General Chemistry
In the world of chemical kinetics, the study of factors that affect the speed of chemical reactions, two important theories to remember are collision theory and transition state theory. Collision theory proposes that the speed of chemical reactions depends on three factors: the number of molecular collisions, the frequency of molecules being hit with enough energy to break existing bonds, and the proper alignment of molecules to break existing bonds when they collide. The reaction rate can be calculated using the equation Z x f, where Z represents the number of collisions per second, and f is the fraction of collisions that successfully break existing bonds.
Transition state theory focuses on the peak of the reaction, where molecules have the highest energy. It proposes the existence of a theoretical molecule that is half-product, half-reactant for a short time at the peak of the reaction. This transient molecule, however, cannot be isolated in a laboratory. The overall rate of a reaction is determined by the rate-determining step, the slowest part of the reaction mechanism. Reaction rates can be influenced by various factors, such as temperature, concentration of molecules, the medium in which the reaction takes place, and the presence of a catalyst, which lowers the activation energy for a specific reaction.
Lesson Outline
<ul> <li>Two important theories about chemical kinetics <ul> <li>Collision theory</li> <li>Transition state theory</li> </ul> </li> <li>Collision theory <ul> <li>Three factors affecting the speed of reactions <ul> <li>Number of molecular collisions</li> <li>Energy of collisions to break existing bonds</li> <li>Correct alignment of molecules</li> </ul> </li> <li>Reaction rate equation: Z x f</li> </ul> </li> <li>Activation energy <ul> <li>Energy needed to start a specific reaction</li> <li>Varies for each reaction</li> <li>Reverse reactions require more energy</li> </ul> </li> <li>Transition state theory <ul> <li>Peak of the reaction with highest energy</li> <li>Theoretical half-product, half-reactant molecules</li> <li>Transient nature of transition state molecules</li> </ul> </li> <li>Factors that influence reaction rate <ul> <li>Mechanism and rate-determining step</li> <li>Temperature</li> <li>Concentration of molecules</li> <li>Medium of reaction</li> <li>Catalysts</li> </ul> </li> </ul>
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FAQs
Collision theory is a concept in chemistry that explains how chemical reactions occur and why reaction rates differ for different reactions. It postulates that for a reaction to occur, the reactant molecules need to collide with each other with the correct orientation and a sufficient amount of energy, known as the activation energy.
Transition state rheory, also known as activated complex theory, is an extension of the collision theory. While collision theory focuses on the frequency and energy of molecular collisions, transition state theory delves deeper into the mechanism of the reaction itself. It proposes that reactions occur through the formation of a transient, high-energy configuration called the transition state or activated complex.
Activation energy is the minimum amount of energy required for reactants to form products in a chemical reaction. It acts as an energy barrier that must be overcome for the reaction to proceed. Higher activation energies result in slower reaction rates, as fewer molecules possess the required energy to initiate the reaction. Conversely, lower activation energies lead to faster reaction rates, as a greater number of molecules have the necessary energy to react. This explains why different reactions have different rates, depending on the energy required for the reaction to occur.
The rate-determining step (RDS) is the slowest step in a multi-step chemical reaction, effectively controlling the overall rate of the reaction. In a reaction mechanism, multiple steps occur sequentially, each with its own rate. The RDS is the step with the lowest rate, as it forms a bottleneck in the overall reaction progress. A reaction can only proceed as quickly as its rate-determining step allows. When studying or optimizing chemical reactions, understanding the RDS is crucial, as it holds the key to manipulating the overall reaction rate.
Catalysts are substances that increase the rate of chemical reactions without being consumed in the process. They function by lowering the activation energy of the reaction, effectively reducing the energy barrier for the reaction to occur. Catalysts accomplish this by providing an alternative reaction pathway or a new mechanism with a lower activation energy, often affecting the rate-determining step. As a result, more molecules can successfully collide and react, leading to an increase in the overall reaction rate. Catalysts are widely used in various industries, such as pharmaceuticals and chemical production, to optimize reaction rates and improve efficiency.
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