rate Laws and Rate Constants

A rate law is an equation that connects the rate of a reaction to the rate constant and the reactant concentrations. A rate constant, k, is a constant of proportionality for a specific reaction. The general rate law is typically represented as follows:

The reaction rate for a particular reaction is an important tool for calculating the exact sequence of a reaction. The sequence of a reaction is crucial because it allows us to simply and effectively identify individual chemical processes. Understanding the reaction order rapidly enables us to grasp many aspects of the reaction, such as the rate law, units of the rate constant, half-life, and much more.

It is crucial to note that, while the rate law may be used to predict the order of the reaction, there is no fundamental link between the reaction order and the stoichiometric coefficients in the chemical equation. The most highly studied reaction orders are zero-order, first-order, second-order reactions, and third although fractional-order reactions do exist. If the reaction order with regard to [A] is 2 (s = 2) and [B] is 1 (t = 1), then the concentration of reactant A decreases by a factor of 2, and the quantity of B decreases by a factor of 1. When you have a reaction order of Zero (i.e., s+t=0), it signifies that the concentration of the reactants has no effect on the rate of reaction. The rate would not alter if reactants are removed or added to the mixture. 

Summarizing the rate expressions in terms of concentration, we have

• Rate computed in M/s units
• The concentration of a reactant/product is indicated by brackets around it.
• Instantaneous rate as the rate at a specific point in time
• Instantaneous rate derived from a straight line tangent that intersects the curve at a certain location
• Slopes represent instantaneous rates.

In all the reactions, an important factor called half-life (t1/2) is considered. It is the time necessary for reactant concentration to fall to half of its original value. First-order t1/2 is independent of starting concentrations. At each given moment of response, the half-life is the same. In a first-order reaction, reactant concentrations decline by half in each set of regularly spaced time periods.

Read more about Rate of Reactions

Factors Influencing Reaction Rates

Reaction rates are influenced a lot by external factors, apart from the concentration of the reactants. Some of the most common reasons that affect the rate of reaction are:

• Reactant concentrations 
• Temperature and pressure at which the reaction takes place
• A catalyst's presence
• Solid or liquid reactant and/or catalyst surface area
• Light or dark conditions
• Moisture in the atmosphere
• Choice of solvent(s)

Role of Catalysts in Altering Reaction Kinetics

A catalyst is a material that increases the rate of a chemical reaction without incurring irreversible chemical change. A homogeneous catalyst exists in the same phase as the reacting molecule, while a heterogeneous catalyst exists in a different phase than the reactants. In general, catalysts reduce the overall activation energy of a chemical process.

References

1. https://users.cs.duke.edu/~reif/courses/molcomplectures/Kinetics/KineticsOverview/KineticsOverview.pdf
2. https://www.britannica.com/science/chemical-kinetics
3. Atkins P. and de Paula J., Physical Chemistry (8th ed., W.H. Freeman 2006) p.793 ISBN 0-7167-8759-8
4. Image sources: Chem Libretexts