Name: collision theory and reaction speed
Uploaded: 2020-09-24T15:06:50.000Z
Duration: 2 min 55 s
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Chemistry and Physics

IPC.7

IPC.7.F

Temperature Dependence on Reaction Rate

Las velocidades de reacción química son sensibles a los cambios de temperatura. Un aumento de 10 grados Celsius puede acelerar la velocidad de reacción de 3 a 4 veces, pero:por qué? La ley de velocidad de una reacción define la relación entre la concentración de un reactivo y la velocidad de reacción.Dado que la concentración es independiente de la temperatura, solo la constante de velocidad permanece para influir en la velocidad de reacción dependiendo de la temperatura. Así, en 1889, el químico sueco Svante Arrhenius concluyó que la dependencia de la temperatura de la velocidad de reacción está incluida en la constante de velocidad. La constante de velocidad describe la relación entre la temperatura y los parámetros cinéticos relacionados con la energía de colisión, orientación y activación de las moléculas que reaccionan a través de la ecuación de Arrhenius.A es una constante llamada factor de Arrhenius o factor de frecuencia, e es un factor exponencial que integra la energía de activación medida en julios por mol, la constante de gas y la temperatura en kelvin. La dependencia de la temperatura de los parámetros se puede explicar con el modelo de colisión, que establece que las moléculas que reaccionan deben colisionar con suficiente energía en la orientación correcta para iniciar una reacción química. El factor de frecuencia consta de dos componentes:la frecuencia de colisión y el factor de orientación.La frecuencia de colisión es el número de colisiones moleculares por unidad de tiempo, mientras que el factor de orientación describe la probabilidad de colisiones con una orientación favorable. Aún así, solo una pequeña fracción de las colisiones conduce a una reacción. Esto se debe a que las moléculas que reaccionan tienen que superar una barrera de energía, llamada energía de activación, para transformarse en productos.Solo aquellas moléculas que colisionan con suficiente energía cinética tendrán suficiente energía potencial para doblar, estirar o romper enlaces, para transformarse en un intermedio de alta energía llamado estado de transición o complejo activado. El complejo activado, inestable y de corta duración, pierde energía para formar productos estables, cuya energía total es menor que la de los reactivos. El factor exponencial en la ecuación de Arrhenius representa la fracción de colisiones exitosas que dan como resultado productos.Un aumento de temperatura influye tanto en el factor de frecuencia como en el factor exponencial. A temperaturas elevadas, las moléculas se mueven más rápido, con más fuerza y con energías térmicas más altas, lo que conduce a colisiones más favorables. Por lo tanto, un aumento de temperatura da como resultado una frecuencia más alta y factores exponenciales que conducen a un aumento en la constante de velocidad, lo que se traduce en una velocidad de reacción acelerada.

