Chemistry
Activation Energy
Activation energy is the minimum amount of energy required for a chemical reaction to occur. It represents the energy barrier that must be overcome for the reaction to proceed. Higher activation energy means the reaction is slower, while lower activation energy means the reaction is faster.
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7 Key excerpts on "Activation Energy"
eBook - PDF- Julianne Zedalis, John Eggebrecht(Authors)
- 2018(Publication Date)
- Openstax(Publisher)
Free energy diagrams illustrate the energy profiles for a given reaction. Whether the reaction is exergonic or endergonic determines whether the products in the diagram will exist at a lower or higher energy state than both the reactants and the products. However, regardless of this measure, the transition state of the reaction exists at a higher energy state than the reactants, and thus, E A is always positive. Chapter 6 | Metabolism 249 Watch an animation of the move from free energy to transition state at this (http://openstaxcollege.org/l/ energy_reaction) site. Explain why transitional states are unstable. a. Molecules have relaxed molecular structure with low energy. b. Molecules have strained molecular structure with high energy. c. Molecules have relaxed molecular structure with high energy. d. Molecules have strained molecular structure with low energy. Where does the Activation Energy required by chemical reactants come from? The source of the Activation Energy needed to push reactions forward is typically heat energy from the surroundings. Heat energy (the total bond energy of reactants or products in a chemical reaction) speeds up the motion of molecules, increasing the frequency and force with which they collide; it also moves atoms and bonds within the molecule slightly, helping them reach their transition state. For this reason, heating up a system will cause chemical reactants within that system to react more frequently. Increasing the pressure on a system has the same effect. Once reactants have absorbed enough heat energy from their surroundings to reach the transition state, the reaction will proceed. The Activation Energy of a particular reaction determines the rate at which it will proceed. The higher the Activation Energy, the slower the chemical reaction will be. The example of iron rusting illustrates an inherently slow reaction. This reaction occurs slowly over time because of its high E A .
- Leo J. Malone, Theodore O. Dolter, Steven Gentemann(Authors)
- 2012(Publication Date)
- Wiley(Publisher)
All collisions between reactants do not lead to products, however, since collisions must have the proper orientation and a minimum amount of energy known as the Activation Energy. This is the minimum kinetic energy that colliding molecules must have in order for old bonds to break so that new ones can form. The lower the Activation Energy, the faster the reaction. At maximum impact, colliding molecules form an activated complex, where the bonds rearrange to form products. The Activation Energy is equal to the difference in potential energy between reactants and the activated complex. Chemical reactions take place at a wide range of rates. The chemical reaction in an explosion is almost instantaneous, whereas the reaction in the aging of a fine wine takes years. The rate at which a particular reaction takes place is affected by five factors. They are as follows: 1. The Activation Energy. This depends on the particular reaction. 2. The temperature. The temperature relates to the average kinetic energy of the molecules. Thus, the higher the temperature, the more molecules that will have the minimum kinetic energy (the Activation Energy) for a reaction to occur. The rate of a reaction is also proportional to the frequency of collisions between reacting molecules. Since a higher temperature increases the average velocity of the molecules, collisions become more frequent at higher temperatures, which also increases the rate. 3. The concentration of reactants. Concentration also relates to the frequency of collisions. The more concentrated the reactants, the more frequent the collisions and the faster the rate of the reaction. 4. Particle size. In a heterogeneous reaction, the more surface area for a particular reactant, the more frequent the collisions and the faster the reaction. 297 5. Catalysts. A catalyst provides an alternate reaction pathway with lower Activation Energy which thus increases the rate of the reaction.
eBook - PDFPhysical Chemistry
Statistical Mechanics
- Horia Metiu(Author)
- 2006(Publication Date)
- Taylor & Francis(Publisher)
(b) Perform the same calculation if the energy difference is 90 kcal/mol. Copyrighted Material This conclusion has immediate implications for kinetics. During the reaction the point representing the molecular configuration has to pass over the ridge. Let us denote by C r any configuration on the ridge. If A/(C r )/A/*(A) is very small, only veiy few molecules reach the ridge. This means that only a few molecules react (go over the ridge, into the minimum corresponding to B). Of all the points on the ridge, the transition state has the lowest energy. This means that among the molecules whose configuration is close to the ridge, most are near the transition state; the crossing of the ridge is most likely to take place near the transition state. The transition state is like a mountain pass on a trail: if a choice is available, most people getting into the mountains are likely to go through the lowest mountain pass. While for hikers this is optional (a matter of common sense), for molecules this is dictated by the laws of statistical mechanics. At a given temperature, the higher the difference V(TS) — V(A), the smaller the number of molecules reaching the transition state (see Eq. 12.8) and the smaller the number of molecules that react. You have probably noticed the resemblance of the factor V(TS)-V(A) 6XP L hT . (which appears in Eq. 12.8) to the factor r E i that appears in the Arrhenius equation (Eq. 12.2). If we assume that the rate constant is proportional to the number of molecules that can reach the transition state, then we are led to conclude that E A = V(TS) — V(A) (12.10) This is a breathtaking conclusion: with qualitative, careful thinking, we have come to identify the meaning of the Activation Energy: it is the difference between the energy The Dynamics of a Chemical Reaction 211 of the saddle point (the transition state) and that of the reactant. Our argument is not a proof but it is a powerful conjecture.
