RATE OF REACTION
P R E S E N T E D B Y K E V I N S T A R K ( L I U )
Rate of reaction refers to the speed of a chemical reaction proceeds. It can be expressed as the formed product’s concentration per unit volume in a unit of time. It can be also expressed as the amount (concentration) of reactants consumed during a specific period of time. (Laidler, 2018)
Rate of reaction can be affected by the number of collisions between particles, varies of affective factors below,
- Temperature (the change in kinetic energy between particles indicates the change in temperature, the higher the temperature, the faster a chemical reaction would occur, molecules under a high temperature condition possess more kinetic energy to break apart into atoms to react with other chemicals. Instead a lower temperature leads the particles having less kinetic energy and being ‘lazy’ to move around)
- Surface Area (the change in surface area affects the chance particle collides with each other. Increasing surface area can boost the reaction effectively by making particles more likely to collide. Moderate surface area would have a relatively lower rate of reaction, which consumes more time to be completely reacted)
- Pressure (the change in pressure between particles determines the distance between molecules, adding pressure to reactants reduce the distance between particles and molecules / atoms would be more likely to collide each other. Less pressure would disperses particles and less chance collision would occur)
- Catalyst (the use of catalyst will speed up the reaction but does not react within a chemical reaction)
The collision theory is used to determine the rate of reaction. It states that gas phrase chemical reactions occur when molecules collide with sufficient energy. (LibreTexts, 2017) The rate at which a chemical reaction proceeds is equal to the frequency of effective collisions. The particles must have enough energy for the collision to be successful in producing a reaction. The rate of reaction depends on the rate of successful collisions between reactant particles. The reaction occurs faster as long as the higher success rate of collision there is. (BBC - GCSE Bitesize, 2011) Because atomic or molecular frequencies of collisions can be calculated with some degree of accuracy only for gases, the application of the collision theory is limited to gas-phase reactions.
Activation energy is the minimum amount of energy a reaction or physical transport requires to occur. In a chemical reaction, molecules must collide in order to be reacted, according to the collision theory. At the same point, collision of particles requires a certain amount of kinetic energy that causes the movement of them, which is activation energy. Additionally, the success rate of a collision also depends on the molecular orientation at the time, atom(s) attached on a molecule must face the proper side in order to strike off the atom(s) on another molecule to form a new substance. Both aspects determine the amount of successful collision in a chemical reaction.
The way that can increase the number of collisions:
- Increase in energy level (e.g. increase temperature, add pressure) between particles that are being reacted, therefore more kinetic energy is applied to the reactant particles, the success rate of collisions is increased.
- Increase in particle amount (concentration) can lower the distance between particles and increase the chance of successful collision in a unit time. The more successful collision in a unit time there is, the faster the reaction would occur (higher rate of reaction)
- Increase in surface area can efficiently boosts the reaction by increasing the chance reactant particles meet together and collide.
A catalyst is a substance that speeds up a chemical reaction, but does not react during the reaction process, it provides an alternative route for the reaction with a lower activation energy. acts a role of boosting agent in a chemical reaction. A catalyst can still be chemically recovered after it catalyze the reaction.
Catalysts are also used in industrial process. For example, during the process of producing ammonia with nitrogen gas and hydrogen gas, iron can be an effective catalyst. As a catalyst, the iron (high purity magnetite, Fe3O4) is being grinded into powder (providing more surface area to increase the efficiency to the experiment) and placed in the container with the reactants (hydrogen and nitrogen gases) before the process starts. (Essential Chemical Industry, 2013)
Kevin Stark (Liu)
Year 11 Physics, Condell Park High School
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