• this eliminates a lot of interesting and important reactions but heterogeneous reactions are more complicated to study
  • new variables include the surface area, the rate of mixing

The questions we are likely to probe experimentally are

• 1. what is the rate law (order) for the reaction?
  • how the rate depends on the concentration of reactants?
• 2. how does the reaction depend on temperature?
  • can we evaluate the Energy of Activation for the reaction?
• 3. are there catalysts that affect the rate of this reaction ?

• 1. What are the actual steps in the reaction mechanism?
• 2. Are there chemical modifications worth trying--
  • could we change the rate of reaction to meet our goals?
  • example: the setting time of glue (too short and the glue sets before the parts are assembled; too long and we wait forever to assemble the object.)

• The first consideration is the approximate time scale of the reaction
  • Reactions that take 10 minutes to 4 hours are pretty easy
    • we start the reaction by mixing the reactants
    • we keep this mixture at a constant temperature
    • at regular intervals we remove a sample
    • we analyze the sample by any available method
      • spectrophotometer, titration, pH meter, ...
    • we collect and plot the data (concentration vs. time)
  • Faster reactions need a little more attention
    • the time used to mix the sample needs to be small compared to the reaction time
    • the time spent collecting a sample and analyzing it needs to be a very small fraction of the reaction time
    • for a reaction that takes about a minute, we'd need to perform these tasks in about 1-2 second
  • reactions that take place in less than a second will require some special tricks
    • fast ways to mix the samples
    • there's not enough time to take out samples, so we need to make very quick measurements in the reaction cell itself
    • [in our department we have tools that can let us look at chemical reactions that take place in 1 msec, 1 microsec, 1 nsec, 1 psec and we have reached the femtosecond region.]
  • Slower reactions are often equally difficult to study
    • reactions that occur over months call for very long periods of study
      • important cases would include the aging of paint, corrosion of metals, osteoporosis (the gradual loss of bone as we age), the elimination of persistent pesticides from soil.
    • it's often difficult to keep conditions the same during the reaction.
      • reactions that normally take years or centuries are almost impossible to observe directly.
      • geologists often are faced with such questions
      • radioactive decay and carbon dating fits such a situation

• Experimental Work
  
  This experiment will extend over two weeks. This week some of you will use spectrometers to follow a kinetics experiment. others will use them the following week. The rest of you will study a clock reaction. In the following week, you switch. This allows us to make good use of a limited number of spectrometers. .
  

• We will use a spectrophotometer to follow the progress of a reaction
  • this limits us, in practice, to reactions that involve either a colored reactant or a colored reaction product.
  • this usually limits us to species at very low concentrations
• our spectrophotometers will respond to changes that occur in 1 second, so we can study reactions that take 1-15 minutes by simply leaving the sample in the spectrometer.
• the digital spectrometers are able to report the absorbance to a computer. So, lazy chemists that we are, we can let the machine collect and plot the data
  • This also allows us to handle faster reactions (people need about 5-10 seconds to read and record a date point)
  • This is also a little more reliable (people tend to lose focus and begin making mistakes after they've measured 20-30 data points. They also tend to make more mistakes when they copy and plot the numbers.)
    

You will work with one of the reactions below--

Detailed Experimental Directions will be provided at lab time.:

• reaction #1-- The loss of color from the reaction of Fe (phenathroline)₃2+
  
  • (This is the same species we formed last week, when we measured Fe spectrophotometrically)
    • The loss of color means Fe(phen)₃2+ ----> Fe²⁺ + 3 phen
      • (It is not obvious that all 3 phen are lost, but we will treat it that way)
    • The reaction only occurs when something else binds up the phen; otherwise, it would simply reattach itself to the iron.
    • Two reactions we can study are
      • Fe(phen)₃2+ + 3H⁺ ----> Fe²⁺ + 3 H-phen
      • Fe(phen)₃2+ + Cu²⁺ ---> Fe2+ + Cu(phen)₃2+
      • One reference suggests that both of these reactions occur at the same rate because
        they all depend on the same key step: loss of the first phen by Fe(phen)₃2+
        

