Understanding Slow Chemical Reactions: Mechanisms, Kinetics and Temperature Effects

Introduction to Slow Chemical Reactions

A slow chemical reaction is a type of reaction that progresses at a much slower rate than typical reactions. Common examples include the rusting of iron, the decomposition of organic matter, and food digestion in the human body. These reactions can take days, weeks, or even years to complete, depending on environmental factors such as humidity, salts, and acids.

Mechanisms of Slow Chemical Reactions

The primary example of a slow chemical reaction is the rusting of iron, which occurs when iron reacts with oxygen and moisture to form iron oxide, rust. The overall reaction can be described as:

4Fe 3O? 6H?O → 4Fe(OH)?

Other examples include the decomposition of organic matter and the digestion of food in the human body. In these cases, the reaction rate can be significantly influenced by various factors, such as the concentration of reactants, surface area, and environmental conditions.

Controlling Factors: Arrhenius Equation and Mass Transfer

Pierre Smith has detailed the common controlling factors of chemical kinetics, emphasizing the Arrhenius equation, which applies to mixed or close-contact molecules. However, in many practical scenarios, the rate of a reaction is controlled by mass transfer, which is the rate at which molecules come together.

For instance, the combustion of a gas jet is generally controlled by the rate at which the gas reaches the reaction site through the pipe. Similarly, in a wood fire, the reactions occur more quickly and intensely when air is forced, increasing the amount of oxygen available and thus the combustion rate. In liquid reactions, viscosity can also play a significant role.

The Arrhenius equation, given by:

ln(k) -E?/(RT) ln(A)

where E? is the activation energy, R is the gas constant, T is the temperature, and A is the pre-exponential factor or frequency factor, describes the temperature dependency of each reaction rate coefficient k. This relationship explains that reaction rates generally double for every 10 degrees Celsius increase in temperature.

Kinetic Rate and Temperature Effects

The rate of reaction is determined by the number of effective collisions between reacting particles. Not all collisions are effective, and products are formed only when the colliding particles possess a certain minimum energy called the threshold energy. A common rule of thumb is that reaction rates for many reactions double for every 10 degrees Celsius increase in temperature. The temperature coefficient Q, defined as the ratio of the rate constants at two temperatures, is commonly used.

For a given reaction, the ratio of its rate constant at a higher temperature to its rate constant at a lower temperature is known as its temperature coefficient Q. The value of Q10 is used as the ratio of rate constants that are 10 °C apart. Understanding these principles helps explain why some reactions can be slow, even at room temperature, unless they are endothermic.

References and Further Reading

The concepts described here are further elaborated in various scientific texts. For instance, Laidler and Meiser (1982) discuss the Arrhenius equation and chemical kinetics in detail, while Laidler (1987) provides insight into the study of reaction rates in solution. Other key references include Connors' work on chemical kinetics and Isaacs' comprehensive text on physical organic chemistry.

Key References: Laidler, K. J., Meiser, J. H. (1982). Physical Chemistry. Benjamin/Cummings. ISBN: 0-8053-5682-7. Laidler, K. J. (1987). Chemical Kinetics. Harper Row, 3rd ed. p. 277. ISBN: 0060438622. Connors, K. (1990). Chemical Kinetics: The Study of Reaction Rates in Solution. VCH Publishers, p. 14. ISBN: 978-0-471-72020-1. Isaacs, N. S. (1995). Physical Organic Chemistry. 2nd ed. Harlow: Addison Wesley Longman. ISBN: 9780582218635.