Essay: Rate and Order of a Reaction

Theory
Much of what surrounds us relies on a few, simple chemical concepts, one of them, whose rules dictate my experiment, is the concept of rate and order of chemical reactions. The question in which my work is based is: how these two things, rate and order of reactions, are affected by both a change in concentration and a change in temperature studying the reaction between solutions of potassium iodide and iron (III) chloride?
The rate of reaction refers to the change in concentration of the products and reactants of a chemical equation per time. When talking about the rate of reaction, the factor used to measure it is usually per unit of time which is typically expressed in seconds. As a time passes and a reaction moves along, reactants become products and therefore the concentration of the former decreases while the concentration of the later increases. Using this information, we can derive the following expressions to represent rate of reaction.
rate of reaction=(increase in product concentration)/(time taken) = (decrease in reactant concentration )/(time taken)
Provided that  represents the change in, R represents the concentration of reactants, and P represents the concentration of products, we can simplify the above equations as
rate of reaction=Δ[P]/t = – ∆[R]/t
Though the negative symbol for the last equation that indicates a decrease in concentration, the rate is still expressed as positive. Since it is concentration over time, the units for rate of reaction are mol dm-3 s-1. The graphs for the rate of reaction are not straight lines and can therefore be given only for a particular time period by drawing a tangent and measuring the slope. Since the shape of the curve is not straight, we can see that it is greater (steeper) at the start of the reaction and gradually slows down as time goes past. The reason for it being greater at the beginning is because this is the point where the reactant concentration is at its highest. Because of the difference in rates, the initial rates are the ones that are usually compared as they give insight into the effect of concentration on rate.
There are different ways of measuring the rate of the reaction and choosing the right one depends on each specific case. In most techniques, the concentration is not measured directly but through some sort of indicator of a change in it. Gas syringes or water displacement is used when measuring the change in volume of a gas. A balance can be used when there is a change in mass, quenching when working in titrating a reaction and a conductivity meter is typically employed when there is a change in electrical conductivity through the reaction.
For this particular experiment, it is most convenient to use a colorimeter or spectrometer as there is a change in color. This measures the change in absorption in the visible regions as a sample of the solution is put in a chamber through which a light of specific wavelength passes and its intensity by the reaction components is measured. As a result of the color of the solution increasing, the absorption decreases and a photocell measures this and the information is recorded on the computer or meter. Since the reaction is being read continuously, an absorbance versus time graph can be plotted and consecutively transformed into a concentration versus time graph using a standard curve based on known concentrations.
All of the theories and concepts that make up the study of chemical kinetics are referred collectedly as the kinetic-molecular theory of matter. One of the most important aspects of this is the fact that particles move randomly because of the kinetic energy they hold which means that not all the particles have the same kinetic energy simultaneously. This is directly proportional to the temperature as the average particle kinetic energy increases as the temperature increases.
We can show the average particle kinetic energy in a Maxwell-Boltzmann distribution curve which plots the number of particles (or the probability of one particle) and the kinetic energy of them.
(Source: webchem.net (not available anymore, Internet Archive), retrieved via the Internet Archive)
For particles to collide with each other and react, two things need to be met for the collision to be successful and lead to the creation of a product which are, energy of collision and geometry of collision.
The amount of energy needed for particles to react is called activation energy, Ea. It can be thought of a hill which only particles with enough kinetic energy can go past in the way to going from reactants to products. It can be represented in enthalpy diagrams in different ways for exothermic and endothermic reaction.
The value of activation energy depends on the reaction and it affects the rate of said reaction as well as the speed in which it is completed, with reactions with a higher activation energy taking more time and having a smaller reaction rate.
The second aspect that affects the rate of reaction is the geometry of collision and it requires the particles to have the correct orientation. There is an increased chance of having the right orientation if there are more collisions per second and this, in turn, is increased by a higher pressure (only in the case of a gas reaction), a higher temperature (more energy for the particles to move around and collide), reducing the particle size (a bigger surface area means more contact and therefore, more collisions per second), and a higher concentration (bigger molarity means more particles to collide).
The other property of the reaction that will be determined through the experiment is order of reaction which searches for the relationship between the concentration of the solution and the rate of reaction. To find the order of reaction in respect to one of the compounds on the chemical equation, you must have the rate of reaction of two trials where the concentration of the compound is changed while the rest of the variables remain the same. You must divide the different rates and concentration between each other and calculate the value of exponent that the concentration can be raised to for it to equal the rate of reaction. This must be calculated for each of the components of the reaction to explain the dependence of the reactant concentration on the rate of reaction.