eBook - PDFChemistry
An Industry-Based Introduction with CD-ROM
- John Kenkel, Paul B. Kelter, David S. Hage(Authors)
- 2000(Publication Date)
- CRC Press(Publisher)
The more reactant molecules there are, the more often contact will occur. Therefore, all other things being equal, reaction rate will increase with higher reactant concentrations . The exception to this is when reactions happen on a catalyst surface, such as in a catalytic converter in an automobile, where the restricted catalyst surface area limits the extent of reaction. The molecules must collide with sufficient energy to react (Fig. 14.1b). The kinetic energy of a system is proportional to its temperature. Therefore, the higher the temperature, the higher the velocity of the particles (K.E. mv 2 ). At higher velocities, particles collide more often and a higher fraction of them have enough energy for successful reaction. The amount of energy needed to initiate a reaction is called the Activation Energy, E a . We will discuss Activation Energy later. A rule of thumb is that the reaction rate will double for every 10°C increase in system temperature . Ammonium salts often react endo-thermically, so that the temperature of the reaction vessel actually decreases with the reaction. Yet even this reaction has an Activation Energy barrier that must be overcome in order for the reaction to proceed. This energy input goes to breaking bonds. The molecules must collide with the proper orientation for reaction (Fig. 14.1c). The molecules must meet in such a way that they can react. An analogy can be made between molecules and long-lost friends going to give each other a hug. If they are facing away from each other, the hug will not happen. Nor will it happen if they are beside each other. Only if they have the proper orientation, facing one another, will they be able to hug. So it is with particles colliding. A catalyst provides an alternative pathway for a reaction to occur (Fig. 14.1d). The chemical reactions that we discuss are really summaries of what are often very complex processes. For example, the process of photosynthesis is often summarized with this reaction.
eBook - PDFFundamental Chemical Kinetics
An Explanatory Introduction to the Concepts
- M R Wright(Author)
- 1999(Publication Date)
- Woodhead Publishing(Publisher)
products. The lowest potential energy pathway between reactant and product configurations represents the changes which take place during reaction. The critical configuration lies on this pathway at the configuration with the highest potential energy. It is called the activated complex or transition state, and it must be attained before reaction can occur. The rate of reaction is the rate of change of configuration through the critical configuration. This theory will be developed in Chapters 2 and 3. Transition state theory has proved to be a very powerful tool, vastly superior to collision theory, and only recently challenged by the modem advances in molecular beams and molecular dynamics which look at the microscopic details of a collision, and which can be regarded as a modified collision theory. These topics are developed in Chapters 7 to 10. The development of the experimental technique of molecular beams, and the computational technique of molecular dynamics have revolutionised the science of reaction rates. Kinetics has now moved to an even more fundamental level of study, where it is probing at the very heart of the microscopic details of molecular behaviour and reaction. There are now two branches of kinetics, conventional and microscopic studies, each of which alone constitutes a major aspect of chemistry, but which fully complement each other. Experimental techniques have also made astounding advances throughout the latter half of this century, and these are described in Chapter 7 to 10. Experimental techniques have been forced to advance in order to meet the two complementary needs (a) of being a tool with which to test theoretical advances, and (b) acting as an inspiration to develop theory to account for the experimental results. Sec. 1.2] Concepts involved in modem state to state kinetics 3 For a long time chemists have used the term mechanism to refer to the chemical steps which make up a reaction.
eBook - PDF- F. Xavier Malcata(Author)
- 2023(Publication Date)
- Wiley(Publisher)
Evans and Michael Polanyi, who developed the alterna- tive concept of transition state; this is one of the most important developments in chem- istry during the twentieth century – and essentially postulates existence of a quite unstable chemical intermediate, which can degenerate to product or regenerate reactant instead. The strongest argument behind this theory is again the hypothesized existence of a minimum energy threshold for a chemical reaction to occur (as pioneered by Arrhenius) – the so-called Activation Energy; hence, only reactants with sufficient (kinetic) energy in the first place will be able to convert it into the (potential) energy required for cleavage/establishment of chemical bonds. However, Eyring’s theory provides a more generic framework to understand how reaction rate is affected by processing parameters – including pressure and surface tension. The collision theory and the transition state theory are somehow equivalent to each other, though – in that they lead to basically the same functional form for dependence of a kinetic constant on temperature, within the (relatively narrow) temperature range of practical interest for enzymatic reactions. The most general case of temperature dependence encompasses its effects upon both enzyme-catalyzed reaction and enzyme decay itself; the latter may, however, reveal itself as a much more complex process, due to the topology of the enzyme molecule vis-à-vis with the molecular topology of common substrates.
eBook - PDF- Margaret Robson Wright(Author)
- 2005(Publication Date)
- Wiley-Interscience(Publisher)
The Kassel theory deals explicitly with this process, and imposes a much more severe restriction than does Hinshelwood. Before an activated molecule can react there must be a flow of energy at least " 0 into a 158 THEORIES OF CHEMICAL REACTIONS critical normal mode. Once this has happened the activated molecule has the correct distribution of vibrational energy for it to be able to react. The time lag is the time taken for this to happen. 4.5.11 Critique of the Kassel theory The derivation is much more complicated than those outlined previously, especially when the quantum version is considered. Both versions are a vast improvement on the previous theories, and predict that 1/k 1st versus 1/[A] should be a curve, and that the plot of log e k 1 versus 1/T should be linear. Both predictions are verified by experiment. More advanced theories continue to make further improvements. 4.5.12 Energy transfer in the activation step The change over from first order kinetics, when the reaction step is rate determining, to second order kinetics, when activation is rate determining, depends on the total pressure, the complexity of the reactant and to a lesser extent on the tempe- rature. Effect of pressure. This has been dealt with in detail on a quantitative basis. But on a physical basis it is more likely that activation is faster than the reaction step at high pressures where there is a greater chance of collisions, simply because there are more molecules around. At low pressures there is less chance of collisions occurring, again simply because there are fewer molecules around, and so activation is now slower than reaction. Remember: activation is bimolecular, while reaction is unimolecular. Complexity of the reacting molecule. This reflects the number of atoms in the molecule, which determines the number of normal modes of vibration open to the molecule.
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