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Kinetics: Reaction Rates and Potential Energy Diagrams

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Presentation on theme: "Kinetics: Reaction Rates and Potential Energy Diagrams"— Presentation transcript:

1 Kinetics: Reaction Rates and Potential Energy Diagrams
Honors Chemistry Unit 9- Chapter 17

2 Collision Theory of Reactions
A successful reaction depends upon the effective collisions between molecules. Reactant molecules must collide with enough energy to break bonds Not every collision will result in a reaction. The minimum energy required to break the bonds of the reactants and start a reaction is called the Activation Energy

3 Examples of Activation Energy
Scraping a match on a rough surface to light it. Using that lighted match to light a candle. Lightening strike causes oxygen (O2) to form ozone (O3)

4 Reaction Rates The rate of a reaction can be increased, in general, by increasing the chances for effective collisions between molecules. Likewise, you can slow a reaction down by reducing the chances of collisions. There are four main factors that affect reaction rates.

5 1. Temperature Higher temperature molecules move faster  Molecules will have more energy There will be more effective collisions Molecules collide with the minimum energy (activation energy) more often.

6 2. Concentration Higher Concentration = More particles
More particles = more possible collisions More collisions = faster reaction! For a gas - Increase Pressure Volume decreases  concentration increases.

7 3. Surface Area Particle size determines surface area
Smaller particles  more surface area. More surface area means more area for collisions to occur MORE COLLISIONS = FASTER REACTION!

8 4. Catalysts A catalyst is a substance that increases the rate of a reaction without being used up in the reaction. It increases the rate of the reaction by lowering the activation energy for a reaction Low activation energy means = faster reactions There is a greater chance for an effective collision if the activation energy is lower.

9 Potential Energy Diagrams
In a reaction mixture the reactants and products contain potential energy. This potential energy is also known as enthalpy (symbol H). During a chemical reaction the enthalpy (or Potential Energy ) of the reactants changes as the reactants form new products. The enthalpy change (∆H) for a reaction can be calculated from a potential energy diagram. ∆H = H(products) - H(reactants) ∆H is measured in kilojoules per mole (kJ/ mol)

10 Exothermic Reactions An exothermic reaction releases heat, we can feel the heat given off to the surroundings as the reaction happens. **Energy required to break reactant bonds is less than the energy released when product bonds form. A + B  C + D + Energy/heat The enthalpy change for an exothermic reaction is always negative. (-ΔH)

11 Endothermic Reactions
An endothermic reaction absorbs heat, we can feel a beaker become colder as a reaction proceeds as heat is taken in from the surroundings. **Energy required to break reactant bonds is more than the energy released when product bonds form A + B + Energy/heat  C + D The enthalpy change for an endothermic reaction is always positive. (+ΔH)

12 ΔHrxn = Heat energy released/absorbed in reaction
Reverse reaction Activation energy ΔHrxn = Heat energy released/absorbed in reaction

13 Potential Energy Diagram with Catalyst
With a catalyst – The Activation Barrier is Lowered This Increases Reaction Rate Note - Enthalpy , ΔH Remains the same!

14 Exothermic ∆H = (-) Endothermic ∆H = (+)

15 Exothermic Reaction: Example
The activation energy (Ea) for the forward reaction is shown by (a), what is the value of the forward activation energy? Ea = 200 – 150 = 50 kJ/mole

16 The activation energy (Ea) for the reverse reaction is shown by (b): what is the value?
Ea (reverse) = = 150 kJ/mole

17 The enthalpy change for the reaction is shown by (c): What is the value of ∆H ?
∆H = H products) – H(reactants) = 50 – 150 = kJ/mol

18 Kinetics Movie Clip


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