dependencia de la temperatura en la velocidad de reacción

La teoría de las colisiones
 Los átomos, moléculas o iones deben colisionar antes de que puedan reaccionar entre sí. Los átomos deben estar cerca para formar enlaces químicos. Esta premisa es la base de una teoría que explica muchas observaciones con respecto a la cinética química, incluyendo factores que afectan las velocidades de reacción.
 La teoría de las colisiones se basa en los postulados de que (i) la velocidad de reacción es proporcional a la velocidad de las colisiones de los reactivos, (ii) las especies que reaccionan chocan en una orientación que permite el contacto entre los átomos que se unen en el producto, y (iii) la colisión se produce con energía adecuada para permitir la penetración mutua de las capas de valencia de las especies que reaccionan, de modo que los electrones puedan reorganizarse y formar nuevos enlaces (y nuevas especies químicas).
 Por ejemplo, en una reacción en fase gaseosa entre el monóxido de carbono y el oxígeno, que se produce a alta temperatura y presión, el primer paso es una colisión entre las dos moléculas.
 <img alt="Eq1" src="/files/ftp_upload/11379/11379_Equation_1.png" style="text-align: center; height: 50px;" />
 Sin embargo, podría haber muchas orientaciones relativas posibles diferentes en las cuales las dos moléculas colisionan. Por lo tanto, la orientación de las moléculas en colisión tiene gran importancia en la determinación parcial de la viabilidad de que ocurra una reacción entre ellas.
 En un caso, el lado del oxígeno de la molécula de monóxido de carbono puede colisionar con la molécula de oxígeno. En otro caso, el lado del carbono de la molécula de monóxido de carbono puede colisionar con la molécula de oxígeno. Es más probable que el segundo caso dé lugar a la formación de dióxido de carbono, con un átomo central de carbono unido a dos átomos de oxígeno (O = C = O).
 <img alt="Image1" src="/files/ftp_upload/11379/11379_Image_1.png" style="text-align: center; height: 250px;" />
 Sin embargo, incluso si la colisión tiene lugar en la orientación correcta, la garantía de que la reacción procederá a formar dióxido de carbono es limitada. Esto se debe a que, además de la orientación adecuada, la colisión también debe ocurrir con suficiente energía, llamada energía de activación, para dar lugar a la formación del producto. Cuando las especies reactantes chocan en orientación correcta y con suficiente energía de activación, se combinan para formar una especie inestable llamada complejo activado o estado de transición. Estas especies son de corta duración y generalmente indetectables por la mayoría de los instrumentos analíticos. En algunos casos, las mediciones espectrales sofisticadas pueden observar estados de transición.
 La teoría de las colisiones explica por qué la mayoría de las velocidades de reacción aumentan a medida que aumenta la temperatura; con un aumento de la temperatura, aumenta la frecuencia de las colisiones. Más colisiones significan una velocidad de reacción más rápida, suponiendo que la energía de las colisiones es adecuada.
 Energía de activación
 La energía mínima necesaria para formar un producto durante una colisión entre reactivos se denomina energía de activación (Ea). La diferencia en la energía de activación requerida y la energía cinética proporcionada por las moléculas reactantes en colisión es un factor primario que afecta la velocidad de una reacción química. Si la energía de activación es mucho mayor que la energía cinética promedio de las moléculas, la reacción ocurrirá lentamente, ya que sólo unas pocas moléculas de movimiento rápido tendrán suficiente energía para reaccionar. Si la energía de activación es mucho más pequeña que la energía cinética promedio de las moléculas, una gran fracción de moléculas tendrá la energía suficiente, y la reacción procederá rápidamente.
 Los diagramas de reacción se utilizan ampliamente en cinética química para ilustrar varias propiedades de una reacción de interés. Muestra cómo cambia la energía de un sistema químico a medida que sufre una reacción, convirtiendo los reactivos en productos.
 Por ejemplo, considere el siguiente diagrama de reacción para una reacción exotérmica: A + B → C + D;
 <img alt="Image2" src="/files/ftp_upload/11379/11379_Image_2.png" style="text-align: center; height: 350px;" />
 El diagrama se visualiza de izquierda a derecha. Inicialmente, el sistema sólo consta de reactivos (A + B). Una vez que las moléculas reactantes con suficiente energía colisionan, forman un complejo activado de alta energía o un estado de transición. A continuación, el estado de transición inestable decae para producir productos estables (C + D).
 El diagrama muestra la energía de activación de la reacción, Ea, como la diferencia de energía entre los reactivos y el estado de transición. La diferencia de energía entre los reactivos y los productos corresponde al cambio de entalpía de la reacción (ΔH). En este caso, la reacción es exotérmica (ΔH < 0), ya que produce una disminución en la entalpía del sistema.
 Ecuación de Arrhenius
 La ecuación de Arrhenius, k =Ae−Ea/RT relaciona la energía de activación y la constante de velocidad, <ek, para muchas reacciones químicas.
 En esta ecuación, R es la constante de gas ideal, que tiene un valor de 8,314 J/mol·K, T es la temperatura en kelvin, Ea es la energía de activación en jules por mol, e es la constante 2,7183, y A es una constante llamada factor de frecuencia, que está relacionado con la frecuencia de las colisiones y la orientación de las moléculas que reaccionan. La ecuación de Arrhenius se adapta perfectamente a los postulados de la teoría de las colisiones. El factor de frecuencia, A, refleja qué tan bien que las condiciones de reacción favorecen las colisiones correctamente orientadas entre moléculas reactivas. Una mayor probabilidad de colisiones eficientemente orientadas da como resultado valores mayores para una A y velocidad de reacción más rápidas.
 El término exponencial, e−Ea/RT, describe el efecto de la energía de activación en la velocidad de reacción. Según la teoría cinético molecular, la temperatura de la materia es una medida de la energía cinética promedio de sus átomos o moléculas constituyentes; una energía de activación más baja resulta en una fracción más significativa de moléculas energizadas adecuadamente y una reacción más rápida.
 El término exponencial también describe el efecto de la temperatura sobre la velocidad de reacción. Una temperatura más alta representa una fracción correspondientemente mayor de moléculas que poseen suficiente energía (RT) para superar la barrera de activación (Ea). Esto produce un valor más alto para la constante de velocidad y una velocidad de reacción más rápida.
 <img alt="Image3" src="/files/ftp_upload/11379/11379_Image_3.png" style="text-align: center; height: 280px;" />
 Las distribuciones de energía molecular muestran el número de moléculas con energías que exceden dos energías de activación diferentes a una temperatura dada, y una energía de activación dada a dos temperaturas diferentes.
 Este texto es adaptado de<a href="https://openstax.org/books/chemistry-2e/pages/12-5-collision-theory"> Openstax, Química 2e, Sección 12.5: Teoría de las Colisiones. </a>

