6 Kinetics

1. 6 Kinetics Year 11 DP Chemistry

2. Particle Collisions For a reaction to proceed to products, the reactants must collide with one another. Rate of reaction The rate of a reaction is the rate that reactants disappear and products form. As the reaction proceeds and reactants are used up, the rate decreases. Since the reacting particles must collide to react, increasing the rate of collisions increases the rate of reaction. Rate of reaction: the change in concentration of reactants or products over time

3. This graph shows the decrease in concentration of the reactants in a reaction. The instantaneous rate can be determined at any point along the graph by measuring the gradient which is change in concentration over time Δ[R] Δt The average rate of reaction can be calculated by using the concentration change over the entire reaction. Rate Equations Rate =

4. Some Factors Affecting Reaction Rates • Temperature • Higher temperature means particles have more kinetic energy, leading to an increase in reaction rate. • Concentration • Increasing the concentration of the reactants increases the chance of collisions, which increases the rate of reaction. • Surface area • Fine powders have more surface area than large pieces leading to more collisions. For this reason, fine powders create an explosion hazard when mixed with air • Catalysts • Catalysts lower the activation energy, therefore increasing the rate of reaction

5. Kinetic Theory describes the movement of particles in matter. There are three basic movements: • Vibrational (vibrate) • Rotational (spin) • Translational (move from here to there) Liquids have particles that are still close together, but they can move around. This is why they flow. Solids vibrate around fixed positions as the atoms are strongly attracted holding a definite shape. Gases have much greater movement and all three motions are possible. Kinetic Theory of Matter Source: http://www.saburchill.com/physics/chapters/0098.html

6. Collision Theory states that in order for reactions to occur, successful collisions must take place between reacting particles. To collide, they must be in motion... Kinetic Theory of Gases describes properties that result from the motion of gas particles. Some assumptions of the Kinetic Theory of Gases: A gas consists of a collection of small particles travelling in straight-line motion. The molecules in a gas occupy negligible volume (that is, they are points). Collisions between molecules are perfectly elastic (that is, no energy is gained or lost during the collision, only transferred from one particle to the other). There are no attractive or repulsive forces between the molecules. The average kinetic energy of a molecule is proportional to the absolute temperature ( in Kelvin). (Specifically, 3kT/2, where T is the absolute temperature and k is the Boltzmann constant.) Collision Theory

7. Kinetic Theory - Gases Source: http://isite.lps.org/sputnam/chem_notes/UnitVI_Gases.htm

8. Collision Theory • For a reaction to occur, there must be successful collisions. This means they must have: • Enough energy (>Ea) • Correct orientation http://boomeria.org/chemlectures/rates/rates.html

9. Collision Theory – particle size effect How will the rate of reaction be affected by the change in particle size?

10. Remember, temperature is a measure of the average kinetic energy of the particles. These particles need energy to get over the Activation Energy barrier in order to go to products. Collision Theory – Temperature effect How will the rate of reaction be affected by the increase in temperature? Go to: http://www.ltscotland.org.uk/highersciences/chemistry/animations/collision_theory.asp to see an animation about the effect of temperature

11. Collision Theory – Concentration effect How will the rate of reaction be affected by the increase in concentration?

12. Collision Theory – Pressure effect How will the rate of reaction be affected by the increase in pressure?

13. Particles in a sample of gas do not all have the same kinetic energy at a given temperature. The curve on the right shows that particles can have a wide range of energy values. The particles over the threshold energy, are able to react It also shows that if we increase the temperature, the number of particles over the threshold also increases. Note that the area under the two curves is the same and represents the total number of particles. The diagram on the right shows how the energy distribution curves relate to the enthalpy diagram. Notice that the threshold energy is the same as the Ea. Energy Distribution Curves

14. Gas particles in a container will increase their kinetic energy and their velocity with an increase in temperature. (K.E. = ½ mv2) Maxwell-Boltzmann distribution curves show this relationship. Maxwell-Boltzmann Energy Distribution Curves

15. M-B curves can be thought of in terms of probability: • The highest point in the curve represents the most probable kinetic energy value • The average kinetic energy is slightly higher due to the asymmetric shape of the curve Maxwell-Boltzmann Energy Distribution Curves

16. Recall that catalysts lower the Activation energy of a reaction. This means that more particles are energetic enough to react at the same temperature. See the enthalpy diagram and M-B distribution curve below. Catalysis and energy distribution

17. 6 Kinetics – 2011 Rob Sliderplease share this resource