1. Upcoming Science Club Presentations
Sanofi Biogenius Challenge
Fri. Oct. 23
Noon in Room 229
Justin Sietz – Freelance Hacker
Wed. Oct. 28
3:30pm in Room 220
WEAMS
(Wild and Exotic Animal Medical Society)
3:30pm in Room 229

2. HW Check & Student-Directed Project Research Check-in
While I come around to do these checks, do the following question:
Draw a labelled potential energy graph with all of the following criteria:
1 intermediate E
2 activated complexes
Ea 1 = 100kJ
Ea 2 = 20kJ
ΔH = + 40kJ
Products are C+D
Reactants are A+B
Is this reaction endothermic or exothermic?
Which step is the rate determining step?

3. Factors that Affect Rates of Reaction
Unit 3a: Reaction Rates

4. 5 min Video

5. Factors that affect rate of reactions
Video Talks about:
1. Decrease volume/increase concentration
2. # Particles
3. Temperature
4. Surface Area
5. Catalysts

6. Factors that affect rate of reactions
Video Talks about:
1. Decrease volume/increase concentration
2. # Particles
3. Temperature
4. Surface Area
5. Catalysts
There are 4 main factors:
1. Nature of the Reactants
2. Temperature
3. Concentration and Pressure
4. Catalysts

7. Nature of the Reactants
Since chemical bonds are broken and reformed in a chemical reaction, the nature of the bonds themselves plays an important role in the rate of reaction.
That is, how readily the bonds are broken and formed.
1. A reaction which involves primarily the ionic bonds (exchange of electrons) tends to be very quick.
Reactions that involve ions in solution tend to be rapid.

8. Nature of the Reactants
2. Reactants that involve covalent bonds tend to be very slow, unless they are highly exothermic (combustions).
For example, the decomposition of hydrogen peroxide into hydrogen gas and oxygen gas happens very slowly.
3. The phase of the reactants is also important.
That is, reactants in liquid, solution, or gas form react much more rapidly than solids.

9. Nature of the Reactants
4. The exposed surface area also affects the rate of reaction. The greater the surface area, the faster the reaction.
For example, kindling vs. logs when starting a fire.
5. Stirring also increases the reaction rate because it increases the frequency of collisions.

10. Temperature
In general, if the temperature of a system increases, so does the rate of reaction.
A rule of thumb is that the rate will double with an increase of 10˚C.
Why does the rate increase?
Faster particles (increased kinetic energy)
More and harder collisions
More particles have enough energy (the activation energy) to make successful collisions.
Faster reactions

11. Temperature
Note that increasing the temperature does not change the activation energy.
Potential energy curves will remain unchanged when we raise the temperature of a system.
The curve says nothing about the rate, but if the system is at a higher temperature, more particles would be able to get ‘over the hill’ to make successful reactions.

12. Concentration and Pressure
Increase reactants concentration = Increased rate of reaction
Only if the reactants are all in the same phase (homogeneous chemical reaction)
Why?
because if you place more particles in a given area, collisions will be more frequent.
Think about a crowded room compared to one with few people in it.

13. Concentration and Pressure
GASES
The concentration can be increased by decreasing the volume (putting the gas in a smaller container).
This forces the particles to be closer together, increasing the concentration.
This also increases the pressure.
***Changing the concentration or pressure does not alter the energy of the particles; it just increases the frequency of the collisions.

14. Concentration & Pressure
To summarize, as a general rule of thumb:
As concentration of reactants increases, the rate of reaction also increases (and vice versa)
As the pressure on the reactants increases, the rate of reaction also increases (and vice versa)
There are a few exceptions to these rules:
i.e. – increasing one reactant, but not another (resulting in the limiting reactant getting used up and ending the reaction). This process would still happen faster than if you didn’t increase the concentration of one reactant
i.e. – increasing the pressure to the point where reactants cannot interact anymore
However, for our purposes the rules of thumb will apply