Dependencia de la Temperatura en la Velocidad de Reacción

La teoría de la colisión
Los átomos, moléculas o iones deben colisionar antes de que puedan reaccionar entre sí. Los átomos deben estar muy juntos para formar enlaces químicos. ...

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collision theory and reaction speed

Collision theory explains why reaction speed depends on how particles collide. Atoms, molecules, or ions must come close together before they can react. The theory says the reaction rate depends on the collision rate, the orientation of the colliding species, and whether the collision has enough energy for bonds to change.

A gas-phase reaction between carbon monoxide and oxygen shows why orientation matters. The molecules can collide in more than one way. If the oxygen side of carbon monoxide hits oxygen, the collision is less likely to form carbon dioxide. If the carbon side of carbon monoxide hits oxygen, the arrangement is more favorable for making O = C = O.

Even a correctly oriented collision may not react if it does not have enough energy. The minimum energy needed to form products is called activation energy, or Ea. When reactants collide with the right orientation and enough energy, they form an unstable activated complex, also called a transition state. This state is very short-lived and usually cannot be detected with standard instruments, although some advanced spectral methods can observe it.

Temperature affects reaction rate because higher temperature increases collision frequency. More collisions usually means a faster reaction, as long as the collisions also have enough energy. If the activation energy is much larger than the average kinetic energy of the molecules, only a small number of fast-moving particles can react, so the reaction is slow. If the activation energy is much smaller, many molecules have enough energy, and the reaction is faster.

Reaction diagrams help show how energy changes as reactants become products. In an exothermic reaction such as A + B → C + D, the system starts with reactants, reaches the high-energy transition state, and then forms stable products. The activation energy is the energy gap between the reactants and the transition state. The enthalpy change, ΔH, is the energy difference between the reactants and products, and it is negative for an exothermic reaction.

The Arrhenius equation links activation energy and reaction rate: k = Ae−Ea/RT. In this equation, R is the ideal gas constant, T is temperature in kelvin, Ea is activation energy in joules per mole, e is 2.7183, and A is the frequency factor. The frequency factor reflects how often collisions occur and how well the molecules are oriented. The exponential term shows that higher temperature and lower activation energy increase the rate constant and speed up the reaction.

Collision Theory and Reaction Speed

reaction rate

Measuring Reaction Rates

Measuring reaction rates shows how fast reactants are used up or products are formed over time. The rate of reaction is the change in the amount of a reactant or product per unit time. Because of that, scientists measure how a property changes as the reaction proceeds and connect that change to reactant or product amounts. For reactions that make or use a gas, the rate can be followed by changes in volume or pressure. A rate expression is the mathematical way to write how concentration changes...

measuring reaction rates

Tracking Reaction Kinetics with Light and Pressure

Reaction kinetics can be tracked with polarimetry, spectrometry, and pressure measurements. These methods let scientists follow a reaction without disturbing the sample. They also help determine how reactant or product concentrations change over time. Polarimetry is used for optically active substances. These substances rotate the plane of polarized light, which is called optical rotation. The effect comes from molecular structure. Normal monochromatic light is unpolarized, but after it passes...

concentration and rate law

Reaction Order from Concentration

Reaction order from concentration shows how reactant amounts affect reaction rate. Rate laws, also called differential rate laws or rate equations, are mathematical expressions that connect reaction rate with the concentrations of reactants. For a generic reaction a A + b B → products, the rate law can be written using the reactant concentrations [A] and [B] and a rate constant k. In a rate law, k is the rate constant for one specific reaction at one specific temperature. The exponents m and n...

