 This section is a continuation of the discussion of the factors that affect rates of reactions.  Today we will focus on the 3 rd & 4 th factors that influence the rate of reactions

 If the concentration of the reactants is increased the rate of reaction will also increase if the reactants are all in the same phase.  The rate will increase 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.

 As far as gases go, 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 is 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.  As products are formed from the reactants, the concentration of the reactants decreases, which is why the rate of reaction decreases as a given reaction proceeds.  Recall:

 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

 Example 1: Consider the following reaction that occurs between hydrochloric acid, HCl, and zinc metal: 2HCl(aq) + Zn(s) → H 2 (g) + ZnCl 2 (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

 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

 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? AB

 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. AB

 A catalyst is a substance that provides an alternate reaction mechanism with a lower activation energy.  Catalysts do not get used up in the reaction. That is, they are introduced as a reactant and show up as a product.  If we can lower the activation energy, the number of particles with enough energy to make successful collisions naturally increases.

 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.

 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.  Like a shortcut to school, the particles may or may not use the catalyst because the original route is still an option.  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.

 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 + Cl 2  COCl 2  b. Identify any reaction intermediates. COCl  c. Identify any catalysts. Cl

 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

 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…

 So then…

 a. -20 kJ

 b. +20 kJ

 c. -40 kJ

 d. +20 kJ

 e. +80 kJ

 f. +40 kJ

 g. Step 1 – higher activation energy than step 2

 h. +60 kJ

 i. exothermic – overall ∆H is negative (-20 kJ)