Example:
Experiment A
B Initial
rate
1 1.00M 1.00M 2.43×10-3
2 2.00M 1.00M 1.02×10-3
Order of reaction of A
(2.43×10-3)/(1.02×10-3)=2.38 2.00/1.00=2 2x= 2.38, x  1.00  It is 1st order with respect to A
All of this calculations and results are important to determine how fast a reaction occurs over a period of time and how this is dependent on temperature, concentration of each reactant, pressure, and other aspects surrounding the experiment. This can be quite important in medical sciences such as figuring out the length and effects of certain medicines and treatments. It can also be applied in fields such as nuclear safety and environmental measures.
Method
Materials:
-Computer
-Colorimeter or Spectrometer
-Vernier computer interface (only if using a colorimeter)
-Three 250mL beakers
-Two 100 mL beakers
-Plastic cuvettes (at least 3)
-0.020M potassium iodide, KI, solution
-0.020M iron (III) chloride, FeCl3, solution in 0.10M HCl
-Distilled water
-Three 25 mL graduated cylinders
-Five plastic Beral pipets
-A kettle or a device to boil water
-Bucket or large container
-Thermometer
-Balance
Though this is a relatively safe lab, some chemicals like the hydrochloric acid (though it is not a very high molarity) do require certain measurements. Safety can be guaranteed by wearing protective googles, a lab coat, and gloves. It should be noted that the potassium iodide solution can stain clothing, surfaces, and equipment and so it must be handled with due care.
Procedure:
Part 1: Setting up the Colorimeter
Make sure to wear all protective gear.
Obtain all equipment stated under materials.
Prepare a blank cuvette by filling it ¾ full of distilled water. While handling cuvettes, the user must remember to
Wipe the outside of the cuvette with a lint-free tissue
Handle them by the top edge of the ribbed sides.
Dislodge any bubbles by tapping the cuvette on a hard surface
Position cuvette correctly (light must pass through clear sides)
Connect colorimeter to computer interface and prepare the computer for data collection by opening the file “25 Rate and Order” from the pre-existing Advanced Chemistry with Vernier folder of Logger Pro.
Open the colorimeter lid and insert the blank. Do not forget to close the lid.
Press the < or > button on the Colorimeter and select the 430 nm (blue) wavelength. Proceed to click the CAL button until the red light begins to flash. When the red flashing stops, the calibration is complete.
During the experiment you will conduct 5 solutions per each temperature between KI and FeCl3 using the following volumes.
Solution FeCl3 (mL)
(0.02M) KI (mL)
(0.02M) H2O (mL)
1 20.0 20.0 0.0
2 20.0 10.0 10.0
3 10.0 20.0 10.0
4 15.0 10.0 15.0
5 10.0 15.0 15.0
Conduct Solution 1 at room temperature (25C).
Measure 20.0 mL of FeCl3 solution into a 100 mL beaker.
Measure 20.0 mL of KI solution into a second 100 mL beaker.
Add the 20.0 mL of FeCl3 solution to the beaker of KI solution and swirl the beaker to mix.
Rinse the cuvette with less than 1 mL amount of the reaction mixture, then fill it ¾ full and place it in either the Colorimeter or Spectrometer. Make sure to close the lid if using a Colorimeter.
Click the box titled Collect. Though the default setting is 1 sample per 200 seconds, you may click the Stop box to collect early. Observe the progress of the reaction.
When the data is completely collected, carefully remove the cuvette from the respective device and dispose of its contents and those from the beaker as indicated. Rinse and clean both the cuvette and beaker for the next solution.
Examine the graph collected in the first solution and select a linear region in the first minute of the data collection to analyze. Click on the Linear Regression button to get the equation of the line of best fit and record the slope as the initial reaction rate for solution 1 at temperature 1 in your data table and clear the data.
Repeat the steps from 8 to 11 for the remaining four solutions. It works well to add the distilled water to the KI solution before the FeCl3 solution. Use the same range of the graph to calculate the initial reaction rates for the rest of the trials as you used for the first one.
Fill a water container with hot water big enough to hold the solution and measure the temperature of the water until it reaches 60C and then repeat the whole process described above while placing each beaker with the finished solution on the water bath for a few moments before putting a portion of it into the devices.
Repeat step 13 for the following temperatures: 50C and 45C.
Notes:
The KI might not always be available in liquid form and, if used in powder form, must be converted from milliliters to moles to grams and measured accordingly by using the balance.
Moles of KI= MV= (.02)(.020)= .0004mol .0004molx162.2g/mol= .064g
Moles of KI= MV= (.02)(.010)= .0002mol .0002molx162.2g/mol= .032g
Moles of KI= MV= (.02)(.015)= .0003mol .0003molx162.2g/mol= .048g
This might affect the particle size and therefore the result of the experiment.
The temperature of the water must always remain the same throughout the process, so it is recommended to use the smallest amount of water possible in the beginning so adding more will have a greater effect.
To record the data from the colorimeter, one must bring a USB which can be connected on the port of the computer.
If possible, a single cuvette should be used in order to eliminate errors induced by the slight variations in the absorbance of different