determining order of reaction

Finding Rate Law from Initial Rates

Rate laws connect the speed of a chemical reaction to the concentration of its reactants. The rate constant, k, and the reaction orders are not taken from the chemical equation. They are found experimentally by watching how the rate changes when the reactant concentrations change. One common way to do this is the method of initial rates. In this method, several trials are run with different starting concentrations. The measured rates are then compared to determine the reaction orders and,...

the integrated rate law: the dependence of concentration on time

the-integrated-rate-law-the-dependence-of-concentration-on-time

Reading Reaction Progress with Integrated Laws

Integrated rate laws connect reaction concentration to time. They help show how much reactant or product is present after a certain period. They can also help estimate how long a reaction must run to reach a given extent. A common use is finding how long radioactive material must be stored before its radioactivity drops to a safe level. These equations come from the differential rate law after it is integrated with respect to time. For a first-order reaction, where rate = k[A], the integrated...

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Reaction Half-Life in Kinetics

Reaction half-life is the time needed for half of a reactant to be consumed. After each half-life, half of the remaining reactant is used up. A hydrogen peroxide decomposition example shows this pattern clearly. The concentration drops from 1.000 M to 0.500 M in the first 6.00 hours, then to 0.250 M in the next 6.00 hours, and then to 0.125 M in the third period. Each later 6.00-hour interval removes half of what is left. For first-order reactions, the half-life does not depend on the reactant...

temperature dependence on reaction rate

Collision theory explains why reaction speed depends on how particles collide. Atoms, molecules, or ions must come close together before they can react. The theory says the reaction rate depends on the collision rate, the orientation of the colliding species, and whether the collision has enough energy for bonds to change. A gas-phase reaction between carbon monoxide and oxygen shows why orientation matters. The molecules can collide in more than one way. If the oxygen side of carbon monoxide...

arrhenius plots

Determining Activation Energy from Temperature

The Arrhenius equation links activation energy and the rate constant, k, for chemical reactions. It is written as k = Ae−Ea/RT. In this equation, R is the ideal gas constant with a value of 8.314 J/mol·K, T is the temperature on the kelvin scale, Ea is the activation energy in J/mole, e is the constant 2.7183, and A is the frequency factor. The frequency factor depends on collision frequency and the orientation of reacting molecules. Experimental kinetic data can be used to find the activation...

reaction mechanisms

Elementary Steps in Reaction Mechanisms

Chemical reactions can happen in steps, and a balanced equation does not show the full path at the molecular level. A reaction mechanism, also called a reaction path, gives the step-by-step process that explains how reactants become products. The decomposition of ozone is one example of a reaction that can be described by a sequence of steps. Each step in a mechanism is an elementary reaction. These elementary reactions happen in order, and when they are added together, they give the balanced...

rate-determining steps

Reaction Bottlenecks in Chemical Kinetics

Reaction bottlenecks help explain rate-determining steps in multistep reaction mechanisms. In a mechanism with several elementary steps, one step is often much slower than the others. That slowest step is called the rate-limiting step, or rate-determining step. A reaction cannot move faster than its slowest step, so this step controls the overall reaction rate. The idea is like traffic on a four-lane freeway with one lane closed. Cars can only move as fast as the bottleneck allows. In the same...

catalysis

Catalyst Action in Reaction Rates

Catalysts change the rate of a chemical reaction by giving it a faster path. A catalyst increases reaction rate without being consumed during the process. Its effect is often explained with reaction mechanisms and energy diagrams. For an endothermic reaction, the reactants and products start and end at the same energies in both cases. The uncatalyzed path is a one-step mechanism with one transition state. The catalyzed path follows a different mechanism with two transition states and a lower...

enzymes

Enzyme Action and Inhibition

Enzymes are biological catalysts that speed up many reactions in living organisms. They help control cellular metabolism by lowering the activation energy needed for biochemical reactions. When activation energy drops, reaction rates can increase. Enzyme problems can also lead to serious disease. A genetic abnormality that causes G6PD, or glucose-6-phosphate dehydrogenase, deficiency can disrupt the pathway that supplies NADPH to cells. In red blood cells, this can reduce glutathione and...

key terms and concepts

Learn how collision theory, activation energy, and the Arrhenius equation explain why higher temperature can speed up chemical reactions in high school chemistry.

Collision Theory and Reaction Speed | JoVE K12

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Collision Theory and Reaction Speed

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