Chapter 14: Problem 71
What are the differences between an intermediate and a transition state?
Short Answer
Expert verified
In summary, intermediate states are temporary species with defined geometries and lower energy levels, which can exist and be observed under certain conditions. On the other hand, transition states are high-energy barriers with no defined molecular structure, representing the critical point in the transformation between reactants and products. The main differences between the two lie in their stability, energy levels, existence, and molecular geometry.
Step by step solution
01
Definition
An intermediate state is a temporary species formed during a chemical reaction, which has a well-defined energy minimum. It exists for a discrete period and can further react to form products. On the other hand, a transition state represents the highest energy point along the reaction coordinate, which is between the reactants and products. It is not an actual molecular species but an energy barrier that needs to be crossed for the reaction to proceed.
02
Energy/state of stability
An intermediate has a lower energy than the transition state, indicating a greater stability than the transition state. An intermediate exists at an energy minimum and can be isolated under certain conditions, whereas a transition state exists at an energy maximum (called the activation energy), making it a high-energy, less stable entity that cannot be isolated.
03
Existence
Intermediate states can exist for a certain period and can be observed under specific experimental conditions. In contrast, a transition state exists for an extremely short period and cannot be directly observed due to its fleeting existence during the reaction.
04
Geometry
The geometry of molecules differs between intermediates and transition states. An intermediate has a defined geometry, which can often be determined using various spectroscopic techniques. A transition state, however, does not have a defined molecular structure, as it represents the highest energy point during the transformation from reactants to products.
05
Role in the reaction mechanism
In a reaction mechanism, intermediate states are actual chemical species that participate in reactions and can be involved in multiple steps of a reaction pathway. Transition states, however, represent the energy barriers between intermediates or reactants and products, and their role is to determine the activation energy and reaction rates.
In conclusion, the main differences between intermediates and transition states lie in their stability, energy levels, existence, and molecular structure. While intermediate states are actual chemical species with a defined structure, transition states are high-energy barriers that need to be crossed for a reaction to proceed.
Unlock Step-by-Step Solutions & Ace Your Exams!
-
Full Textbook Solutions
Get detailed explanations and key concepts
-
Unlimited Al creation
Al flashcards, explanations, exams and more...
-
Ads-free access
To over 500 millions flashcards
-
Money-back guarantee
We refund you if you fail your exam.
Over 30 million students worldwide already upgrade their learning with Vaia!
Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Intermediate States
Intermediate states in chemical reactions are fascinating because they act as stepping stones from reactants to products. They are actual chemical species that form temporarily during a reaction process.
They exist at an energy minimum, which makes them relatively stable compared to transition states. Because of their lower energy, intermediates can at times be isolated and studied under certain experimental conditions.
The presence of intermediates can provide insights into the mechanism of a reaction and the path it takes from start to completion.
They exist at an energy minimum, which makes them relatively stable compared to transition states. Because of their lower energy, intermediates can at times be isolated and studied under certain experimental conditions.
The presence of intermediates can provide insights into the mechanism of a reaction and the path it takes from start to completion.
- Intermediates have defined molecular geometries, making it possible to analyze them using spectroscopic techniques.
- Unlike transition states, intermediates can be stable enough to detect and study.
- Their stability is due to their lower energy compared to transition states.
Transition States
Transition states are critical yet elusive parts of a chemical reaction. They represent the pinnacle of energy that reactants must overcome to form products.
These are not tangible species that can be captured or observed directly, because they exist at the highest energy point along the reaction coordinate.
Despite their fleeting nature, understanding transition states is essential for grasping reaction kinetics.
These are not tangible species that can be captured or observed directly, because they exist at the highest energy point along the reaction coordinate.
Despite their fleeting nature, understanding transition states is essential for grasping reaction kinetics.
- They occur at the peak energy level, known as the activation energy, that must be crossed for a reaction to proceed.
- Lacking a defined structure, transition states are theoretical constructs used to explain the maximum energy barrier.
Activation Energy
Activation energy is the energy barrier that needs to be overcome for a chemical reaction to proceed. It is the minimum energy required for reactants to transform into products.
Imagine activation energy as a hill that molecules need to climb over to reach the other side.
The height of this energy "hill" determines how fast the reaction happens. The higher the activation energy, the slower the reaction because more energy is needed for the process.
Imagine activation energy as a hill that molecules need to climb over to reach the other side.
The height of this energy "hill" determines how fast the reaction happens. The higher the activation energy, the slower the reaction because more energy is needed for the process.
- Activation energy can be influenced by catalysts, which lower this energy barrier, allowing reactions to proceed faster and with less energy input.
- Understanding activation energy is crucial for controlling reaction rates and designing efficient chemical processes.
Reaction Coordinate
The reaction coordinate is a conceptual path that depicts the progression of a chemical reaction from reactants to products.
It is a helpful tool for visualizing the changes in energy throughout the course of the reaction.
By plotting the energy of molecules at various stages, this coordinate helps in illustrating both the intermediate and transition states.
It is a helpful tool for visualizing the changes in energy throughout the course of the reaction.
By plotting the energy of molecules at various stages, this coordinate helps in illustrating both the intermediate and transition states.
- It often appears as a graph showing energy on the y-axis and progress of the reaction on the x-axis (reaction pathway).
- Peaks on this graph represent transition states, while valleys indicate intermediates.