• Bromine is a colored species and it is easy to determine its concentration.
• We can adjust concentrations so this reaction occurs in 3-10 minutes, making it a good candidate for study
• We will have time to vary concentrations and determine the reaction order.

reaction #3-- The oxidation of Bromcresol Green

• Bromcresol green is a common acid base indicator; it is strongly colored and easily measured with a spectrophotometer
  • the base form is blue and the acid form is yellow
  • We will keep the pH >6.0 so the dye will remain in the blue (base) form.
• The dye reacts with bleach (hypochlorite ion) and turns yellow.
• We will monitor the course of this reaction with the Spectronic20.
  
• This is a true experiment, meaning we (instructor and students) don't know exactly what is happening and may need to find the conditions necessary to get meaningful results. (More on this later*.)

Remarks on Spectrophotometry studies of kinetics

• Spectrophotometry limits the concentrations we can study.
  • We usually need very low concentrations of strongly colored species
  • We need to limit the Absorbance to 0.02 to 1.0
  • we need to start the experiment near the larger concentration
    • The molar extinction coefficient of Fe(phen)₃2+ is about 11000
    • So the concentration is about 0.8/11000 = 7 x 10^-5 Molar
  • We can vary concentrations by a factor of 2 to 4 fold
    • varying concentration more makes the spectrophotometry less reliable
    • (we can use higher concentrations if we select wavelengths where the species
      absorbs less efficiently.)
• In most cases, the other reactants are present in higher concentrations
  • as a result, they change very little over the course of a reaction run.
  • To see the effect of concentration, repeat the run with different starting concentrations.
    • *The bromcresol green experiment may be an exception
    • the dye concentration is low to keep in the spectrometer's useful range
    • the hypochlorite concentration must also be kept low or the reaction
      is too fast to study.
    • this means that hypochlorite concentration decreases significantly during a run.
      therefore two reactants (dye and hypochlorite) are changing simultaneously
      too little bleach and the reaction stops early
    • too much bleach and the reaction is over before we get it into the machine
      The reference I'm using in planning this experiment doesn't seem to recognize
      this situation.
      

Clock Reactions

• The other type of Kinetics experiment uses simple glassware and is more dramatic to observe.
• The reaction is between I- (iodide ion) and S₂O₈2- (peroxydisulfate ion)
  • The reaction produces free Iodine, I₂ and SO₄2-, sulfate ion
  • The I₂ actually combines with I- to form the I₃- ion
    • We add starch, and there is a colorful reaction between starch and I₃-
    • the color is dark blue
• if we simply mixed I- and peroxydisulfate, the starch would immediately turn blue
  • it might get a bit darker as more I₂ is produced, but we wouldn't really see much
• the trick is to also add a fixed amount of thiosulfate,
  • thiosulfate reacts with I₂ preventing it from reacting with starch
  • as long as thiosulfate remains, there will be no I₂ collecting
  • as long as thiosulfate remains, the solution remains colorless
  • but, as soon as the thiosulfate is gone, I₂ collects, and the solution goes blue/black
    • the color change is abrupt-- 1-2 seconds, typically
  • This gives us one data point for this reaction
  • it tells us when a fixed amount of I₂ had been produced.
• This is an initial reaction method
  • we keep thiosulfate << other reactants
  • primary reaction will occur for 10-60 minutes
    • but the thiosulfate will be gone in 15-60 seconds
    • so we haven't used more that 2-4% of the starting materials
    • since concentrations have changed very little, the rate of reaction has stayed constant
• rate = [ ] / [time when color changes]
  • To find the order of the reaction-- we repeat using different starting concentrations
• To find the effect of temperature, we repeat a run at higher and lower temperatures.
• 
  

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