Unit 9 Reaction Rates

1. Unit 9 Reaction Rates

2. Reaction Rates • Some chemical reactions go fast, some slowly rapidly, some reactions go to completion, some seem to get stuck halfway, and still others do not seem to occur at all. This unit is about the factors that control the speed and outcome of a reaction.

3. Reaction Rates • Take a look at the following exothermic rxns. : • P4 (s) + 5 O2 (g) → P4O10 (s) • 4 Fe (s) + 3 O2 (g) → 2 Fe2O3 (s) • 2 H2 (g) + O2 (g) → 2 H2O (l) Reaction 1: Ignites spontaneously; phosphorus kept under water Reaction 2: Rust process is slow; may not be noticeable for years Reaction 3: BOOM!!! But, hydrogen & oxygen may be stored indefinitely until sparked.

4. What determines how a rxn. starts? How fast it proceeds? • It’s all about…kinetics. • To determine the speed or rate of a reaction, you would measure the increase in product concentration over time, or the decrease in reactant concentration over time: • rate = decrease in reactant concentration time interval or • rate = increase in product concentration time interval

5. What you choose to measure depends mostly on what is easiest to measure. For this reaction: • CO (g) + NO2 (g) → CO2 (g) + NO (g) colorless red-brown colorless colorless

6. Factors that affect reaction rates: Concentration • If you add HCl to a solution of sodium thiosulfate, the thiosulfate ion decomposes to produce fine particles of solid sulfur, which makes the solution cloudy: S2O32– (aq) + 2 H1+ (aq) → S (s) H2SO3 (aq) • If the S2O32– solution is twice as concentrated, the mixture becomes cloudy twice as fast. The result is not surprising, but can we explain it in molecular terms?

7. Factors that affect reaction rates: Concentration • Anything that would increase the number of collisions would be likely to increase the rate of the reaction. • This is called, collision theory of reaction rates. • From collision theory, you predict that increasing concentration increases reaction rate. (more collisions)

8. Factors that affect reaction rates: Concentration • Anything that brings the components of a reaction into better contact is likely to increase the number of successful collisions, and thus the rate or speed of the reaction: • increasing the number of reactant molecules in the mixture (increasing concentration) • increasing the surface area of a solid reacting with a liquid or gas • stirring to increase contact between a solid and a liquid

9. Increasing concentration will not always affect reaction rate. • Example: CO (g) + NO2 (g) → CO2 (g) + NO (g) • Adding more CO does not change the rate • But, add more NO2 makes it go faster • Why? • The overall speed of the reaction is determined by the slowest step, called the rate determining step. For the reaction of NO2 and CO, the mechanism that has been suggested is • (1) NO2 + NO2 → NO3 + NO slow • (2) NO3 + CO → NO + CO2 fast

10. Factors that affect reaction rates: Temperature • Collisions alone are not enough to cause a chemical change. • Most collisions are ineffective and don’t form a new product. • “To be effective, a collision must involve a certain amount of energy.”

11. “To be effective, a collision must involve a certain amount of energy.”

12. Can the activation energy be lowered? • You bet! This is how a catalyst affects a reaction.

13. Catalyst • The catalyst is involved in the reaction mechanism but is not changed in the overall reaction. • Reaction can proceed at a lower temperature than would be required for the uncatalyzed reaction mechanism.

14. Real World Application of Catalyst • Enzymes are proteins that act as biological catalyst. • Photosynthesis

15. Equilibrium • Colorless CO gas reacts with red-brown NO2 gas to produce colorless CO2 and NO gases: • CO (g) + NO2 (g) → CO2 (g) + NO (g) colorless red-brown colorless colorless 1) rxn. proceeds, color fades, stops fading, rxn. is stuck 2) Rxn. is at equilibrium

16. When is a system at equilibrium? • When its macroscopic properties –– temperature, concentration, color, pressure, or any other measurable property of the system –– are constant. • Equilibrium can only be achieved in a closed system. • We recognize that any process is at equilibrium when its observable properties become constant, and we explain equilibrium as the balance of two continuous opposing processes.

17. LeChatlier’s Principle • In 1888 a French chemist, Henri LeChatelier, summarized observations of how changing conditions affect equilibrium in a rule known as LeChatelier’s Principle: • If a system at equilibrium is changed in a way that upsets the equilibrium, processes occur that minimize the disturbance and return the system to equilibrium.

18. Consider if you will… • 2 NO2 (g) ↔ N2O4 (g) • But...what if you add NO2 gas? • 1) NO2 is consumed and more N2O4 will be produced • 2) Two rates come back into balance, equilibrium is restored. • 3) NO2 you added has been consumed and more N2O4 has formed • We say that this reaction “shifted right” ; more product produced

19. Same reaction…different scenario • 2 NO2 (g) ↔ N2O4 (g) • But… what if you remove NO2 • 1) Forward rate slows down • 2) New equilibrium, the concentration of N2O4 is lower than before, some of the NO2 you removed has been replaced • We say this reaction “shifted left”; more reactant

20. What else can we do to affect equilibrium? • Change the temperature • Change the volume. • What would this look like for temp? • When you think of temp, think energy. • 2 NO2 (g) ↔ N2O4 (g) + energy (exothermic rxn) • Cooling rxn. removes energy- LeChatlier’s predicts rxn. will shift right to replace energy • Heating rxn.- shifts rxn to the left using some of the added energy.

21. What would this look like for volume? • 2 NO2 (g) ↔ N2O4 (g) • Increase volume (decrease pressure)- shifts rxn. toward more gas molecules to restore pressure • Decrease volume (increase pressure)- shifts rxn. toward fewer gas molecules to relieve pressure

22. Hey! What about a catalyst? • Adding a catalyst does not affect the equilibrium. • Lowers activation energy, allows rxn. to achieve equilibrium sooner • The Rule of Thumb (LeChatlier’s Principle predicts…) • equilibrium shifts away from whatever you add • equilibrium shifts toward whatever you remove

23. Law of Mass Action • We can describe any state of dynamic equilibrium quantitatively with an expression known as the law of • mass action. • Imagine for a moment the hypothetical reaction • aA + bB ↔ cC + dD