15. Concentration and Pressure
Example 1: Consider the following reaction that occurs between hydrochloric acid, HCl, and zinc metal:
2HCl(aq) + Zn(s) → H2 (g) + ZnCl2 (aq)
Will this reaction occur fastest using a 6 M solution of HCl or a 0.5 M solution of HCl? Explain.
6M  increasing concentration means there will be more collisions between particles

16. Concentration and Pressure
Example 2: Again consider the reaction between hydrochloric acid and zinc.
How will increasing the temperature affect the rate of the reaction? Explain.
Increasing temp. means more molecules will have sufficient energy to collide, and there will be more collisions  a faster reaction

17. Concentration and Pressure
Example 3: Based on the following kinetic energy curves, which reaction will have a faster rate, A or B? Explain.
Also, which reaction, A or B, would benefit most in terms of increased rate if the temperature of the system were increased?
A B

18. Concentration and Pressure
Reaction B will have a faster rate, since it has more particles that are likely to have sufficient energy for a reaction to occur.
Reaction A would then benefit more from a temperature increase, since it has more particles below the threshold energy.
A B

19. Catalysts
A catalyst is a substance that provides an alternate reaction mechanism with a lower activation energy.
Substance that speeds up a reaction without being used up
Speeds up reaction by giving the reaction a new path.
The new path has a lower activation energy.
More molecules have this energy.
The reaction goes faster.
Inhibitor- a substance that slows down a reaction (raises the activation energy).

20. Catalysts Solid line = uncatalyzed reaction
Dotted line = catalyzed reaction
Notice that for the diagram above, the ΔH = 15 J for both the catalyzed and uncatalyzed reaction. Therefore, the heat of reaction is independent of the pathway (Hess’s Law).
However, the activation energy for the uncatalyzed reaction is 25 kJ, while it is only 10 kJ for the catalyzed reaction.

21. Catalysts
Although a catalyst can be introduced, it only offers an alternate pathway for the reaction. Therefore, some particles may not use the catalyst and react as if it is not there.
Since we know from previous science courses that a catalyst speeds up a reaction, we can conclude that the lower the activation energy, the faster the reaction.

22. Catalysts
Example 4: Phosgene, COCl2, one of the poison gases used during World War I, formed from a free radical* form of chlorine (non- diatomic Cl) and carbon monoxide. The mechanism is thought to proceed by:
a. Write the overall reaction equation.
b. Identify any reaction intermediates.
c. Identify any catalysts.

23. Catalysts
Example 4: Phosgene, COCl2, one of the poison gases used during World War I, formed from a free radical* form of chlorine (non-diatomic Cl) and carbon monoxide. The mechanism is thought to proceed by:
a. Write the overall reaction equation.
CO + Cl2  COCl2
b. Identify any reaction intermediates.
COCl
c. Identify any catalysts. Cl

24. *Side note: “Free Radicals”
Atoms, molecules or ions that have unpaired valence electrons or an open shell (non-full/non-octet valence shell) configuration
Very unstable
Play a role as catalysts, commonly in biochemical processes
The free radical Cl is produced by subjecting diatomic chlorine to ultraviolet (UV) light:
Cl2  2Cl or
This is not something we will be studying in further detail, so no need to worry about it. I just wanted to explain why Cl was in that form in that mechanism

25. Putting it all together: More Complex Energy Curves
Example 5: We have typically been simplifying our potential energy curves somewhat; for multistep reactions, potential energy curves are more accurately shown with multiple peaks. Each peak represents the activated complex for an individual step.
Consider the PE curve for a two-step reaction on the next slide…

26. Example 5
So then…

27. Example 5
a. -20 kJ

28. Example 5
b. +20 kJ

29. Example 5
c. -40 kJ

30. Example 5
d. +20 kJ

31. Example 5
e. +80 kJ

32. Example 5
f. +40 kJ

33. Example 5
g. Step 1 – higher
activation energy
than step 2

34. Example 5
h. +60 kJ

35. Example 5
i. exothermic – overall
∆H is negative (-20 kJ)